Submitted Successfully!
To reward your contribution, here is a gift for you: A free trial for our video production service.
Thank you for your contribution! You can also upload a video entry or images related to this topic.
Version Summary Created by Modification Content Size Created at Operation
1
César Sequeira
 -- 3438 2022-07-01 12:32:34 |
2 format
Jason Zhu
 + 1 word(s) 3439 2022-07-04 05:56:47 | |
3 format
Jason Zhu
 -1 word(s) 3438 2022-07-07 04:19:24 | |
4 format
Jason Zhu
 Meta information modification 3438 2022-07-08 03:49:35 |

Video Upload Options

Do you have a full video?
Send video materials Upload full video

Confirm

Are you sure to Delete?
Yes No
Cite
If you have any further questions, please contact Encyclopedia Editorial Office.
Sequeira, C. Application of Mixed Potential Theory to Leaching. Encyclopedia. Available online: https://encyclopedia.pub/entry/24740 (accessed on 17 May 2024).
Sequeira C. Application of Mixed Potential Theory to Leaching. Encyclopedia. Available at: https://encyclopedia.pub/entry/24740. Accessed May 17, 2024.
Sequeira, César. "Application of Mixed Potential Theory to Leaching" Encyclopedia, https://encyclopedia.pub/entry/24740 (accessed May 17, 2024).
Sequeira, C. (2022, July 01). Application of Mixed Potential Theory to Leaching. In Encyclopedia. https://encyclopedia.pub/entry/24740
Sequeira, César. "Application of Mixed Potential Theory to Leaching." Encyclopedia. Web. 01 July, 2022.
Application of Mixed Potential Theory to Leaching
Edit
This entry is adapted from the peer-reviewed paper 10.3390/reactions3020023

Leaching is a central unit operation in the hydrometallurgical processing of minerals, which often occurs by means of electrochemical reactions. Application of mixed potential theory to explain the kinetics of oxidative and reductive leaching processes is a useful concept in explaining observed results. Native metals, selected oxides, and most base metal sulfides are electron-conducting phases. For these minerals, leaching may take place by normal corrosion, passivation or galvanic couple mechanisms, which provide individual electrode kinetics enabling the calculation of mixed potentials and overall reaction kinetics. 

Leaching mechanism dissolution galena gold silver pyrite sphalerite chalcopyrite mixed potential theory

1. Leaching of Galena

The electrochemical dissolution of galena in aqueous solutions has received considerable attention. The electrical conductivity of galena is unusually high, and galena is recognized as having both n- and p-type conductivity. Eadington and Prosser [1] have demonstrated that the oxidation of galena is noticeably influenced by the semiconductor type. For non-stoichiometric PbS, the rate of oxidation was much higher for sulfur-rich (p-type semi-conductor) specimens than for lead-rich (n-type) specimens [2][3][4]. In aqueous solution, lead forms a large array of sparingly soluble or insoluble compounds; these materials can form surface films which mask the original electronic properties of the solid [5]. Galena reacts anodically according to the overall reaction
 Paul et al. [5] suggested a mechanism for the dissolution of galena in acidic perchlorate solutions. The overall anodic reaction is given in Equation. Tafel slopes fail to support this simple two-electron- transfer reaction. Therefore, the oxidation of galena is assumed to take place according to sequential one-electron-transfer steps,
The Butler–Volmer treatment for steady-state conditions (d[S−]/dt=0) yields the following equation for the net anodic current
where [Pb2+]S is the concentration of lead ions at the electrode surface, k⃗  and k are the respective forward and reverse rate constants for Reaction, k′ is the specific rate of Reaction and the other terms are analogous to those previously defined. At sufficiently positive values of E or at low [Pb2+], the expression for ia reduces to
Conversely, when the [Pb2+] is high and E is low, 
Both cases cited above, along with an additional condition when [Pb2+] approaches zero, have been verified experimentally. Therefore, the proposed mechanism involving consecutive one-electron-transfer reactions is an accurate description of the anodic dissolution of galena. The initial dissolution reaction explained by Equation is soon superceded by diffusion rate control due to the products of reaction on the galena surface.
Chloride solution chemistry is of especial interest to lead metallurgy because of the high solubility of PbCl2 at elevated temperatures [6][7][8]. Galena will react anodically, as illustrated by the half-cell reaction, Equation. This reaction may be promoted by the coupled cathodic reduction of ferric chloride
 resulting in the overall stoichiometry 
The lead released to the solution forms a series of lead chloride complex ions, PbCl2−yy, depending on the amount of excess Cl− in solution. Fuerstenau et al. [8] and Dutrizac [7] have shown that the reaction kinetics follow Fick’s diffusion equation for isometric particles:
Dutrizac [7] has shown a similar shift in rate control, finding kp to be independent of FeCl3 for concentrations greater than 0.1 M in the absence of added NaCl. In the higher FeCl3 concentration range, kp decreased linearly with increased PbCl2 concentration. These results suggest the rate-controlling step is the inward pore diffusion of ferric chloride through the growing sulfur product layer at low rates of reaction and outward pore diffusion of lead chloride at high rates of reaction. For the latter, the calculated activation energy was 16 kJ/mol (4 kcal per mol) consistent with the pore-diffusion model. Nitrate, sulphate, and fluosilicate solutions chemistry are also of particular interest for leaching of lead [9].

2. Leaching of Gold and Silver

Leaching is the primary means for the recovery of precious metals from ore. Hydrometallurgy also represents an alternative for the recovery of metals from industrial metallic waste for purposes of recycling. For 100 years, cyanidation has been the principle method for gold (Au) and silver (Ag) extraction worldwide [10][11][12][13][14]. The leaching mechanism was recognized by early investigators to be electrochemical in nature. Moreover, dissolution rates have been observed to display unique kinetic characteristics, shifting from one kinetic regime to another with only slight variations in concentration of oxidant or complexing agent in solution [15][16]. In the case of conventional Au and Ag extraction recovery, O2 in air is the oxidant and CN−the complexing agent.
Kudryk and Kellogg [17] published an important fundamental paper establishing the electrochemical nature of gold cyanidation and explained the abrupt shift in kinetic regimes on the basis of mixed potentials.
In the plateau regions, the current is independent of voltage, indicative of solution boundary-layer diffusion. The plateau values at 0.005%, 0.01%, and 0.0175% KCN are controlled kinetically by CN− diffusion. The plateau region of the cathodic branch is a region where oxygen diffusion is rate controlling. The mixed potential is the potential that a solid metallic particle will assume when exposed to the leaching solution and, for a given solution composition, corresponds to the point of intersection of the anodic and cathodic branches. The mixed potential and its associated mixed current establish the rate of leaching.
As %KCN is increased, the rate of reaction moves from CN− diffusion control to O2 diffusion control. Kudryk and Kellogg also showed the reverse to be true for a fixed %KCN and changing %O2. The kinetics are well described by the Fick diffusion equation when the mixed potential falls within either the CN− or O2 plateau regions. To explain the kinetic results in the transition region, the electrokinetic reactions that establish the anodic and cathodic branches for all concentrations must be determined. It can be shown that the experimental curves measured by Kudryk and Kellogg are not totally diffusion overpotential limited, but require mixed diffusion plus surface charge-transfer kinetics to explain the shape of the curves.
Equation describes the overall anodic dissolution of Au in cyanide solutions.
where DCN is the diffusion coefficient for CN−, (CN¯¯¯¯¯−) is the bulk solution concentration, (CN−) is the surface concentration, α is the diffusion boundary thickness, F is the Faraday constant, δ is the thickness of the boundary layer (which varies between 2 and 0.009 cm, depending on the speed and method of agitation), αa is the transfer coefficient and Ea is the voltage for the anodic reaction. Assuming steady-state conditions, Equations may be combined, giving 
where ka=KCNF/2δ and Ba=αaF/RT. The constant kawas evaluated from average plateau values for the 0.005, 0.01, and 0.0175% KCN anodic curves. In this region, the current density is the limiting current density, ida, and is equal to ka(CN¯¯¯¯¯−), providing a means to evaluate ka. The value of Ba was determined from one point on the rising portion of the 0.0175% KCN curve.
Oxygen discharge is known to proceed through a series of reactions including peroxide intermediates. In basic solutions, simplified steps may be represented by
 the total reaction being 
In the cyanidation process, the effective transfer of electrons will fall in the range of two to four electrons per mole of O2, depending on the net rate of Reaction. The cathodic O2 discharge curves of Kudryk and Kellogg are also adequately described by a mixed diffusion plus surface charge-transfer process.
where ic is the cathodic current density, D(O2) is the diffusion coefficient for oxygen, and n is the effective number of electrons transferred. Equations may be combined to give
where kc=DO2nF/δ and Bc=αcF/RT. The dashed curve was calculated using Equation. The mixed potential E and current i correspond to the conditions Ea=Ec=E and ia=ic=i. Equations may be combined to give 
where
The calculated rates may be determined from the electrochemical data. The rate of reaction R (mg cm−2 h−1) is directly proportional to the mixed current density such that Equation becomes
where Rda and Rdc are the limiting rates of dissolution corresponding to the limiting current densities ida and idc. Ideally, D and Ba/Bc values may be determined from the electrochemical data alone. A good correlation is obtained for the electrochemical data of Kudryk and Kellogg. Measured rates of dissolution fall somewhat below those predicted from anodic and cathodic polarization data alone. This may be attributed in part to the oxidation of CN− in the presence of oxygen. Normally dissolution data, rather than electrochemical data, will be available and used to evaluate D and BaBc.

3. Leaching of Pyrite

Pyrite, the most common sulfide mineral, dissolves by the following reactions:

The overall reaction can be written in terms of the half-reactions for the oxidation of pyrite and the reduction of ferric ions and oxygen. Holmer and Crundwell [18] investigated each of these half-reactions and investigated the overall reaction. These measurements showed that the kinetics of the half-reaction for the anodic dissolution of pyrite is given by:

 and that the half-reaction for the reduction of ferric ions is given by:

Here E is the mixed potential, after the corrosion theory of Wagner and Traud (1938) [19]. The order of the reaction of—0.5 with respect to H+ for the anodic reaction indicates that reaction consists of some steps, possibly involving adsorbed hydroxide ions, prior to the rate-determining step. Considering that every electron donated by the anodic dissolution of pyrite is instantaneously accepted by the reduction of ferric ions, then τFeS2=τFe and 

The substitution of this equation for the mixed potential into Equation yields the rate equation for the oxidative dissolution of pyrite by ferric ions. When [Fe2+] is greater than 0.001 M, i.e., KFeS2[H+]−1/2 is much less than K′Fe[Fe2+], then this equation is given by

This equation predicts that the rate of dissolution of pyrite is one-half order in ferric ions and negative one-half order in H+. Results of recent investigations of the rate of dissolution of pyrite by ferric ions are correctly described by the electrochemical mechanism derived by Holmes and Crundwell [18]. They also measured the cathodic reduction of oxygen on pyrite, and showed that the order of reaction with respect to H+ is 0.14. Using this result, they derived an expression for the rate of dissolution of pyrite in the presence of oxygen, which is given as follows:

This result is again related to the electrochemical mechanism of Holmes and Crundwell [18], and was confirmed by Rimstidt and Vaughan [20]. The effect of the concentrations of ferric and ferrous ions on the mixed potential of pyrite showed that the slope of the mixed potential is 0.059 mV/decade, which is in good agreement with the theoretical value. If KFeS2[H+]−1/2 is much less than K′Fe[Fe2+], then the predicted value of the slope of the mixed potential is—0.059 mV/decade, which is close to the experimentally determined value of—0.056 mV/decade. Thus, once again, the experiments of the mixed potential confirm the theory.

4. Leaching of Sphalerite

Sphalerite (ZnS) is a semiconductor with a wide band-gap, which is another way of saying that it is an electrical insulator. The valence band is comprised of electron orbitals that are of bonding character, and the conduction band is comprised of electron orbitals that are of non-bonding character. In order for this material to be dissolved by an oxidative mechanism, electrons must be removed from the bonding orbitals, that is, from the valence band. The positive charge that arises from the removal of an electron from the valence band is referred to as a hole. As a result, for dissolution to occur bonding electrons must be removed from the valence orbital by the oxidant in solution.
The presence of iron that substitutes for zinc atoms in the sphalerite lattice results in a d-orbital band within the band-gap of sphalerite. The iron d-orbitals of this band are of bonding character. This means that the removal of an electron from this band (injection of a hole) also results in dissolution of the solid. The iron impurity and its associated d-orbital band have two consequences for the dissolution of sphalerite: the d-orbital band presents a narrow localized band with which the transfer of electrons is energetically more favorable than it is with the valence band, and the d-orbital band “pins” the Fermi level so that changes in the interfacial potential occur on the solution side of the interface rather than on the solid side.
Using the principles above it is easy to derive that the dissolution of sphalerite by ferric ions occurs according to the reaction:

 whose rate of reaction can be described by the following kinetic expression:

Crundwell [21], using principles of quantum electrochemistry, derived the following expressions for the rates of the anodic half-reaction and the cathodic half-reaction

where Nd is the concentration of iron in the sphalerite (mol Fe/mol Zn). Eliminating E between Equations, it can be obtained the rate of the surface reaction:

where αa and αc have been assumed to be equal to one-half, in accordance with a rate-determining step based on the transfer of charge.
An interesting feature is the effect of light on the rate of dissolution. Electrons can be excited from bonding orbitals at the surface by shining light on them if the energy of the light matches the band gap of the semiconductor.
Clearly, ultraviolet light increases the rate of reaction, which serves to confirm that the rate determining step in the dissolution of this mineral is the transfer of charge at the mineral surface. If the rate-determining step was not the transfer of charge, there would have been no effect as a result of shining ultraviolet light on the particles during dissolution.

5. Leaching of Chalcopyrite

Chalcopyrite is not only the most abundant of the copper sulfides, but also the most stable, making it recalcitrant to hydrometallurgical processes. Hence, hydrometallurgical processing of chalcopyrite continues to be an attractive area of research due to the vaguely understood surface chemistry of the mineral in different aqueous media [22][23][24][25][26][27][28][29][30][31][32][33].
Different routes for hydrometallurgical treatment of chalcopyrite can be followed. These include thermal treatment prior to leaching, direct leaching, and direct electrochemical leaching.
Ammoniacal solutions are attractive and effective lixiviants that form stable amine complexes with some base metal cations while rejecting iron. Leaching of chalcopyrite in ammoniacal solutions in the presence of an oxidant is possible due to the stabilization of copper (I) and copper (II) by ammonia at elevated pH levels. In oxygenated ammonia solutions, it has been suggested that chalcopyrite dissolves according to Equation:
Chalcopyrite is characterized by very slow leaching kinetics and this has been strongly linked to the formation of a passive film on its surface. Researchers have not reached a consensus on the actual composition or degree of stability of this passive surface film. Fe2O3 or its hydrated form Fe2O3xH2O has been reported to be the surface product of chalcopyrite oxidation. Moyo [25] studied the electrochemical oxidation of chalcopyrite in ammonia-ammonium sulphate solutions and reported significant amounts of ferrous iron (3–20%) in the product film. The researchers postulated that a ferrous iron intermediate was formed, which can be readily oxidized to ferric, and the oxidation can be achieved even if only traces of dissolved oxygen are present. If ferrous iron is formed from chalcopyrite dissolution, it is expected that its further reactions would be affected by the solution conditions, hence the different surface effects as observed under the carbonate and sulphate ammonium salts. Asselin [34] reported that Fe (II) ammines are thermodynamically stable only under reducing conditions and concluded that they were unlikely to be formed if oxygen is present. He presented quasi-equilibrium Pourbaix diagrams for the Fe−NH3−H2O system. According to the diagram, Fe(OH)3 is the species present at noble potentials across all pH ranges, while Fe(OH)2 is present only above pH 11.
Many other processes have been used for the leaching of chalcopyrite and the chemistry of iron for solutions limited for the chalcopyrite as leaching process and other ones assisted leaching of chalcopyrite, as cited above [22][23][24][25][26][27][28][29][30][31][32][33][34]. At present, based on tests which were done in ammonium sulphate, carbonate, and perchlorate solutions, the leaching of chalcopyrite can be synthesized as follows:
  • The number of electrons (seven) transferred per mole copper during anodic oxidation is similar for the ammonia-ammonium sulphate and ammonia-ammonium carbonate solutions;
  • Ammonia-ammonium perchlorate solutions promote a five-electron transfer/copper reaction, possibly forming elemental sulfur on the mineral surface;
  • Ammonium sulphate leaching results in the formation of a Fe-oxyhydroxide layer with low sulfur on the mineral surface;
  • Ammonia-ammonium carbonate solutions resulted in marginal accumulation of iron on the mineral surface, but no formation of a layer was observed;
  • Ammonium perchlorate leaching results in the formation of a Fe-oxyhydroxide layer with moderate sulfur on the mineral surface;
  • The surface product was largely amorphous (90%) and significantly more porous (9–12 times) than unleached chalcopyrite. The observed morphology of the surface product suggests that it is formed through secondary precipitation rather than as part of the chalcopyrite dissolution mechanism;
  • Surface abrasion allows for the removal of the surface product, leading to improved leaching recoveries;
  • The abraded surface product from the small particles leaching experiment contained no sulfur, while surface products found on the stationary block of mineral contained small quantities of sulfur.
It has thus been shown that a relatively unstable iron product forms on the surface of chalcopyrite through secondary reactions to the faradaic oxidation reaction. This product may contain small percentages of sulfur but, regardless of the system, the majority of leached sulfur reports to the solution. Choice of ammonium salt and the hydrodynamic environment of leaching influence the presence or absence, as well as the nature, of the surface product. It appears that the formation of surface products in turn influences the reaction mechanism of chalcopyrite dissolution, and the two aspects need to be studied in conjunction. Lack of space does not allow further consideration of many works provided in the literature [35][36][37][38][39][40][41][42][43][44][45][46][47][48][49][50][51][52][53][54][55][56].
Nevertheless, it seems important to discuss, even briefly, future directions and opportunities in electrochemical separation and hydrometallurgy processes that can point to pathways for future implementation. One of these directions is defining consistent selectivity metrics. In fact, there has often been a lack of a consistent framework in reporting selectivity metrics in electrochemical separations, which can hamper the development of new metal recycling technologies. It can be seen that numerous metrics has been used (separation factor, selectivity factor, selectivity coefficient, selectivity rate), without a uniform approach; for instance, in the case of selective Li recovery, whereas some articles defined selectivity coefficient (KLi/M) without accounting for feed solution concentration, other studies defined the same term (selective coefficient) as the ratio of adsorbed Li to other metals divided by the ratio of amount of Li to other metal ions in the feed solution. Another direction is the speciation control. Considering chemical transformations within the liquid medium is critical for hydrometallurgical processes, especially complex aqueous or organic speciation of various metals. The dilute concentration of metals in waste streams makes it difficult to capture target ions selectively and efficiently. In this regard, speciation control can be leveraged for the effective separation and recovery from liquid streams; provided distinct chemical properties given by control of speciation offers a path for discrimination between metals with similar properties. Some ligands, such as cyanide and thiosulfate, bind strongly to metals and alter the electrochemical behavior of metals; their reduction potentials are controlled depending on their speciation and electrolytic condition (e.g., concentration of ligand, pH). For example, thermodynamic analysis has been carried out to identify potential levels where selective deposition of dilute concentrations of silver or gold from concentrated copper-containing solutions occurs—the potentials turned out to be shifted to cathodic direction as the concentration of thiosulfate increased [56], providing a way of modulating selectivity via speciation and potential control. In addition, some weakly binding ligands can exhibit the same speciation effect by having high concentration. These ligands, either strong or weak, indeed have been employed as lixiviants in conventional hydrometallurgical processes—therefore, the electrochemical recovery can take benefits from former leaching processes in relation to the speciation chemistry, and the judicious selection of appropriate lixiviants and electrochemical methods allows for a continuous, integrated approach for selective separations.
Recent advances in the synthesis of highly tailored redox interfaces offer a promised platform for the future development of electrochemical hydrometallurgy, especially for metals that have largely negative reduction potentials and thus are not easily electrodeposited. However, many studies reported so far feature-selective metal recovery in synthetic binary mixtures. Even though the studies in simple mixtures enable quick evaluation of new recovery methods and provide important information, these results can often be limited when inferring performance under realistic conditions—multicomponent, complex mixtures with a broader range of simultaneous competing targets. In parallel, the design of counter electrodes, types of electrical stimuli, optimizing electrochemical parameters, etc., can significantly improve the separation and energy efficiency. Combined efforts for engineering of heterogeneous electrode interfaces, electrolytes, fluid dynamic behavior, and operational parameters are needed.

References

  1. Eadington, P.; Prosser, A. Oxidation of lead sulfide in aqueous suspension. IMM Trans. Sect. C 1969, 78, 74–82.
  2. Pridmore, D.F.; Shuey, R.T. The electrical conductivity of galena, pyrite and chalcopyrite. Am. Mineral. 1976, 61, 248–259.
  3. Gurin, G.; Titov, K.; Ilyin, Y.; Tarasov, A. Induced polarization of disseminated electronically conductive minerals: A semi-empirical model. Geophys. J. Int. 2015, 200, 1555–1565.
  4. Dusabermariya, C.; Qian, W.; Bagaragaga, R.; Faruwa, A.; Ali, M. Some experiences of resistivity and induced polarization methods on the exploration of sulfide: A review. J. Geosci. Environ. Prot. 2020, 8, 68–92.
  5. Paul, R.L.; Nicol, M.J.; Diggle, J.W.; Saunders, A.P. The electrochemical behaviour of galena (lead sulphide)-1. Anodic dissolution. Electrochim. Acta 1978, 23, 625–633.
  6. Baba, A.A.; Adekola, F. Comparative analysis of the dissolution kinetics of galena in binary solutions of HCl/FeCl3 and HCl/H2O2. Int. J. Miner. Metall. Mater. 2011, 18, 1–9.
  7. Dutrizac, J.E. The dissolution of galena in ferric chloride media. Metall. Trans. B 1986, 17, 5–17.
  8. Fuerstenau, M.C.; Chen, C.C.; Han, K.M.; Palmer, B.R. Kinetics of galena dissolution in ferric chloride solutions. Metall. Trans. B 1986, 17, 415–423.
  9. Chen, A.A. Kinetics of Leaching Galena Concentrates with Ferric Fluosilicate Solution. Master Eng. Thesis, University of British Columbia, BC, Canada, 1992.
  10. Azizi, A. Gold cyanidation revisited-Kinetic & electrochemical studies of gold-sulfidic ore mixed/multilayer fixed beds. Ph.D. Thesis, University Laval, Quebec City, QC, Canada, 2011.
  11. Wilkominsky, I.; Rojas, N.; Balladares, E. Gold and silver cyanidation from a residue produced by leaching dead-roasted copper cohite metal. Can. Metall. Q. 2010, 49, 29–37.
  12. Medina, D.; Anderson, C.G. A review of cyanidation treatment of copper-gold ores and concentrates. Materials 2020, 10, 897.
  13. Azizi, A.; Ghardrahmati, R. Optimizing and evaluating the operational factors affecting the cyanide leaching circuit of the Aghdareh gold processing plant using a CCD model. Proc. R. Soc. A 2015, 471, 20150681.
  14. Birich, A.; Stopic, S.; Friedrich, B. Kinetic investigation and dissolution behaviour of cyanide alternative gold leaching. Sci. Rep. 2019, 9, 7191.
  15. Sabir, S. Silver Hydrometallurgy: Recovery and Recycling; Nova Publishers: Riyadh, Saudi Arabia, 2017.
  16. Crundwell, F.K. The dissolution and leaching of minerals. Mechanisms, myths and misunderstandings. Hydrometallurgy 2013, 139, 132–148.
  17. Kudryk, V.; Kellog, H.H. Mechanism and rate-controlling factors in the dissolution of gold in cyanide solution. JOM 1954, 6, 541–548.
  18. Holmes, P.R.; Crundwell, E.K. The kinetics of the oxidation of pyrite by ferric ions and dissolved oxygen: An electrochemical study. Geochim. Cosmochim. Acta 2000, 64, 263–274.
  19. Sequeira, C.A.C. High Temperature Corrosion: Fundamentals and Engineering; John Wiley & Sons: Hoboken, NJ, USA, 2019.
  20. Rimstidt, J.D.; Vaughan, D.J. Pyrite oxidation: A state-of-the-art assessment of the reaction mechanism. Geochim. Cosmochim. Acta 2003, 67, 873–880.
  21. Crundwell, F.K. Effect of iron impurity in zinc sulfide concentrate on the rate of dissolution. AICHE J. 1988, 34, 1128–1134.
  22. Lu, Z.Y.; Jeffrey, M.I.; Lawson, F. An electrochemical study of the effect of chloride ions on the dissolution of chalcopyrite in acidic solutions. Hydrometallurgy 2000, 56, 145–155.
  23. Hua, X.; Zheng, Y.; Xu, Q.; Lu, X.; Cheng, H.; Zou, X.; Song, Q.; Ning, Z.; Free, M.L. Leaching mechanism and electrochemical oxidation on the surface of chalcopyrite in ammonia-ammonium chloride solution. J. Electrochem. Soc. 2018, 165, E466–E476.
  24. Tanne, C.K.; Schippers, A. Electrochemical investigation of chalcopyrite (bio) leaching residues. Hydrometallurgy 2019, 187, 8–17.
  25. Moyo, T. An Electrochemical and Leach Study of the Oxidative Dissolution of Chalcopyrite in Ammoniacal Solutions. Ph.D. Thesis, University of Cape Town, Rondebosch, Cape Town, South Africa, 2016.
  26. Eghbalnia, M. Electrochemical and Raman Investigation of Pyrite and Chalcopyrite Oxidation. Ph.D. Thesis, University of British Columbia, Vancouver, Canada, 2012.
  27. Asgari, K.; Hassanzadeh, A.; Nazari, S.; Kakylabad, A.B.; Hosseinzadeh, M. Effect of externally adding pyrite and electrical current on galvanic leaching of chalcopyrite concentrate. Physicochem. Probl. Miner. Process. 2021, 57, 106–120.
  28. Liu, Q.; Chen, M.; Zheng, K.; Yang, Y.; Feng, X.; Li, H. In situ electrochemical investigation of pyrite assisted leaching of chalcopyrite. J. Electrochem. Soc. 2018, 165, H813–H819.
  29. Solis Marcial, O.J.; Nájera Bastida, A.; Banuelos, J.E.; Valdés Martinez, O.U.; Luevano, L.A.; Serrano Rosales, B. Chacopyrite leaching kinetics in the presence of methanol. Int. J. Chem. Reactor Eng. 2019, 17, 20190081.
  30. Arena, F.A.; Suegama, P.H.; Bevilaqua, D.; dos Santos, A.L.A.; Fugivara, C.S.; Benedetti, A.V. Simulating the main stages of chalcopyrite leaching and bioleaching in ferrous ions solution: An electrochemical impedance study with a modified carbon paste electrode. Miner. Eng. 2016, 92, 229–241.
  31. Peng, T.; Liao, W.; Wang, J.; Miao, J.; Peng, Y.; Gu, G.; Wu, X.; Qiu, G.; Zeng, W. Bioleaching and electrochemical behavior of chalcopyrite by a mixed culture at low temperature. Front. Microbiol. 2021, 12, 663757.
  32. Sequeira, C.A.C.; Santos, D.M.F.; Chen, Y.; Anastassakis, G. Chemical metathesis of chalcopyrite in acidic solutions. Hydrometallurgy 2008, 92, 135–140.
  33. Sequeira, C.A.C.; Santos, D.M.F. Transient film formation on chalcopyrite in acidic solutions. J. Appl. Electrochem. 2010, 40, 123–131.
  34. Asselin, E. Thermochemistry of the Fe, Ni and Co-NH3-H2O system as they relate to the Caron process: A review. Min. Metall. Process. 2011, 28, 169–175.
  35. Sequeira, C.A.C. Electrohydrometallurgical Recovery of Cadmium and Nickel from Spent Batteries. In Mineral Processing and the Environment; Gallios, G.P., Matis, K.A., Eds.; NATO ASI Series 2: Environment Kluwer; Academic Publishers: Dordrecht, The Netherlands, 1998; Volume 43, pp. 129–142.
  36. Brito, P.S.D.; Patricio, S.; Rodrigues, L.F.; Santos, D.M.F.; Sequeira, C.A.C. Electrodeposition of Zn-Mn alloys from recycling battery leach solutions in the presence of amines. In The Sustainable World-WIT Transactions on Ecology and the Environment; Brebbia, C.A., Ed.; WIT Press, Wessex, Institute of Technology: Southampton, UK, 2010; Volume 142, pp. 367–378.
  37. Sousa, N.R.; Borges, P.M.R.; Magueijo, V.M.; Brito, P.S.D.; Sequeira, C.A.C. Electrolytic reactors for the recovery of cadmium from leaching solutions. Key Eng. Mater. 2002, 230–232, 416–419.
  38. Rademan, J.A.M.; Lorenzen, L.; Van Deventer, J.S.J. The leaching characteristics of Ni-Cu matte in the acid-oxygen pressure leach process at Impala platinum. Hydrometallurgy 1999, 52, 231–252.
  39. Nikkhou, F.; Xia, F.; Yao, X.; Adegoks, I.A.; Gu, Q.; Kimpton, J.A. A flow-through reaction cell for studying minerals leaching by in-situ synchrotron powder X-ray diffraction. Minerals 2020, 10, 990.
  40. Sander, M.; Hofstetter, T.B.; Gorski, C.A. Electrochemical analyses of redox-active iron minerals: A review of nonmediated and mediated approaches. Environ. Sci. Technol. 2015, 49, 5862–5878.
  41. Gow, R.N.V. Spectroelectrochemistry and Modeling of Enargite (Cu3AsS4) Reactivity under Atmospheric Conditions. Ph.D. Thesis, University of Montana, Missoula, Butte, MT, USA, 2015.
  42. Yessengaziyev, A.; Kenzhaliyev, B.; Berkinbayeva, A.; Sharipov, R.; Suleimenov, E. Electrochemical extraction of Pb and Zn from a collective concentration using a sulfur-grafite electrode as a cathode. J. Chem. Technol. Metall. 2017, 52, 975–980.
  43. Pugaev, D.; Nicol, M.; Senanayake, G. The mechanisms of the passivation of sulfide minerals in oxidative leaching processes. In Proceedings of the 6th Southern African Base Metals Conference, Phalaborwa, South Africa, 18–20 July 2011; pp. 39–48.
  44. Moreno-Saldaña, S.I.; Martinez-Gómez, V.J.; Valle-Cervantes, S.; Lucho-Chigo, R.; Rojas-Montes, J.C.; Fuentes-Aceituno, J.C.; Pérez-Garibay, R. Analysis of galena leaching and maximum electrodeposition capacity of Pb using an electrochemical cell. JOM 2021, 73, 1353–1361.
  45. Chaerun, S.K.; Putri, E.A.; Mubarok, M.Z. Bioleaching of indonesian galena concentrate with an iron- and sulfur-oxidizing mixotrophic bacterium at room temperature. Front. Microbiol 2020, 11, 557548.
  46. Zhang, Z.; Liu, B.; Wu, M.; Sun, L. An electrochemical method to investigate the effects of compound composition on gold dissolution in thiosulfate solution. Green Proc. Synth. 2020, 9, 496–502.
  47. Sun, C.B.; Zhang, X.L.; Kou, J.; Xing, Y. A review of gold extraction using noncyanide lixiviants: Fundamentals, advancements, and challenges toward alkaline-sulfur containing leaching agents. Int. J. Miner. Metall. Mater. 2020, 27, 417–431.
  48. Sanchez-Ortiz, W.; Aldana-González, J.; Monh, T.L.; Romero-Romo, M.; Mejia-Caballero, I.; Ramirez-Silva, T.; Arce-Estrada, E.M.; Mugica Álvarez, V.; Palomar-Pardavé, M. A deep eutectic solvent as leaching agent and electrolytic bath for silver recovery from spent silver oxide batteries. J. Electrochem. Soc. 2021, 168, 016508.
  49. Reyes-Sandoval, E.; Fuentes-Aceituno, J.C. A study of the metallic silver dissolution with the MEA-NH3-Cu system. Rev. Matéria 2018, 23, e-12004.
  50. Tanne, C.; Schippers, A. Electrochemical investigation of microbially and galvanically leached chalcopyrite. Hydrometallurgy 2021, 202, 105603.
  51. Ahmed, M.; Hussein, I.A.; Onawole, A.T.; Saad, M.-A.; Khaled, M. Electrochemical removal of pyrite scale using green formulations. Sci. Rep. 2021, 11, 4796.
  52. Ma, Y.; Yang, Y.; Gao, X.; Fan, R.; Chen, M. The galvanic effect of pyrite enhanced (bio)leaching of enargite, Cu3 As S4. Hydrometallurgy 2021, 202, 105613.
  53. Lundstrom, M. Chalcopyrite Dissolution in Cupric Chloride Solutions; Helsinki University of Technology: Helsinki, Finland, 2009.
  54. Dizer, O.; Rogozhnikov, D.; Karimov, K.; Kuzas, E.; Suntsov, A. Nitric acid dissolution of tenantite, chalcopyrite and sphalerite in the presence of Fe (III) ions and FeS2. Materials 2022, 15, 1545.
  55. Tafoya-Medina, N.A.; Chuck-Hernandez, C.; Medina, D.I. Study of the electrooxidation of a zinc concentrate. Materials 2021, 14, 2868.
  56. Alonso, A.R.; Lapidus, G.T.; González, I. A strategy to determine the potential interval for selective silver electrodeposition from ammoniacal thiosulfate solutions. Hydrometallurgy 2007, 85, 144–153.
More
©Text is available under the terms and conditions of the Creative Commons-Attribution ShareAlike (CC BY-SA) license; additional terms may apply. By using this site, you agree to the Terms and Conditions and Privacy Policy.
Upload a video for this entry
Information
Subjects: Electrochemistry
Contributor MDPI registered users' name will be linked to their SciProfiles pages. To register with us, please refer to https://encyclopedia.pub/register :
César Sequeira
View Times: 665
Revisions: 4 times (View History)
Update Date: 08 Jul 2022
Table of Contents
    1000/1000
    Hot Most Recent
    Feedback