Sustainable Recovery of Lead from Secondary Waste in Chloride Medium: A Review

Abstract

Environmental regulations on lead recycling are becoming increasingly stringent, prompting the search for sustainable alternatives to conventional high-temperature processes. Hydrometallurgical methods in chloride media have emerged as a viable option for recovering lead from mining and urban wastes, including lead anode corrosion residues, zinc leaching residues, and spent lead–acid batteries. This study reviews the key conditions for lead recovery in chloride media, highlighting the variables that optimize lead dissolution, and the potential challenges associated with these processes. The findings indicate that efficient lead recovery requires high chloride concentrations, with acidity playing a critical role depending on the relative concentrations of lead and sulfate in the solution. When lead and sulfate concentrations are similar, stable lead–chloride complexes form within a pH range of 0 to 6.0. However, at higher sulfate concentrations, the pH range narrows significantly to 0 to 2.0, necessitating a more acidic environment for effective dissolution. Chloride media offer a distinct advantage through the formation of stable lead–chloride complexes, whose stability is influenced by chloride concentration, sulfate concentration, pH, and redox potential. Moreover, this approach provides a sustainable alternative that could integrate seawater into industrial processes, particularly in regions facing water scarcity.

1. Introduction

Lead is a heavy metal pollutant and is hazardous to human health [1]. It naturally occurs in minerals such as galena (PbS), anglesite (PbSO₄), and cerussite (PbCO₃) and is also present in various mining and urban residues. These include corrosion products from lead anodes in copper and manganese electrowinning [2], zinc leaching residues [3,4], and urban waste such as spent lead–acid batteries [5,6], solder materials [7], and solar panels [8]. The improper management of these wastes can result in significant environmental harm, such as soil contamination [9]. Moreover, prolonged exposure to lead represents a serious health hazard, including poisoning and damage to the central nervous system, particularly affecting children and adults [10].

Lead recovery processes are increasingly gaining importance due to the possibility of reusing the metal from processes and waste, the possibility of performing the process with a reduced consumption of raw materials and energy, and obtaining products of satisfactory quality. Lead recovery from minerals has traditionally been carried out through pyrometallurgical methods, which are effective but have significant drawbacks, including high energy consumption, the release of polluting gases, and lead dust emissions, posing environmental and occupational health risks [11]. Pyrometallurgical processes are accompanied by the formation of waste gases (metal dust and particles, volatile organic compounds with dioxins, SO_2, and other acidic gases), wastewater with metal compounds, as well as waste and residues from the process (sludge, residues rich in migratory metals, filter dust, and slag). Due to the content of certain elements (lead, arsenic, and antimony) and the possibility of their migration into the environment, the handling of the above waste streams must be strictly controlled. Despite these limitations, a portion of lead recycling from mining and urban wastes continues to rely on these methods [12].

In Chile, international and local environmental regulations, such as the Basel Convention, ratified in 1992—which prohibits the transboundary movement of hazardous waste, including spent lead–acid batteries and lead-containing materials [13]—and the Extended Producer Responsibility (EPR) Law 20920 [14], which mandates that producers take responsibility for the waste management of their products at the end of their life cycle, are driving the search for alternative methods of lead recycling. These regulations pose a significant challenge to the sustainable management of mining and urban wastes with high lead contents. In Chile, lead production is minimal, and official data on production and waste generation are scarce or restricted due to regulatory and confidentiality concerns. Historically, lead waste has been stabilized and disposed of underground rather than recovered. The recent EPR Law aims to improve traceability, but no consolidated public statistics are available.

Hydrometallurgical processes using chloride media emerge as sustainable alternatives. These methods use solutions with sodium chloride (NaCl), calcium chloride (CaCl₂), and magnesium chloride (MgCl₂), among other salts [15,16,17]. Future applications could include seawater or brine waste, which is particularly relevant for water-scarce regions in Chile [18].

Lead-containing mining and urban waste materials commonly consist of compounds such as metallic lead (Pb), lead oxides (PbO, PbO₂), and lead sulfate (PbSO₄), the latter being one of the significant components in some wastes [19,20,21]. Research highlights the refractory nature of PbSO₄, characterized by its chemical stability and low solubility in common acids such as hydrochloric acid (HCl), nitric acid (HNO₃), and acetic acid (CH₃COOH) [22,23]. Lead recovery from residues containing PbSO₄ in a chloride medium has emerged as a promising approach, primarily due to the ability of chloride ions to form stable lead–chloride complexes and their impact on altering the surface morphology of reaction products [15,24,25].

This review focuses on lead recovery in chloride media, drawing from key studies available in databases such as Scopus and Web of Science. Given the relevance of historical research in understanding fundamental mechanisms, both classic and recent experimental studies were considered, prioritizing those providing direct evidence of lead recovery in this medium. This article explores the chemistry of chloride-based solutions, the solubility of lead chloride (PbCl₂), and the formation of lead–chloride complexes. Additionally, it reviews the key variables influencing lead recovery from PbSO₄ in chloride media derived from various residues, with the aim of providing a comprehensive overview of these studies and guiding future research in this field.

2. Chemistry of Chloride Solutions

Chloride solutions offer significant advantages in the leaching of waste materials and ores. Studies indicate that the improved recovery of lead from PbS and compounds such as PbSO₄ is due to the formation of complexes between chloride ions and lead, as well as changes in the surface morphology of the reaction product [15,24]. Research on the leaching of copper and lead sulfide ores in chloride media suggests that the increased metal extraction can be attributed to the effects of chloride ions on solvation. This interaction between solvent and solute particles enhances proton activity, reducing the need for acid additions to maintain low pH values [26,27].

Majima and Awakura [28] reported that the activity of hydrochloric acid (HCl) increases significantly with the rise in NaCl and CaCl₂ salt concentrations. In a subsequent study, Majima and colleagues [29] concluded that proton activity improves according to the following sequence of salts: MgCl₂ > CaCl₂ > SrCl₂ > LiCl > NaCl > KCl.

A second interpretation suggests that the enhanced dissolution rate of lead from galena ore in the presence of high chloride concentrations could be due to the specific adsorption of chloride ions or the formation of surface complexes on the mineral [30].

3. Lead Chloride Solubility

In hydrometallurgical processes, it is essential that the metal of interest exhibits adequate solubility in the leaching solution to minimize the precipitation of compounds and prevent potential passivation phenomena [27]. Understanding the maximum amount of the element that can dissolve in a chloride medium is therefore critical.

Winand [31] reports that the saturated solubility of PbCl₂ is 0.99 g per 100 g of water at 20 °C and 3.34 g per 100 g of water at 100 °C. However, the solubility of PbCl₂ can increase significantly through complexation reactions with excess chloride ions, which may originate from hydrochloric acid or alkaline and alkaline earth chlorides. Holdich and Lawson [32] studied the solubility of aqueous PbCl₂ solutions and indicated that this compound is less soluble in 2 M CaCl₂, 4 M NaCl, and 6.2 M HCl solutions due to the common ion effect. However, at concentrations higher than those mentioned, lead solubility increases sharply due to the enhanced chloride activity (Figure 1a), facilitated by the formation of soluble lead–chloride complexes, which will be discussed in the following section. They also examined lead solubility at 25 °C and boiling temperature, using NaCl and CaCl₂ (Figure 1b), and found that the maximum lead solubility at the boiling point, with a total chloride concentration of 6.35 M, was 0.26 M with NaCl and 0.15 M with CaCl₂. This indicates that chloride ion activity is considerably higher in NaCl solutions compared to CaCl₂ at elevated chloride concentrations [32].

Thus, the solubility of PbCl₂ is directly proportional to the activity of chloride ions in the solution. Solubility limits are critical in these systems, as they play a key role in determining the maximum process yield when leaching lead concentrates, residues, or complex mixtures containing various lead compounds.

4. Lead–Chloride Complex

Lead exhibits oxidation states of 0, +2, and +4. The plumbic ion (Pb⁴⁺), corresponding to the +4 oxidation state, is known to be unstable in aqueous solutions and tends to reduce to the plumbous ion (Pb²⁺), which has an oxidation state of +2 [33]. Therefore, lead is commonly found in the +2 oxidation state, particularly in the form of chloride complexes in aqueous solutions. The equilibria and stability constants for lead–chloride complexes are provided in the following equations [34,35]:

$$\mathrm{P}{\mathrm{b}}^{2+}+{\mathrm{C}\mathrm{l}}^{-}\leftrightarrow \mathrm{P}\mathrm{b}\mathrm{C}{\mathrm{l}}^{+}{\mathrm{K}}_1=12.59$$

(1)

$$\mathrm{P}\mathrm{b}\mathrm{C}{\mathrm{l}}^{+}+{\mathrm{C}\mathrm{l}}^{-}\leftrightarrow \mathrm{P}\mathrm{b}\mathrm{C}{\mathrm{l}}_2{\mathrm{K}}_2=14.45$$

(2)

$$\mathrm{P}\mathrm{b}\mathrm{C}{\mathrm{l}}_2+{\mathrm{C}\mathrm{l}}^{-}\leftrightarrow \mathrm{P}\mathrm{b}\mathrm{C}{\mathrm{l}}_3^{-}{\mathrm{K}}_3=0.398$$

(3)

$$\mathrm{P}\mathrm{b}\mathrm{C}{\mathrm{l}}_3^{-}+{\mathrm{C}\mathrm{l}}^{-}\leftrightarrow \mathrm{P}\mathrm{b}\mathrm{C}{\mathrm{l}}_4^{2-}{\mathrm{K}}_4=0.089$$

(4)

Studies on the stability and reactivity of lead–chloride complexes and their anionic chloroplumbate derivatives highlight the existence of a wide range of additional complexes, formed by chloride ions or chlorine gas [25,36,37]. Below are some equations for stable complexes in aqueous solutions:

$$\mathrm{P}\mathrm{b}\mathrm{C}{\mathrm{l}}_4^{2-}+{\mathrm{C}\mathrm{l}}^{-}\leftrightarrow {\mathrm{P}\mathrm{b}\mathrm{C}\mathrm{l}}_5^{3-}\mathsf{\Delta }{G}_{r,298}=-3.4 \mathrm{k}\mathrm{J}/\mathrm{m}\mathrm{o}\mathrm{l}$$

(5)

$${\mathrm{P}\mathrm{b}\mathrm{C}\mathrm{l}}_5^{3-}+{\mathrm{C}\mathrm{l}}^{-}\leftrightarrow {\mathrm{P}\mathrm{b}\mathrm{C}\mathrm{l}}_6^{4-}\mathsf{\Delta }{G}_{r,298}=-53.3 \mathrm{k}\mathrm{J}/\mathrm{m}\mathrm{o}\mathrm{l}$$

(6)

$${\mathrm{P}\mathrm{b}\mathrm{C}\mathrm{l}}_3^{-}+{\mathrm{C}\mathrm{l}}_2\leftrightarrow {\mathrm{P}\mathrm{b}\mathrm{C}\mathrm{l}}_5^{-}\mathsf{\Delta }{G}_{r,298}=-80.2 \mathrm{k}\mathrm{J}/\mathrm{m}\mathrm{o}\mathrm{l}$$

(7)

$${\mathrm{P}\mathrm{b}\mathrm{C}\mathrm{l}}_4^{2-}+{\mathrm{C}\mathrm{l}}_2\leftrightarrow {\mathrm{P}\mathrm{b}\mathrm{C}\mathrm{l}}_6^{2-}\mathsf{\Delta }{G}_{r,298}=-154.2 \mathrm{k}\mathrm{J}/\mathrm{m}\mathrm{o}\mathrm{l}$$

(8)

$${\mathrm{P}\mathrm{b}\mathrm{C}\mathrm{l}}_5^{-}+{\mathrm{C}\mathrm{l}}^{-}\leftrightarrow {\mathrm{P}\mathrm{b}\mathrm{C}\mathrm{l}}_6^{2-}\mathsf{\Delta }{G}_{r,298}=-5.1 \mathrm{k}\mathrm{J}/\mathrm{m}\mathrm{o}\mathrm{l}$$

(9)

Freza et al. [25] reported that Equations (5) and (6) represent electrically stable complexes, while Equation (7) describes an exothermic and spontaneous process in aqueous solutions, despite requiring the overcoming of specific activation barriers. In contrast, Equations (8) and (9) yield PbCl₆²⁻ as the stable product in solution.

To gain deeper insights into the nature of lead–chloride solutions, speciation diagrams of lead in the presence of chloride ions were constructed using the software “Make Equilibrium Diagrams Using Sophisticated Algorithms” (MEDUSA 32 bit vers. 16 December 2010) [38]. Figure 2a,b illustrate the fraction of lead chemical species as a function of total chloride concentration, under the conditions of 0.001 M Pb²⁺ and 25 °C and a solution potential of 0.5 V. Figure 2a, developed at pH 1.0, reveals that, at chloride concentrations below 0.5 M, the predominant species is the PbCl⁺ (Equation (1)), whereas, at chloride concentrations exceeding 2.0 M, the stable complex is PbCl₄²⁻ (Equation (4)). Figure 2b, constructed at pH 7.0, shows that at total chloride concentrations below approximately 3.2 M, lead predominantly exists as a solid PbOHCl. However, at higher chloride concentrations, PbCl₄²⁻ becomes the dominant species. These results highlight that neutral pH can negatively impact lead recovery when total chloride concentrations are below 3.2 M.

Figure 3a,b depict the fraction of lead chemical species as a function of pH, constructed at a Pb²⁺ concentration of 0.001 M, a temperature of 25 °C, and a total chloride concentration of 3.0 M. Figure 3a, at a solution potential of 0.5 V, indicates that within the pH range of 0 to 6.7, the distribution of the complexes remains constant, with PbCl₄²⁻ being the predominant species. However, at higher pH values, lead undergoes a series of precipitation reactions. Beyond pH 8.0, the alkaline hydrolysis of lead occurs, resulting in the formation of solid compounds such as Pb(OH)₂, PbO_2, or Pb(OH)₄. Figure 3b, at a solution potential of 1.2 V, demonstrates that in the pH range of 0 to 4.0, the stable complex is PbCl₆²⁻, as described in Equations (8) and (9). Creighton and Woodward [36] noted that this complex is predominant in concentrated hydrochloric acid solutions.

5. Effect of Lead–Chloride Complex Formation on the Reduction Potential

The formation of complexes between a metal ion and a ligand can significantly affect the value of the standard reduction potential [27,39]. Equation (10) shows the reduction reaction of Pb²⁺ to Pb⁰ and its standard reduction potential.

$${\mathrm{P}\mathrm{b}}_{\left(\mathrm{a}\mathrm{c}\right)}^{2+}+2{\mathrm{e}}^{-}\leftrightarrow {\mathrm{P}\mathrm{b}}_{\left(\mathrm{s}\right)}^0{E}^0=-0.13\mathrm{V}$$

(10)

The presence of complexing ligands, such as chloride, modifies the reduction potential of lead. This effect can be interpreted as a decrease in the activity of free Pb²⁺ ions, which hinders their reduction to metallic lead and requires a more negative potential, as shown in Figure 4, extracted from the study of Lin and co-workers [9]. Furthermore, their findings indicate that the reduction potential between lead–chloride complexes and metallic lead becomes more negative as the concentration of chloride ions increases. Consequently, higher chloride concentrations necessitate more negative potentials, implying greater energy consumption for reducing lead in a chloride medium.

6. Lead Recovery from PbSO₄ in Chloride Medium from Waste

The dissociation of PbSO₄ in aqueous solutions yields plumbous ions (Pb²⁺) and sulfate ions (SO₄²⁻), as shown in Equation (11).

$${\mathrm{P}\mathrm{b}\mathrm{S}{\mathrm{O}}_4}_{\left(\mathrm{s}\right)}\to {{\mathrm{P}\mathrm{b}}^{2+}}_{\left(\mathrm{a}\mathrm{c}\right)}+{\mathrm{S}{\mathrm{O}}_4^{2-}}_{\left(\mathrm{a}\mathrm{c}\right)}{\mathrm{K}}_{\mathrm{p}\mathrm{s}}=2\times {10}^{-8}$$

(11)

The solubility of PbSO₄ is limited due to its low solubility product (K_PS). However, in chloride media, lead can form soluble complexes, as described by Equations (1) through (9).

Table 1. Lead recovery from lead sulfate through agitated leaching in chloride media.

PbSO4 Source	Pb Content (wt, %)	Leaching Agent	Concentration	Temperature	pH	Particle Size	Stirring Speed, RPM	Density Pulp	Leaching Time	Pb Recovery, %	Reference
Synthetic	68.3	CaCl2	444 g/L	80 °C	-	12.5 μm	125	303 g/L	240 min	95	[15]
		MgCl2	380 g/L						60 min	99
Zn residue	25	NaCl	300 g/L	30 °C	2.0	150 μm	-	20 g/L	30 min	92	[40]
		HCl
	15.5	NaCl	300 g/L	95 °C	-	68 μm	250	50 g/L	10 min	98.90	[3]
		HCl	30 mL/L
	12.4	NaCl	300 g/L	37 °C	1.0	P80 = 118 μm	400	25 g/L	30 min	89.43	[35]
		HCl
	14.4	NaCl	300 g/L	70 °C	2.0	P80 = 88 μm	700	25 g/L	10 min	78.94	[34]
		HCl
	2.6	CaCl2	400 g/L	80 °C	1.0	<106 μm	500	143 g/L	45 min	93.79	[4]
		HCl
	14.1	CaCl2	350 g/L	45 °C	2.0	-	-	63 g/L	120 min	85.78	[21]
	6.2	NaCl	175 g/L	25 °C	-	P50 = 9.6 μm	120	50 g/L	-	72	[17]
		HCl	0.1 M
	3.4	NaCl	350 g/L	room	-	P90 = 150 μm	300	10%	90 min	75.72	[41]
		HCl	2 M
Concentrate	5.2	NaCl	233 g/L	90 °C	-	-	500	100 g/L	600 min	99.50	[42]
		H2SO4	0.05 M
		O2	0.03 m3/h
Battery paste	71.1	CaCl2	400 g/L	90 °C	1.0	<150 μm	500	33 g/L	120 min	99.20	[23]
		Fe2+	5 g/L
		HCl

summarizes the leaching conditions for PbSO₄ in chloride media under agitated conditions, which have resulted in the highest lead recoveries. Although these studies may also target the recovery of other elements of interest, the focus here is exclusively on lead recovery.

Table 1 illustrates the wide range of lead concentrations found in lead-bearing residues, varying from 3.4% in certain zinc residues to 71.1% in lead-rich wastes, such as spent battery paste. The collected results demonstrate the effectiveness of chloride media in lead recovery, with efficiencies ranging from 72% to 99%. Due to the complexity of comparing the various experimental conditions, the key variables influencing lead recovery are described below.

6.1. Effect of Chloride Salts and Their Concentration

The main chloride salts studied were MgCl₂, CaCl_2, and NaCl. Sinadinovic and coworkers [15] compared the use of MgCl₂ and CaCl₂ in the dissolution of synthetic PbSO₄ (Figure 5a,b) and reported that MgCl₂ enhances lead recovery and reduces leaching time compared to CaCl₂. The use of MgCl₂ results in the formation of soluble magnesium sulfate (MgSO₄), whereas CaCl₂ leads to the formation of a porous calcium sulfate (CaSO₄) layer on the surfaces of PbSO₄ particles. This porous layer, observed through microphotography, was suggested to decrease the reaction kinetics [15]. Similarly, Xie and co-workers [21] highlighted that CaCl₂ is employed to precipitate CaSO₄ and promote the progression of the reaction.

NaCl has been extensively studied as a chloride source for lead recovery due to its low cost. Research indicates that increasing salt concentration or chloride activity (MgCl₂, CaCl_2, and NaCl) enhances the formation of lead–chloride complexes, thereby improving lead recovery [3,15,34]. However, the application of these salts is constrained by their solubility limits: NaCl saturates at approximately 350 g/L [3], CaCl₂ at around 665 g/L [15], and MgCl₂ at nearly 542 g/L. Wang et al. [4] noted that excessively high salt concentrations increase system viscosity, rendering the process economically unfeasible at an industrial scale.

Furthermore, the prolonged leaching of PbSO₄ with NaCl has been shown to result in sulfate accumulation in the solution, adversely affecting lead recovery [40,43]. Raghavan et al. [40] confirmed this phenomenon in a lead leaching system containing 300 g/L of NaCl and sodium sulfate (Na₂SO₄) concentrations ranging from 20 to 500 g/L (Figure 6). They found that sulfate concentrations exceeding 200 g/L significantly hinder lead recovery due to the common ion effect or the formation of insoluble precipitates.

6.2. Influence of Acid Type and pH

PbSO₄ is stable in common acids such as HCl, HNO₃, and CH₃COOH [9] but dissolves in saturated chloride solutions acidified with either HCl or H₂SO₄ acid [35].

Lead recovery in a 300 g/L NaCl solution at pH 2.0, adjusted with HCl or H₂SO₄, was found to be 78.94% and 77.78%, respectively, showing comparable results with both acids [34]. Consequently, the positive effect of chloride ions introduced by HCl and the negative effect of sulfate ions from H₂SO₄ at this pH are negligible. For economic reasons, Behnajady and co-workers [34] continued their studies with H₂SO₄.

However, Silwamba et al. [17] demonstrated that a higher concentration of sulfate, whether originating from H₂SO₄ or other sources, can adversely affect lead recovery. This conclusion was further corroborated by their use of speciation diagrams constructed with MEDUSA software [38] under conditions of 3.0 M total chloride concentration and a solution potential of 0.5 V (Figure 7). Figure 7a illustrates that when the concentrations of lead and sulfate are equal, pH has little to no effect on lead speciation within the range of 0 to 5.0. Additionally, they extended their analysis by constructing speciation diagrams to evaluate scenarios with varying sulfate concentrations. As shown in Figure 7b, when the sulfate concentration exceeds that of lead, pH becomes a determining factor in lead speciation. At pH values above 2.0, lead precipitates as PbSO₄, whereas at pH values below 2.0, lead predominantly remains in solution as the PbCl₄²⁻ complex. On the other hand, Ruşen et al. [3] reported significant improvements in lead recovery by directly adding 10–30 mL/L of HCl to a leaching system with 300 g/L of NaCl at 95 °C.

6.3. Effect of Temperature

It is widely recognized that increasing temperature enhances element recovery in leaching systems. Figure 5 illustrates the positive effect of temperature on lead recovery, showing a marked increase as the temperature rises. However, Sinadinovic et al. [15] noted that as the solution cools after leaching, lead precipitates as PbCl₂ when low chloride concentrations are present. Chmielewski et al. [42] observed this phenomenon upon cooling a leached solution containing 118 g/L NaCl.

6.4. Effect of Reaction Time

Ruşen et al. [3] suggested controlling the reaction time within a range of 5 to 15 min. Their study, conducted at 95 °C with 300 g/L NaCl, reported a gradual decrease in lead extraction after 10 min, which they attributed to the precipitation of dissolved lead. This behavior was also observed by Nuñez et al. [44] who proposed it as a possible reversal of the chemical reaction.

In contrast, Behnajady et al. [34] examined reaction times of 1 and 8 h and demonstrated that, at 55 °C with 300 g/L NaCl, lead recovery increased significantly from 35.93% to 75.38% after 8 h. Notably, this study was conducted at a lower temperature compared to that of Ruşen et al.

Furthermore, Xing et al. [23] investigated lead recovery from spent lead battery paste using a leaching solution containing 400 g/L CaCl₂, 5 g/L Fe²⁺, and a pH of 1.0 at 90 °C. They concluded that the optimal leaching time under these conditions was 2 h. These results suggest that the leaching time required to achieve optimal lead recovery is influenced by the working temperature and solution composition.

6.5. Effect of Pulp Density

Numerous studies have demonstrated that lead recovery in chloride media is inversely proportional to pulp density [4,21,35], identifying it as a highly significant parameter in the leaching process [34].

Raghavan et al. [40] investigated the effect of pulp density, ranging from 10 to 300 g of solid sample per liter of solution, using NaCl. As shown in Figure 8, an increase in pulp density results in decreased lead recovery, attributed to the limited solubility of PbCl₂ [34]. Similarly, Wang et al. [4] and Xie et al. [21] emphasized that lower pulp densities reduce suspension density and system viscosity, thereby minimizing mass transfer resistance at the liquid–solid interface and enhancing the leaching reaction.

Farahmand et al. [35] further recommended employing low slurry densities to optimize lead recovery. However, the use of CaCl₂ has been reported as effective for improving lead recovery, even at pulp densities exceeding 63 g/L [4,15,21].

6.6. Effect of Stirring Speed

The study by Wang et al. [4] demonstrated that lead recovery increases with agitation speed up to 400 RPM, beyond which it plateaus, as shown in Figure 9. Similarly, Behnajady et al. [34] observed that at agitation speeds above 500 RPM, lead recovery remains largely unchanged, indicating that further increases in agitation speed have a negligible effect.

7. Kinetic Studies of the Dissolution of Lead Sulfate in Chloride Media

Leaching kinetics is relevant for process design, optimization, and control [27], and it is influenced by various variables, such as those discussed above. Studies on the kinetics of PbSO₄ leaching in chloride media have predominantly employed the “Shrinking-Core Model” (SCM). This model proposes that the reaction begins at the external surface of the solid particle and progresses inward, leaving behind a fully converted and inert solid layer, resulting in the gradual shrinking of the unreacted core during the process [45]. The primary rate-controlling mechanisms considered include diffusion through the liquid film (Equation (12)), diffusion through the porous solid layer (Equation (13)), and chemical reaction control (Equation (14)).

$$\mathsf{\alpha }=\mathrm{k}·\mathrm{t}$$

(12)

$$1-{\displaystyle \frac{2}{3}}\mathsf{\alpha }-{(1-\mathsf{\alpha })}^{{\displaystyle \raisebox{1ex}{$2$}\!\left/ \!\raisebox{-1ex}{$3$}\right.}}=\mathrm{k}·\mathrm{t}$$

(13)

$$1-{(1-\mathsf{\alpha })}^{{\displaystyle \raisebox{1ex}{$1$}\!\left/ \!\raisebox{-1ex}{$3$}\right.}}=\mathrm{k}·\mathrm{t}$$

(14)

In these equations, α represents the fraction converted, t the leaching time (min), and k is the kinetic constant (min⁻¹).

A summary of PbSO₄ leaching kinetics is presented in

Table 2. Lead recovery from lead sulfate through agitated leaching in chloride media.

PbSO4 Source	Leaching Medium	Temperature, °C	Rate-Determining Step	Activation Energy, kJ/mol	Reference
Synthetic	CaCl2	25–80	Mixed -Diffusion product layer	31.0	[15]
			-Chemical reaction	69.0
	MgCl2	25–80	Mixed -Diffusion	15.0
			-Chemical reaction	46.0
Intermediate residue	NaCl	25–80	Diffusion	12.4	[46]
Zn residue	CaCl2	35–65	Diffusion	17.6	[21]
	NaCl HCl	25–60	Diffusion product layer	16.7	[41]
	NaCl	45–90	Mixed control	13.4	[16]
	CaCl2
	NaClO3

. The primary chloride salts investigated include MgCl₂, CaCl_2, and NaCl, with activation energy (Ea) values ranging from 12.41 to 69.0 kJ/mol. Sinadinovic et al. [15] fitted their experimental data to models governed by chemical reaction or diffusion control, but, without fully validating their results, proposed a mixed-control mechanism for PbSO₄ leaching with CaCl₂. This mechanism involves diffusion through the porous layer due to the formation of calcium sulfate (CaSO₄) on the PbSO₄ surface and chemical reactions occurring on the surface. The presence of CaSO₄ was identified through EDS analysis and confirmed by microphotographs of PbSO₄. Similarly, leaching with MgCl₂ also exhibits a mixed-control mechanism, where both liquid–film diffusion and surface chemical reactions contribute to the rate-controlling step. To analyze this behavior, the authors applied the additive time law and introduced a mathematical correlation between experimental leaching rate data and time. Additionally, they used the Nelder–Mead optimization algorithm to minimize errors between experimental and modeled data. They also analyzed the impact of temperature and stirring speed on the leaching rate, fitting their results to the Arrhenius equation.

In contrast, Geidarov et al. [46] identified liquid–film diffusion as the rate-determining step for PbSO₄ leaching in NaCl. They demonstrated that the reaction is zero-order with respect to lead concentrations within the range of 0 to 120 g/L and reported a linear correlation between the leaching rate constants and NaCl concentration at 25 °C, with constants of 0.043, 0.079, 0.111, and 0.145 min⁻¹ for 50, 130, 200, and 270 g/L NaCl, respectively. However, a recent study by Houshmand et al. [41] using NaCl with HCl identified a different mechanism, where the rate-determining step is diffusion through a porous film. They did not employ experimental techniques to identify the phases formed after leaching; instead, their study relied entirely on theoretical analysis using the shrinking core model.

Xie et al. (2019) [21] identified diffusion as the rate-controlling step when using CaCl₂. By analyzing leaching rates at 35–65 °C and fitting the data to shrinking-core equations, they derived kinetic constants and confirmed the diffusion-controlled process.

Both Houshmand et al. [41] and Zhang et al. [16] determined activation energies using the Arrhenius equation, reinforcing the importance of temperature-dependent kinetic studies in PbSO₄ leaching.

The differences in rate-controlling steps observed for similar chloride media highlight the complexity of PbSO₄ leaching mechanisms, indicating the need to continue research in this area.

8. Lead Concentration Methods from Chloride Media

After lead leaching in chloride media, different concentration methods can be applied to recover lead efficiently. One of the most studied approaches is electrowinning, as demonstrated in the PLACID process [5]. This method employs selective membrane cells, allowing lead deposition (Equation (15)), while oxygen evolution occurs at the anode (Equation (16)). The use of proton-permeable membranes maintains the separation of anodic and cathodic electrolytes, facilitating the regeneration of HCl for leaching. With current densities around 1200 A/m², pure lead powder is obtained, which is later compacted for smelting.

$${\mathrm{P}\mathrm{b}\mathrm{C}\mathrm{l}}_{2\left(\mathrm{a}\mathrm{c}\right)}+2{\mathrm{e}}^{-}\leftrightarrow {\mathrm{P}\mathrm{b}}_{\left(\mathrm{s}\right)}^0+2{\mathrm{C}\mathrm{l}}_{\left(\mathrm{a}\mathrm{c}\right)}^{-}$$

(15)

$$2{\mathrm{H}}_2\mathrm{O}\leftrightarrow {\mathrm{O}}_2+4{\mathrm{H}}^{+}+2{\mathrm{e}}^{-}$$

(16)

Xing et al. [23] further investigated lead electrowinning from CaCl₂-FeCl₂ leachates of spent lead–acid batteries, finding that optimal conditions (10 g/L Pb, 200 A/m²) resulted in 99.6% purity, a current efficiency of 96.3%, and an energy consumption of 85.9 kWh per ton of lead.

Lead cementation using aluminum has also been explored [17,35]. Farahmand et al. [35] determined that at 40–50 °C, the process follows a chemically controlled model, whereas at higher temperatures, diffusion control dominates with an activation energy of 26.8 kJ/mol. Silwamba et al. [17] demonstrated that adding zero-valent aluminum during leaching significantly enhances Pb removal, even at low chloride concentrations (Equation (17)).

$${3\mathrm{P}\mathrm{b}}^{2+}+{2\mathrm{A}\mathrm{l}}^0\to 2{\mathrm{P}\mathrm{b}}^0+2{\mathrm{A}\mathrm{l}}^{3+}$$

(17)

Another effective method is PbCl₂ crystallization through cooling [47,48]. Xie et al. [47] reported that cooling a PbCl₂-saturated solution to 15 °C for 3 h resulted in 82% crystallization with 99.5% purity. Xing et al. [48] applied this technique to a chloride leachate from bismuth refining slags, achieving a 58.9% lead crystallization yield at 20.6 g/L Pb concentration.

Beyond industrial applications, fundamental studies have explored chloride crystallization mechanisms and conditions. Holdich and Lawson [32] examined Pb solubility in chloride solutions containing CaCl₂ and NaCl acidified with HCl and CuCl₂, where lead was introduced as PbCl₂. They found that in acidic solutions with 5–10 M chloride, up to two-thirds of the lead could precipitate as fine PbCl₂ crystals (>99% purity) by cooling from boiling temperature to 20 °C, with no evidence of other insoluble double salts. Veintemillas et al. [49] investigated PbCl₂ crystal growth in boiling aqueous solutions, obtaining isometric PbCl₂ crystals up to 1 cm in size after three weeks using continuous lead chloride feeding into a boiling KNO₃-H₂O system. Li et al. [50] introduced an alternative approach using the ionic liquid Aliquat 336 (methyltrioctylammonium chloride) to simultaneously remove and recover Pb²⁺ from wastewater at room temperature, inducing the formation of rod-like PbCl₂ crystals at the organic–aqueous interface.

These concentration methods provide versatile pathways for lead recovery in chloride media, with the choice of method depending on process conditions, selectivity, energy consumption, and final product requirements.

9. Implications of Future Sustainable Recovery of Lead from Secondary Waste

The recovery of lead from mining and municipal waste has been investigated through various hydrometallurgical methods. These include sulfur removal treatments using sodium hydroxide (NaOH), sodium carbonate (Na₂CO₃), or ammonium carbonate ((NH₄)₂CO₃), followed by lead leaching with fluoroboric acid (HBF₄) or fluorosilicic acid (H₂SiF₆). Such processes, often applied to lead battery pastes [51,52,53], face significant challenges during electrowinning due to the emission of toxic gases, including hydrogen fluoride (HF) and silicon tetrafluoride (SiF₄) [22,54]. Alternative approaches involve thermal sulfur removal, followed by leaching in media such as ammonium acetate, which have been applied to materials like electrolytic manganese anode slime [2,55].

The direct leaching of mining and urban wastes for lead recovery has been studied in various environments. Acidic media, such as nitric acid, have been extensively investigated [7,56], along with acid chloride systems, which are reviewed in this article, and alkaline media employing reagents like ammonium sulfate ((NH₄)₂SO₄), ammonium hydroxide (NH₄OH), and NaOH [57,58].

Additionally, organic media, including urea with acetic acid, sodium citrate with citric acid, and ethylenediaminetetraacetic acid (EDTA), have been explored as biodegradable and environmentally friendly reagents [20,59,60,61]. However, carboxylic acids are more expensive than inorganic acids such as HCl or H₂SO₄ and are less readily available due to lower global production volumes [62,63].

Chloride-based media have demonstrated their effectiveness and cost-efficiency as leaching agents for copper recovery from ores. In Chile, seawater and chloride solutions have been utilized in mining operations, such as the Michilla plant, for leaching copper from secondary ores. This approach has been motivated by the need to identify alternative water sources due to the declining availability of freshwater [64]. In contrast, lead recycling in Chile is predominantly carried out through pyrometallurgical processes aligned with North American technical guidelines [65]. While effective, these processes present notable challenges, including the emission of lead-containing gases and particulates, soil contamination, and significant occupational health risks.

The review summarized in Table 1 underscores the effectiveness of chlorinated media as a viable alternative for lead recovery from various waste streams. A prominent example is the PLACID process, developed and patented by the R&D team at Técnicas Reunidas in Spain. This process involves the direct leaching of lead paste from end-of-life batteries using a hydrochloric acid (HCl) solution diluted in heated NaCl brine [5].

It is well established that increasing the concentration of chloride salts enhances lead recovery, making this a critical variable to consider. However, caution must be exercised when approaching the solubility limit, as this can increase the system’s viscosity, complicating subsequent processing steps [4]. Among the chloride salts studied, NaCl, CaCl₂, and MgCl₂ are the most prominent. While CaCl₂ may lead to the precipitation of CaSO₄, MgCl₂ has been reported to reduce lead leaching time by generating soluble MgSO₄ [15,21].

NaCl has received significant attention due to its low cost and accessibility. However, studies suggest that the prolonged leaching of PbSO₄ with NaCl can adversely affect lead recovery [40,43] by increasing the sulfate concentration in the solution. When sulfate concentration exceeds lead concentration, pH becomes a critical variable. At pH values above 2.0, lead precipitates as PbSO₄, whereas at lower pH levels, the PbCl₄²⁻ complex predominates (Figure 7b) [17]. Conversely, when lead and sulfate concentrations are balanced in solution, pH has little impact within the range of 0 to 6.0, and the PbCl₄²⁻ complex remains predominant (Figure 7a). Another study found no significant effect of using HCl or H₂SO₄ at pH 2.0 with 300 g/L NaCl on lead recovery. At this pH, neither the addition of HCl to enhance chloride activity nor the addition of H₂SO₄, which introduces sulfate ions, showed a measurable impact [34].

Another important factor frequently overlooked is the measurement of solution potential. This parameter offers valuable insights into the formation of additional complexes, such as PbCl₆²⁻, which predominates at a solution potential of 1.2 V and pH values between 0 and 4.0, when the total chloride concentration is 3 M (Figure 3b).

The primary advantage of using chloride media for lead recovery from waste materials lies in the formation of stable lead–chloride complexes. The predominance of specific complexes in solution depends on key factors, including chloride concentration, sulfate concentration, pH, and solution potential. Lead leaching times in chloride media are typically short, with maximum recovery achieved within 15 to 120 min [3,23]. Temperature plays a pivotal role and requires careful regulation. Studies indicate that allowing the post-leaching solution to cool at low chloride concentrations, such as 118 g/L NaCl, can lead to PbCl₂ precipitation, potentially hindering complete lead recovery [15,42]. Agitation speeds of 600 to 700 RPM are generally enough to facilitate lead recovery [4].

Finally, pulp density is another significant variable due to the limited solubility of PbCl₂. Studies recommend limiting pulp densities to below 25 g/L when employing NaCl to maximize lead recovery [35,40].

10. Conclusions

This review provides a comprehensive evaluation of the conditions required for efficient lead recovery from diverse mining and urban wastes in chloride media, positioning it as a sustainable solution to the environmental challenges posed by conventional recycling methods. Chloride media demonstrate significant benefits, including accelerated leaching kinetics, the stabilization of lead–chloride complexes, and high recovery rates under optimized conditions of elevated chloride concentrations and low pulp densities.

The studies analyzed underscore the importance of chloride concentration, sulfate content, pH, and redox potential in driving lead dissolution and complex formation. They also reveal potential challenges, such as sulfate accumulation and PbCl₂ precipitation, which must be carefully managed to maintain high recovery efficiencies.

Future research should focus on diversifying lead sources, performing detailed kinetic and thermodynamic analyses, and investigating dissolution mechanisms in chloride media. These efforts will establish chloride-based hydrometallurgical processes as a sustainable and efficient alternative to conventional lead recycling methods.