Metallic Copper (Cu[0]) Obtained from Cu2+-Rich Acidic Mine Waters by Two Different Reduction Methods: Crystallographic and Geochemical Aspects

  Citation: Sánchez-España, J.; Ilin, A.; Yusta, I. Metallic Copper (Cu[0]) Obtained from Cu2+-Rich Acidic Mine Waters by Two Different Reduction Methods: Crystallographic and Geochemical Aspects. Minerals 2022,12, 322. https://doi.org/ 10.3390/min12030322 Academic Editors: Radoslaw Pomykala, Barbara Tora and Katerina Adam Received: 31 January 2022 Accepted: 1 March 2022 Published: 4 March 2022 Publisher’s Note: MDPI stays neutral with regard to jurisdictional claims in published maps and institutional affiliations. Copyright: © 2022 by the authors. Licensee MDPI, Basel, Switzerland. This article is an open access article distributed under the terms and conditions of the Creative Commons Attribution (CC BY) license (https:// creativecommons.org/licenses/by/ 4.0/). minerals Article Metallic Copper (Cu[0]) Obtained from Cu2+-Rich Acidic Mine Waters by Two Different Reduction Methods: Crystallographic and Geochemical Aspects Javier Sánchez-España 1,* , Andrey Ilin 2and Iñaki Yusta 2 1Mine Waste and Environmental Geochemistry Research Group, Department of Geological Resources (IGME-CSIC), Calera 1, Tres Cantos, 28760 Madrid, Spain 2Department of Geology, University of the Basque Country (UPV/EHU), Faculty of Science and Technology, Apdo. 644, 48080 Bilbao, Spain; andrey[email protected] (A.I.); [email protected] (I.Y.) *Correspondence: [email protected] Abstract: The recovery of valuable metals from different types of wastes has become of prime strategic interest given the scarcity of primary critical raw materials at international scale. Implementation of new methods or refinement of classical techniques with modern technological advances is, therefore, an active research field. Mine wastes are of special interest because their high metal concentrations make them environmentally harmful and economically profitable at the same time. In this study, we evaluated two different methods of Cu recovery from extremely acidic mine waters seeping from wastes and abandoned mines in SW Spain. Through a series of different batch experiments, we compared the method efficiency and crystallographic properties of elemental copper (Cu[0]) obtained by reduction of Cu 2+ ions by (1) chemical reduction using ascorbic acid at different environmental conditions of pH (1.50–3.95), temperature (25–80 ◦ C) and ascorbic acid concentration (10 mM to 0.1 M), and (2) classical cementation method with scrap iron at pH 1.50 and 25 ◦ C. Our study demonstrates that the precipitation of Cu[0] can take place at pH 3.95 and low AA concentrations (0.1 M), resulting in large ( µ m-scale), perfectly developed crystals of copper with pseudoprismatic to acicular habit after 24 h of aging, likely through formation of a transient compound consisting in Cu 2+ -ascorbate and/or cuprite (Cu 2 O) nanocolloids. Reduction experiments at higher AA concentrations (0.1 M) showed faster precipitation kinetics and resulted in high-purity (>98%) copper suspensions formed by subrounded nanoparticles. The AA method, however, yielded very low recovery rates (15–25%) because of the low pH values considered. The cementation method, which produced tree-like aggregates formed by sub-micron crystals arranged in different directions, proved to be much more efficient (>98% recovery) and cost-effective. Keywords: copper; chemical reduction; acidic mine waters; secondary raw materials; metal recovery 1. Introduction Copper is of paramount importance in modern technology, being widely used in micro/nanoelectronics [ 1 ]. Copper nanoparticles and nanowires are also of major interest for their use in diverse fields, such as catalysis, optoelectronics, and liquid phase or gas sensors [ 2 – 4 ]. Copper nanoparticles are also used to produce conductive inks for electrodes of silicon crystal solar cells and in the printed electronic industry [ 5 ]. Other potential applications of copper nanoparticles aim to take advantage of their known antibacterial properties in the food or medical industry, or to produce lubricating substances for the fabrication of diverse composite materials [ 2 ]. Copper nanoparticles have been traditionally obtained by reduction of Cu 2+ ions in solution, which does not require expensive equipment and allows the control of particle size and morphology through variation of certain parameters [ 2 ]. This method usually requires an alkaline pH to be efficient, since the electrode potential of AA is much lower at alkaline medium (+80 mV compared to Minerals 2022,12, 322. https://doi.org/10.3390/min12030322 https://www.mdpi.com/journal/minerals Minerals 2022,12, 322 2 of 18 +166 mV at pH 4 [ 2 ]). At present, copper nanoparticles may be synthesized by different reduction methods using microwaves, as well as sonochemical, electrochemical, thermal, and chemical techniques, whereas copper nanowires are currently produced by different methods such as electrochemical deposition/electroplating (commonly the easiest and cheapest technique, and highly efficient at very acidic pH), pulsed laser deposition, or chemical vapor deposition [ 3 , 6 – 8 ]. In chemical reduction, the use of several reductants in early times, such as hydrazine or sodium hydrophosphate, which were seen to be either highly toxic (e.g., hydrazine) or technically inefficient (e.g., sodium hydrophosphate), has been progressively replaced by other reducing agents, such as ascorbic acid and other organic reductants such as glucose [ 2 , 3 , 5 ]. The antioxidant properties of AA (which prevent the resulting copper from undergoing oxidation in oxygenated environments) make this compound advantageous with respect to other reductants [ 3 ]. The control of size and morphology of Cu[0] and Cu 2 O nanocrystals and nanostructures through the use of different organic and inorganic additives, with or without stabilizing agents, has also become a common practice in materials chemistry to produce nanomaterials with very specific conductive properties (e.g., superconductors) [9]. In the context of a high demand of raw materials for the modern industry and the implementation of circular economy concepts at global scale, the recovery of valuable and critical metals from mine wastes has become of growing interest [ 10 ]. Acidic mine waters seeping from mine wastes (e.g., tailings, waste piles) or abandoned mine portals or open pits, usually referred to as acid mine drainage (AMD), is among the major environmental problems of the metal and coal mining industry, due to its severe acidity and high concentration of sulfate and many toxic metals [ 11 , 12 ]. At the same time, however, these acidic liquors may contain valuable metals, such as Cu, Zn, Co, or Ni, at such high concentrations that can make them a resource of secondary raw materials more than a waste [ 13 – 15 ]. The acidic mine waters of the Iberian Pyrite Belt, in SW Spain, with AMD leachates exceeding concentrations of 440 ppm (mg/L) [ 16 ] and with pit lakes showing extreme contents up to 1350 mg/L Cu 2+ [ 17 – 19 ], probably represent a paradigmatic case at international level. Although its benefit at industrial scale by modern techniques is still utopic, the recovery of copper by the classical cementation method using scrap iron was widely used in the region during the 19th and 20th centuries [ 20 ], being also applied at present in some AMD sites and pit lakes in other mining districts of the world [21]. Despite the plethora of studies on different methods of copper production from solutions of variable chemistry and origin, there is an apparent lack of studies comparing the efficiency of different methods potentially applicable to the recovery of copper from AMD, and both the structural/crystallographic properties of the obtained copper by reduction of Cu 2+ ions, and the geochemical controls on Cu[0] crystallization, are still not sufficiently well known. Any Cu 2+ -reduction method needs to be cost-effective and must provide a high-quality product in order to be profitable and implementable at industrial scale. In the present study, we aimed to evaluate the efficiency of Cu 2+ reduction in Cu-rich, acidic mine waters by (1) chemical reduction using ascorbic acid (a widely known reducing agent in Cu synthesis studies), and (2) classical cementation method using scrap iron as reducing agent. Both methods were tested at low pH conditions (<4.0) under oxygenic atmosphere and without any stabilizer, since the main goal of our investigation was to evaluate possible on-site treatments in abandoned mines without the need of complex or expensive infrastructure. We were especially interested in contrasting the geometry and size of the copper particles and/or crystals formed by these two different methods, for which we used a combination of X-ray diffraction and scanning electron microscopy. Geochemical modeling based on solution chemistry was also used to evaluate the distribution and molar abundance of Cu aqueous species in solution and the saturation state of the acidic solutions with respect to elemental copper and other copper compounds. The practical implications are discussed. Minerals 2022,12, 322 3 of 18 2. Materials and Methods 2.1. Field Work and Sampling Water samples (1 L each) were taken from three different locations in the province of Huelva (SW Spain), a region severely affected by acid mine drainage (AMD) pollution [ 16 ]. These acidic mine waters included a leachate from a big waste rock pile in the Riotinto mine (TI), an effluent from an abandoned mine portal in Tharsis mine (CA) and an outflow from an acidic pit lake in a big open pit in San Telmo mine (ST). The pH and Cu concentration of these samples are provided in Table 1. A common feature of these three mine waters was their high acidity (pH 2.5–2.9). They differed, however, as regards to their oxidation state (Eh = 570–750 mV), as well in sulfate and metal concentration, including dissolved copper (23–183 mg/L Cu). Table 1. Selected chemical parameters of the three samples of acidic mine water used in this study. Sample pH Eh SO4Fe Cu mV mg/L mg/L mg/L TI 2.6 570 23,300 1290 183 CA 2.5 615 3700 1250 124 ST 2.9 750 3350 125 23 These waters were all filtered on site through 0.45 µ m pore size, nitrocellulose membrane filters (Millipore) and stored at 4 ◦ C during transport to the laboratory. Once in the IGME Geochemistry lab, the waters were used for different experiments aimed at reducing the dissolved Cu 2+ ions present in the acidic liquors to obtain metallic, zero-valence copper (Cu[0]). We used two different reduction methods, namely: (1) reduction with ascorbic acid (C 6 H 8 O 6 ) as the reducing agent, and (2) classic cementation method using scrap metallic iron (Fe[0]) as the reducing agent. To assess the effect of pH and temperature on crystallization kinetics, as well as on the crystallinity and particle size of metallic copper, the first method (ascorbic acid) was conducted at varying pH (1.50–3.95), temperature (25 ◦ C to 80 ◦ C), and ascorbic acid concentration (10 mM to 0.1 M C 6 H 8 O 6 ). The second method was solely conducted at pH 1.50 and 25 ◦ C. Both reduction methods are briefly described below. 2.2. Batch Experiments of Cu2+ Reduction 2.2.1. Reduction with Ascorbic Acid The selected acidic mine waters (500 mL in volume) were poured in glass flasks and placed in a magnetic stirring stage set at 400 rpm, where they were partially neutralized by dropwise addition of 1 M NaOH to increase the pH of the waters to values around 3.90–3.95 (Figure S1 in Supplemental Material). The goal of this partial neutralization was to remove all dissolved ferric iron in solution (Table 1) so that the experiments could be conducted without the presence of any colloids of Fe(III) minerals, such as schwertmannite, which are often formed in these solutions upon pH increase [ 22 – 25 ]. This pH limit, in addition, ensured the absence of aluminum precipitates that form in AMD systems at pH > 4.0 [26,27] . The titrated solutions were filtered in a vacuum filtration unit (Millipore) using 0.45 µm pore size filters to remove all Fe(III) precipitates (Figure S1). The partly neutralized (pH ~3.95), acidic mine waters were subsequently titrated by dropwise addition of an ascorbic acid solution (1 M, pH 2.5). This ascorbic acid solution had been previously prepared by dilution of pure (reagent-grade), dry crystalline ascorbic acid (C 6 H 8 O 6 ) from Merck in ultrapure distilled water (Millipore). The addition of ascorbic acid (AA) to the MilliQ water resulted in an immediate acidification, and the increase of the AA concentration in the solution tended to attain an apparent equilibrium pH around 2.5 for AA concentrations greater than 1 g/L (Figure S2 in Supplemental Materials). This acidification was provoked by rapid deprotonation of AA and partial conversion to ascorbate through Reaction (1) [28]: Minerals 2022,12, 322 4 of 18 C6H8O6→C6H7O6−+ H+(1) Both AA and ascorbate can effectively reduce Cu(II) and Fe(III) ions, though AA reactions with transition metals such as Fe(III) or Cu(II) have rate constants that differ by up to several orders of magnitude, with the reactions of ascorbate usually being faster [ 28 ]. In any case, in this study we refer to ‘ascorbic acid’ to consider the total amount of the protonated form, AA, and the deprotonated form (ascorbate) present in the reducing solutions. The pH of the solutions was continuously measured during the reduction experiments with previously calibrated CRISON pH meters, and visual changes of color and turbidity were recorded and annotated. When a metallic copper precipitate was apparently formed (based on color change or turbidity formation), the solutions were immediately filtered through 0.45 µ m pore size filters and the obtained solids were washed several times by MilliQ water, and air-dried at ambient temperature for further morphological, mineralogical, and crystallographic analysis (Figure 1). Minerals 2022, 11, x FOR PEER REVIEW 4 of 19 acidification was provoked by rapid deprotonation of AA and partial conversion to ascorbate through reaction (1) [28]: C 6 H 8 O 6 → C 6 H 7 O 6− + H + (1) Both AA and ascorbate can effectively reduce Cu(II) and Fe(III) ions, though AA reactions with transition metals such as Fe(III) or Cu(II) have rate constants that differ by up to several orders of magnitude, with the reactions of ascorbate usually being faster [28]. In any case, in this study we refer to ‘ascorbic acid’ to consider the total amount of the protonated form, AA, and the deprotonated form (ascorbate) present in the reducing solutions. The pH of the solutions was continuously measured during the reduction experiments with previously calibrated CRISON pH meters, and visual changes of color and turbidity were recorded and annotated. When a metallic copper precipitate was apparently formed (based on color change or turbidity formation), the solutions were immediately filtered through 0.45 µm pore size filters and the obtained solids were washed several times by MilliQ water, and air-dried at ambient temperature for further morphological, mineralogical, and crystallographic analysis (Figure 1). Figure 1. (a) Natural Cu-rich aqueous solution from AMD-affected acidic stream mixed with 10 mM AA (pH = 3.95, T = 25 °C) and containing an undetermined transient Cu complex; (b) Metallic copper precipitate formed on the filter after filtration of the solution shown in (a) after 24 h; (c,d) Reflected light microscope photomicrographs (at 500× magnification) taken on the filter shown in (b), showing a network of small acicular microcrystals of metallic copper (Cu[0]) along with sparse large translucent crystals of gypsum which could not be avoided (due to high Ca and SO 4 concentration in the AMD waters). The above explained experiments at pH 3.95 and 25 °C were repeated at conditions of lower pH (1.5–2.5) and higher temperature (40–80 °C), as well as at different AA concentration (10 mM to 0.1 M) to investigate the effect of these parameters on the kinetics of Cu[0] crystallization (see Figure S3 in Supplemental Material). All experiments were conducted under oxygenic conditions. Figure 1. ( a ) Natural Cu-rich aqueous solution from AMD-affected acidic stream mixed with 10 mM AA (pH = 3.95, T = 25 ◦ C) and containing an undetermined transient Cu complex; ( b ) Metallic copper precipitate formed on the filter after filtration of the solution shown in ( a ) after 24 h; ( c , d ) Reflected light microscope photomicrographs (at 500 × magnification) taken on the filter shown in ( b ), showing a network of small acicular microcrystals of metallic copper (Cu[0]) along with sparse large translucent crystals of gypsum which could not be avoided (due to high Ca and SO 4 concentration in the AMD waters). The above explained experiments at pH 3.95 and 25 ◦ C were repeated at conditions of lower pH (1.5–2.5) and higher temperature (40–80 ◦ C), as well as at different AA concentration (10 mM to 0.1 M) to investigate the effect of these parameters on the kinetics of Cu[0] crystallization (see Figure S3 in Supplemental Material). All experiments were conducted under oxygenic conditions. Minerals 2022,12, 322 5 of 18 2.2.2. Reduction with Scrap Iron A second set of experiments was performed to obtain Cu[0] by the classical cementation method using iron as reducing agent. Instead of scrap iron we used clean nuts and screws (previously sonicated and rinsed with ethanol and MilliQ water to remove any dust or contamination from the hardware surfaces) which were submersed in the Cu 2+ - containing acidic liquors (Supplemental Figure S4). These experiments were conducted at conditions of pH = 1.5 (after previous adjustment of the acidic solutions with H 2 SO 4 1 M) and T = 25 ◦ C, with intermittent shaking and in the presence of oxygen, until the pH and Eh values (which were continuously measured) were fully stabilized (the time elapsed in these trials was between 80 and 90 min). The pH for these experiments was adjusted to pH 1.5 following previous research which report a high efficiency of Cu[0] cementation at these very acidic conditions [ 4 , 20 ] and also to avoid the presence of solid-phase Fe(III) (e.g., minerals such as jarosite) which may also form in these waters at pH > 2.0 [ 16 , 17 ]. When the solutions had been stabilized, the iron pieces appeared fully coated by a reddish-brown colored film, and abundant blackish particles had been formed in the aqueous solution (Figure S4). These black, elongated particles were separated from the aqueous solution by vacuum filtration as discussed above. The filters were also air-dried at ambient T and were studied by XRD and SEM-EDX, as described below. 2.3. X-ray Diffraction, Optical Microscopy and Electron Microscopy The metallic copper precipitates were analyzed by different techniques at the SGIker facilities of the University of the Basque Country (UPV/EHU). Cu[0]-coated filters were firstly examined under reflected light optical microscope Nikon Eclipse 50i/POL using a LV-UEPI Universal Epi Illuminator. X-ray diffraction (XRD) was carried out on a PANalytical X’Pert Pro diffractometer (Malvern Panalytical Ltd., Almelo, The Netherlands), with graphite monochromator, a programmable divergence slit, and a solid-state PIXcel detector. The working conditions were 40 kV and 40 mA, Cu K α radiation and a continuous scan range of 2–90 ◦ 2 θ . Samples were finely grinded in an agate mortar, dispersed in ethanol and placed on zero-background Si holder. The goniometer was calibrated with a silicon standard and diffraction data processed using X’Pert High Score software with ICDD database. The obtained precipitates were further studied by Scanning Electron Microscopy (SEM). A JEOL JSM-7000F field emission electron microscope (JEOL Ltd., Tokyo, Japan) coupled to Energy Dispersion X-ray Spectroscope (EDX) was used. Working conditions included 20 kV beam voltage, 1 nA beam current, 10 mm working distance, vacuum < 8.35 ×10−4Pa . Samples were examined in both secondary electron (SE) and backscattered electron (BSE) modes with acquisition time in range of 60 to 100 s at every point of interest. Prior to SEM analysis, samples were mounted on C stab, passed through plasma cleaning process (3 min), and metalized with carbon, avoiding introducing any additional metal to the system. 2.4. Metal Concentration in the Aqueous Solutions Metal concentrations in the aqueous solutions before and after the experiments were also conducted to quantify Cu removal during Cu(II) reduction. These analyses were performed by Inductively Coupled Plasma Atomic Emission Spectrometry (ICP-AES) in a Varian Vista-MPX spectrometer (Agilent Technologies Inc., San Clara, CA, USA). 2.5. Geochemical Modeling The geochemical software package PHREEQCI (Version 3.0.5-7748, US Geological Survey, Reston, VA, USA [ 29 ]) was used for (1) modeling of copper ionic species in solution at varying pH and redox conditions (Eh), and (2) calculation of saturation indices (SI) of selected Cu-bearing minerals. All the calculations were conducted using the MINTEQA2 thermodynamic database (Version 4.0, USEPA, Washington, DC, USA) [ 30 ]. The saturation indices of selected copper-containing mineral phases were calculated using the correspond- Minerals 2022,12, 322 6 of 18 ing solubility products (log K sp values) included in the MINTQA2.V4 database, as well as the known composition of the acidic waters and the experimental conditions. 3. Results 3.1. Reduction of Cu2+ Ions with Ascorbic Acid The first batch reduction experiments were conducted at low AA concentration (10 mM), pH = 3.95 and T = 25 ◦ C (Figure 1). At these conditions, the solutions turned to a yellowish color after a few minutes (Figure 1a). The color of the solution did not change after filtration through 0.45 and 0.1 µ m pore size. Therefore, this color was interpreted to be diagnostic of the formation of either a transient colloidal (nano-sized) compound or an ionic complex containing Cu 2+ or Cu + ions chelated with ascorbate anion. No crystallization was observed for hours, and these solutions were left to evolve overnight at ambient temperature (25 ◦ C). After 24 h, the solution had turned colorless, and a pinkish precipitate had been accumulated in the flask bottom. This solid, which was removed by vacuum filtration (Figure 1b) and subsequently studied by optical microscopy and XRD, consisted almost exclusively in metallic copper (Cu[0]) that was present in the form of elongated, reddish, micrometric crystals of up to 50 µ m, sometimes visible to the naked eye as small reflections (Figure 1c), mingled up with much smaller equidimensional crystals (average around 2 ×2µm). The XRD analyses revealed the major peaks diagnostic of metallic copper (e.g., facets such as 111, 200, 220, or 311), although traces of cuprite (Cu 2 O) could be also deduced by the presence of tiny peaks at around 38 ◦ 2 θ , corresponding to the most intense 111 facets of this mineral (Figure 2). XRD analyses of the precipitates formed in other batch experiments at higher AA concentrations revealed that metallic copper was always the final and stable mineral product resulting from the interaction of the Cu 2+ ions with the added AA, though some tiny peaks suggestive of some minor presence of cuprite could be observed in some samples (Figure 2). Due to the use of AA in the experiments, which is a wellknown antioxidant compound preventing Cu[0] oxidation [ 2 – 5 ], we believe that these trace amounts of Cu 2 O represent remnants of a transitional cuprite nanocolloidal precursor, rather than the result of post-experimental Cu[0] oxidation. Minerals 2022, 11, x FOR PEER REVIEW 6 of 19 The saturation indices of selected copper-containing mineral phases were calculated using the corresponding solubility products (log Ksp values) included in the MINTQA2.V4 database, as well as the known composition of the acidic waters and the experimental conditions. 3. Results 3.1. Reduction of Cu2+ Ions with Ascorbic Acid The first batch reduction experiments were conducted at low AA concentration (10 mM), pH = 3.95 and T = 25 °C (Figure 1). At these conditions, the solutions turned to a yellowish color after a few minutes (Figure 1a). The color of the solution did not change after filtration through 0.45 and 0.1 µm pore size. Therefore, this color was interpreted to be diagnostic of the formation of either a transient colloidal (nano-sized) compound or an ionic complex containing Cu2+ or Cu+ ions chelated with ascorbate anion. No crystallization was observed for hours, and these solutions were left to evolve overnight at ambient temperature (25 °C). After 24 h, the solution had turned colorless, and a pinkish precipitate had been accumulated in the flask bottom. This solid, which was removed by vacuum filtration (Figure 1b) and subsequently studied by optical microscopy and XRD, consisted almost exclusively in metallic copper (Cu[0]) that was present in the form of elongated, reddish, micrometric crystals of up to 50 µm, sometimes visible to the naked eye as small reflections (Figure 1c), mingled up with much smaller equidimensional crystals (average around 2 × 2 µm). The XRD analyses revealed the major peaks diagnostic of metallic copper (e.g., facets such as 111, 200, 220, or 311), although traces of cuprite (Cu2O) could be also deduced by the presence of tiny peaks at around 38°2θ, corresponding to the most intense 111 facets of this mineral (Figure 2). XRD analyses of the precipitates formed in other batch experiments at higher AA concentrations revealed that metallic copper was always the final and stable mineral product resulting from the interaction of the Cu2+ ions with the added AA, though some tiny peaks suggestive of some minor presence of cuprite could be observed in some samples (Figure 2). Due to the use of AA in the experiments, which is a wellknown antioxidant compound preventing Cu[0] oxidation [2–5], we believe that these trace amounts of Cu2O represent remnants of a transitional cuprite nanocolloidal precursor, rather than the result of post-experimental Cu[0] oxidation. Figure 2. XRD patterns of Cu[0] obtained in the different batch experiments conducted in this study: (I) large Cu[0] crystals formed by reduction of Cu(II)-rich solutions with ascorbic acid 10 mM at pH 3.95 and 25 °C (TI; bottom), (II) Cu[0] nanoparticles formed by reduction with ascorbic acid 0.1 M at pH 3.95 and 80 °C (CA, ST; center), and (III) Cu[0] nanoparticles formed by reduction with Fe[0] (scrap iron) at pH 1.5 and 25 °C (ST; top). Figure 2. XRD patterns of Cu[0] obtained in the different batch experiments conducted in this study: (I) large Cu[0] crystals formed by reduction of Cu(II)-rich solutions with ascorbic acid 10 mM at pH 3.95 and 25 ◦ C (TI; bottom), (II) Cu[0] nanoparticles formed by reduction with ascorbic acid 0.1 M at pH 3.95 and 80 ◦ C (CA, ST; center), and (III) Cu[0] nanoparticles formed by reduction with Fe[0] (scrap iron) at pH 1.5 and 25 ◦C (ST; top). Minerals 2022,12, 322 7 of 18 The experiments conducted at pH 3.95 and AA concentrations of 0.1 M showed much faster precipitation kinetics. A dense brownish-red-colored suspension containing Cu[0] particles was formed around 5–10 min after the addition of AA (Figure S3). The reduction experiments carried out at higher T (40–80 ◦ C) did not show a significant change in the precipitation kinetics as compared to the ones conducted at room temperature, though the reddish-brown suspensions were denser in the former case at comparable pH values. On the other hand, reduction experiments performed at low pH (1.5–2.5) did not produce either metallic copper or any other mineral product after 1 h reaction time. This is consistent with the increased electrode potential of AA at very low pH [ 2 ], which could have significantly decreased the reactions kinetics. The lower intensity of major Cu[0] peaks (e.g., 111, 200, 220) in the XRD patterns obtained in the reduction experiments at higher AA concentration, along with faster precipitation kinetics suggested lower crystallinity degree and/or size of the Cu[0] particles (Figure 2). This hypothesis was further confirmed by the SEM-EDX studies reported in the next section. 3.2. Textural and Crystallographic Properties Revealed by Electron Microscopy Under SEM, the Cu[0] crystals formed in the slow crystallization experiments at low AA concentration appeared with diverse morphology ranging from acicular-elongated, pseudoprismatic to isomorphic (e.g., cubes, truncated cubes, cuboctahedra), tetrahedral or truncated tetrahedra, and trapezoidal polyhedra (Figures 3and 4). The latter ranged in size between 2–7 µ m (width) and 1–4 µ m (length), with a very good correlation between both dimensions (Figure 4d,e). On the other hand, pseudoprismatic and acicular-elongated crystals showed widths of around 1–4 µ m but ranged largely in length between 2 and 50 µ m (Figure 4b), sometimes appearing as very thin bent needle-like crystals ( Figure 3a–e ). Some bigger crystals showed much smaller crystallites (with varying shapes and diameters smaller than 200 nm) showing oriented growth in some of their edges or faces (e.g., Figure 3f), suggesting heterogeneous, surface-controlled crystal growth. The EDX analyses of these crystals yielded compositions of almost pure, electrolytic copper (>99% Cu ; not shown). In contrast to the slow kinetics–low AA concentration experiments, the fast kinetics– high AA concentration experiments produced a precipitate formed by nanometric, subrounded particles of metallic copper ranging in size between 40 and 200 nm (Figure 5). These SEM observations were in agreement with the lower crystallinity degree suggested by the XRD spectra of these samples, as described above (Figure 2). The discrete nanoparticles coalesced together in larger aggregates (with diameter of up to few microns) that could contain tens to hundreds of smaller particles (Figure 5). 3.3. Reduction of Cu2+ Ions with Scrap Iron The reduction experiments using scrap iron at 25 ◦ C and pH 1.5 resulted in the simultaneous formation of black particles in the acidic solution and a brownish-red-colored coating on the metallic pieces as the reaction proceeded (Figure S4). A high density of gas bubbles was also observed in these experiments, consistent with the production of a gaseous phase (likely hydrogen [H 2 ], as discussed below). Filtration of the solution after 30 min of reaction time produced a dense, black-colored precipitate (Figure S4). SEM examination of these solids revealed that these were formed by almost pure (>99%) metallic copper. In this case, metallic copper was generally present as pine-tree-like forms composed of arrays of submicron crystallites growing in diverse directions (Figure 6a–e). Less abundant forms of metallic copper produced by the cementation method included pseudo-spherical aggregates composed by coalesced, nanometric particles (<100 nm, Figure 6f) and diverse encrusting aggregates showing porous growths as boxwork-type or thin curved films (Figure 6e). Minerals 2022,12, 322 8 of 18 Minerals 2022, 11, x FOR PEER REVIEW 8 of 19 Figure 3. Scanning electron microscopy (SEM) images of the Cu crystals obtained in batch reduction experiments with Cu(II)-rich acidic solutions (pH 3.95, 25 °C) and ascorbic acid 10 mM. Pictures (a– f) display different textural and crystallographic features in a single sample (same sample as in Figure 1b,c). Figure 3. Scanning electron microscopy (SEM) images of the Cu crystals obtained in batch reduction experiments with Cu(II)-rich acidic solutions (pH 3.95, 25 ◦ C) and ascorbic acid 10 mM. Pictures ( a – f ) display different textural and crystallographic features in a single sample (same sample as in Figure 1b,c). Minerals 2022,12, 322 9 of 18 Minerals 2022, 11, x FOR PEER REVIEW 9 of 19 Figure 4. Geometric characteristics of different types of Cu[0] crystals formed in ascorbic acid reduction experiments (10 mM AA, pH 3.95, 25 °C) as deduced from the SEM pictures shown in Figure 3: (a,b) Pseudoprismatic-elongated-acicular crystals, (c) Isomorphic polyhedra, (d,e) Tetrahedraltrapezoidal polyhedra. In contrast to the slow kinetics–low AA concentration experiments, the fast kinetics– high AA concentration experiments produced a precipitate formed by nanometric, subrounded particles of metallic copper ranging in size between 40 and 200 nm (Figure 5). These SEM observations were in agreement with the lower crystallinity degree suggested Figure 4. Geometric characteristics of different types of Cu[0] crystals formed in ascorbic acid reduction experiments (10 mM AA, pH 3.95, 25 ◦ C) as deduced from the SEM pictures shown in Figure 3 : ( a , b ) Pseudoprismatic-elongated-acicular crystals, ( c ) Isomorphic polyhedra, ( d , e ) Tetrahedraltrapezoidal polyhedra. Minerals 2022,12, 322 16 of 18 The high yield of Cu recovery in these assays (>98%) in short reaction times of around 90 min, along with the very low cost of the reducing agent, makes this mechanism much more efficient and cheaper than the ones based in chemical reduction using AA or any other reducing agent, such as glucose or hydrazine hydrate. The high purity of metallic copper aggregates (>99% Cu[0]) obtained by this technique and the very little requirements (e.g., good performance even at very low pH, no need for expensive chemical reagents, no need for high temperature or intense stirring, etc.) makes cementation the easiest and simplest method of Cu recovery from Cu-rich mine waters and their acidic leachates. 5. Conclusions The current practices in different industrial sectors focused on Cu recovery from different wastes include Cu 2+ reduction by ascorbic acid as a feasible and apparently efficient technology. At the same time, there is an important need to develop cost-effective strategies of metal recovery in abandoned mine sites within the framework of circular economy and mine waste recycling [ 10 , 15 ]. An obvious question, therefore, emerges about the convenience of ascorbic acid-based procedures to recover Cu from acidic mine waters containing high concentrations of this metal as dissolved loadings. We provide here, for the first time, a comparison of metallic copper products obtained by two different reduction methods from Cu-rich acid mine waters. The present work has shown that high-purity metallic copper can also be obtained from highly acidic, metal-rich mine drainage solutions by reduction of dissolved Cu 2+ by ascorbic acid. This procedure may result in the formation of either (i) well-developed large crystals or (ii) much smaller rounded particles of sub-micron to nanometric diameter, depending on the crystallization kinetics and saturation state, which have been observed to vary with the ascorbic acid concentration used, as well as with other parameters such as pH or temperature. Compared to previous studies [ 2 – 9 , 34 – 36 ], we have identified a surprisingly high diversity of crystal geometries (e.g., from acicular to polyhedral) and sizes (from <200 nm to >50 µ m) when using low concentrations of AA and slow crystallization kinetics. However, this procedure shows a low efficiency as regards to metal recovery rates (<25%) at the pH conditions ( 1.50–3.95 ) and AA concentrations (10 mM–0.1 M) considered in this study. This method presents some other disadvantages, such as the need of a previous partial neutralization to increase the Cu-reduction kinetics, or the undesirable removal of Cu 2+ by adsorption on Al colloids formed in these waters at pH > 4.0. These two limitations leave a narrow operational pH window for on-site treatments in abandoned mine sites. In any case, the possibility of controlling crystal geometry or particle size as a function of varying experimental conditions (e.g., pH, T, [Cu/AA] ratio) could make this procedure still interesting if further improvements are conducted to increase its efficiency. In contrast, the classical cementation method of Cu 2+ reduction by scrap iron, historically applied in the IPB and other mine districts [ 20 , 21 ], offers an easier, cheaper, and more effective option of Cu recovery from mine waters. The metallic copper produced by this method is also highly crystalline, although commonly forms aggregates of smaller crystallites than those formed by the slow kinetics ascorbic acid method. The big advantages of the cementation method are: (i) it can be applied at very low pH (thus, with no need of alkalization), and (ii) it does not require expensive reagents. Therefore, this method could be used in combination with some other techniques in a multi-metal recovery scheme aimed at obtaining secondary raw materials or at least at defraying remediation treatment costs. Supplementary Materials: The following supporting information can be downloaded at: https:// www.mdpi.com/article/10.3390/min12030322/s1, Figure S1: Pictures on partial neutralization of AMD water prior to Cu 2+ reduction experiments; Figure S2: Plot of pH vs ascorbic acid concentration; Figure S3 : Cu 2+ reduction experiments at higher pH and ascorbic acid concentration; Figure S4 : Pictures showing one of the experiments of Cu 2+ reduction using iron (Fe[0]) metallic pieces; Figure S5 : Eh/pH diagrams for the Cu-O-H system using different thermodynamic databases. Minerals 2022,12, 322 17 of 18 Author Contributions: Conceptualization, J.S.-E.; Methodology, J.S.-E.; Validation, J.S.-E., I.Y. and A.I.; Formal Analysis, J.S.-E., I.Y. and A.I.; Investigation, J.S.-E., I.Y. and A.I.; Resources, J.S.-E. and I.Y.; Data Curation, J.S.-E. and A.I.; Writing—Original Draft Preparation, J.S.-E.; Writing—Review and Editing, J.S.-E., I.Y. and A.I.; Project Administration, J.S.-E.; Funding Acquisition, J.S.-E. All authors have read and agreed to the published version of the manuscript. Funding: This research was funded by the Spanish Ministry of Science and Innovation through grant number CGL2016-74984-R. Acknowledgments: We acknowledge the technicians and staff at the Geochemistry Laboratory at IGME-CSIC for their kind help and support during the experimental work. We also thank the personnel at the SGIker facilities at UPV/EHU for their support during the XRD and SEM analyses. We are grateful to two anonymous reviewers who helped improve this manuscript. Conflicts of Interest: The authors declare no conflict of interest. The funders had no role in the design of the study; in the collection, analyses, or interpretation of data; in the writing of the manuscript, and in the decision to publish the results. References 1. Ghosh, S. Electroless copper deposition: A critical review. Thin Solid Films 2019,669, 641–658. [CrossRef] 2. Saikova, S.V.; Murasheva, K.S.; Vorobyev, S.A.; Kochmarev, K.Y.; Karimov, E.E.; Eremina, A.D.; Mikhlin, Y.L. Synthesizing highly concentrated hydrosols of copper nanoparticles via reduction by ascorbic acid in the presence of gelatose. Chem. Sust. Dev. 2013 , 21, 403–409. 3. Andal, V.; Buvaneswari, G. Effect of reducing agents in the conversion of Cu 2 O nanocolloid to Cu nanocolloid. Eng. Sci. Technol. 2017,20, 340–344. [CrossRef] 4. Kumar, S.; Kumar, V.; Sharma, M.L.; Chakarvarti, S.K. Electrochemical synthesis of metallic micro-rose having petals in nanometer dimensions. Superlattice Microstruct. 2008,43, 324–329. [CrossRef] 5. Tsai, C.-Y.; Chang, W.-C.; Chen, G.-L.; Chung, C.-H.; Liang, J.-X.; Ma, W.-Y. A study of the preparation and properties of antioxidative copper inks with high electrical conductivity. Nanoscale Res. Lett. 2015,10, 357. [CrossRef] [PubMed] 6. Virk, H.S. Fabrication and characterization of metallic copper and copper oxide nanoflowers. Pak. J. Chem. 2011 ,1, 148–154. [CrossRef] 7. Virk, H.S.; Kishore, K.; Balouria, V. Fabrication of copper nanowires by electrodeposition using anodic alumina and polymer templates. J. Nano Res. 2010,10, 63–67. [CrossRef] 8. Sarkar, J.; Khan, G.G.; Basumallick, A. Nanowires: Properties, applications and synthesis via porous anodic aluminum oxide template. Bull. Mater. Sci. 2007,30, 271–290. [CrossRef] 9. Zhang, D.-F.; Zhang, H.; Guo, L.; Zheng, K.; Han, X.-D.; Zhang, Z. Delicate control of crystallographic facet-oriented Cu 2 O nanocrystals and the correlated adsorption ability. J. Mater. Chem. 2009,19, 5220–5225. [CrossRef] 10. Lottermoser, B.G. Recycling, Reuse and Rehabilitation of Mine Wastes. Elements 2011,7, 405–410. [CrossRef] 11. Nordstrom, D.K.; Alpers, C.N. Geochemistry of acid mine waters. In The Environmental Geochemistry of Mineral Deposits, Part A; Processes, Techniques, and Health Issues; Plumlee, G.S., Logsdon, M.J., Eds.; Society of Economic Geologists: Littleton, CO, USA, 1999; Volume 6, pp. 133–156. 12. Geller, W.; Klapper, H.; Schultze, M. Natural and anthropogenic sulphuric acidification of lakes. In Acidic Mining Lakes–Acid Mine Drainage, Limnology and Reclamation; Geller, W., Klapper, H., Salomons, W., Eds.; Springer: Berlin/Heidelberg, Germany, 1998; pp. 3–14. 13. Barthen, R.; Sulonen, M.L.K.; Peräniemi, S.; Jain, R.; Lakaniemi, A.M. Removal and recovery of metal ions from acidic multi-metal mine water using waste digested activated sludge as biosorbent. Hydrometallurgy 2022,207, 105770. [CrossRef] 14. Isosaari, P.; Sillampää, M. Use of Sulfate-Reducing and Bioelectrochemical Reactors for Metal Recovery from Mine Water. Sep. Purif. Rev. 2014,46, 1–20. [CrossRef] 15. León, R.; Macías, F.; Cánovas, C.; Pérez-López, R.; Ayora, C.; Nieto, J.M.; Olías, M. Mine waters as a secondary source of rare earth elements worldwide: The case of the Iberian Pyrite Belt. J. Geochem. Exp. 2021,224, 106742. [CrossRef] 16. Sánchez-España, F.J.; López Pamo, E.; Santofimia, E.; Aduvire, O.; Reyes, J.; Barettino, D. Acid Mine Drainage in the Iberian Pyrite Belt (Odiel river watershed, Huelva, SW Spain): Geochemistry, Mineralogy and Environmental Implications. App. Geochem. 2005 , 20, 1320–1356. [CrossRef] 17. Sánchez España, J.; López Pamo, E.; Santofimia, E.; Diez-Ercilla, M. The acidic mine pit lakes of the Iberian Pyrite Belt: An approach to their physical limnology and hydrogeochemistry. App. Geochem. 2008,23, 1260–1287. [CrossRef] 18. López-Pamo, E.; Sánchez-España, J.; Santofimia, E.; Diez-Ercilla, M.; Reyes, J. Cortas Mineras Inundadas de la Faja Pirítica: Inventario e Hidroquímica; Instituto Geológico y Minero: Madrid, Spain, 2009; 279p. 19. Sánchez-España, J.; Diez, M.; Santofimia, E. Mine pit lakes of the Iberian Pyrite Belt: Some basic limnological, hydrogeochemical and microbiological considerations. In Acidic Pit Lakes: The Legacy of Coal and Metal Surface Mines; Geler, W., Ed.; Springer: Berlin/Heidelberg, Germany, 2013; pp. 315–342. Minerals 2022,12, 322 18 of 18 20. Taylor, J.H.; Whelan, P.F. The leaching of cupreous pyrites and the precipitation of copper at Rio Tinto, Spain. Trans. Inst. Min. Metall. 1943,52, 35–96. 21. Tucci, N.J.; Gammons, C.H. Influence of copper recovery on the water quality of the acidic Berkeley Pit lake, Montana, U.S.A. Environ. Sci. Technol. 2015,49, 4081–4088. [CrossRef] 22. Bigham, J.M.; Nordstrom, D.K. Iron and Aluminum Hydroxysulfates from Acid Sulfate Waters. In Sulfate Minerals: Crystallography, Geochemistry, and Environmental Significance. Rev. Mineral. Geochem. 2000,40, 351–403. [CrossRef] 23. Sánchez-España, J. The behavior of iron and aluminum in acid mine drainage: Speciation, mineralogy, and environmental significance. In Thermodynamics, Solubility and Environmental issues; Letcher, T., Ed.; Elsevier: Amsterdam, The Netherlands, 2007; pp. 137–150. 24. Sánchez-España, J.; López-Pamo, E.; Santofomia, E.; Reyes, J.; Martín-Rubí, J.A. The removal of dissolved metals by hydroxysulphate precipitates during oxidation and neutralization of acid mine waters, Iberian Pyrite Belt. Aquat. Geochem. 2006 ,12, 269–298. [CrossRef] 25. Sánchez-España, J.; Yusta, I.; López, G. Schwertmannite to jarosite conversion in the water column of an acidic mine pit lake. Mineral. Mag. 2012,76, 2659–2682. [CrossRef] 26. Nordstrom, D.K. The effect of sulphate on aluminum concentrations in natural waters: Some stability relations in the system Al2O3-SO3-H2O at 298 K. Geochim. Cosmochim. Acta 1982,46, 681–692. [CrossRef] 27. Sánchez-España, J.; Yusta, I.; Burgos, W.D. Geochemistry of dissolved aluminum at low pH: Hydrobasaluminite formation and interaction with trace metals, silica and microbial cells under anoxic conditions. Chem. Geol. 2016,441, 124–137. [CrossRef] 28. Shen, J.; Griffiths, P.T.; Campbell, S.J.; Utinger BKalberer, M.; Paulson, S.E. Ascorbate oxidation by iron, copper and reactive oxygen species: Review, model development, and derivation of key rate constants. Sci. Rep. 2021 ,11, 7417. [CrossRef] [PubMed] 29. Parkurst, D.L.; Appelo, C.A.J. User’s guide to PHREEQC (version 2)—A computer program for speciation, batch-reaction, one-domensional transport, and inverse geochemical calculations. In U.S. Geological Survey Water-Resources Investigations Report 99-4259; U.S. Geological Survey: Denver, CO, USA, 1999; 312p. 30. Allison, J.D.; Brown, D.S.; Novo-Gradac, J. MINTEQA2/PRODEAFA2, A Geochemical Assessment Model for Environmental Systems: User Manual Supplement for Version 4.0; U.S. Environmental Protection Agency: Athens, GA, USA, 1999. 31. Sánchez-España, J.; Yusta, I. Low-crystallinity products of trace-metal precipitation in neutralized pit-lake waters without ferric and aluminous adsorbent: Geochemical modelling and mineralogical analyses. Mineral. Mag. 2015,79, 781–798. [CrossRef] 32. Takeno, N. Atlas of Eh-pH diagrams: Intercomparison of thermodynamic databases. In Geological Survey of Japan Open File Report No. 419; National Institute of Advanced Industrial Science and Technology: Tokyo, Japan, 2005; 285p. 33. Mystkowski, E.M. The oxidation of ascorbic acid in the presence of copper. Biochem, J. 1942,36, 494–500. [CrossRef] 34. Hacisevkí, A. An overview of ascorbic acid biochemistry. J. Fac. Pharm. Ank. 2009,38, 233–255. 35. Goçalves, T.A.; Botelho, A.B.; de Moraes, V.T.; Espinos, D.C.R. Study of pH Influence in the Synthesis of Copper Nanoparticles Using Ascorbic Acid as Reducing and Stabilizing Agent. In Proceedings of the TMS 2020 149th Annual Meeting & Exhibition, San Diego, CA, USA, 23–27 February 2020; The Minerals, Metals & Materials Society, Ed.; The Minerals, Metals & Materials Series; Springer Nature: Cham, Switzerland, 2020; pp. 1547–1557. 36. Liu, Q.M.; Yasunami, T.; Kuruda, K.; Okido, M. Preparation of Cu nanoparticles with ascorbic acid by aqueous solution reduction method. Trans. Nonferrous Met. Soc. China 2012,22, 2198–2203. [CrossRef] 37. Songping, W. Preparation of fine copper powder using ascorbic acid as reducing agent and its application in MLCC. Mater. Lett. 2017,61, 1125–1129. [CrossRef] 38. Sánchez-España, J. Crystallization in acidic media: From nanoparticles to macrocrystals. Semin. Soc. Esp. Mineral. 2017 ,13, 15–34. 39. Jacukowicz-Sobala, I.; Stanisławska, E.; Baszczuk, A.; Jasiorski, M.; Kociołek-Balawejder, E. Size-Controlled Transformation of Cu 2 O into Zero Valent Copper within the Matrix of Anion Exchangers via Green Chemical Reduction. Polymers 2020 ,12, 2629. [CrossRef] 40. Ford, W.E. Dana’s Textbook of Mineralogy: With an Extended Treatise on Crystallography and Physical Mineralogy, 4th ed.; CBS Publishers: New Delhi, India, 2006; 851p. 41. Ünalero˘glu, C.; Mert, Y.; Zümreo˘glu-Karan, B. Synthesis and characterization of copper ascorbate. Synth. React. Inorg. Met.-Org. Chem. 2001,31, 1531–1543. [CrossRef] 42. Stefanowicz, T.; Osinska, M.; Napieralska-Zagozda, S. Copper recovery by the cementation method. Hydrometallurgy 1997 ,47, 69–90. [CrossRef]