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Comparison of chalcocite dissolution in the sulfate, perchlorate, nitrate, chloride, ammonia, and cyanide systems
Minerals Engineering, Vol. 7, No. 1, pp. 99-103, 1994
0892-6875194 $6.00+0.00 © 1993 Pergamon Press Ltd
Printed in Great Britain
TECHNICAL NOTE COMPARISON OF CHALCOCITE DISSOLUTION IN THE SULFATE, PERCHLORATE, NITRATE, CHLORIDE, AMMONIA, AND CYANIDE SYSTEMS
W.W. FISHER Department of Metallurgical and Materials Engineering, University of Texas at El Paso, E1 Paso, Texas 79968, U.S.A. (Received 5 August 1993; accepted 9 August 1993)
ABSTRACT The dissolution of chalcocite in aqueous solution is dramatically influenced by the presence or absence of ligands that form stable cuprous complexes. Three types of behavior have been observed, bz systems that do not contain a ligand that forms strong cuprous complexes, the dissolution occurs in two steps, but the rate of the first step is very slow and is the same for each system. In systems that contain a ligand that forms strong cuprous complexes, the dissolution also occurs in two steps, but the rate of the first step is 60 to 90 times greater than in non-complexing systems and depends on the ligand. In the cyanide system, dissolution occurs in the absence o f oxidation at a rate that is 6 to 9 times that in the ammonia or chloride systems and more than 570 times that in noncomplexing systems.
Keywords Dissolution; chalcocite; sulfate; chloride; perchlorate; nitrate; ammonia; cyanide
INTRODUCTION The dissolution of chalcocite has been studied in a variety of chemical leaching systems as illustrated in Table 1. This summary does not include electrochemical (anodic dissolution) studies. Three of these studies [7, 12, 141 have been performed on the same mineral sample over the same range of particle size. Fisher, Henderson and Flores [7] investigated chalcocite dissolution in oxygenated, acid solutions containing sulfate and chloride [7] as the predominant anions. Aguayo [12] studied chalcocite dissolution in oxygenated, basic solution containing ammonium sulfate and ammonia [12]. Shantz and Fisher [14] examined chalcocite dissolution in the absence of oxygen or any other oxidant in basic solution containing cyanide. Each of these investigations used sized chacocite obtained from one large specimen of chalcocite, and therefore, the data from these three investigations can be used to directly compare the leaching response of a single chalcocite sample in several chemical systems.
EXPERIMENTAL The Chalcocite Sample The material used in each study came from a single piece of chalcocite donated by the Phelps Dodge Corporation from their New Cornelia Mine at Ajo, Arizona. The original sample was almost pure 99
100
Technical Note
chaleocite bounded on two or more sides by quartz with some visible quartz inclusions. The massive chalcoeite was broken into small piece.s with a hammer. Most of the quartz was rejected at this point by hand sorting. The small pieces of chalcocite were crushed and ground with an agate mortar and pestle. The ground material was sized with a series of Tyler screens to yield the size fractions -0.297 +0.210, 0.210 +0.177, -0.177 +0.149, -0.149 +0.105 and -0.105 +0.074 ram.
TABLE 1 Chemical systems used for chalcocite dissolution i
Aqueous System
Sample Type
Method
Source
Ferric sulfate
Natural, particles
Rolling bottles
1
Ferric sulfate
Synthetic, disk
Rotating disk
2
Ferric sulfate
Synthetic, disk
Rotating disk
3
Ferric sulfate
Synthetic, particles
Stirred reactor
4
Ferric chloride
Synthetic, particles
Stirred reactor
5
Sulfuric acid/oxygen
Natural, concentrate
Stirred reactor
6
Sulfuric acid/oxygen
Natural, particles
Stirred reactor
7
Perchloric acid/oxygen
Natural, particles
Shaking reactor
8
Sulfuric acid/oxygen
Synthetic, particles
Stirred reactor
9
Hydrochloric acid/oxygen
Natural, particles
Stirred reactor
7
Hydrochloric acid/chlorine
Natural, particles
Column percolation
10
Ammonia/oxygen
Natural, concentrate
Stirred reactor
11
Ammonia/oxygen
Natural, particles
Stirred reactor
12
Ammonia peroxy-disulfate/oxygen
Synthetic, disk
Rotating disk
13
Cyanide
Natural, particles
Stirred reactor
14
Pyridene hydrochloride/oxygen
Natural, particles
Stirred reactor
15
EDTA/oxygen
Synthetic, particles
Stirred reactor
16
EDTA/oxygen
Natural, particles
Stirred reactor
17
An x-ray diffraction analysis of the -0.297 +0.210 mm fraction from the work of Fisher, Henderson and Flores [7] identified the minerals present as chalcocite and quartz. Emission spectrographic analysis of the same fraction showed the major elements to be copper and silicon with iron, aluminum, calcium and silver in minor amounts and magnesium, manganese, arsenic, tungsten, cobalt and molybdenum in trace amounts. Chemical analysis of this size fraction showed the material to contain 71.52% copper. The amount of copper in the material used in the studies by Shantz and Fisher [14] and Aguayo [12] was respectively 76.29 and 78.60 percent. The difference in copper content between the three studies is due to the amount of insoluble residue which was identified as quartz. Since the leaching results are reported as percent extraction, the difference in initial copper content has no effect on the comparison of results.
Technical Note
101
Equipment All leaching experiments were carried out using 2.5 liters of solution and 25 g of sized chalcocite. A 2.5 liter, acrylic plastic reactor [7, 14] was used for all experiments except those in the ammonia/oxygen system. The ammonia/oxygen study was performed in a 1 gallon, stainless steel reactor using 2.5 liters of solution [12]. Details of the experimental procedures are given in these references.
RESULTS AND DISCUSSION Table 2 lists the conditions for the six chemical systems to be compared in this paper. For all of the systems, the temperature was 303 K, the oxygen or inert gas pressure was 0.086 MPa, the particle size was .-0.210+0.177 ram, and the percent solids in the slurry was 1.0%. All of the tests being compared were carried out at a stirring speed such that agitation was not a variable.
TABLE 2 Experimental conditions for different chemical systems
Chemical System
Ligand Conc., M
H + Cone., M
pH
Source
Sulfate/oxygen
0.176
0.10
- 1
7
Nitrate/oxygen
0.352
0.10
- 1
New
Perchlorate/oxygen
0.352
0.10
- 1
New
Chloride/oxygen
0.50
0.35
< 1
7
Ammonia/oxygen
0.87
9.5-10.5
12
Cyanide
0.33
12
14
Figure 1 compares the leaching curves for each dissolution system. The single curve for sulfate, nitrate, and perehlorate is an average of the results from each of the tests. The difference between corresponding sulfate, nitrate, and perchlorate data points is less than the size of the circle representing each data point. The curves in Figure 1 appear to fall into three groups. The curves for sulfate, nitrate, and perchlorate represent systems in which there is no ligand present to form complexes with cuprous ion. The dissolution rate is essentially the same for each of these non-complexing systems. When there is no ligand present to complex cuprous ion, the reaction occurs in two steps with synthetic covellite formed as an intermediate product by: Cu2S + J,~O2 + 2H + -. CuS + Cu +2 + H20
(1)
CuS + 202 -. Cu +2 + SO4"2
(2)
The first step is very slow, and therefore, very long leaching time.s or high temperature are required to reach the seeond stage. The ammonia and chloride curves represent systems in which there is a ligand present that strongly complexes cuprous ion. When a ligand such as chloride is present, copper is again extracted in two distinct stages, but the first stage reaction is given by [7]:
102
Technical Note
2Cu2S + tAO2 + 2H + + 6CI- + 2CuS + 2CUC13-2 + H20
(3)
The first stage reaction with ammonia as the complexing ligand is given by:
(4)
2Cu2S + tAO2 + 2H + + 4NH 3 + 2CuS + 2Cu(NH3)2+ + H20
The cuprous complex from Eqs. 3 or 4 may then be oxidized in solution to form the cupric complex. The second stage reaction in the chloride or ammonia system occurs by Eq. (2). The rate of dissolution in the first stage is independent of oxygen pressure in both systems [7, 12]. The difference in first stage rate between the chloride and ammonia systems illustrated in Figure 1 is due to the difference in stability of the two complexes and the ligand concentration. The rate in the second stage is independent of the ligand present, because the reaction is controlled by an electrocbemieal step [7, 12].
60
o
x 114 o
•--Sulfate,
1
2
Nitrate, Perchlorata
3 Time,
4
5
102 m i n u t e s
Fig. 1 Chalcocite dissolution in the sulfate, nitrate, perchlorate, chloride, ammonia, and cyanide systems. The cyanide curve represents a system in which the formation of the cuprous complex is energetic enough to extract cuprous ion from the chalcocite lattice without oxidation as shown by [14]: Cu2S + 8CN--, 2Cu(CN)4 -3 + S-2
(5)
To verify that oxidation is necessary for dissolution in non-cyanide systems, tests were run in the sulfate [7], chloride [7], and ammonia [12] systems with a nitrogen (inert) atmosphere in place of oxygen. A small amount of copper dissolves immediately, and then, no further dissolution occurs. The very limited initial dissolution is attributed to surface oxidation of the chalcocite particles.
Technical Note
103
CONCLUSIONS The dissolution of chalcocite in aqueous solution is dramatically influenced by the presence or absence of ligands that form stable cuprous complexes. Three types of behavior have been observed. In systems that do not contain a ligand that forms strong cuprous complexes, the dissolution occurs in two steps, but the rate of the first step is very slow and is the same for each system. In systems that contain a ligand that forms strong cuprous complexes, the dissolution also occurs in two steps, but the rate of the first step is 60 to 90 times greater than in non-eomplexing systems and depends on the ligand. In the cyanide system, dissolution occurs in the absence of oxidation at a rate that is 6 to 9 times that in the ammonia or chloride systems and more than 570 times that in non-eomplexing systems.
REFERENCES .
2.
.
4.
5.
°
7. 8. 9.
10. I1. 12. 13. 14. 15. 16. 17.
Sullivan, J.D., Chemistry of Leaching Chalcocite, U.S. Bur. Mines Tech. Paper, TP473, 24, (1930). Thomas, G., Ingraham, T.R. & MacDonald, R.J.C., Kinetics of Dissolution ofSyntbetic Digenite and Chaleoeite in Aqueous Acidic Ferric Sulphate Solutions, Can. Met. Quart., 6, 281-292 (1967). Mulak, W. & Niemiec, J., Kinetics of Cu2S Dissolution in Acidic Solutions of Ferric Sulphate, Rocz. Chem., 42, 775-781 (1968). Marcantonio, P., Kinetics of Dissolution of Chalcocite in Ferric Sulfate Solutions, Ph.D. Thesis, University of Utah, (1976). King, J.A., Burkin, A.R. & Ferreira, R.C.H., Leaching of Chalcocite by Acidic Ferric Chloride Solutions. In: A.R. Burkin (Editor), Leaching and Reduction in Hydrometallurgy, Inst. Min. Metall., London, 36-45 (1975). Warren, I.H., A Study of Pressure Leaching of Chalcopyrite, and Covellite, Aust. J. Appl. Sci., 9, 36-51 (1958). Fisher, W.W., Henderson, J.A. & Flores, F.A., Comparison of Chalcocite Dissolution in the Oxygenated, Aqueous Sulfate and Chloride Systems, Minerals Engineering, 5, 817-834 (1992). Peters, E. & Loewen, F., Pressure Leaching of Copper Minerals in Perchloric Acid Solutions, Met. Trans., 4, 5-14 (1973). Mao, M.H. & Peters, E., Acid Pressure Leaching of Chalcocite, In: K. Osseo-Asare and J.D. Miller (Editors), Research, Development and Plant Practice in Hydrometallurgy, AIME, New York, 243-260 (1983). Groves, R.D. & Smith, P.B., Reactions of Copper Sulfide Minerals with Chlorine in an Aqueous System, U.S. Bur. Mines Rept. Invest., RI7801, 10, (1973). Kuhn, M.C., Arbiter, N. & Kling, H., Anaconda's Arbiter Process for Copper, CIMBull., 67, 62-73 (1974). Aguayo, S., A Kinetic Study of Chalcocite Dissolution in the Low-Pressure Oxygen-Ammonia System, M.S. Thesis, University of Arizona, (1978). Mulak, W. & Niemiec, J., Kinetics of the Oxidation of Synthetic Chalcocite in Aqueous Ammonium Peroxydisulphate Solutions, Rocz. Chem., 42, 775-781 (1968) Shantz, R. & Fisher, W.W., The Kinetics of the Dissolution of Chalcocite in Alkaline Cyanide Solution, Met. Trans., 8B, 253-259 (1977). Pilarczyk, M., Strzelecki, H. & Grzybkowski, W., Recovery of Copper from Chalcocite by Means of the Pyridene and Chloride Ion Donor Systems, HydrometaUurgy, 11,247-251 (1983). Duda, L.L. & Bartecki, A., Dissolution of Cu2S in Aqueous EDTA Solutions Containing Oxygen, HydrometaUurgy, 8, 341-354 (1982). Konishi, Y., Kathoh, M. & Asai, S., Leaching Kinetics of Copper from Natural Chalcocite in Alkaline Na4EDTA Solutions, Met. Trans., 22B, 295-303 (1991).
0892-6875194 $6.00+0.00 © 1993 Pergamon Press Ltd
Printed in Great Britain
TECHNICAL NOTE COMPARISON OF CHALCOCITE DISSOLUTION IN THE SULFATE, PERCHLORATE, NITRATE, CHLORIDE, AMMONIA, AND CYANIDE SYSTEMS
W.W. FISHER Department of Metallurgical and Materials Engineering, University of Texas at El Paso, E1 Paso, Texas 79968, U.S.A. (Received 5 August 1993; accepted 9 August 1993)
ABSTRACT The dissolution of chalcocite in aqueous solution is dramatically influenced by the presence or absence of ligands that form stable cuprous complexes. Three types of behavior have been observed, bz systems that do not contain a ligand that forms strong cuprous complexes, the dissolution occurs in two steps, but the rate of the first step is very slow and is the same for each system. In systems that contain a ligand that forms strong cuprous complexes, the dissolution also occurs in two steps, but the rate of the first step is 60 to 90 times greater than in non-complexing systems and depends on the ligand. In the cyanide system, dissolution occurs in the absence o f oxidation at a rate that is 6 to 9 times that in the ammonia or chloride systems and more than 570 times that in noncomplexing systems.
Keywords Dissolution; chalcocite; sulfate; chloride; perchlorate; nitrate; ammonia; cyanide
INTRODUCTION The dissolution of chalcocite has been studied in a variety of chemical leaching systems as illustrated in Table 1. This summary does not include electrochemical (anodic dissolution) studies. Three of these studies [7, 12, 141 have been performed on the same mineral sample over the same range of particle size. Fisher, Henderson and Flores [7] investigated chalcocite dissolution in oxygenated, acid solutions containing sulfate and chloride [7] as the predominant anions. Aguayo [12] studied chalcocite dissolution in oxygenated, basic solution containing ammonium sulfate and ammonia [12]. Shantz and Fisher [14] examined chalcocite dissolution in the absence of oxygen or any other oxidant in basic solution containing cyanide. Each of these investigations used sized chacocite obtained from one large specimen of chalcocite, and therefore, the data from these three investigations can be used to directly compare the leaching response of a single chalcocite sample in several chemical systems.
EXPERIMENTAL The Chalcocite Sample The material used in each study came from a single piece of chalcocite donated by the Phelps Dodge Corporation from their New Cornelia Mine at Ajo, Arizona. The original sample was almost pure 99
100
Technical Note
chaleocite bounded on two or more sides by quartz with some visible quartz inclusions. The massive chalcoeite was broken into small piece.s with a hammer. Most of the quartz was rejected at this point by hand sorting. The small pieces of chalcocite were crushed and ground with an agate mortar and pestle. The ground material was sized with a series of Tyler screens to yield the size fractions -0.297 +0.210, 0.210 +0.177, -0.177 +0.149, -0.149 +0.105 and -0.105 +0.074 ram.
TABLE 1 Chemical systems used for chalcocite dissolution i
Aqueous System
Sample Type
Method
Source
Ferric sulfate
Natural, particles
Rolling bottles
1
Ferric sulfate
Synthetic, disk
Rotating disk
2
Ferric sulfate
Synthetic, disk
Rotating disk
3
Ferric sulfate
Synthetic, particles
Stirred reactor
4
Ferric chloride
Synthetic, particles
Stirred reactor
5
Sulfuric acid/oxygen
Natural, concentrate
Stirred reactor
6
Sulfuric acid/oxygen
Natural, particles
Stirred reactor
7
Perchloric acid/oxygen
Natural, particles
Shaking reactor
8
Sulfuric acid/oxygen
Synthetic, particles
Stirred reactor
9
Hydrochloric acid/oxygen
Natural, particles
Stirred reactor
7
Hydrochloric acid/chlorine
Natural, particles
Column percolation
10
Ammonia/oxygen
Natural, concentrate
Stirred reactor
11
Ammonia/oxygen
Natural, particles
Stirred reactor
12
Ammonia peroxy-disulfate/oxygen
Synthetic, disk
Rotating disk
13
Cyanide
Natural, particles
Stirred reactor
14
Pyridene hydrochloride/oxygen
Natural, particles
Stirred reactor
15
EDTA/oxygen
Synthetic, particles
Stirred reactor
16
EDTA/oxygen
Natural, particles
Stirred reactor
17
An x-ray diffraction analysis of the -0.297 +0.210 mm fraction from the work of Fisher, Henderson and Flores [7] identified the minerals present as chalcocite and quartz. Emission spectrographic analysis of the same fraction showed the major elements to be copper and silicon with iron, aluminum, calcium and silver in minor amounts and magnesium, manganese, arsenic, tungsten, cobalt and molybdenum in trace amounts. Chemical analysis of this size fraction showed the material to contain 71.52% copper. The amount of copper in the material used in the studies by Shantz and Fisher [14] and Aguayo [12] was respectively 76.29 and 78.60 percent. The difference in copper content between the three studies is due to the amount of insoluble residue which was identified as quartz. Since the leaching results are reported as percent extraction, the difference in initial copper content has no effect on the comparison of results.
Technical Note
101
Equipment All leaching experiments were carried out using 2.5 liters of solution and 25 g of sized chalcocite. A 2.5 liter, acrylic plastic reactor [7, 14] was used for all experiments except those in the ammonia/oxygen system. The ammonia/oxygen study was performed in a 1 gallon, stainless steel reactor using 2.5 liters of solution [12]. Details of the experimental procedures are given in these references.
RESULTS AND DISCUSSION Table 2 lists the conditions for the six chemical systems to be compared in this paper. For all of the systems, the temperature was 303 K, the oxygen or inert gas pressure was 0.086 MPa, the particle size was .-0.210+0.177 ram, and the percent solids in the slurry was 1.0%. All of the tests being compared were carried out at a stirring speed such that agitation was not a variable.
TABLE 2 Experimental conditions for different chemical systems
Chemical System
Ligand Conc., M
H + Cone., M
pH
Source
Sulfate/oxygen
0.176
0.10
- 1
7
Nitrate/oxygen
0.352
0.10
- 1
New
Perchlorate/oxygen
0.352
0.10
- 1
New
Chloride/oxygen
0.50
0.35
< 1
7
Ammonia/oxygen
0.87
9.5-10.5
12
Cyanide
0.33
12
14
Figure 1 compares the leaching curves for each dissolution system. The single curve for sulfate, nitrate, and perehlorate is an average of the results from each of the tests. The difference between corresponding sulfate, nitrate, and perchlorate data points is less than the size of the circle representing each data point. The curves in Figure 1 appear to fall into three groups. The curves for sulfate, nitrate, and perchlorate represent systems in which there is no ligand present to form complexes with cuprous ion. The dissolution rate is essentially the same for each of these non-complexing systems. When there is no ligand present to complex cuprous ion, the reaction occurs in two steps with synthetic covellite formed as an intermediate product by: Cu2S + J,~O2 + 2H + -. CuS + Cu +2 + H20
(1)
CuS + 202 -. Cu +2 + SO4"2
(2)
The first step is very slow, and therefore, very long leaching time.s or high temperature are required to reach the seeond stage. The ammonia and chloride curves represent systems in which there is a ligand present that strongly complexes cuprous ion. When a ligand such as chloride is present, copper is again extracted in two distinct stages, but the first stage reaction is given by [7]:
102
Technical Note
2Cu2S + tAO2 + 2H + + 6CI- + 2CuS + 2CUC13-2 + H20
(3)
The first stage reaction with ammonia as the complexing ligand is given by:
(4)
2Cu2S + tAO2 + 2H + + 4NH 3 + 2CuS + 2Cu(NH3)2+ + H20
The cuprous complex from Eqs. 3 or 4 may then be oxidized in solution to form the cupric complex. The second stage reaction in the chloride or ammonia system occurs by Eq. (2). The rate of dissolution in the first stage is independent of oxygen pressure in both systems [7, 12]. The difference in first stage rate between the chloride and ammonia systems illustrated in Figure 1 is due to the difference in stability of the two complexes and the ligand concentration. The rate in the second stage is independent of the ligand present, because the reaction is controlled by an electrocbemieal step [7, 12].
60
o
x 114 o
•--Sulfate,
1
2
Nitrate, Perchlorata
3 Time,
4
5
102 m i n u t e s
Fig. 1 Chalcocite dissolution in the sulfate, nitrate, perchlorate, chloride, ammonia, and cyanide systems. The cyanide curve represents a system in which the formation of the cuprous complex is energetic enough to extract cuprous ion from the chalcocite lattice without oxidation as shown by [14]: Cu2S + 8CN--, 2Cu(CN)4 -3 + S-2
(5)
To verify that oxidation is necessary for dissolution in non-cyanide systems, tests were run in the sulfate [7], chloride [7], and ammonia [12] systems with a nitrogen (inert) atmosphere in place of oxygen. A small amount of copper dissolves immediately, and then, no further dissolution occurs. The very limited initial dissolution is attributed to surface oxidation of the chalcocite particles.
Technical Note
103
CONCLUSIONS The dissolution of chalcocite in aqueous solution is dramatically influenced by the presence or absence of ligands that form stable cuprous complexes. Three types of behavior have been observed. In systems that do not contain a ligand that forms strong cuprous complexes, the dissolution occurs in two steps, but the rate of the first step is very slow and is the same for each system. In systems that contain a ligand that forms strong cuprous complexes, the dissolution also occurs in two steps, but the rate of the first step is 60 to 90 times greater than in non-eomplexing systems and depends on the ligand. In the cyanide system, dissolution occurs in the absence of oxidation at a rate that is 6 to 9 times that in the ammonia or chloride systems and more than 570 times that in non-eomplexing systems.
REFERENCES .
2.
.
4.
5.
°
7. 8. 9.
10. I1. 12. 13. 14. 15. 16. 17.
Sullivan, J.D., Chemistry of Leaching Chalcocite, U.S. Bur. Mines Tech. Paper, TP473, 24, (1930). Thomas, G., Ingraham, T.R. & MacDonald, R.J.C., Kinetics of Dissolution ofSyntbetic Digenite and Chaleoeite in Aqueous Acidic Ferric Sulphate Solutions, Can. Met. Quart., 6, 281-292 (1967). Mulak, W. & Niemiec, J., Kinetics of Cu2S Dissolution in Acidic Solutions of Ferric Sulphate, Rocz. Chem., 42, 775-781 (1968). Marcantonio, P., Kinetics of Dissolution of Chalcocite in Ferric Sulfate Solutions, Ph.D. Thesis, University of Utah, (1976). King, J.A., Burkin, A.R. & Ferreira, R.C.H., Leaching of Chalcocite by Acidic Ferric Chloride Solutions. In: A.R. Burkin (Editor), Leaching and Reduction in Hydrometallurgy, Inst. Min. Metall., London, 36-45 (1975). Warren, I.H., A Study of Pressure Leaching of Chalcopyrite, and Covellite, Aust. J. Appl. Sci., 9, 36-51 (1958). Fisher, W.W., Henderson, J.A. & Flores, F.A., Comparison of Chalcocite Dissolution in the Oxygenated, Aqueous Sulfate and Chloride Systems, Minerals Engineering, 5, 817-834 (1992). Peters, E. & Loewen, F., Pressure Leaching of Copper Minerals in Perchloric Acid Solutions, Met. Trans., 4, 5-14 (1973). Mao, M.H. & Peters, E., Acid Pressure Leaching of Chalcocite, In: K. Osseo-Asare and J.D. Miller (Editors), Research, Development and Plant Practice in Hydrometallurgy, AIME, New York, 243-260 (1983). Groves, R.D. & Smith, P.B., Reactions of Copper Sulfide Minerals with Chlorine in an Aqueous System, U.S. Bur. Mines Rept. Invest., RI7801, 10, (1973). Kuhn, M.C., Arbiter, N. & Kling, H., Anaconda's Arbiter Process for Copper, CIMBull., 67, 62-73 (1974). Aguayo, S., A Kinetic Study of Chalcocite Dissolution in the Low-Pressure Oxygen-Ammonia System, M.S. Thesis, University of Arizona, (1978). Mulak, W. & Niemiec, J., Kinetics of the Oxidation of Synthetic Chalcocite in Aqueous Ammonium Peroxydisulphate Solutions, Rocz. Chem., 42, 775-781 (1968) Shantz, R. & Fisher, W.W., The Kinetics of the Dissolution of Chalcocite in Alkaline Cyanide Solution, Met. Trans., 8B, 253-259 (1977). Pilarczyk, M., Strzelecki, H. & Grzybkowski, W., Recovery of Copper from Chalcocite by Means of the Pyridene and Chloride Ion Donor Systems, HydrometaUurgy, 11,247-251 (1983). Duda, L.L. & Bartecki, A., Dissolution of Cu2S in Aqueous EDTA Solutions Containing Oxygen, HydrometaUurgy, 8, 341-354 (1982). Konishi, Y., Kathoh, M. & Asai, S., Leaching Kinetics of Copper from Natural Chalcocite in Alkaline Na4EDTA Solutions, Met. Trans., 22B, 295-303 (1991).