Chalcopyrite Leaching in Ferric Sulphate: The Effect of Fe₃O₄-CuFeS₂ Galvanic Couple on the Cu Dissolution

Abstract

Galvanic interactions present alternative strategies to achieve a more efficient Cu dissolution from CuFeS₂. The present work studied the interaction between chalcopyrite–magnetite (CuFeS₂-Fe₃O₄) in acidified ferric sulphate Fe₂(SO₄)₃-H₂SO₄ at a solution pH of 1.8 and a temperature of 25 or 50 °C. The addition of Fe₃O₄ to CuFeS₂ forms a galvanic couple, which positively impacts the dissolution of Cu. The results showed that the presence of Fe₃O₄ led to high and fast Cu dissolution rates and decreased significantly the activation energy, from 83 to 57 kJ/mol. In addition to that, the solid residues revealed that CuFeS₂ dissolution produced intermediate Cu-S-rich phases: CuS, Cu₂S and Cu₅FeS₄, which appeared to envelop CuFeS₂, had no observable intermediate phase while in the presence of Fe₃O₄. The results showed that 94% of Cu could be recovered after 5 h of leaching at 50 °C at a Fe₃O₄/CuFeS₂ ratio of 4:1 and a 460 mV Ag/AgCl solution potential. The findings of this study present an option for efficient Cu dissolution from CuFeS₂ in ferric sulphate at atmospheric pressure.

1. Introduction

Copper (Cu) has been used by humans for thousands of years, and its discovery marked a significant milestone in developing the world’s civilization. Recently, it has been listed among the critical metals that provide significant contributions towards a more sustainable and low-carbon economy [1]. A major source of Cu is chalcopyrite (CuFeS₂), a copper sulphide mineral, which accounts for around 70%–80% of the total global production of Cu. The pyrometallurgy route produces about 85% of the worldwide Cu through a reverberatory furnace or flash smelting. However, the pyrometallurgical treatment appears unattractive for metal extraction due to high sulphur dioxide (SO₂) emission and decreased high-grade mineral ore bodies [2,3]. The hydrometallurgical operation, therefore, seems to promote adequate process economics based on its high metal selectivity and environmental considerations.

The leaching or dissolution process is a vital operation in the hydrometallurgy process. As a result, its efficiency has a major impact on the various downstream processes and, mainly, on the economic and technical aspects of mineral hydrometallurgy. Efficient leaching of chalcopyrite concentrates at ambient pressure still remains a challenge because of the slow dissolution kinetics of this mineral in most leaching media [4,5]. The slow and incomplete dissolution is attributed to forming a passive layer on the mineral surface, regardless of the method used, whether it is performed via a chemical or bioleaching processes [6,7]. The nature and composition of this barrier remain subject to controversy [8,9,10]. The most often suggested candidates to form the passivating layers are metal-deficient phases, elemental sulphur and jarosite [11]. In order to enhance the Cu dissolution from the CuFeS₂ mineral, different methods and techniques have been proposed from the various laboratory studies. These methods include using a sulphate–chloride solution, microorganisms, fine/ ultra-grinding and the addition of silver ions [5,12].

CuFeS₂ dissolution could be discussed based on the Ellingham or Eh–pH diagram, which predicts the predominance area of susceptible pieces or phases based on their equilibrium condition for all possible redox reactions (Equations (1)–(4)). The increase of potential on the mineral surface leads to different oxidation reactions, which promote the formation of intermediate phases: bornite (Cu₅FeS₄), covellite (CuS) and chalcocite (CuS₂). According to the Eh–pH diagram, successful copper dissolution from the mineral could be achieved after raising the redox potential above 0.44 V. In addition to that, all intermediate-formed phases are susceptible to dissolving in acidic media at a potential above 0.5 V.

CuFeS₂ + 2Fe₂(SO₄)₃ = CuSO₄ + 5FeSO₄ + 2S          ∆G = −16.3

(1)

5CuFeS₂ + 4Fe₂(SO₄)₃ ⇌ Cu₅FeS₄ + 12FeSO₄ + 6S        ∆G = −68.7

(2)

2CuFeS₂ + 2Fe₂(SO₄)₃ ⇌ Cu₂S + 6FeSO₄ + 3S          ∆G = −30.8

(3)

CuFeS₂ + Fe₂(SO₄)₃ ⇌ CuS + 3FeSO₄ + S            ∆G = −18.6

(4)

The strategy to overcome passivation and improve metal recovery during the hydrometallurgical treatment of the CuFeS₂ mineral is applying and using sulphur-oxidizing microorganisms, which promotes the oxidation of elemental sulphur. The second method, on the other hand, involves galvanic interactions between CuFeS₂ and other minerals. These minerals usually have an associated/symbiotic relationship with chalcopyrite and include phases such as pyrite, bornite, magnetite and Arsenopyrite [3]. Due to the difference in the potential between the two minerals, a galvanic effect may occur, making the lower potential phase act as an anode, leading to enhanced oxidation and improved kinetics. The covalent character of most sulphide minerals provides non-localization of charge, resulting in appreciable intrinsic electronic conductivity. FeS₂ is very common in the CuFeS₂ mineral ore matrix. Hence, these two phases are in galvanic contact, which leads to the rapid dissolution of Cu from FeCuS₂, while FeS₂ is galvanically protected. Particular interest has been given to the galvanic couple CuFeS₂-FeS₂, and numerous experimental works [13,14,15,16] have been conducted to understand and provide more insight into this alternative route of CuFeS₂ dissolution. In the galvanic setup of CuFeS₂,-FeS₂ pyrite represents the cathode on which the reduction of the active oxidant in the system takes place while the anode’s (CuFeS₂) dissolution is significantly enhanced.

The Galvanox^TM process makes use of the FeS₂-CuFeS₂ galvanic couple [13,14,15] and reports appreciable results at a Cu recovery of 98% with a FeS₂-to-CuFeS₂ (Py:Cp) ratio: roughly 2:1 to 4:1, with the solution potential varying from 440 mV (Ag/AgCl) to 470 mV or higher and the temperature in the range of 70 to 80 °C [17]. The presence of a galvanic couple led to a higher degree of Cu dissolution from CuFeS₂ in the presence of FeAsS in an acidic medium culture at a pH of 2 and a temperature of 30 °C [3].

Magnetite (Fe₃O₄) could also be found as an existing impurity in the CuFeS₂ ore matrix, especially in the case of carbonatite [18,19]. Being a spinel, the magnetite (Fe²⁺Fe³⁺₂O₄) in an aqueous phase could represent a substantial oxidant source. To the author’s knowledge, one recent study reported the existence of galvanic interactions between CuFeS₂ and Fe₃O₄ [20]. These interactions might also serve as another alternative for Cu dissolution from the CuFeS₂ since oxide minerals are generally more ionic in nature and usually have higher resistivities than sulphide minerals. In addition, the problem with the direct application of the Eh–pH diagram is the possible formation of metastable intermediates, which may be attributed to slow chemical kinetics. However, there is no way to predict the presence of intermediates from thermodynamics alone. It is, therefore, appropriate to review the kinetic of CuFeS₂ in Fe₂(SO₄)₃ to better understand the dissolution and phase transition.

Although there have been many studies on mineral additives that can promote the dissolution of chalcopyrite, many types of minerals related to chalcopyrite have not been studied. This work compares the dissolution of Cu from CuFeS₂ with and without the addition of Fe₃O₄ in (H₂SO₄-Fe₂(SO₄)₃) at a pH value of 1.8, at temperatures of 25 and 50 °C. The study highlights the effect of Fe₃O₄ and evaluates optimum conditions (temperature, potential and Fe₃O₄/CuFeS₂ ratio) for maximum Cu dissolution. Lastly, the study reviews intermediate phase conversion and elaborates on the dissolution based on the Cu-S phases.

2. Materials and Methods

2.1. Material

Solutions of the desired pH were prepared using analytical reagent-grade sulphuric acid (H₂SO₄ 98% A.C.E.), ferric sulphate (Fe₃(SO₄) • H₂O ACE) and deionized water (<5.0 µs/cm). The redox potential (Ag/AgCl) measurements were determined in reference to the Standard Hydrogen Electrode (SHE).

2.2. Chalcopyrite and Magnetite Sample Characterization

The CuFeS₂ sample was obtained as a concentrate from the Phalaborwa Copper Mining Company (Phalaborwa, South Africa). The chalcopyrite sample was dried in an oven at 50 °C for seven days before sub-sampling. Homogenization, in accordance with the soil sampling protocol by the U.S. Environmental Protection Agency [21], was conducted. Approximately 0.5 kg of concentrate was sub-sampled and dried for approximately 20 min at 105 °C. A grain size of less than n 200 microns (<200 µm) was used as a dissolution feed.

The samples were analysed for mineral composition and bulk chemistry using X-ray analysis—X-ray diffraction (XRD) (RigakuUltimaIV with PDXL analysis software (Japan)) and XRF (Rigaku-ZSX Primus II with SQX analysis software (Japan)), respectively. An XRD Rigaku Ultima IV was operated at 40 kV and 30 mA. The Integrated X-ray Powder Diffraction Software (PDXL) (Ver.2.0–2.8.X, Rigaku, Japan) was used for mineral phases, and the instrument’s detection limit was 2%. Data were recorded over the range 5° ≤ 2θ ≤ 95° at a scan rate of 0.5 degrees/min and a step width of 0.01°. The elemental composition (XRF) powder method was carried out on a Rigaku-ZSX Primus II in conjunction with the S.Q.X. analysis software (Rigaku, Japan), operating at 4 kW, 60 kV and 150 mA. The results are illustrated in Figure 1.

2.3. Leaching Media and Tests

CuFeS₂ was dissolved in an acidified ferric sulphate solution (H₂SO₄-Fe₂(SO₄)₃), with and without Fe₃O₄ addition. A ratio of Fe₃O₄/CuFeS₂ varying from 1:1 to 4:1 was used to evaluate the effect of Fe₃O₄ addition on Cu dissolution. The solution was obtained by mixing Fe₂(SO₄)₃ with H₂O (water) and H₂SO₄. An initial Fe content (Fe³⁺ = 0.05 MFe) was used for all tests. The media was agitated for 12 h before using it. Dissolution tests were performed at atmospheric conditions at 25 °C or 50 °C. The media pH was measured and maintained at 1.8 with periodic addition of H₂SO₄ (98%). At the same time, the solution oxidation–reduction potential (O.R.P.) was allowed to evolve throughout the dissolution test and was recorded periodically (every 10 min).

Additional dissolution tests were conducted at 50 °C at 400, 430 and 460 mV, and these solution potentials were maintained throughout the dissolution test by periodic addition of hydrogen peroxide in the slurry. This was intended to evaluate the potential dependency of the varying Fe₃O₄-CuFeS₂ ratio.

A pulp density of 10% solid was used in all tests, and kinetic information was obtained by periodic withdrawal (every 20 min) of a 5 mL sample for chemical analysis of total Cu using atomic absorption flame spectrometry (AAFS, Thermo Scientific ICE 3000 series, Cambridge, UK). The redox potential of the leaching solution was measured with a platinum electrode using a saturated Ag/AgCl (3 M KCl) electrode as the reference.

Solid leached residues were analysed for mineral composition using XRD analysis (Rigaku Ultima IV) operating at 40 kV and 30 mA, and PDXL analysis software was used; the instrument’s detection limit was 2%. Data were recorded over the range 5° ≤ 2θ ≤ 95° at a scan rate of 0.5°/min and a step width of 0.01°.

3. Results

3.1. Cu Dissolution Curve and Recovery

All the dissolution curves were asymptotic and characterized by different stages (Figure 2a–c and Figure 3a–c): the first stage (0–60 min) was more rapid than the second stage (60–360 min), and the third stage was the plateau with no Cu dissolution (360–720 min). The rapid stage could be attributed to the fast Cu withdrawal at the mineral surface. At the same time, the second one could correspond to the dissolution of the transient Cu phases formed, which would lead to the third stage with no further dissolution (passivation) [2]. It is worth highlighting that the dissolution conducted without Fe₃O₄ displayed the lowest Cu recovery, as opposed to the mixed CuFeS₂-Fe₃O₄ sample, which exhibited high Cu recovery at a much-prolonged first stage.

3.2. Cu Dissolution Recovery at Free O.R.P.

Figure 2 shows the Cu dissolution at free O.R.P. at 25 (Figure 2a) or 50 °C (Figure 2b) for the different samples (CuFeS₂ and Fe₃O₄-CuFeS₂). Figure 2c,d also show the O.R.P. evolution during the dissolution. At both experimental temperatures (25 and 50 °C), it was observed that the Cu recovery of the mixed sample (Fe₃O₄-CuFeS₂) was higher than that of the pure CuFeS₂ concentrate sample without added Fe₃O₄. At 25 °C (Figure 2a), only 15% Cu was dissolved from the pure concentrate after 300 min, while 20, 24, 27 and 32% Cu were observed for the mixed sample (Fe₃O₄-CuFeS₂), at a Fe₃O₄/CuFeS₂ ratio of 1, 2, 3 and 4, respectively. At 50 °C (Figure 2b), much more Cu was dissolved—47, 51, 57 and 60% Cu were dissolved from the mixed sample, at Fe₃O₄/CuFeS₂ ratios of 1, 2, 3 and 4, respectively, while only 21% Cu was dissolved for the pure concentrate (CuFeS₂).

A sharp drop (Figure 2c,d) in the solution potential (from 348–336.7 mV) was observed for the CuFeS₂ concentrate sample, while all mixed (Fe₃O₄-CuFeS₂) sample O.R.P. curves displayed a stagnant/static region at the early stage of the dissolution and dropped thereafter. The rapid drop in solution potential observed for the pure concentrate (without Fe₃O₄ addition) could be due to the rapid consumption of ferric (Fe³⁺), leading to Cu withdrawal and favouring the formation and increase in Fe²⁺ (Equation (1)). The static region observed for the Fe₃O₄-CuFeS₂ sample could be related to the galvanic interaction between CuFeS₂ and Fe₃O₄. This stationary state corresponded to the first linear Cu dissolution stage observed on the Cu recovery curve. It could, therefore, be responsible for the high rate and Cu recovery observed for the mixed Fe₃O₄-CuFeS₂ sample.

The results from the present investigation show that the Fe₃O₄ addition positively affects the dissolution of CuFeS₂ rate and recovery. In addition, the presence of Fe₃O₄ appears to some extent to stabilize the solution O.R.P., since a margin drop-in solution potential is observed in the sample where Fe₃O₄ was absent.

Figure 2. Cu dissolution at uncontrolled O.R.P. and 25 or 50 °C. (a) Dissolution at 25 °C at various ratios of magnetite; (b) Dissolution 50 °C at various ratios of magnetite; (c) Dissolution at 25 °C at various ratios of magnetite; (d) Dissolution at 50 °C at various ratios of magnetite.

Figure 2. Cu dissolution at uncontrolled O.R.P. and 25 or 50 °C. (a) Dissolution at 25 °C at various ratios of magnetite; (b) Dissolution 50 °C at various ratios of magnetite; (c) Dissolution at 25 °C at various ratios of magnetite; (d) Dissolution at 50 °C at various ratios of magnetite.

3.3. Effect of Solution Potential

The effect of potential was evaluated at 50 °C. Figure 3 shows the Cu dissolution at 400, 430 and 460 mV solution potentials. At all solution O.R.P.s (Figure 3a–c), it was observed that increasing the (Fe₃O₄/CuFeS₂) Mag/Ch ratio (from 1:1to 4:1) led to a significant increase in the copper recovery and extraction rate. However, only a smaller difference, between 3:1 and 4:1 (Mag/Ch), was identified. The solution potential appears to play a vital role during the Cu dissolution from the CuFeS₂. It was observed that an increase in the solution O.P.R. enhanced the dissolution. More Cu dissolved at 400 mV, much more at 430 and close to complete dissolution was obtained at 460 mV. The present results showed an optimum O.R.P. value of 460 mV, at which 96 and 94% Cu could be achieved after 300 min at a Mag/Ch ratio of 4:1 and 3:1, respectively. These results support previous studies where high Cu recoveries were obtained at low redox potentials) [12,22]. Monitoring and controlling the solution potential appear to be key factors for increasing the dissolution rate in CuFeS₂ leaching systems. Most previous work agreed on a narrow potential range of 400–450 mV (Ag|AgCl) where the Cu dissolution yield was at the maximum.

Figure 3. Cu dissolution at controlled O.R.P. (400, 430 or 460 mV Ag/AgCl). (a) Dissolution at 400 mV and 50 °C; (b) Dissolution at 430 mV and 50 °C; (c) Dissolution at 460 mV and 50 °C.

Figure 3. Cu dissolution at controlled O.R.P. (400, 430 or 460 mV Ag/AgCl). (a) Dissolution at 400 mV and 50 °C; (b) Dissolution at 430 mV and 50 °C; (c) Dissolution at 460 mV and 50 °C.

3.4. Effect of Temperature

The temperature was found to affect both the rate and recovery positively. Consequently, its effect was evaluated using the shrinking core model, as shown in Equation (5) when surface chemical reaction controls.

1 − (1 − a)^1/3 = k′·t

(5)

where a is the copper ion concentration, k′ corresponds to the reaction rate constant and t means the leaching time. The Arrhenius equation (i.e., Equation (6)) was used to calculate the activation energy of chalcopyrite leaching:

k′= Ae ^−(Ea/RT)

(6)

where A is the pre-exponential factor, Ea is the activation energy (J·mol⁻¹), R is the universal gas constant (J·mol⁻¹·K⁻¹) and T is the absolute temperature in Kelvin (K). Figure 4a shows the fitting of the lnk′ against 1/T (Arenhius plot) according to the Cu extraction data collected, and Figure 4b shows the effect of Fe₃O₄ addition on the CuFeS₂ Ea values.

The CuFeS₂ concentrate sample dissolution had the highest Ea (around 80 kJ/mol). Its high value was consistent with the literature values [23]. The activation energy decreased with the addition of Fe₃O₄ (Figure 4b). At a Fe₃O₄/CuFeS₂ ratio of 1:1, the Ea value was 70 kJ·mol⁻¹. This value went down to 65 when the Fe₃O₄ content was double (Fe₃O₄/CuFeS₂ = 2) and decreased further to 61 kJ·mol⁻¹ and 56 kJ·mol⁻¹ when the Fe₃O₄/CuFeS₂ ratio increased respectively to 3 and 4. The decrease in the Ea value with the addition of Fe₃O₄ could suggest that the dissolution process of CuFeS₂ becomes much easier in the presence of Fe₃O₄. Although the Ea value substantially decreased from 84 to 56 kJ·mol⁻¹, all the Ea values were above 40 kJ·mol⁻¹, implying that all dissolution was controlled by the same process. The dissolution process was chemically controlled at the mineral surface [24].

3.5. Solid Residue Analysis

Figure 5 summarises the mineralogical composition of the solid residues at 50 °C and 460 mV. Figure 5a displays the residue mineral content of the CuFeS₂ without adding Fe₃O₄, while Figure 5b presents the mineral phase of the mixed Fe₃O₄-CuFeS₂ at a Mag/Ch ratio of 4:1.

In both cases, the CuFeS₂ main peaks (2 theta = 29.46, 48.93 and 57.95) appeared to decrease in intensity compared to the feed sample. The dissolution of CuFeS₂ without Fe₃O₄ (Figure 5a) displayed the presence of a Cu-rich phase (bornite (Bo), chalcocite (Cx) and covellite (Co)). The presence of these Cu-rich intermediates appeared to support earlier investigation [25], which underlined the preferential dissolution of Fe over Cu, leading to the formation of a Fe-deficient chalcopyrite mineral (defect chalcopyrite structure Cu1-xFe1-yS2-z) to which the present identified species could be a part.

On the other hand, the mixed Fe₃O₄-CuFeS₂ sample residues remained dominated by the Fe₃O₄ (Figure 5b), and the Fe₃O₄ content increased due to the CuFeS₂ dissolution. These results showed that the Fe₃O₄ content increased from 66.4% in the feed sample (Mag/Ch = 4:1) to 86.2% after 5 h of dissolution (Figure 5d). This supports the presence of the galvanic couple between CuFeS₂ and Fe₃O₄, as CuFeS₂ dissolves more rapidly, while Fe₃O₄ becomes galvanically protected. This finding also confirms that most of the Fe₃O₄ could be recycled if the process is used. In addition to that, the Fe₃O₄-CuFeS₂ mixed residue sample did not reveal the presence of any intermediate phase. This could suggest that the dissolution of Cu from the CuFeS₂ sample takes place by mineral phase intermediates. In contrast, no intermediate phase is formed in galvanic interaction, and the Cu dissolves directly from CuFeS₂, according to Equation (1).

Thermodynamically, bornite is the most favourable phase to be formed (Equation (2)), followed by chalcocite (Equation (3)) and, lastly, covellite (Equation (4)). These Cu-S phases could react further into new Cu-S-richer phases or leach to promote Cu recovery. CuS is more likely to form from Cu₅FeS₄ (Equation (5)) than Cu₂S (Equation (6)). Similarly, [26,27] also reported the presence of CuS due to a transient species during the ferric leaching of Cu₅FeS₄. The intersection points in Figure 5c between the various intermediate phases could suggest, on the one hand, the dissolution competition between phases. On the other hand, it could suggest the transformation from one phase into the other. In that sense, I1, I2 and I3 observed between CuS/Cu₅FeS₄, CuS/Cu₂S and CuS/CuFeS₂ suggest the competition between both dissolution reactions and, since CuS is refractory, Cu₅FeS₄, Cu₂S and CuFeS₂ dissolve by transformation into CuS according to Equations (7)–(9). This explains the increase in the CuS content (Figure 5a) and its cumulative behaviour (Equation (11)), as observed in the XRD analysis. Meanwhile, I4 strictly represents the direct conversion of CuFeS₂ to Cu₂S without any Cu²⁺ withdrawal, according to Equation (11).

Cu₅FeS₄ + Fe₂(SO₄)₃ ⇌ 4CuS + 3FeSO₄ + Cu²⁺          ∆G = −22.4

(7)

Cu₂S + Fe₂(SO₄)₃ ⇌ CuS + 2FeSO₄ + Cu²⁺            ∆G = −4.35

(8)

CuFeS₂ + Fe₂(SO₄)₃ ⇌ CuS + 3FeSO₄ + S            ∆G = −18.6

(9)

CuS + Fe₂(SO₄)₃ ⇌ 2FeSO₄ + S₀ + Cu²⁺            ∆G = 2.12

(10)

2CuFeS₂ + 2Fe₂(SO₄)₃ ⇌ Cu₂S + 6FeSO₄ + 3S          ∆G = −30.8

(11)

4. Conclusions

This work investigated the dissolution of CuFeS₂ in acidic Fe₂(SO₄)₃ in the presence of Fe₃O₄. The results show the positive impact of Fe₃O₄ on Cu dissolution from CuFeS₂. It was observed that Fe₃O₄ enhanced the Cu dissolution, which was highly correlated to temperature. The results showed that 94% of Cu could be recovered after 5 h at 50 °C at a Fe₃O₄/CuFeS₂ ratio of 4:1. The activation energy considerably decreased in the presence of Fe₃O₄. It was found to be 57 kJ/mol, suggesting a chemically controlled process through surface reaction. The solid residues revealed that CuFeS₂ dissolution produced intermediate Cu-rich phases: CuS, Cu₂S and Cu₅FeS₄, which appeared to envelop CuFeS₂, while in the presence of Fe₃O₄, no intermediate phase was observed. The findings of this study present an option for efficient Cu dissolution from CuFeS₂ in ferric sulphate.

5. Patents

Data from this work have been used to claim the novelty of the work, and a patent was granted in 2022.