Evaluation of Silver Recovery from High Sulphur Mining Waste Using the Thiourea-Oxalate System

1. Introduction

Tailings are mining residues originating from the ore beneficiation process, where valuable minerals were extracted, leaving behind mineral phases such as iron sulfides, silicates, silica, and minerals containing toxic metals, which are of low economic value [1]. Unfortunately, this type of waste had no environmental regulation during its open-pit disposal, leaving it exposed to weathering factors and microbial activity (e.g., the presence of Thiobacillus ferrooxidans). This exposure results in the oxidation of sulfide-type mineral phases (mainly pyrite) leading to the generation of acid (H⁺) and, as a consequence, the dissolution of heavy metals that accompany the mineral phases of low economic value. The leachates obtained directly contaminate soil and water [2,3,4,5,6,7].

Based on the above, in this type of waste, it is important to evaluate the acid-base balance, concentration of toxic metals, as well as to propose alternatives that help to totally or partially reduce future acid generation. And due to the content of metallic values of commercial interest (copper, lead, zinc, gold and silver) use these wastes as raw material in hydrometallurgical processes for the dissolution of metals of interest.

This implies a thorough characterization to help identify the mineral phases that provide the metal of interest, analyze the degree of liberation since the metals of interest such as gold and silver can be found in carbonaceous, argillaceous, arsenosulfide, sulfide, sulfide, sulfosalt and refractory mineral matrices [8,9,10,11,12], necessitating pretreatments such as intensive grinding, roasting, autoclave oxidation, or bioleaching [13,14]. However, these processes often increase operational costs, limiting their economic feasibility.

On the other hand, it is well known that cyanidation is the most widely used process for the recovery of metals such as gold and silver. However, the environmental problems caused by cyanidation have encouraged the search for new leachants to replace sodium cyanide. Proposed alternatives include chlorides [15], thiosulfate (S₂O₃²⁻) [16], thiocyanate (SCN⁻) [17], and thiourea (TU, CS(NH₂)₂) [18]. Similarly, oxidants such as manganese dioxide (MnO₂), monopersulfate compounds (HSO₅), ozone, hydrogen peroxide (H₂O₂), and ferric ion (Fe³⁺) have been studied, the latter being the most commonly used in TU leaching (Eq. 1) [19,20,21,22].

$$\mathrm{A}\mathrm{g}+2\mathrm{C}\mathrm{S}({\mathrm{N}\mathrm{H}}_2{)}_2+{\mathrm{F}\mathrm{e}}^{3+}=\mathrm{A}\mathrm{g}(\mathrm{C}\mathrm{S}({\mathrm{N}\mathrm{H}}_2{)}_2{)}_2^{+}+{\mathrm{F}\mathrm{e}}^{2+}$$

(1)

Acid thioureation (TU in acid medium) has proven to be the most attractive alternative for the treatment of silver ores [23] due to its high efficiency and low environmental toxicity [20,24,25]. The use of TU involves the formation of formamidine disulfide (FDS, (NH₂)₂CSSC(NH₂)₂²⁺) (Eq. 2) [26], which acts as an oxidant in the presence of dissolved oxygen. However, FDS is unstable and decomposes into TU, cyanamide (CH₂N₂), and elemental sulfur (S⁰) (Eq. 3) [27,28]. The formation of S⁰ is undesirable because it causes mineral surface passivation, which decreases the leaching efficiency of valuable metals [26,29,30].

$$2\mathrm{C}\mathrm{S}({\mathrm{N}\mathrm{H}}_2{)}_2=({\mathrm{N}\mathrm{H}}_2{)}_2\mathrm{C}\mathrm{S}\mathrm{S}\mathrm{C}({\mathrm{N}\mathrm{H}}_2{)}_2^{2+}+{2\mathrm{e}}^{-} \mathrm{E}°=0.42\mathrm{V}$$

(2)

$$\left({\mathrm{N}\mathrm{H}}_2{)}_2\mathrm{C}\mathrm{S}\mathrm{S}\mathrm{C}\right({\mathrm{N}\mathrm{H}}_2{)}_2=\mathrm{C}\mathrm{S}({\mathrm{N}\mathrm{H}}_2{)}_2+\mathrm{N}{\mathrm{H}}_2\mathrm{C}\mathrm{N}+{\mathrm{S}}^0$$

(3)

According to the study by D. Calla et al. (2016, 2019), TU oxidation is potentiated by an increase in the redox potential (about 0.5 V vs SHE) caused by the dissolved iron (Fe³⁺) and copper (Cu²⁺) themselves. Moreover, the ferric ion alone promotes the oxidation of TU (Eq. 4), negatively affecting silver leaching [19,31]:

$$2CS({NH}_2{)}_2+2{Fe}^{3+}=2{Fe}^{2+}+({NH}_2{)}_{2}CSSC({NH}_2{)}_2^{2+}$$

(4)

Another negative aspect of the use of thiourea is the ease of forming complexes as metal ions (Eq. 5 to 7) which implies adding additional TU so as not to mismatch the TU needed in the complex with silver.

$${\mathrm{F}\mathrm{e}}^{3+}+\mathrm{C}\mathrm{S}({\mathrm{N}\mathrm{H}}_2{)}_2=[\mathrm{F}\mathrm{e}\mathrm{C}\mathrm{S}({\mathrm{N}\mathrm{H}}_2{)}_2{]}^{3+}$$

(5)

$${\mathrm{F}\mathrm{e}}^{3+}+2\mathrm{C}\mathrm{S}({\mathrm{N}\mathrm{H}}_2{)}_2=[\mathrm{F}\mathrm{e}\left(\mathrm{C}\mathrm{S}\right({\mathrm{N}\mathrm{H}}_2{)}_2{)}_2{]}^{3+}$$

(6)

$$6\mathrm{C}\mathrm{S}({\mathrm{N}\mathrm{H}}_2{)}_2+2{\mathrm{C}\mathrm{u}}^{2+}=2\mathrm{C}\mathrm{u}(\mathrm{C}\mathrm{S}(\mathrm{N}{\mathrm{H}}_2{)}_2{)}_2^{+}+{(\mathrm{N}\mathrm{H}}_2{)}_2\mathrm{C}\mathrm{S}\mathrm{S}\mathrm{C}{(\mathrm{N}\mathrm{H}}_2{)}_2^{2+}$$

(7)

In attention to the problems of TU oxidation Chandra et al. (2005) and D. Calla et al. (2020, 2021) have implemented the use of the organic ligand oxalate (Ox, C₂O₄²⁻) as a way to decrease the activity of copper and iron ions (Eq. 8 to 11) thus allowing the availability of thiourea for silver dissolution [32,33,34].

$$\mathrm{C}{\mathrm{u}}^{2+}+2{\mathrm{C}}_2{\mathrm{O}}_4^{2-}=\mathrm{C}\mathrm{u}{\left({\mathrm{C}}_2{\mathrm{O}}_4\right)}_2^{2-} \log \mathrm{K}=10.23$$

(8)

$${\mathrm{F}\mathrm{e}}^{3+}+2{\mathrm{C}}_2{\mathrm{O}}_4^{2-}=\mathrm{F}\mathrm{e}{\left({\mathrm{C}}_2{\mathrm{O}}_4\right)}_2^{-} \log \mathrm{K}=13.81$$

(9)

$${\mathrm{F}\mathrm{e}}^{3+}+3{\mathrm{C}}_2{\mathrm{O}}_4^{2-}=\mathrm{F}\mathrm{e}{\left({\mathrm{C}}_2{\mathrm{O}}_4\right)}_3^{3-} \log \mathrm{K}=18.6$$

(10)

$${\mathrm{F}\mathrm{e}}^{3+}+2{\mathrm{C}}_2{\mathrm{O}}_4^{2-}+2{\mathrm{e}}^{-}=\mathrm{F}\mathrm{e}{\left({\mathrm{C}}_2{\mathrm{O}}_4\right)}_2^{3-} \log \mathrm{K}=16.25$$

(11)

Based on the aforementioned issues, this research conducted a characterization and leaching study of mine tailings to evaluate their use as a value-added raw material for silver leaching in a Thiourea-Oxalate medium. Consequently, the study examines the acid drainage potential through acid-base balance assessments, silver (Ag), copper (Cu), iron (Fe), and arsenic (As) leaching profiles as a function of TU and Ox concentrations, as well as the direct influence of pH and temperature on this leaching process.

2. Materials and Methods

2.1. Chemical and Mineralogical Characterization

The tailings sample used in this study was provided by the mining company 'El Espiritu', which is located in the mining district of Zimapan, Hidalgo, Mexico [35]. From the total sample, a quartering process was performed to obtain five kilograms of a representative sample. Subsequently, particle size analysis was conducted using Tyler series sieves, following the methodology outlined by [36].

Each retained fraction was processed through digestion. For this purpose, one gram of each fraction was weighed separately and mixed with 20 mL of aqua regia (HCl and HNO₃ in a 3:1 ratio) for 60 minutes. The digestion solutions, along with the samples from the leaching experiments, were analyzed using Inductively Coupled Plasma Optical Emission Spectroscopy (ICP-OES) with a Perkin Elmer Model 8300 instrument.

The neutralization potential (NP) was determined using the modified acid-base accounting test (PM-ABA) established in the standard [37] and suggested by A. Ruiz-Sánchez [5]. For the determination of NP, one gram of the representative sample was placed on a watch glass, and 0.25 mL of 25% v/v HCl was added to observe its degree of effervescence (caused by the reaction of calcium carbonate with HCl), which was categorized as none, low, moderate, or strong.

Next, 2 grams of the sample were placed in a 250 mL Erlenmeyer flask containing 90 mL of distilled water, and magnetic stirring was set at a speed of 400 rpm. Based on the recorded effervescence level, volumes of 1 N HCl were added at the times indicated in Table 1. After 22 hours of stirring, the pH was measured. If the pH was greater than 2.5, additional HCl was added until a pH value of 2 was achieved. However, if the pH measured after 24 hours was less than 2, the measurement was repeated due to the possibility of excessive HCl addition at the beginning of the procedure. Finally, the titration was completed with 0.1 N sodium hydroxide (NaOH).

The NP value (g CaCO₃/kg) was calculated based on the stoichiometry of the reaction in Equation 12, which indicates a proportion of 1.37 g of CaCO₃ per gram of HCl consumed when neutralizing one kilogram of mine tailings.

$${CaCO}_{3\left(s\right)}+2HCL\to {CaCl}_2+H_{2}O+C{\mathrm{O}}_2$$

(12)

The acid potential (AP) was determined based on studies conducted by Petersen, (1969), Edwards et al. (1982) and A. Ruiz-Sanchez et al. (2022) [5,38,39], which approximate AP using the pyrite content in mine tailings through acid digestion. First, 0.2 g of the representative sample was digested with aqua regia (HCl and HNO₃ in a 3:1 ratio). Subsequently, another 0.2 g of the representative sample was digested with hydrochloric acid (HCl) to remove all iron except that from pyrite. The difference in iron concentrations was attributed to the iron from pyrite.

Additionally, 0.2 g was digested with nitric acid (HNO₃) to dissolve all forms of iron, including pyrite and sulfur. According to the standard [40], the difference in readings between digestion with HNO₃ and HCl corresponds to pyrite content. The AP value (g CaCO₃/kg) was determined from the pyrite content and the stoichiometry of the reaction (Eq. 13), which indicates a ratio of 1.67 g CaCO₃ per gram of FeS₂ for every kilogram of mine tailings.

$$4Fe{S}_2+8CaC{O}_3+15O_2+{6H}_{2}O\to 4Fe{\left(OH\right)}_{3\left(s\right)}+8S{O}_4^{2-}+{8Ca}^{2+}+{8CO}_2$$

(13)

Characterization of the mine tailings sample was performed using X-ray Diffraction (XRD) with an EQUINOX 2000 X-ray diffractometer using Co-Kα1 radiation (1.789010 Å) at 30 mA, 20 kV, and 220 V. The powdered samples were embedded in epoxy resin to form a pellet, which was polished to a mirror finish and analyzed via Scanning Electron Microscopy with Energy Dispersive Spectroscopy (SEM-EDS) using a JEOL JSM-6610LV system. More details on sample characterization can be found in a preliminary study by Erick Muñoz et al. 2023 [41].

2.2. Leaching Tests

Leaching experiments were designed to evaluate the dissolution behavior of silver, copper, iron, and arsenic from the tailings under varying conditions.

The following analytical-grade reagents were used for the experiments: Thiourea (CH₄N₂S, 99%, Sigma-Aldrich), Potassium Oxalate (K₂C₂O₄, 99.5%, Analytyka), Sulfuric Acid (H₂SO₄, 98%, Jalmek), and Deionized Water (1 μS/cm).

Based on the granulometric and metal content analysis (section 3.1), mine tailings with a particle size (r₀) of -53 +37 μm were selected. Unless otherwise specified, all experiments lasted 60 minutes. The initial TU and Ox concentrations were based on prior research D. Calla et al. (2016, 2019, 2020, 2021), I. Chandra et al. (2005), Xue-yi Guo et al. 2020 and Erick Muñoz et al. 2022 [19,31,32,33,34,42,43,44]. Reference parameters were set as follows: [TU]= 0.2 M, [Ox]= 0.2 M, 10 g sample, pH= 2, T= 30°C, and stirring speed= 600 rpm.

The leaching experiments used a Thermo Scientific HP88857190 heating plate, equipped with a 0.5 L reactor and a stirring system integrated with an IKA EW 20 motor and a Teflon propeller. The pH and temperature values were recorded using a pH/ATC electrode connected to a Thermo Scientific Orion potentiometer.

The pH value (1, 1.5, or 2) was adjusted in 0.5 L of the leaching solution (TU-Ox) before stirring was initiated at a constant speed of 600 rpm (determined in a prior study not included in this work). The temperature was set to 20, 30, or 35°C. Once the solution reached the target temperature, 20 g of mine tailings were added per liter of leaching solution. The moment the mineral was added was considered the start of the leaching process (t = 0 minutes). Every 10 minutes, approximately 10 mL of the leaching liquor was carefully sampled, filtered, and analyzed via ICP-OES to quantify the concentrations of Ag, Cu, Fe, and As.

The leaching percentage was calculated using Eq. 14 [45,46]:

$${\mathrm{X}}_{\mathrm{m}}={\displaystyle \frac{{\mathrm{C}}_{\mathrm{E}}\ast \mathrm{V}}{\mathrm{M}\ast \mathrm{x}}}\mathrm{x}100\mathrm{\%}$$

(14)

Where CE is the concentration of the element measured by ICP-OES (mg/L) in the leaching reactor, V is the total volume of solution (L), x is the mass fraction of the metal (%), and M is the mass of tailings used during leaching (g).

3. Results and Discussion

3.1. Granulometric Analysis and Metal Content

Figure 1 illustrates the percentage of mine tailings retained on each sieve during the granulometric analysis. As shown, 29.62% of the sample has a particle size greater than 150 μm, while 25.87%, 15.77%, 13.85%, and 6.33% correspond to -150 +105 μm, -105 +74 μm, -74 +53 μm, and -53 +37 μm, respectively.

Additionally, due to the presence of fine particles smaller than 37 μm (approximately 8.5% of the total sample), it is recommended to securely confine this type of waste. These fine particles pose a risk of being transported by wind, potentially reaching humans through inhalation, which can lead to respiratory diseases such as silicosis [47].

The elemental composition for each particle size range is presented in Table 2. As observed, the silver and copper content tends to increase as the particle size decreases, suggesting that silver phases are more likely to be liberated in smaller particles. Conversely, the composition of Fe and As remains nearly identical for particles within the -150 +105 μm to -74 +53 μm size range, with the highest concentration observed in the 37 μm sieve fraction. These results suggest that Fe and As originate from the same mineral phase, most likely arsenopyrite (FeAsS), which is confirmed by SEM in Section 3.3.

The variation in the metal contents shown in Table 2 is evidence that mine tailings are heterogeneous, for this reason in the leaching tests it is necessary to consider an overall composition ($\mathrm{\%}\overline{\mathrm{w}}{\mathrm{t}}_{\mathrm{j}}$) as a reference to help calculate the leaching percentages. For this, equation (15) reported in the work of Ruiz-Sánchez et al. (2022) was used where n corresponds to each of the 6 particle sizes, $\mathrm{\%}{\mathrm{w}}_{\mathrm{j}}$ is the percentage of mass retained in each fixed particle size n (Figure 1) present in the tailing, and $\mathrm{w}{\mathrm{t}}_{\mathrm{j}}$ represents the metallic composition for each particle size (Table 2), as a result an average concentration of 50 g/t, 1180 g/t, 77600 g/t and 5280 g/t was obtained for Ag, Cu, Fe and As, respectively.

$$\mathrm{\%}\overline{\mathrm{w}}{\mathrm{t}}_{\mathrm{j}}=\sum _{\mathrm{j}=1}^{\mathrm{n}}\left({\displaystyle \frac{\mathrm{\%}{\mathrm{w}}_{\mathrm{j}}}{100}}\right)\left(\mathrm{w}{\mathrm{t}}_{\mathrm{j}}\right)$$

(15)

Another important aspect about the particle size distribution and metallic composition (Table 3) is that about 72% and 52% of the total silver and copper are in particles smaller than 74 μm, so for a hydrometallurgical process of silver dissolution it is not necessary to perform an additional milling. Therefore, it is necessary to consider this type of waste as a value-added raw material to establish a complete hydrometallurgical process for silver dissolution.

Although the silver leaching process proposed in this study should ideally be conducted using the tailings sample as collected from the deposit site, it is evident that such processing complicates conversion calculations for the target metal. This is because the mass balance would require digesting the entire solid residue after each leaching experiment. Therefore, in the present study, only the -53 +37 μm fraction was selected, as it contains a nearly uniform silver concentration.

3.2. Neutralization Potential (NP) and Acid Potential (AP)

As a result of adding 0.25 mL of 25% v/v HCl for the NP calculation, a strong effervescence was observed, indicating a significant presence of alkaline species such as calcite (CaCO₃) and microcline (Al_1.03K_0.986Na_0.014Si_2.97O₈) identified in section 3.3. Accordingly, the neutralization potential was determined to be 116.5 g CaCO₃/kg (Table 5). The pyrite content determined by chemical methods showed comparable results with a minimal difference of approximately 1%. This discrepancy could be attributed to the higher aggressiveness of aqua regia, which can dissolve a greater amount of iron, or to the heterogeneity of the sample. Nevertheless, it is important to note that the sample contains at least 8% pyrite, which resulted in a calculated acid potential of 140 g CaCO₃/kg (Table 5).

The NP/AP ratio was 0.83 (Table 5). According to the standard [37], the tailings are classified as potential acid drainage generators. This highlights the importance of studying metal dissolution from this type of residue and suggests a way to utilize these residues as raw materials. Additionally, the addition of neutralizing agents such as CaCO₃, Ca(OH)2, CaO, Al(OH)₃, Na₂CO₃, or NaOH is recommended to mitigate their environmental impact and enhance their industrial reuse potential [48].

Table 4. Pyrite content analysis in tailings using chemical methods.

Iron content%
Aqua regia	HCL	Pyritic iron	Pyrite content (%)
4.3	0.6	3.8	8.1
HNO3	HCL	Pyritic iron	Pyrite content (%)
3.9	0.6	3.3	7.1

Table 4. Pyrite content analysis in tailings using chemical methods.

Iron content%
Aqua regia	HCL	Pyritic iron	Pyrite content (%)
4.3	0.6	3.8	8.1
HNO3	HCL	Pyritic iron	Pyrite content (%)
3.9	0.6	3.3	7.1

Table 5. Neutralization potential and acid potential relationship.

Parameter	Value
NP (g CaCO3 /kg)	116.5
AP (g CaCO3 /kg)*	140
Relación NP/AP	0.83
*Estimation considering 8.1% FeS2 obtained by chemical method.

Table 5. Neutralization potential and acid potential relationship.

Parameter	Value
NP (g CaCO3 /kg)	116.5
AP (g CaCO3 /kg)*	140
Relación NP/AP	0.83
*Estimation considering 8.1% FeS2 obtained by chemical method.

3.3. X-Ray Diffraction Analysis (XRD) and Scanning Electron Microscopy Analysis (SEM)

Figure 2 presents the diffractogram of the tailings sample along with the corresponding mineral identification using Match 3 software. The peaks identified in the diffractogram correspond to calcite [(CaO₃) (96-101-0963)] and microcline [(Al_1.03K_0.986Na_0.014Si_2.97O₈) (96-900-5304)], pyrite [(FeS₂) (96-901-0596)], sphalerite [(Zn_0.66Fe_0.34S) (96-101-1234)], argentite [(Ag₂S) (96-101-1338)], and polybasite [(Ag₃₁As_0.203CuSb_3.797S₂₂) (96-901-3300)]. These results suggest that silver leaching may not be fully efficient since part of the silver is present in the form of polybasite, a sulfosalt that is difficult to oxidize [49]. Additionally, it is worth mentioning that the mineral phases identified in this study have also been reported in other research exploring similar tailings [50,51], albeit from a different area within the deposit.

The spot analyses using energy dispersive spectroscopy (EDS) confirmed the abundant presence of calcite (CaCO₃, Figure 3a) and feldspar (KAlSi₃O₈, Figure 3b), the latter associated with microcline. This characterization also identified mineral phases such as sphalerite (ZnS, Figure 4a) and arsenopyrite (FeAsS, Figure 4b). EDS analysis further confirmed the presence of pyrite (FeS₂, Figure 5a) and wüstite (FeO, Figure 5b) as iron-contributing mineral phases. Consequently, the presence of sulfides suggests that pyrite, sphalerite, and arsenopyrite may lead to acid drainage. This occurs because the exposure of these sulfides to moisture and an oxidizing agent, such as atmospheric air, promotes the generation of acid drainage (Equations 16 to 23) [52,53,54].

$${\mathrm{F}\mathrm{e}\mathrm{S}}_2+{\displaystyle \frac{7}{2}}{\mathrm{O}}_2+{\mathrm{H}}_2\mathrm{O}\to {\mathrm{F}\mathrm{e}}^{2+}+2{\mathrm{S}\mathrm{O}}_4^{2-}+2{\mathrm{H}}^{+}$$

(16)

$${\mathrm{F}\mathrm{e}}^{2+}+{\displaystyle \frac{1}{4}}{\mathrm{O}}_2+{\mathrm{H}}^{+}\to {\mathrm{F}\mathrm{e}}^{3+}+{\displaystyle \frac{1}{4}}{\mathrm{H}}_2\mathrm{O}$$

(17)

$${\mathrm{F}\mathrm{e}}^{3+}+3{\mathrm{H}}_2\mathrm{O}\to {\mathrm{F}\mathrm{e}\left(\mathrm{O}\mathrm{H}\right)}_{3\left(\mathrm{s}\right)}+3{\mathrm{H}}^{+}$$

(18)

$${\mathrm{F}\mathrm{e}\mathrm{S}}_2+{\displaystyle \frac{15}{4}}{\mathrm{O}}_2+{\displaystyle \frac{7}{3}}{\mathrm{H}}_2\mathrm{O}\to {\mathrm{F}\mathrm{e}\left(\mathrm{O}\mathrm{H}\right)}_{3\left(\mathrm{s}\right)}+2{\mathrm{S}\mathrm{O}}_4^{2-}+4{\mathrm{H}}^{+}$$

(19)

$$\mathrm{F}\mathrm{e}\mathrm{A}\mathrm{s}\mathrm{S}+{\displaystyle \frac{7}{2}}{\mathrm{O}}_2+4{\mathrm{H}}_2\mathrm{O}\to {\mathrm{F}\mathrm{e}\left(\mathrm{O}\mathrm{H}\right)}_{3\left(\mathrm{s}\right)}+{\mathrm{S}\mathrm{O}}_4^{2-}+{\mathrm{H}}_2\mathrm{A}\mathrm{s}{\mathrm{O}}_4^{-}+4{\mathrm{H}}^{+}$$

(20)

$$4\mathrm{F}\mathrm{e}\mathrm{A}\mathrm{s}\mathrm{S}+11{\mathrm{O}}_2+6{\mathrm{H}}_2\mathrm{O}\to 4{\mathrm{F}\mathrm{e}}^{2+}+4{\mathrm{H}}_3\mathrm{A}\mathrm{s}{\mathrm{O}}_3+4{\mathrm{S}\mathrm{O}}_4^{2-}$$

(21)

$$2{\mathrm{H}}_3\mathrm{A}\mathrm{s}{\mathrm{O}}_3+{\mathrm{O}}_2\to 2\mathrm{H}\mathrm{A}\mathrm{s}{\mathrm{O}}_4^{2-}+4{\mathrm{H}}^{+}$$

(22)

$$2{\mathrm{H}}_3\mathrm{A}\mathrm{s}{\mathrm{O}}_3+{\mathrm{O}}_2\to 2\mathrm{H}\mathrm{A}\mathrm{s}{\mathrm{O}}_4^{-}+4{\mathrm{H}}^{+}$$

(23)

If sufficient Fe³⁺ ions are present at the deposition site, they act as an oxidizing agent [55], enabling the oxidation of other sulfides, such as sphalerite (Equation 24) [56]:

$$ZnS+{2Fe}^{3+}\to {Zn}^{2+}+{2Fe}^{2+}+S^0$$

(24)

On the other hand, for the mineral phases contributing copper and silver, EDS analyses confirmed the presence of polybasite ((Ag,Cu)₁₆Sb₂S₁₁, Figure 6a) and a silver sulfide (argentite), both phases previously identified by XRD (Figure 6b). However, the micrographs reveal a natural association between polybasite and argentite, similar to that observed with other sulfides such as galena, pyrite, and arsenopyrite [57,58].

3.4. Thiourea Concentration

Figure 7a shows the percentage of silver leaching. As can be seen, the percentage of silver leaching increases with higher TU concentration. The maximum value of silver leaching at 60 min is 52.96%, 65.05%, and 82.40% for 0.03 M, 0.1 M, and 0.2 M, respectively. However, the results show that the presence of TU is important since in its absence (0 M TU) about 15% is leached.

A tentative explanation for the low percentage of silver leaching with only oxalate is due to the fact that Ag(I)-Ox complexes has lower stability (lower log K values) which means that the presence of Ox^2- ligands do not promote the increase in solubility since the only oxidant in the medium is dissolved oxygen around 8 mg/L. Instead, the increase in silver leaching rates is due to the fact that Ag(I)-TU complexes are more stable (higher log K) compared to Ag(I)-Ox complexes as shown in Equations 25 to 27.

In addition, the presence of TU not only helps the stability of Ag(I)-TU complexes but also promotes the formation of formamidine disulfide ((CS(NH₂)₂)₂²⁺), as shown in reactions 28 and 29. The species CS(NH₂)₂)₂²⁺ being one more oxidizing species together with oxygen to promote the oxidation of silver, which in the case of Ag₂S can be represented by the reactions.

$${\mathrm{A}\mathrm{g}}^{+}+{\mathrm{O}\mathrm{x}}^{2-}\to \mathrm{A}\mathrm{g}\left({\mathrm{O}\mathrm{x}}^{-}\right) \log \mathrm{K}=3.02$$

(25)

$${\mathrm{A}\mathrm{g}}^{+}+2\mathrm{C}\mathrm{S}{\left({\mathrm{N}\mathrm{H}}_2\right)}_2\to \mathrm{A}\mathrm{g}{{\left(\mathrm{C}\mathrm{S}{\left({\mathrm{N}\mathrm{H}}_2\right)}_2\right)}_2}^{+} \log \mathrm{K}=10.35$$

(26)

$${\mathrm{A}\mathrm{g}}^{+}+3\mathrm{C}\mathrm{S}{\left({\mathrm{N}\mathrm{H}}_2\right)}_2\to \mathrm{A}\mathrm{g}{{\left(\mathrm{C}\mathrm{S}{\left({\mathrm{N}\mathrm{H}}_2\right)}_2\right)}_3}^{+} \log \mathrm{K}=12.87$$

(27)

$$\mathrm{C}\mathrm{S}{\left({\mathrm{N}\mathrm{H}}_2\right)}_2+{\mathrm{H}}^{+}\to {\mathrm{H}\mathrm{C}\mathrm{S}{\left({\mathrm{N}\mathrm{H}}_2\right)}_2}^{+} \log \mathrm{K}=1.44$$

(28)

$${2\mathrm{H}\mathrm{C}\mathrm{S}{\left({\mathrm{N}\mathrm{H}}_2\right)}_2}^{+}\to {{\left(\mathrm{C}\mathrm{S}{\left({\mathrm{N}\mathrm{H}}_2\right)}_2\right)}_2}^{+}+{2\mathrm{H}}^{+}+{2\mathrm{e}}^{-} \log \mathrm{K}=14.2$$

(29)

$${\mathrm{A}\mathrm{g}}_2\mathrm{S}+6\mathrm{C}\mathrm{S}{\left(\mathrm{N}{\mathrm{H}}_2\right)}_2+{{\displaystyle \frac{1}{2}}\mathrm{O}}_{2\left(\mathrm{a}\mathrm{c}\right)}+{2\mathrm{H}}^{+}\leftrightarrow 2\mathrm{A}\mathrm{g}{{\left[\left(\mathrm{C}\mathrm{S}{\left(\mathrm{N}{\mathrm{H}}_2\right)}_2\right)\right]}_3}^{+}+{\mathrm{S}}^0+{\mathrm{H}}_2\mathrm{O}$$

(30)

$${\mathrm{A}\mathrm{g}}_2\mathrm{S}+4\mathrm{C}\mathrm{S}{\left(\mathrm{N}{\mathrm{H}}_2\right)}_2+{{\left(\mathrm{C}\mathrm{S}{\left({\mathrm{N}\mathrm{H}}_2\right)}_2\right)}_2}^{+}\leftrightarrow 2\mathrm{A}\mathrm{g}{{\left[\left(\mathrm{C}\mathrm{S}{\left(\mathrm{N}{\mathrm{H}}_2\right)}_2\right)\right]}_3}^{+}{+\mathrm{S}}^0$$

(31)

As for copper (Figure 7b), it is observed that in the absence of TU the percentage of copper leaching was around 5.22% at 20 minutes. This percentage does not increase probably because in the leaching solution equilibrium is reached for the formation of copper oxalate as shown by the Pourbaix type diagram in Figure 8 and described by equation 32. The increase of at least 5% more in silver dissolution can be attributed to the formation of TU-Cu(II) complexes that promote the solubility of copper in the form of solid copper oxalate. While the increase in copper solubility from solid copper oxalate is not as significant due to the higher log K (Equation 33) for the Ox-Cu(II) complexes compared to the TU-Cu(II) complexes whose thermodynamic constants are shown in equations 34 to 37.

$${\mathrm{C}\mathrm{u}}^{2+}+{\mathrm{O}\mathrm{x}}^{2-}\to \mathrm{C}\mathrm{u}{\left(\mathrm{O}\mathrm{x}\right)}_{\left(\mathrm{s}\right)} \log \mathrm{K}=11.19$$

(32)

$${Cu}^{2+}+{2Ox}^{2-}\to Cu{{\left(Ox\right)}_2}^{2-} \log \mathrm{K}=10.23$$

(33)

$${Cu}^{2+}+CS{\left({NH}_2\right)}_2\to Cu{\left(CS{\left({NH}_2\right)}_2\right)}^{2+} \log \mathrm{K}=0.8$$

(34)

$${Cu}^{2+}+2CS{\left({NH}_2\right)}_2\to Cu{{\left(CS{\left({NH}_2\right)}_2\right)}_2}^{2+} \log \mathrm{K}=0.8$$

(35)

$${Cu}^{2+}+3CS{\left({NH}_2\right)}_2\to Cu{{\left(CS{\left({NH}_2\right)}_2\right)}_3}^{2+} \log \mathrm{K}=0.9$$

(36)

$${Cu}^{2+}+4CS{\left({NH}_2\right)}_2\to Cu{{\left(CS{\left({NH}_2\right)}_2\right)}_4}^{2+} \log \mathrm{K}=1.1$$

(37)

The analysis of iron (Figure 9a) and arsenic (Figure 9b) leaching shows similar behaviors, suggesting that the leaching process for these metals originates from the oxidation of arsenopyrite. This is deduced from the dissolution percentage ratios of Fe and As at 60 minutes in the presence of 0.2 M TU, with values of 19.44% and 25.68% (As/Fe ratio of 1.32), which is numerically close to the 1.34 ratio derived from the molecular weights of As and Fe in the arsenopyrite chemical formula.

On the other hand, the behavior during the first 20 minutes of Fe and As leaching percentages does not show a clear trend, possibly due to sample heterogeneity, the natural association of arsenopyrite with other mineral phases, or the occlusion of arsenopyrite within other mineral phases present in the tailings. However, it is evident that TU presence promotes arsenopyrite dissolution, as oxalate stabilizes thiourea, preventing its oxidation by arsenopyrite itself or by the Fe (III) resulting from arsenopyrite oxidation, as suggested by Ke Li et al. (2024) in their study on arsenopyrite oxidation using Fe (III) ions with thiourea. Based on Ke Li et al.’s work, the arsenopyrite leaching process can be described by reaction 38 [60].

$$FeAsS+{{\displaystyle \frac{3}{2}}O}_{2\left(ac\right)}+{2Ox}^{2-}\leftrightarrow {Fe{\left(Ox\right)}_2}^{-}+{{AsO}_3}^{3-}{+S}^0$$

(38)

3.5. Oxalate Concentration

To evaluate the impact of oxalate ion concentration on the leaching of Ag, Cu, Fe, and As, the TU concentration was fixed at 0.2 M, along with the parameters defined in the methodology (see Section 2.2: Leaching System). Figure 10a shows the silver percentages obtained with and without the presence of oxalate. In the absence of oxalate, the maximum silver dissolution percentage reached was 25.77%. However, it is evident that the presence of oxalate, ranging from 0.0012 M to 0.3 M L⁻¹, promotes an increase in silver leaching percentages up to 87.70%.

The benefit of using oxalate in silver leaching is that it prevents the decomposition of TU by arsenopyrite. In the absence of oxalate, as reported by Ke Li et al. (2024), arsenopyrite catalyzes the decomposition of thiourea into formamidine disulfide, an unstable species that easily decomposes. Consequently, TU is depleted, and the leaching process reaches equilibrium within 5 minutes [60].

In addition to the above analysis, it is also observed that the increase in the oxalate concentration from 0.2M to 0.3M does not significantly benefit the dissolution of silver in the form of Ag(CS(NH₂)₂)₃⁺. This is probably due to the fact that the Ag(I)-Tu complexes have a larger solubility window in potential and pH as shown in the Pourbaix diagram (Figure 11); which does not occur for the Ag(I)-Ox system where the formation of the Ag(Ox)⁺ complex (Figure 12) is at the limit of the maximum pH value in the experiments of this study.

According to Figure 11 and Figure 12, both TU and oxalate ions can form complexes with silver. This suggests a synergistic interaction that explains the improved leaching percentages within a shorter time frame. This complex interaction between silver, thiourea and oxalate suggests it has also been reported by Xiyun Yang et al. (2011) [61] when using thiourea with thiocyanate. For the above mentioned, it is considered that the best leaching conditions are the use of 0.2M of Tu and Ox, respectively.

The copper leaching profiles (Figure 10b) show that in the absence of oxalate ions, the maximum copper concentration (30.63%) is achieved, which simultaneously implies maximum dissolution of silver associated with copper, specifically from polybasite. Given that the Ag/Cu mass ratio in polybasite is approximately 1.64, it can be inferred that about 50% of the total leached silver (87.7%) originates from polybasite. Additionally, it is observed that increasing the oxalate concentration to 0.3 M results in a decrease in dissolved copper due to the formation of solid copper oxalate (Figure 8), as demonstrated by Ruiz-Sánchez et al. (2020) in the chalcopyrite leaching process using organic ligands [62].

Consequently, the competition between Cu ions and Ag ions for the formation of complexes with thiourea is reduced.

Figure 13a,b show leaching profiles for iron and arsenic with similar trends and relationships to the results in Figure 9, supporting the assertion that these metals in aqueous form originate from arsenopyrite oxidation. Interestingly, in the absence of oxalate, the dissolution of copper and iron follows an almost linear trend, suggesting rapid and preferential dissolution of arsenopyrite. This behavior indicates that, without oxalate ions, arsenopyrite catalyzes the decomposition of thiourea into formamidine disulfide, an unstable species that leads to cyanamide (CH₂N₂) and elemental sulfur (S⁰) [27]. Consequently, the silver phases present are not significantly oxidized, resulting in a silver recovery of only about 20% (Figure 10).

3.6. pH Effect

Figure 14a presents the silver leaching percentages as a function of the initial pH of the leaching solution. The results indicate that acidic pH levels enhance the dissolution of silver phases, achieving 88%, 87%, and 86% silver recovery within the first 5 minutes for pH values of 1, 1.5, and 2, respectively. This behavior is attributed to the precipitation of copper as copper oxalate (Figure 14b) on the surface of silver phases at pH = 1, which creates a transport resistance for reactants toward the reaction surface. This phenomenon is reflected in the slower increase in silver leaching profiles for pH = 1 and 1.5 during the first 5 minutes.

Iron (Figure 15a) and arsenic (Figure 15b) leaching experienced a notable increase as the pH decreased. For Fe, leaching percentages of 29.48% and 21.16% were observed at pH = 1 and 1.5, respectively. In the case of As, leaching percentages reached 27.23% and 17.55%. Lowering the pH triggers significant changes in species chemistry, favoring the formation of soluble species such as Fe(Ox)⁺, Fe(Ox)₂⁻ and Fe(Ox)₃³⁻, as illustrated in the species distribution diagram (Figure 16).

3.7. Temperature Effect

Figure 17a shows the silver leaching profiles as a function of temperature. The results indicate that increasing the temperature enhances the dissolution of silver phases. This effect suggests chemical reaction control, as raising the temperature from 20°C to 35°C allows for the recovery of 90% of the available silver. Additionally, it is observed that higher temperatures accelerate equilibrium attainment, requiring only 5 minutes of operation at 35°C. This effect implies that higher temperatures improve the dissolution kinetics of silver but simultaneously promote the rapid decomposition of thiourea, preventing the leaching of the remaining 10% of silver. In fact, regardless of the operating temperature, it was found that after 45 minutes, thiourea concentration had decreased by more than 70% from its initial value (0.2 M), which was already in excess relative to the stoichiometry of reactions 31 and 32.

The leaching profiles for copper (Figure 17b), iron and arsenic (Figure 18a y 18b) were also enhanced by increasing the temperature.

4. Conclusions

• The TU-Ox system represents a suitable option for recovering precious metals present in tailings. This approach adds value to these residues, which are rich in sulfides and may potentially generate acid drainage in the future.
• The most critical variables for optimizing this process are the temperature and the concentrations of thiourea and oxalate.
• The TU-Ox system demonstrated improvements in the leaching process, particularly in the presence of arsenopyrite—a mineral phase that catalyzes the decomposition of thiourea.
• Oxalate helps stabilize thiourea, ensuring its availability to form complexes with silver. However, increasing oxalate concentration can lead to the formation of solid copper oxalate, which creates resistance to the transport of reagents and products.
• Higher temperatures enhance silver leaching, suggesting that the process is chemically controlled. However, elevated temperatures also accelerate thiourea decomposition.