Recovery of Copper from Pregnant Leach Solutions of Copper Concentrate Using Aluminum Shavings

Abstract

Copper is one of the most used metals today due to its wide range of applications. Traditionally, this metal has been primarily extracted through pyrometallurgical methods, which presents several environmental and energy-related drawbacks. An alternative is hydrometallurgy, which has achieved acceptable copper extraction rates. However, this process has not found widespread industrial application due to operational challenges and the complexity associated with the selective recovery of copper ions from the Pregnant Leach Solution (PLS), especially due to the coexistence of copper and iron ions, complicating the efficient separation of both metals. In this work, the use of aluminum shavings as a cementation agent is proposed, analyzing variables such as the initial shaving concentration (2.5, 5, 10, 15, and 20 g/L), the agitation speed (0, 200, and 400 rpm), and a temperature of 20, 30, and 40 °C. The results demonstrated selective copper cementation, achieving a 100% recovery in 30 min under stirring conditions of 400 rpm. The analysis performed using X-ray Diffraction (XRD) and Scanning Electron Microscopy (SEM) revealed the formation of solid phases such as metallic copper (Cu), aluminum hydroxide [Al(OH)₃], and elemental sulfur (S). Additionally, it was observed that the iron ion concentration remained constant throughout the experiment, indicating a high selectivity in the process. The kinetic analysis revealed that the reaction follows a first-order model without stirring. An activation energy of 62.6 kJ/mol was determined within the experimental temperature range of 20–40 °C, confirming that the process fits the chemical reaction model. These findings provide a deeper understanding of the system’s behavior, highlighting its feasibility and potential for industrial-scale applications.

1. Introduction

Currently, high-grade copper reserves are being depleted, and the available reserves primarily contain low-grade copper. This has led to a shift in processing technologies, replacing traditional pyrometallurgical methods with hydrometallurgical processes more suitable for treating lower-quality ores. Copper is primarily found in nature in two forms: sulfides and oxides. While oxidized ores can be directly leached to transfer copper to the aqueous phase, the leaching of sulfide ores presents a greater challenge due to the refractory nature of copper to chemical reagents. The composition and purity of the solution obtained during the leaching process significantly depend on the mineralogical structure of the ore and the reagents used. It is important to note that leaching solutions generated using strong acids tend to contain a higher number of impurities, which can complicate subsequent metal recovery stages. Leaching solutions with a low impurity content can be obtained using weak acids, bases, and salts [1].

In this sense, chalcopyrite is one of the most complex minerals to be processed by hydrometallurgical means, due to the great passivity it presents in environmental conditions. To avoid this, several proposals have been made, such as the use of strong oxidizing agents, catalysts, and organic solvents, among others. The use of organic solvents in conjunction with hydrogen peroxide has obtained acceptable results of 90% copper extraction, but there is still a need for further research and resolution of some drawbacks to be able to take the process to an industrial scale. One of the main drawbacks is the existence of a co-dissolution of iron and copper, economically and technically complicating the recovery of copper from the pregnant solution in the hydrometallurgical process and limiting the subsequent techniques that are usually used, such as solvent extraction and/or electrowinning [2]. A technique that has been used for centuries in hydrometallurgy for the purification of pregnant solutions in leaching is cementation [3,4,5,6], which is a heterogeneous galvanic process where the noblest metal ion is reduced with a metal with greater activity, depositing on its surface [7] as an electrochemical precipitation.

$$M_1^{+Z_1}+{\displaystyle \frac{Z_1}{Z_2}}{M}_2^0\leftrightarrow {\displaystyle \frac{Z_1}{Z_2}}{M}_2^{+Z_2}+M_1^0$$

(1)

This involves the displacement of a half-reaction from a more active metal, $M_2$,

$$M_2^0\leftrightarrow M_2^{+Z_2}+z_2{e}^{-}$$

(2)

by a more noble one, $M_1$.

$$M_1^{+Z_1}+z_1{e}^{-}\leftrightarrow M_1^0$$

(3)

In Equation (1), each term represents the species involved in the cementation reaction, expressed in a general stoichiometric form:

$M_1^{+Z_1}$ is the cation of the more noble metal $M_1$ with oxidation state $+Z_1$. This is the species that will be reduced and deposited as metallic $M_1^0$.

$M_2^0$ is the metallic form of the more active (less noble) metal $M_2$, which serves as the reductant in the process.

${\displaystyle \frac{Z_1}{Z_2}}$ is the stoichiometric coefficient that balances the charge transfer between the two metals, ensuring electron conservation during the redox process.

$M_2^{+Z_2}$ is the cation of the more active metal $M_2$, generated after the oxidation of $M_2^0$.

$M_1^0$ is the metallic deposition of the more noble metal $M_1$, formed on the surface of $M_2$ during the cementation process.

Overall, Equation (1) describes the redox displacement reaction where the ions of the nobler metal $M_1$ are reduced and precipitate as metallic copper (or another noble metal), while the more active metal $M_2$ undergoes oxidation and passes into the solution as cations.

On the other hand, copper can be cemented by various metals, including Zn, Fe, and Al. The one that has shown the greatest efficiency is zinc, but it has the drawback of the co-cementation of iron and copper. The process of recovering copper with iron has been used for more than a century, showing great efficiency that has been the most used industrially. A disadvantage of this process is the excessive consumption of iron, sometimes requiring up to 260% stoichiometrically [6]. In the case of aluminum, it has been observed that it cements copper with an efficiency of 100% in alkaline solutions at temperatures above 50 °C [7]. On the other hand, it has been found that, when the cupric ion is in acidic conditions, the cementation kinetics are very slow and inefficient [8]. In this sense, Karavasteva [9] attributes this behavior to the formation of a layer of aluminum oxide on its surface, limiting the contact of the reactants, substantially inhibiting the cementation of copper. In this regard, thermodynamics predicts that both the copper and iron cementation reactions are spontaneous, with cell potentials of 2 V and 1.63 V, respectively [10,11].

$${2Al}^0+3{Cu}^{2+}⟶{2Al}^{3+}+{3Cu}^0{E}^0=2.001\mathrm{V}$$

(4)

$${Al}^0+{Fe}^{3+}⟶{Al}^{3+}+{Fe}^0{E}^0=1.63\mathrm{V}$$

(5)

Therefore, in this work, the cementation of copper is studied using aluminum shavings as a reducing agent of a pregnant solution in the leaching of chalcopyrite concentrate in an acidic medium and hydrogen peroxide, using methanol as a stabilizer of the cuprous ion. The copper ion is then considered to be present in the pregnant solution as a cuprous ion, increasing the cell potential to 2.196 V, which means that it tends to happen more easily (Equation (6)).

$${3Cu}^{+}+{Al}^0⟶{3Cu}^0+{Al}^{3+}{E}^0=2.196\mathrm{V}$$

(6)

Copper cementation with aluminum and zinc [12,13,14,15,16] and iron [17,18] has been studied by different researchers, and this topic is still widely investigated.

2. Materials and Methods

2.1. Experimental Development

The experimental aqueous solution used comes from the leaching of a chalcopyrite concentrate with ${\mathrm{H}}_2\mathrm{S}{\mathrm{O}}_4$, ${\mathrm{H}}_2{\mathrm{O}}_2$, and methanol at 40 °C for 3 h, following the methodology used by Solís-Marcial and Lapidus [19,20].

Table 1. Parameters of pregnant solutions of copper and iron sulfate with methanol used in the experimental tests before and after copper cementation.

Stage	Concentration of Cu [g/L]	Concentration of Fe [g/L]	pH	Redox Potential vs. (SHE) (mV)
Initial	53.5	51.7	1.9	637
Final	0.5	51	1.7	−337

shows the parameters of the pregnant solution used in the experimental tests before and after copper cementation.

The concentration of copper and iron was determined by the atomic absorption spectrophotometry technique with a Bunsen Spectrophotometer (Madrid, Spain). The pH and ORP values of the solution were measured by a pH-meter NS 17 Mcra-Elwro (Wrocław, Poland) with a combined glass–calomel electrode type ERH-I I Hydromet (Stroze, Poland). The aluminum used was primarily rinsed with distilled water and acetone; its form was shavings with an average size of 3 cm, being obtained by mechanizing an aluminum bar on a lathe at the Institute. The cementation tests were carried out by placing 100 mL of the PLS mentioned above, varying the amount of the cementing agent, stirring speed, and temperature maintained with heating plate. Samples were taken at established times and analyzed through Atomic Absorption Spectroscopy for copper and iron. At the end of the experiments, the residues were filtered and dried in the environment to be analyzed with X-Ray Diffraction (XRD), which was conducted in a Bruker AXS (Bruker, Billerica, MA, USA) (Advanced X-Ray solutions) D8 X-ray diffractometer with a Cu Kα (λ-1.5406 Å) radiation. To characterize the formed copper and the solid phases of aluminum remaining in the solid, a mineralogical characterization study was determined using a Carl Zeiss Sigma VP automated Scanning Electron Microscope instrument (SEM), equipped with a Bruker Quantax X-Flash 6130 Energy Dispersive X-ray (EDX) detector.

2.2. Thermodynamic Analysis

Chalcopyrite can be oxidized by hydrogen peroxide mixed with sulfuric acid and methanol solutions according to the following reactions [19,20]:

$$CuF{eS}_2+3H^{+}+{{3/2H}_{2}O}_{2\left(aq\right)}+R\to C{u}^{+}R+F{e}^{2+}+{2S}^0+3H_{2}O$$

(7)

where $R$ is the methanol.

Nevertheless, in concentrated sulfated solutions, copper and iron form several soluble complexes (NIST (2004)) [21]:

$$C{u}^{2+}+S{O}_4^{2-}\to CuS{O}_{4\left(aq\right)}{log}_{10}k=2.1$$

(8)

$$F{e}^{2+}+H^{+}+S{O}_4^{2-}\to FeHS{O}_4^{+}{log}_{10}k=3.068$$

(9)

$$F{e}^{2+}+S{O}_4^{2-}\to FeS{O}_{4\left(aq\right)}{log}_{10}k=2.25$$

(10)

$$F{e}^{2+}+2S{O}_4^{2-}\to Fe(S{O}_4{)}_2^{2-}{log}_{10}k=-7.6$$

(11)

$$F{e}^{2+}+H^{+}+S{O}_4^{2-}\to FeHS{O}_4^{2+}+e^{-}{log}_{10}k=-8.552$$

(12)

$$F{e}^{2+}+S{O}_4^{2-}\to FeS{O}_4^{+}+e^{-}{log}_{10}k=-8.98$$

(13)

When the aqueous complexes of copper and iron are in contact with metallic aluminum, a series of precipitation and complexation reactions occur (NIST):

$${Cu}^{2+}+2e^{-}\to C{u}_{\left(S\right)}{log}_{10}k=11.395$$

(14)

$${Fe}^{2+}+2e^{-}\to F{e}_{\left(S\right)}{log}_{10}k=-16.097$$

(15)

$${Al}_{\left(s\right)}\to A{l}^{3+}+3e^{-}{log}_{10}k=84.3$$

(16)

$${Al}^{3+}+S{O}_4^{2-}\to AlS{O}_4^{+}{log}_{10}k=3.5$$

(17)

$${Al}^{3+}+2S{O}_4^{2-}\to Al(S{O}_4{)}_2^{-}{log}_{10}k=5$$

(18)

$${Al}^{3+}+H^{+}+S{O}_4^{2-}\to AlHS{O}_4^{2+}{log}_{10}k=2.488$$

(19)

To better understand the phases of the copper and iron solutions and, particularly, the solubility zones and precipitation, species distribution diagrams in the presence of sulfate ions were constructed, and the thermodynamic information necessary to construct the predominance diagrams (Eh versus pH) for the three metals considered in sulfated solutions was obtained from the NIST database 46, version 8.0 (NIST, 2004) and was integrated into the MEDUSA Eq-Calcs_32 software suite (Making Equilibrium Diagrams Using Simple Algorithms) [22,23]. The diagrams presented below were constructed considering a solution with a composition of 1.0 M sulfate, 0.8 M Cu(II), 0.8 M Fe(II), and 1.0 M Al(III), which corresponds to experimental conditions.

The diagrams for each metal are shown in Figure 1a (Fe), b (Cu), and c (Al), where the conditions of pH and solution potential of the soluble, metallic, and oxide species can be identified. The predominance diagram for iron shows soluble species within a pH range of 0 to 6 and potential from −0.5 to 0.8 V. Outside this area, metallic iron and goethite are present. For copper, there is only one soluble species at potentials above 0.3 V and pH values below 4; around this region, solid species are found. For aluminum, there are three aqueous complexes at potentials below −1.6 V and pH values from 0 to 4. Around this area, solid species such as Al and $Al(OH{)}_3$ are present. The final conditions of the cemented solution were pH 1.7 and potential −337 mV vs. SHE, which explains why the iron is not cemented and the copper is, given that the commented conditions favored the formation of the ferric ion, metallic copper, and aluminum ion.

Additionally, the presence of metallic aluminum could cement (reduce) the copper ions. This characteristic is especially important, from an environmental perspective for copper, as it is toxic and can cause serious problems if copper ions are present. On the other hand, thermodynamic diagrams predict the cementation stage of aqueous copper sulfate that occurs at a potential 0.3 V with the onset of hydrogen gas formation.

The predominance diagram for sulfur (Figure 2) shows the stable species across a range of pH and electrochemical potentials. Elemental sulfur $\left[S_{\left(s\right)}\right]$ is stable within a potential window of approximately 0 to 0.2 V and pH values between 1 and 5. At higher oxidizing potentials and acidic conditions (above ~0.3 V and pH < 2), sulfur is found as bisulfate (${HS{O}_4}^{-}$), while at slightly higher pH values (pH > 2), sulfate (${S{O}_4}^{2-}$) becomes the dominant species. In more reducing conditions (below 0 V), hydrogen sulfide ($H_{2}S$) is the predominant species at low pH, and at higher pH values (above ~5) the $H{S}^{-}$ ion is more stable.

3. Results

Cementation Test

To better illustrate the experimental design and the parameters investigated, Figure 3 provides a schematic overview of the study. The central diagram depicts the reactor configuration, including temperature control, agitation, and pH monitoring. Around it, the plots highlight the influence of critical variables on cementation kinetics: the effects of the initial metal concentration, stirring speed, and temperature. Complementary characterization was performed using XRD, confirming the crystalline phases of the precipitates, and using SEM-EDS, which revealed the morphology and distribution of the cemented copper on the substrate. This figure integrates both the experimental setup and the main results, providing a concise visualization of the methodology and findings.

A key variable in the cementation process is the initial amount of reducing metal, since this variable can directly influence the economic viability of the process. Figure 4 shows the percentage of cemented copper as a function of time, varying the aluminum concentration (5, 10, 15, 20, 25, and 30 g/L) without agitation. The graph shows an increasing linear behavior in the degree of cementation as time elapses, for all concentrations studied. In addition, a directly proportional relationship between the amount of aluminum shavings and the efficiency of copper cementation is evident. This phenomenon could be explained by the fact that, as the amount of aluminum increases, the active surface available for the reaction also increases. According to the work of Djokic [7], the experimental data fit acceptably to zero-order kinetics. In fact, complete cementation is achieved with concentrations of 25 and 30 g/L after 120 min of experimentation.

The stirring speed is another variable that significantly influences cementation kinetics, due to its impact on mass transfer phenomena. To this effect, three experimental runs were carried out at room temperature, with a concentration of 20 g/L of aluminum shavings and stirring speeds of 0, 200, and 400 rpm. The results are presented in Figure 5. It is observed that the curves present a non-linear behavior, having higher cementation kinetics with the increase in the agitation speed. In the case of 400 rpm, 100% copper precipitated is obtained after 30 min. This behavior may indicate a possible change in the process control stage.

The shrinking core model is used for kinetic analysis, considering that is the mechanism for the cementation of copper with aluminum, in which the chemical reaction controls the process under the assumption of a spherical geometry of the particles. According to the results shown in Figure 6, this model adequately fits the experimental data obtained under stirring conditions with a correlation coefficient higher than 0.9, which is acceptable, where the increase in stirring speed diminished the mass transfer problems. On the other hand, in the absence of stirring, the system exhibits zero-order kinetics, suggesting that the control of the process is not dominated by the aluminum concentration. Furthermore, it is noted that cementation presents inherent challenges related to mass transfer, particularly in systems with a high reactant concentration. This behavior highlights the need to optimize operating conditions, such as the agitation intensity and particle size, to maximize process efficiency.

Another fundamental variable that significantly influences the cementation process is the temperature [4]. To evaluate its effect, mainly when the process control mechanism is the chemical reaction at the solid–liquid interface, experiments were carried out at 20, 30, and 40 °C, with a concentration of 5 g/L of the leaching agent and without agitation of the system. The results obtained, shown in Figure 7, show a clear tendency to increase the reaction rate with temperature, in accordance with the kinetic theory of heterogeneous processes. A significant increase in the cementation fraction is observed, going from 0.25 at 20 °C to 0.7 at 30 °C, and reaching 0.9 at 40 °C.

Figure 8 shows the plots corresponding to the fit of the above data using the shrinking core model, controlled by the superficial reaction. Based on this figure, it can be argued that, with increasing temperature, the diffusion resistance at the boundary layer of the solid–liquid interface is reduced, facilitating the transport of cupric ions to the surface of the cementitious metal. This effect is especially relevant in unstirred systems, where mass transfer is often the limiting factor in reaction kinetics.

An activation energy of 62.6 kJ/mol was determined in the temperature range of 20–40 °C (Figure 9), determined by Arrhenius equation. Constants were obtained from the slope in Figure 8 for each temperature, coinciding in that the reaction rate is predominantly controlled by a chemical reaction mechanism within this thermal range. This value is consistent with systems where cementation kinetics are governed by the charge transfer phenomenon at the solid–liquid interface, rather than being limited by the diffusion of reactive species through a boundary layer. The activation energy was determined using the Arrhenius method by calculating rate constants at 20 °C, 30 °C, and 40 °C. A plot of ln(k) versus 1/T was constructed, and the slope of the regression line was used to obtain a value of 62.6 kJ/mol.

From a thermodynamic perspective, the increase in temperature not only accelerates the kinetics of the process, but also directly influences the redox potential of the cementation reaction, modifying the electrochemical potential difference between the cupric ion and the reducing metal. This effect can translate into a greater efficiency in the deposition of metallic copper on the surface of the cementing agent, favoring the formation of a more homogeneous and adherent film. Furthermore, the increase in temperature could improve the purity of the precipitate by minimizing the co-deposition of impurities, thus optimizing the morphology, density, and compactness of the recovered copper, which is crucial for its subsequent processing and industrial application.

The precipitated copper’s phase crystallinity and structural analysis were determined using X-Ray Diffraction (XRD) via the powder diffraction technique. Figure 10 presents the diffraction pattern of Cu, in which the main peaks found at 2θ values of 43.28°, 50.40°, and 74.81° correspond to (111), (200), and (220) planes, respectively, confirming the cubic lattice of copper, without other significant peaks. All the diffraction peaks agree with the standard pattern for the pure face-centered cubic phase of copper nanoparticles (JCPDS No. 040836) [10,11].

The solid residues were characterized using Scanning Electron Microscopy coupled with energy-dispersive X-ray spectroscopy (SEM-EDS), shown in Figure 11. The analysis confirms that copper precipitates (1), elemental sulfur deposits (2), and aluminum hydroxides (3) were observed on the surface of the Cu particle. Metallic aluminum (4) was detected adjacent to the main particle. Those phases are determined based on the percentages in the EDS microanalysis spectrum obtained, where peaks corresponding to oxygen (O), sulfur (S), copper (Cu), and aluminum (Al) are identified, indicating the coexistence of these phases within the sample.

Figure 12 presents the characterization of solid residues of the cementation of the pregnant solution with a 5 g/L aluminum concentration at 40° C, using SEM-EDS at 300× magnification. Various mineral phases were identified based on morphology and elemental composition. The particles labeled as (1) correspond to copper precipitates, while region (2) is associated with elemental sulfur deposits. Region (3) shows structures consistent with aluminum hydroxides. The bright particles marked as (4) exhibit a strong secondary electron reflectance, typical of metallic aluminum, which is confirmed by the EDS spectrum.

The EDS microanalysis reveals a predominant presence of oxygen (61.26%), followed by aluminum (33.49%), with lower concentrations of sulfur (1.59%) and copper (3.66%). This suggests the coexistence of aluminum oxides and hydroxides with traces of elemental copper and sulfur compounds. The strong oxygen and aluminum signals support the formation of oxidized surface products on the metallic particles, likely a result of the leaching cementing process (Equation (20)). In this case, the elemental sulfur comes from the sulfate ion present in the leaching pregnant solution being reduced by the aluminum shavings.

$$S{O}_4^{-2}+2A{l}^0+2H_{2}O+2H^{+}\to S^0+2Al(OH{)}_3$$

(20)

The elemental sulfur detected may originate from the reduction of the sulfate ion present in the cementing solution. Thermodynamic analysis, based on the Gibbs energy of formation (

Table 2. Gibbs energy of formation of the compounds presented in Equation (20).

Compound	$∆\mathit{G}(\mathbf{k}\mathbf{J}/\mathbf{m}\mathbf{o}\mathbf{l})$
${Al}^0$	0.0
${SO}_4^{-2}$	−744.5
$H^{+}$	0.0
$H_{2}O$	−292.72
$Al(OH{)}_3$	−1306.0
$S^0$	0.0

), and calculating the Gibbs energy of reaction by summing the Gibbs energy of the products minus the Gibbs energy of the reactants (Equation (21)) yield a value of −1282.06 kJ/mol, indicating that the reaction is thermodynamically spontaneous.

$${\mathsf{\Delta }G}_{reaction}^0=\sum {\upsilon }_i{\mathsf{\Delta }G}_{iProducts}^0-\sum {\upsilon }_j{\mathsf{\Delta }G}_{jreacts}^0$$

(21)

where

$${\mathsf{\Delta }G}_{reaction}^0=Gibbsfreeenergyofreaction.$$

$${\mathsf{\Delta }G}_{iProduts}^0=Gibbsfreeenergyofproductsforcompundi.$$

$${\mathsf{\Delta }G}_{jReactants}^0=Gibbsfreeenergyofreactantsforcompundj.$$

$${\upsilon }_{i}and{\upsilon }_j=stoichiometriccoefficientforiandjcompounds,respectively.$$

Replacing

$${\mathsf{\Delta }G}_{reaction}^0=-2612.0{\displaystyle \raisebox{1ex}{$\mathrm{k}\mathrm{J}$}\!\left/ \!\raisebox{-1ex}{$\mathrm{m}\mathrm{o}\mathrm{l}$}\right.}+1059.06{\displaystyle \raisebox{1ex}{$\mathrm{k}\mathrm{J}$}\!\left/ \!\raisebox{-1ex}{$\mathrm{m}\mathrm{o}\mathrm{l}$}\right.}=-1282.06{\displaystyle \raisebox{1ex}{$\mathrm{k}\mathrm{J}$}\!\left/ \!\raisebox{-1ex}{$\mathrm{m}\mathrm{o}\mathrm{l}$}\right.}$$

Additionally, under the final conditions of the cementation experiments (pH 1.7 and –337 mV vs. SHE), the formation of H₂S is thermodynamically favored (Figure 2), although the proximity to the stability field of elemental sulfur may allow for the partial precipitation of solid sulfur, as confirmed by the XRD and SEM analysis.

4. Conclusions

The study of copper cementation from Pregnant Leach Solutions of a chalcopyrite concentrate using aluminum shavings allowed us to identify the key factors influencing process efficiency. It was demonstrated that aluminum concentration and stirring speed significantly impact cementation kinetics, with higher aluminum amounts and increased stirring rates enhancing copper recovery by increasing the reactive surface area and improving mass transfer.

• The shrinking core model is adequately fitted to the experimental results. The lower stage is dependent on the operation conditions. In the absence of stirring, the limit step is the diffusion of the reactants to particles and the observed kinetics follow a zero-order model, indicating that the process control is not dominated by reactant diffusion but rather by the availability of the cementing agent.
• Additionally, an activation energy of 62.6 kJ/mol was determined within the temperature range of 20 to 40 °C, suggesting that the chemical stage of the process predominantly controls the reaction rate. Increasing the temperature accelerated the cementation kinetics and enhanced the process efficiency by improving copper deposition on the aluminum surface, reducing the impurity content in the final precipitate.
• X-Ray Diffraction analysis confirmed the formation of metallic copper with a face-centered cubic crystalline structure, verifying the purity of the precipitated copper. Overall, the results highlight the feasibility of using aluminum shavings as a cementing agent for copper recovery from Pregnant Leach Solutions, achieving a high efficiency and selectivity. However, further studies are recommended to assess process stability on a larger scale and under variable operating conditions to optimize its industrial application.
• Through morphological and compositional analysis using SEM-EDX, the presence of copper precipitates with varying degrees of surface oxidation was confirmed. Particles with a high oxygen and aluminum content were identified, suggesting the formation of aluminum hydroxides or oxides, likely as secondary products from the leaching reactions.
• The detection of elemental sulfur on the surface of copper particles indicates the possible partial oxidation of sulfur species or residues from the sulfidation process. The brightest particles in the micrographs exhibit a spectral profile characteristic of metallic aluminum, indicating that some aluminum remained unreacted during the experimental process. From a thermodynamic standpoint, this observation aligns with the stability domains predicted in the Pourbaix diagram for sulfur, where elemental sulfur is favored under the experimental conditions (pH 1.7 and –337 mV vs. SHE), being close to the stability window of S⁰. This proximity may have allowed the partial precipitation of sulfur, rather than its full conversion into soluble sulfide species such as H₂S or HS⁻. The information obtained through SEM-EDX complements the XRD results, as it allows for a spatial visualization of elemental distribution and confirms the coexistence of multiple solid phases at the microscale.