Transformation and immobilization of sedimental galena (PbS) by phosphate from surface runoff in simulated storm suspensions

Abstract

Sedimental galena (PbS) is an important sink for lead in natural waters. Resuspension of PbS caused by storm disturbance can lead to its oxidative dissolution, resulting in higher ecological risks. Surface runoff resulting from intensive storms can also carry phosphate into receiving water bodies. It is hypothesized that phosphate from surface runoff can regulate Pb concentration and speciation during storm events. To test the hypothesis, the dissolution and transformation of PbS were investigated in the absence and presence of orthophosphate under different pH (5–8) and dissolved oxygen (0–8.4 mg L⁻¹) conditions to simulate the behaviors of suspended PbS. The results indicated that the kinetics of PbS dissolution was mainly controlled by pH while dissolved oxygen played a minor role when orthophosphate was absent. In the presence of orthophosphate (0.5 mg-P L⁻¹), the soluble lead concentration as high as 990 ppb resulting from PbS dissolution decreased immediately to ND (<5.1 μg L⁻¹), except at pH 5, due to the formation of pyromorphite. A saturation index (SI) greater than 16.16 was required to initiate pyromorphite precipitation. The results suggested that phosphate, which is often associated with eutrophication, could sequester soluble lead and reduce associated ecological risks resulting from sedimental PbS dissolution in storm events.

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Water impact

Sedimental galena (PbS) is an important sink for Pb in natural waters but its oxidative dissolution during storm events can release Pb. This study investigates the kinetics of PbS dissolution in simulated storm suspensions and explores whether phosphate in surface runoff can immobilize released Pb. It was found that the kinetics of PbS oxidative dissolution was mainly regulated by pH and soluble Pb decreased immediately after phosphate addition due to pyromorphite precipitation. Although phosphate is often associated with eutrophication, it could sequester soluble Pb and reduce ecological risks resulting from sedimental PbS dissolution in storm events.

1. Introduction

Sediments resulting from weathering and erosion are the habitats of many aquatic organisms.^1,2 However, heavy metal pollution caused by anthropogenic activities has negative impacts on the quality of sediments and has become a threat to ecosystems and human health.^3–6

Sediments are typically anoxic and reduced species such as sulfide are commonly present. The concentrations of heavy metals in the sediments could surpass those in the overlying water by several orders of magnitude because metal ions can quickly precipitate with sulfide to form low solubility metal sulfides, which are stable and considered as the major sink for heavy metals in anoxic sediments.^7–9 It has been suggested that the bioavailability and toxicity of heavy metals in sediments can be predicted using the molar ratio of simultaneous extraction metals (SEM) to acid volatile sulfides (AVS).^3,7,10 If [SEM]/[AVS] is less than 1, stable metal sulfides can form and reduce the acute toxicity for benthic organisms. On the other hand, if [SEM]/[AVS] is greater than 1, no sufficient sulfide is available and metal ions can become bioavailable, thereby increasing the toxicity to sensitive benthic species.^4,11 It should be noted that the dissolution of metal sulfides is also affected by water chemistry such as pH, dissolved oxygen (DO) and other complexing agents.^12–15 Therefore, the influences of water chemistry on this process should be considered in addition to the [SEM]/[AVS] ratio.

The dissolution of metal sulfides can occur due to resuspension of sediments resulting from turbulence or bioturbation, leading to the release of toxic metal ions into oxic water columns.^14,16,17 For example, lead levels in water and suspended sediment can fluctuate by 10-fold during the dredging of contaminated sediment.^17,18 A significant release of lead from contaminated creek and marine sediments due to resuspension has also been reported.^19,20 Such events can elevate lead levels in water that cause adverse effects on the metabolism, survival, growth, development and reproduction of fish, wildlife and invertebrates.²¹

In areas affected by human activities, phosphate can be easily flushed into water bodies by surface runoff during storm events.^22–25 The total phosphate level in the first flush runoff can be up to several mg L⁻¹.^26–28 In typical surface waters, phosphate is considered as the limiting nutrient regulating algae growth and the main contributor to the eutrophication of water bodies.^29–31 The presence of phosphate, however, may lead to the formation of low solubility phosphate minerals that sequester toxic metal ions. For example, addition of phosphate has been considered as an effective remediation technology to reduce the mobility of lead in soils and drinking water by forming pyromorphites including Pb₅(PO₄)₃Cl and Pb₅(PO₄)₃OH according to the following reactions: 5 Pb²⁺ + 3 PO₄³⁻ + Cl⁻ → Pb₅(PO₄)₃Cl(s); 5 Pb²⁺ + 3 PO₄³⁻ + OH⁻ → Pb₅(PO₄)₃OH(s).^32–36

Given the increasing frequency and intensity of storm events driven by climate change, disturbances of anoxic sediments and the flushing of phosphate through surface runoff are expected to occur more often. It is therefore hypothesized that phosphate transported by surface runoff may sequester lead ions released from the oxidative dissolution of sedimental PbS during such storm events. The coupled processes of PbS dissolution and pyromorphite precipitation could ultimately control the concentration of toxic lead ions in the water column. However, these interactions have not yet been thoroughly investigated.

The objectives of this study are to verify the above hypothesis and explore the fate of sedimental PbS during storm events. Specifically, the kinetics of PbS dissolution under different DO and pH conditions is examined. The influences of phosphate on the immobilization of lead and transformation of lead species during PbS dissolution are explored. Overall, the insights into the potential influences of phosphate-containing surface runoff on the lead level and the fate and transformation of sedimental PbS in storm suspensions are provided in this study.

2. Materials and methods

2.1 Chemicals and solution preparation

PbS (99.9%) was purchased from Sigma-Aldrich (USA). KCl (>99.0%), NaHCO₃ (>99.7%), CaCl₂·2H₂O (>99.0%) and MgSO₄·7H₂O (>98.0%), all purchased from Merck (Germany), were used to prepare experimental solutions. Organic buffers including 2-(N-morpholino)ethanesulfonic acid (MES) (>99.5%, Merck, Germany) and 3-morpholinopropane-1-sulfonic acid (MOPS) (>99.5%, Merck, Germany) were used for pH control as they do not form complexes with Pb ions.³⁷ 10 mM MES was used to buffer the pH for experiments at pH 5 and 6. 5 mM MOPS was used for those at pH 7 and 8. 1 M HCl and 1 M NaOH (Honeywell Fluka, USA) were used for pH adjustments. 69% HNO₃ (J.T. Baker, USA) was used to acidify samples. Combined reagents composed of ascorbic acid (>99.0%, Merck, Germany), H₂SO₄ (Honeywell Fluka, USA), K(SbO)C₄H₄O₆·½H₂O (99–103%, Acros Organics, USA) and (NH₄)₆Mo₇O₂₄·4H₂O (81–83%, J.T. Baker, USA) were used in the ascorbic acid method for phosphate measurement. KH₂PO₄ (>99.5%, Nacalai Tesque, Japan) was used to prepare the stock orthophosphate solution (100 mg-P L⁻¹). All solutions were prepared using deionized water generated from a PURELAB Classical system (ELGA, UK).

2.2 Experimental apparatus and methods

The schematic and photograph of the batch aeration setup used in this study are shown in Fig. S1. 1000 mL amber serum bottles were used as the reactors. Synthetic freshwater (Table S1) was prepared as the experimental solution.³⁸ To simulate different DO concentrations of the water column, different partial pressures of pure N₂ and O₂ were mixed using a mass flow controller (MFC, Protec Instrument Corp., USA). The air mixture was passed through an O₂ sensor (PrimaX® I Gas Transmitter, MSA, The Netherlands) before passing through ultra-pure water to minimize the decrease in the solution volume caused by contacting dry gas. The air mixture was purged into the experimental solution with a flow velocity of 3 L min⁻¹ until the DO concentration reached the target level. The flow velocity was then adjusted to 1 L min⁻¹ to keep the desired partial pressure of O₂ in the headspace to retain a steady DO concentration during the experiments.

After the addition of 100 mg L⁻¹ of PbS, the solution was stirred with a magnetic bar at 200 rpm to suspend the particles. The experiments were conducted at different pH (5, 6, 7 and 8) and DO concentrations (0, 3, 5 and 8.4 mg L⁻¹) for 72 h. The PbS concentration and experimental period employed were to simulate the sulfide level in the sediments^17,18 and the suspension period in storm events.³⁹ Different concentrations of orthophosphate (0.25, 0.50, 1 and 2 mg-P L⁻¹) were added to simulate the surface runoff with different phosphate levels. A 9 mL sample was withdrawn and filtered through a 0.22 μm pore size nylon syringe filter at designated time intervals. 4.5 mL of the filtrate was acidified to a pH < 4 with 69% HNO₃ and stored at 4 °C before soluble Pb(II) analysis. The other 4.5 mL filtrate was used for orthophosphate measurement. A polynomial regression was used to fit the results of soluble Pb(II) release versus time. The initial rate of PbS dissolution was determined from the initial slope of the polynomial regression curve at time zero. After the 72 h experiment, the solids in the reactors were collected by filtering the solution through a 0.22 μm pore size nylon membrane and dried at 80 °C in an oven for 24 h. The solids were stored in a dry cabinet prior to scanning electron microscopy (SEM) analysis. All experiments were conducted in duplicate and the error bar represents the range of duplicate data.

2.3 Analytical methods

The surface morphology of the solids was analyzed using a Field-Emission Scanning Electron Microscope (FE-SEM, ZEISS ΣIGMA Essential, Germany). A Brunauer–Emmett–Teller surface area analyzer (Micromeritics ASAP 2420, USA) and Particle Size Analyzer (CILAS 1090L, France) were used to determine the specific surface area and particle size distribution of PbS, respectively. The Pb concentration was measured using an inductively coupled plasma optical emission spectrometer (ICP-OES, Agilent 700 series, USA), with a method detection limit (MDL) of 5.1 μg L⁻¹. The solution pH was measured using a pH meter (SUNTEX sp2100, Taiwan) equipped with an Ag/AgCl electrode (Sensorex SG200C, Taiwan) calibrated using pH 4, 7, and 10 standard buffer solutions. The DO concentration was measured using a DO meter (Thermo Scientific, USA) equipped with a membrane electrode calibrated using air saturation and zero DO solutions. The orthophosphate concentration was measured using the ascorbic acid method following the Standard Method 4500-P E.⁴⁰

3. Results and discussion

3.1 Kinetics of PbS dissolution under different DO and pH conditions

The PbS used in this study possessed a specific surface area of 0.9 m² g⁻¹ and a mean size of 5.9 μm (Fig. S2). The dissolution of PbS under different DO and pH conditions was first examined to establish the dissolution rate law of PbS and the results are shown Fig. 1. The initial dissolution rates calculated from the initial slopes of the soluble Pb(II) vs. time curves are summarized in Table S2. For the DO levels investigated (DO = 0–8.4 mg L⁻¹), the released soluble Pb(II) levels were similar, particularly in the first 5 min (Fig. 1(a)). This result suggested that the dissolution of PbS was not significantly affected by the DO level, especially at the initial stage. For the pH investigated (pH 5–8), the rate of PbS dissolution increased with decreasing pH, implying that pH was the dominant factor regulating PbS dissolution (Fig. 1(b)). The strong influence of pH and the weak influence of DO on the dissolution of PbS were consistent with the results reported in the literature.^41–43

Fig. 1 Effects of (a) DO (pH fixed at 7) and (b) pH (DO fixed at 8.4 mg L⁻¹) on the dissolution of PbS.

It has been proposed that the dissolution of PbS proceeds via the following steps:⁴¹

1. Protonation of the PbS surface due to the adsorption of H⁺:

≡PbS + 2H⁺ → ≡PbSH₂²⁺ (1)

2. In the absence of DO, ≡PbSH₂²⁺ dissolves directly to release Pb²⁺:

≡PbSH₂²⁺ → Pb²⁺ + H₂S (2)

3. In the presence of DO, O₂ adsorbs onto ≡PbSH₂²⁺, followed by the oxidation of surface S²⁻ to SO₄²⁻ and the release of Pb²⁺:

≡PbSH₂²⁺(s) + 2O₂ → ≡PbSH₂²⁺–2O₂ (3)

≡PbSH₂²⁺–2O₂ → Pb²⁺ + SO₄²⁻ + 2H⁺ (4)

Based on the results in Fig. 1, it can be deducted that the rate of PbS dissolution is mainly controlled by proton attacks (eqn (1) and (2)), while that induced by DO attacks (eqn (3) and (4)) played a minor role. The initial dissolution rate can then be fitted to eqn (5) considering the overall influence of [H⁺] on the adsorption of H⁺ to form ≡PbSH₂²⁺ and its subsequent dissolution:

r = d[Pb(II)]/dt = k[H⁺]ⁿ (5)

where r is the initial dissolution rate of PbS (M min⁻¹), [Pb(II)] is the soluble Pb(II) concentration (M), t is the reaction time (min), k is the apparent rate constant (M¹⁻ⁿ min⁻¹), [H⁺] is the proton concentration (M), and n is the reaction order with respect to [H⁺].

Based on the least-square regression, k and n were determined to be 1.08 × 10⁻⁵ M^0.67 min⁻¹ and 0.33, respectively. Consequently, the initial dissolution rate law of PbS can be described as follows:

r = 1.08 × 10⁻⁵ [H⁺]^0.33 (6)

The obtained reaction order with respect to [H⁺] was slightly smaller than those determined by Cama et al.⁴⁴ (n = 0.56 at pH 3 and 25 °C), De Giudici et al.⁴⁵ (n = 0.30–0.60 at pH 1.2–5.8 and 25–75 °C), and Zhang et al.⁴⁶ (n = 1.0 at pH 0.43–2.45, 25–75 °C and 1 M NaCl) under more acidic conditions. It has been reported that the reaction order with respect to [H⁺] could vary at different pH ranges due to different governing mechanisms, in which the diffusion-controlled kinetics dominates at pH < 3 and both surface reaction and diffusion-controlled kinetics are important at a higher pH, such as the range employed in this study.⁴⁵ To verify the rate law, additional experiments were conducted at pH = 5.5, 6.5 and 7.5 in the presence of 8.4 mg L⁻¹ of DO. The results showed that the modeled initial rates agreed well with the measured initial rates (Fig. 2), suggesting that eqn (6) can be used to describe the rate of PbS dissolution observed in this study.

Fig. 2 Comparison between modelled and measured initial dissolution rates.

The dissolution rate of PbS decreased gradually with increasing time. De Giudici and Zuddas⁴⁷ used atomic force microscopy to investigate the changes in surface microtopography during PbS dissolution. They found that PbS dissolution could result in the formation of protrusions, which were composed of secondary minerals such as lead oxydihydroxide (Pb₂O(OH)₂) and lead oxide (PbO) under their experimental conditions. The accumulation of protrusions could attenuate the PbS surface reactivity and hinder PbS dissolution. Visual MINTEQ 3.1 was used to determine the potential mineral supersaturation at the end of our experiments (Tables S3–S6). The results showed that no supersaturation was observed at pH 5 and 6 (Tables S3 and S4). At pH 7, hydrocerussite (Pb₃(CO₃)₂(OH)₂) and cerussite (PbCO₃) were slightly supersaturated with a saturation index of 0.72 and 0.25, respectively (Table S5). At pH 8, Pb₃(CO₃)₂(OH)₂ is mostly supersaturated, followed by Pb(OH)₂ and PbCO₃, with an SI of 1.08, 0.51 and 0.03, respectively. These results indicated that it was thermodynamically possible to precipitate these secondary minerals at the end of the experiments. The potential changes in surface morphology due to the formation of secondary minerals as examined by SEM are discussed later.

3.2 Effects of orthophosphate addition on PbS dissolution

It has been shown that the first flush in the agricultural area reaches the receiving water bodies in the first few hours of the rainfall events.^48,49 Therefore, following the experiments conducted under different DO and pH conditions, 0.5 mg-P L⁻¹ of orthophosphate was added at 180 min to simulate the influences of phosphate-containing surface runoff on the release of soluble Pb(II), and the results are shown in Fig. 3.

Fig. 3 Effects of orthophosphate addition (0.5 mg-P L⁻¹) on the dissolution of PbS in the presence of different (a) DO (pH fixed at 7) and (b) pH (DO fixed at 8.4 mg L⁻¹). Orthophosphate was introduced at 180 min.

For the experiments with different DO, soluble Pb(II) decreased immediately to ND (<5.1 μg L⁻¹) for all DO levels after orthophosphate addition (Fig. 3(a)). The period of ND lapsed for about 720 min for DO = 3.0, 5.0 and 8.4 mg L⁻¹. For DO = 0 mg L⁻¹, soluble Pb(II) remained as ND till the end of the experiment. It has been proposed that PbS could be passivated by pyromorphite grown on its surfaces so that dissolution of PbS was hindered.⁵⁰ Visual MINTEQ 3.1 was used to examine the degree of supersaturation with respect to different lead minerals after orthophosphate addition (Tables S8–S10). The solutions were highly supersaturated with respect to chloropyromorphite (Pb₅(PO₄)₃Cl) (SI = 19.86–20.88) and hydroxypyromorphite (Pb₅(PO₄)₃OH) (SI = 8.54–9.56). The solutions were also supersaturated with respect to Pb₃(PO₄)₂ (SI = 4.65–5.26), Pb₃(CO₃)₂(OH)₂ (SI = 4.65–5.26), PbCO₃ (SI = 0.53–0.73), PbHPO₄ (SI = 0.29–0.49) and Pb(OH)₂ (SI = 0.01–0.21). Given the much higher SI, the two pyromorphites were more thermodynamically favorable to precipitate.

For the experiments conducted at different pH levels, soluble Pb(II) decreased immediately to ND for all pH conditions except pH 5. The period of ND for pH = 6, 7 and 8 also lapsed for about 720 min, followed by a gradual increase of soluble Pb(II), in which the rate of increase was higher at a lower pH. Visual MINTEQ 3.1 was used to examine the degree of supersaturation with respect to different lead minerals after orthophosphate addition (Tables S11–S14). At pH < 6, only lead phosphate minerals, including Pb₅(PO₄)₃Cl, Pb₅(PO₄)₃OH and Pb₃(PO₄)₂ were supersaturated. At pH > 6, in addition to the three lead phosphate minerals, the solutions were also supersaturated with the respective lead carbonate minerals, i.e., Pb₃(CO₃)₂(OH)₂ and PbCO₃. However, the SI was much higher for the two pyromorphites, in which the SI values for Pb₅(PO₄)₃Cl and Pb₅(PO₄)₃OH were 10.99–19.46 and 3.84–8.14, respectively. It should be noted that the gradual increase of soluble Pb(II) after the period of ND under different DO and pH conditions could be due to two reasons: 1. the orthophosphate was completely consumed, therefore no orthophosphate was available to suppress the release of soluble Pb(II) and 2. the residual orthophosphate was not high enough to induce precipitation, in which a meta-saturated status with respect to pyromorphites may exist. This was further examined using different concentrations of orthophosphate shown in the following section.

3.3 Effects of orthophosphate concentrations on PbS dissolution

Different concentrations of orthophosphate (0.25–2.0 mg-P L⁻¹) were added at 180 min to simulate the entry of surface runoff with different strengths. The calculations using Visual MINTEQ 3.1 indicated that the solutions were highly supersaturated with respect to the two pyromorphites, Pb₅(PO₄)₃Cl and Pb₅(PO₄)₃OH, after orthophosphate was introduced (Tables S15–S18).

The profiles of soluble Pb(II) as a function of time are shown in Fig. 4. After orthophosphate addition, soluble Pb(II) decreased immediately to ND (<5.1 μg L⁻¹) for all orthophosphate levels, and then a gradual increase was observed. The periods of ND lapsed for orthophosphate = 0.25, 0.5, 1, and 2 mg-P L⁻¹ were 360, 720, 1440 and 2880 min, respectively. The variations of the orthophosphate concentration as a function of time for 0.25, 0.5 and 1.0 mg-P L⁻¹ were also determined and the results are shown in Fig. S3. For the period when soluble Pb(II) was ND, there was always a sufficient level of orthophosphate in the solution. For example, when 0.25 mg-P L⁻¹ was introduced at 180 min, during the period of ND soluble Pb(II) (180–360 min), the orthophosphate concentration was greater than 0.10 mg-P L⁻¹. The orthophosphate concentration continued to decrease and a gradual increase of soluble Pb(II) was then observed. The same trend applied when 0.5 and 1.0 mg as P L⁻¹ of orthophosphate were introduced (Fig. S3(b) and (c)). Therefore, it can be concluded that the gradual increase of soluble Pb(II) after ND could be attributed to the complete consumption of orthophosphate.

Fig. 4 Effects of orthophosphate concentrations on the dissolution of PbS. Orthophosphate was introduced at 180 min. pH 7, DO = 8.4 mg L⁻¹.

3.4 Changes in surface morphology

The SEM images of PbS before and after 60 min of the dissolution experiment without the addition of orthophosphate under different DO and pH conditions are shown in Fig. 5. Before the experiment, PbS displayed a cubic structure (Fig. 5(a)), which was consistent with the characteristic shape of PbS reported in the literature.^51,52 The SEM images of PbS after dissolution under DO = 0, 5 and 8.4 mg L⁻¹ at pH 7 are shown in Fig. 5(b)–(d) and those under DO = 8.4 mg L⁻¹ at pH 5 and 8 are shown in Fig. 5(e) and (f), respectively. Based on the Visual MINTEQ 3.1 analysis, PbCO₃ and Pb₃(CO₃)₂(OH)₂ were found to be supersaturated in these experiments except at pH < 6. The highest SI values with respect to PbCO₃ and Pb₃(CO₃)₂(OH)₂ were 0.42 and 1.24, respectively. PbCO₃ possesses a prismatic shape and Pb₃(CO₃)₂(OH)₂ possesses a hexagonal plate shape.^53,54 No newly formed particles with these shapes were observed in the SEM images, suggesting that PbCO₃ and Pb₃(CO₃)₂(OH)₂ did not precipitate and a meta-saturated status existed for these two lead carbonate minerals.⁵⁵ These SEM results imply that when the sedimental PbS is disturbed, the dissolution of PbS could release substantial soluble Pb(II) which is more bioavailable than particulate Pb for a certain period of time and poses a higher ecological risk to the aquatic environment.

Fig. 5 SEM images of PbS (a) before the experiment and after the experiment for (b) DO = 0 mg L⁻¹, pH 7, (c) DO = 5 mg L⁻¹, pH 7, (d) DO = 8.4 mg L⁻¹, pH 7, (e) DO = 8.4 mg L⁻¹, pH 5, and (f) DO = 8.4 mg L⁻¹, pH 8. Reaction time = 60 min.

The SEM images of PbS after the dissolution experiments in the presence of orthophosphate under different DO and pH conditions are shown in Fig. 6 and S4. The two pyromorphites (Pb₅(PO₄)₃Cl or Pb₅(PO₄)₃OH) are needle-shaped crystals with a hexagonal structure.^35,56,57 Needle-shaped crystals were observed when DO > 0 and pH > 6 for all concentrations of orthophosphate employed (0.25–2 mg-P L⁻¹). It was not possible to distinguish between Pb₅(PO₄)₃Cl and Pb₅(PO₄)₃OH due to their similar crystal shapes.^54,58 However, the much higher SI for Pb₅(PO₄)₃Cl (Tables S7–S18) leads us to believe that the needle-shaped crystals observed were more likely to be Pb₅(PO₄)₃Cl. It should be noted that since needle-shaped pyromorphites were not observed at pH 6 with an SI = 16.16 (Table S12), the solution exhibited a meta-saturated status, indicating that a higher degree of supersaturation would be required to initiate the precipitation. Further confirmation of the mineralogy of the needle-shaped crystals using XRD, unfortunately, was not possible due to limited amounts of the minerals in the samples.

Fig. 6 SEM images of PbS after the experiment for (a) DO = 0 mg L⁻¹, pH 7, P = 0.5 mg L⁻¹, (b) DO = 8.4 mg L⁻¹, pH 7, orthophosphate = 0.5 mg-P L⁻¹, (c) DO = 8.4 mg L⁻¹, pH 5, orthophosphate = 0.5 mg-P L⁻¹, (d) DO = 8.4 mg L⁻¹, pH 8, orthophosphate = 0.5 mg-P L⁻¹, (e) DO = 8.4 mg L⁻¹, pH 7, orthophosphate = 0.25 mg-P L⁻¹ and (f) DO = 8.4 mg L⁻¹, pH 7, orthophosphate = 2 mg-P L⁻¹. Reaction time = 72 h.

Based on the SEM observations and the soluble Pb(II) measurements, it can be deduced that orthophosphate in surface runoff during storm events could sequester soluble Pb(II) resulting from sedimental PbS dissolution and reduce the associated ecological risks due to pyromorphite precipitation.

4. Conclusions

The effects of orthophosphate on PbS dissolution and transformation under different DO and pH conditions during simulated storm suspensions were investigated in this study. In the absence of orthophosphate, the rate of PbS dissolution was mainly controlled by proton attacks and increased with decreasing pH. After orthophosphate addition, the soluble Pb(II) resulting from PbS dissolution decreased immediately to ND (<5.1 μg L⁻¹) except at pH 5 due to the precipitation of pyromorphites. The soluble Pb(II) then gradually increased because of the consumption of orthophosphate. A higher concentration of orthophosphate could prolong the period of ND soluble Pb(II), in which an SI greater than 16.16 was required to initiate the precipitation of pyromorphites. The results of this study implied that sedimental PbS resuspended in storm suspensions could dissolve and elevate the soluble Pb(II) levels in the water column. However, phosphate-containing surface runoff resulting from the storm, especially in agricultural areas, could sequester soluble Pb(II) by forming low solubility pyromorphites to reduce Pb-associated ecological risks, although phosphate is generally considered as the main nutrient contributing to eutrophication.

Several limitations still exist given the complexity of natural waters. For example, the dynamic roles of natural organic matter (NOM) in PbS dissolution and pyromorphite precipitation was not considered. NOM is known to adsorb on the mineral surfaces to protect them from proton or oxygen induced dissolution,^59,60 as well as to initiate ligand-induced dissolution and metal complexation to enhance the solubility of the minerals.^9,61 In addition, other metals such as Cu and Cd that were not investigated in this study may compete and consume orthophosphate.⁶² Therefore, the influences of NOM and other metals on the dissolution and transformation of PbS should be further investigated. Finally, field studies in water bodies experiencing lead contamination and surface runoff especially in agricultural areas could be conducted to verify the findings of this study.

Author contributions

Yi-Pin Lin: conceptualization, formal analysis, writing – original draft and review & editing, funding acquisition, and supervision. Ze-Xuan Tan: methodology, investigation, and formal analysis.

Conflicts of interest

There are no conflicts to declare.

Data availability

All data included in this study are available upon request by contact with the corresponding author.

Supplementary information (SI): synthetic freshwater composition. The initial dissolution rates of PbS under different DO and pH. Lead mineral saturation index at PbS = 100 mg L⁻¹, pH 5–8, DO = 8.4 mg L⁻¹ and t = 60 min. Lead mineral saturation index right after phosphate addition (t = 180 min): PbS = 100 mg L⁻¹, pH 7, DO = 0–8.4 mg L⁻¹ and 0.5 mg-P L⁻¹ phosphate. Lead mineral saturation index right after phosphate addition (t = 180 min): PbS = 100 mg L⁻¹, pH 5–8, DO = 8.4 mg L⁻¹ and 0.5 mg-P L⁻¹ phosphate. Lead mineral saturation index right after phosphate addition (t = 180 min): PbS = 100 mg L⁻¹, pH 7, DO = 8.4 mg L⁻¹ and 0.25–2 mg-P L⁻¹ phosphate. The schematic and photograph of the experimental setup. Particle size distribution of PbS. The variations of the orthophosphate concentration as a function of time for the addition of 0.25, 0.5 and 1.0 mg L⁻¹ as P of orthophosphate. SEM images of PbS after the experiment for (a) DO = 3.0 mg L⁻¹, pH 7, P = 0.5 mg L⁻¹, (b) DO = 5.0 mg L⁻¹, pH 7, P = 0.5 mg L⁻¹, (c) DO = 8.4 mg L⁻¹, pH 6, P = 0.5 mg L⁻¹, (d) DO = 8.4 mg L⁻¹, pH 7, P = 0.5 mg L⁻¹, (e) DO = 8.4 mg L⁻¹, pH 7, P = 0.5 mg L⁻¹ and (f) DO = 8.4 mg L⁻¹, pH 7, P = 1 mg L⁻¹. Reaction time = 72 h. P indicates orthophosphate added at 180 min. See DOI: https://doi.org/10.1039/d5ew00329f.

Acknowledgements

This work was supported by the NTU Research Center for Future Earth from The Featured Areas Research Center Program (113 L901002) and the Excellence Research Program - Core Consortiums (NTU-CC-114 L892601), within the framework of the Higher Education Sprout Project by the Ministry of Education (MOE) in Taiwan.

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