Environmental Geochemistry of Sulfide Oxidation - American

Chapter 7

Downloaded by NORTH CAROLINA STATE UNIV on November 9, 2013 | http://pubs.acs.org Publication Date: December 20, 1993 | doi: 10.1021/bk-1994-0550.ch007

Solid-Phase Alteration and Iron Transformation in Column Bioleaching of a Complex Sulfide Ore 1

2

LasseAhonen and Olli H. Tuovinen 1

Geological Survey of Finland, SF-02150 Espoo, Finland Department of Microbiology, Ohio State University, 484 West 12th Avenue, Columbus, OH 43210-1292

2

The objective of the work was to characterize solid-phase changes and Fe(III) precipitation during biological leaching of a sulfide ore which contained chalcopyrite, pentlandite, pyrite, pyrrhotite, and sphalerite. The leaching experiments were carried out using bench-scale column reactors which were inoculated with acidophilic Fe- and S-oxidizing thiobacilli. Experimental factors included inoculation, pH, temperature, flood and trickle leaching, aeration, particle size, and mineralogical composition. Secondary solid phases, viz. covellite, jarosites, and elemental S, were detected in biologically active columns. Dissolved ferric iron data were pooled from all experiments and compared with solubility curves calculated for jarosites and ferric hydroxides. The data suggested that ferric-iron solubility was controlled by jarosites.

In chemical and biological leaching systems, dissolved ferric iron is an important redox component and is reduced to Fe by reaction with sulfide minerals (1,2). The reoxidation of ferrous iron is very slow under abiotic leaching conditions, unless strong oxidizing chemicals are used. Iron-oxidizing thiobacilli (Thiobacillus ferrooxidans), on the other hand, oxidize Fe -species at relatively fast rates in the range of pH 1-4 (3,4). Leptospinllum ferrooxidans, an iron-oxidizing chemolithotroph, is commonly found in consortia with Τ ferrooxidans and other acidophilic thiobacilli (5-7). By ferrous iron oxidation, bacteria help maintain high redox-potential values favorable to oxidative dissolution of sulfide minerals (8,9). Thiobacilli also can oxidize directly iron sulfides, resulting in elevated ferric iron concentrations in leach solutions (6,10). The oxidative-dissolution reactions of sulfide minerals may produce or consume acid. The bacterial oxidation of Fe-disulfide (e.g., pyrite) and hydrolysis of ferric iron are acid-producing reactions: 2+

2+

3+

2

4FeS + 150 + 2H 0 -> 4Fe + 8S0 " + 4H 2

2

2

+

4

0097-6156/94/0550-0079$06.00/0 © 1994 American Chemical Society In Environmental Geochemistry of Sulfide Oxidation; Alpers, C., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 1993.

(1)

80

ENVIRONMENTAL GEOCHEMISTRY OF SULFIDE OXIDATION 3+

Fe + 3H 0 -» Fe(OH) + 3H 2

+

(2)

3

Acid is consumed in the bacterial oxidation of ferrous iron: 2+

+

3+

4Fe + 0 + 4H -> 4Fe + 2H 0 2

(3)

2

The oxidation of Fe-monosulfide is also an acid-consuming reaction if the hydrolysis of ferric iron is excluded from the net reaction: +

3+

4FeS + 90 + 4H -> 4Fe + 4S0


2

2 4

+ 2H 0

(4)

2

Laboratory experiments have demonstrated that the oxidation of non-stoichiometric natural pyrrhotite is characterized by the formation of elemental sulfur (11): +

3+

4Fe _ S + (3-3x)0 + (12-12x)H -> (4-4x)Fe + 4S° + (6-6x)H 0 (1

x)

2

2

(5)

Contact of leach solutions with carbonate minerals as well as with some silicate phases causes acid consumption. The mineralogical composition of the ore and factors which influence the pH of the leach solution have major influences on ferric iron solubility. At pH range 1.5-3, precipitation of ferric iron in sulfate environments occurs largely in the form of jarosites. Because the Fe^/Fe * couple constitutes the major redox system in leach solutions, a decrease in the relative concentration of dissolved ferric iron also causes a decline in the redox potential and thereby in the oxidative dissolution of sulfide minerals (2,8,9). The dissolution of sulfide phases may lead to the precipitation of secondary minerals (e.g. jarosites, elemental S). The formation of such solid-phase reaction zones may adversely influence the contact of sulfide-mineral surfaces with leach solution and bacteria, replenishment of 0 at the reaction site, and fluxes of dissolution products (metals, sulfur species). The shrinking-core model describes this situation as a parabolic process, resulting in gradually declining leaching rates because of the lack of proper surface contact or because of an increasing distance between the solution front and the reactive surface (2,12,13). The formation of solid-phase reaction zones has not been characterized well in biological leaching systems, although there is circumstantial evidence for their formation. The scope of the work was to characterize changes in solid phases during the bacterial leaching of a complex sulfide ore and to examine transformations influencing the solubility of ferric iron in leach solutions. The data presented in this paper are a part of a program which was carried out to evaluate the feasibility of biological leaching techniques to recover metals from a complex sulfide ore. Most leaching experiments were relatively long-term, usually lasting in the range 400-700 days, and represented mostly conditions which were relevant to heap or in-situ leaching systems. Data previously published as a part of this research program were concerned with the effect of temperature on the bacterial leaching in column reactors (14) and the effect of Ag -addition in column bioleaching experiments in an attempt to stimulate chalcopyrite oxidation (15). Additional results on temperature effects, using this ore material in separate shake-flask and column bioleaching experiments, have also been published previously (9,16). 2

f

2

+

In Environmental Geochemistry of Sulfide Oxidation; Alpers, C., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 1993.

7.

Column Bioleaching of a Complex Sulfide Ore

AHONEN & TUOVINEN

Materials and Methods The ore material used in this study was obtained from a Cu-Co-Zn mine in Outokumpu, eastern Finland. The mineralogy of this deposit has been previously described (17,18). The main sulfide minerals were chalcopyrite (CuFeS ), sphalerite (ZnS), Co-pentlandite ((Ni,Co,Fe) S ), pyrite (FeS ), and pyrrhotite (Fe^S). Several different mixtures of the sulfide ore and gangue minerals were tested in the study. The gangue material was mostly carbonate- and graphite-containing quartzite. In addition, a mixture of skarn (containing diopside and tremolite) and serpentinite, all typical of the deposit, were used. Both pyrite-rich and pyrrhotite-rich samples were mixed with gangue material. The bacteria used in this work were originally enriched from mine-water samples and used as a mixed culture. The culture was capable of oxidizing ferrous iron and sulfur compounds as the sole source of energy under acidic conditions; therefore, the organisms were designated as acidophilic thiobacilli. Cultures were grown in a mineral salts solution (3.0 mM (NH ) S0 , 2.3 mM K HP0 , and 1.6 mM MgS0 - 7H 0 adjusted to initial pH 2.5 with H S0 ) which was supplemented with the sulfide-ore material as the sole substrate for energy. The experiments were carried out with glass-column reactors. The columns contained either 600-800 g or 12 kg of ore sample. The smaller columns (1:1 solid/liquid ratio) were 50 cm high with an internal diameter of 9 cm. In one column leaching experiment with a high grade ore sample, 400 g of ore material was used and the solid/liquid ratio was 1:2. The larger columns (4:1 solid/liquid ratio) were 100 cm high with an internal diameter of 11 cm. Two different small column types were used. In the first one, a fritted glass base plate supporting the ore was placed at the bottom of the percolator and the ore material was kept submerged in the leach solution (flood leaching). Oxygen supply was ensured by recirculating the leach solution with side-arm airlift. In the second type, the fritted glass base plate was lifted above the level of the leach solution. The solution passing through the ore sample was recirculated either with side-arm airlift or with a peristaltic pump (trickle leaching). After initial experiments, the column design was modified by inserting an open port below the fritted glass base to equalize pressure changes occurring during recirculation of the leach solution. The larger column type (12 kg ore) approximated trickle-leaching conditions: the ore sample was placed on a perforated base plate at the bottom of the column and the leach solution was recirculated with a peristaltic pump. The initial acid consumption was satisfied with sulfuric acid. Where necessary, sodium hydroxide was used to neutralize excessive acid production. Leach-solution samples removed for chemical analyses were replaced with equivalent volumes of sterile mineral-salts solution. Evaporation losses were compensated for by adding sterile, distilled water. The oxidative dissolution of metals from sulfide minerals in the leaching experiments was routinely monitored by atomic absorption spectroscopy (AAS). Sulfate formed in the oxidation was not determined because of excessive amounts of sulfuric acid added for pH control. 2


9

4

2

8

2

4

2

2

4

4


4

2

81

82

ENVIRONMENTAL GEOCHEMISTRY OF SULFIDE OXIDATION

Partial elemental composition of the ore samples was determined by AAS and by X-rayfluorescencespectrometry (XRF). Solid residues were either dissolved in a HC1HN0 -mixture and analyzed by AAS or they were analyzed directly by XRF. Total S in solid residues was determined by combustion and IR-detection (LECO IR 32H), elemental S by iodometric titration, and sulfate indirectly by AAS after precipitation of BaS0 . Mineralogical alterations were examined by X-ray diffraction (XRD) and ore microscopy in solid samples collected from column-leaching systems. The concentration of dissolved ferric iron in actual leach solutions was calculatedfromthe measured redox potential and total soluble iron data, on the basis of the Nernst equation. Only FeS0 , the predominant species in the pH range of the experiments, was accounted for in the calculation of the relative proportion of ferric iron. Figure 1 shows the theoretical distribution of the major aqueous species of ferric iron in acidic leach solutions. The speciation was calculated without charge balance using the PHREEQE program (19). The calculations were based on the equilibrium constants presented by Nordstrom et al. (20) and on the Debye-Huckel formula for the activity coefficient corrections. The calculations were carried out using equimolar concentrations of ferric iron and sulfate (0.1 M). The actual concentrations have a negligible influence on the distribution of the species. Consequently, changes in iron concentration due to precipitation were not considered. 3

4

+


4

Results and Discussion Evolution of H S was sometimes evident in the columns during pre-leaching stages, suggesting non-oxidative leaching of pyrrhotite: 2

+

2+

Fe . S + 2H -> (l-3x)Fe + 2xFe* + H S (1 x)

(6)

2

Microscopic examination of leach residues revealed the presence of covellite (CuS) on pyrrhotite surfaces (Figure 2). Covellite formation was favored at pH >2.5; Le. under conditions where ferric iron precipitated. At pH 10 g/liter) were dissolved during these time courses. In experiments conducted at pH 5 systems (23) and was, therefore, unlikely to exist under the experimental conditions used in the present study. Solubility product data are not available for schwertmannite. The published log values for ferric hydroxide-type precipitates range between -36.6 and -43.7. The lower log values are more characteristic for X-ray amorphous phases and the higher values are approaching those of crystalline Fe(m)-oxyhydroxides. For example, the log value for goethite is -43 (25). Thermodynamically, goethite formation is the preferred product in these leaching systems. However, goethite was not detected by XRD analysis, suggesting that jarosite formation was kinetically favored. Figure 5 shows dissolved Fe concentrations as a function of pH values of leach solutions pooled from the experimental results. Dissolved sulfate concentrations were not systematically analyzed because additional sulfate was introduced in the form of sulfuric acid which was used in pH adjustments. The total amount of sulfate added for the pH adjustment averaged approximately 0.5 M in concentration. The complete oxidative dissolution of sulfides in the ore material would produce an additional 0.5 M concentration of sulfate. Figure 5 also shows the solubility curves of ferric hydroxide in the range of log -37 to -43, based on a sulfate concentration of 0.1 M. Most experimental points are 0


3

4 2

6

+

+

+

+

+

4

+

+

+

3

+

4

8

6

4

5


8

2


7.


AHONEN & TUOVINEN

Figure 3. Polished section of corroded pyrrhotite (FeS) surface, surrounded by a dark rim of sulfur (S) and loosely associated Fe(III) precipitate (Fe ppt). Pegs of pentlandite ((Ni,Fe,Co)çS ) protrude from pyrrhotite, suggesting that pentlandite solubilization was slower than that of pyrrhotite. The sample was taken from a column bioleaching experiment (contact time 122 days). Bar, 50 um. Adapted from ref. 10. 8

Figure 4. Polished section showing chalcopyrite (CuFeS ) network within pyrite (FeS ) matrix. Pyrite is more disintegrated than chalcopyrite, and corrosion seems to propagate along phase boundaries. The sample was taken from a column bioleaching experiment after 90 days contact time. Bar, 50 um. 2

2


85

86

ENVIRONMENTAL GEOCHEMISTRY OF SULFIDE OXIDATION 40

within the Fe(HI) concentration and pH range defined by values of ΙΟ'^-ΙΟ* (Figure 5), indicating that ferric hydroxide precipitation may control iron concentrations at higher pH-values. The solubility curves for Na-, Κ-, H 0-, and NH -jarosites are shown in Figure 6. The jarosite/solution equilibria can be presented with the following equations (equilibrium constants derived from refs. 26-28): 3

+

4

+

NaFe (S0 ) (OH) + 6H * Na + 3Fe* + 2S0 3

4

2

6

2 4

+ 6H 0

(8)

2


log Κ = -5.3 +

2

KFe (S0 ) (OH) + 6H+ * K + 3Fe* + 2S0 ' + 6H 0 3

4

2

6

4

(9)

2

log Κ = -9.2 +

3

2

H OFe (S0 ) (OH) + 5H * 3FC * + 2S0 * + 7H 0 3

3

4

2

6

4

2

3+

2

(10) log Κ = -5.4

+

+

NH Fe (S0 ) (OH) + 6H * NH + 3Fe + 2S0 " + 6H 0 4

3

4

2

6

4

4

(11)

2

log Κ = -6.3 The equilibrium for Na-jarosite/ferric hydroxide can be defined as +

2

NaFe (S0 ) (OH) + 3H 0 * Na + 3Fe(OH) + 2S0 " + 3H 3

4

2

6

2

3

+

(12)

4

log Κ = -14.3 The equilibrium constant for this reaction was inferred from the solubility products of Na-jarosite (log = -5.3) and ferric hydroxide, 3

Fe(OH) * Fe * + 30H*

(13)

3

log ^

= -39

and the dissociation constant of water (log 1^ = -14.0). The equilibrium condition in dilute solutions (activity of H 0 = 1) is expressed in logarithmic form as 2

+

2

log[Na ] + 21og[S0 ] - 3pH = -14.3

(14)

4

+

2

Assuming activities [Na ] = 0.002 and [S0 ] = 0.01, it can be calculated that ferric hydroxide and Na-jarosite are in equilibrium at pH 2.5. This pH value shifts to higher pH values with increased concentrations of Na and S0 \ The equilibrium between goethite (log = -43) and Na-jarosite predicts goethite to be the stable phase at all positive pH values, in contrast to the XRD data of samples from leach columns. 4

+

2

4

Conclusions The initial mechanism in pyrrhotite leaching was concluded to involve non-oxidative dissolution, producing sulfide ion and HjS with subsequent oxidation to elemental S or precipitation as a secondary Cu-sulfide. Pentlandite was more recalcitrant than pyrrhotite. The leaching of pyrite was oxidative and strongly catalyzed by bacteria.



7.


AHONEN & TUOVINEN

ι

1

1

1

1

1

1

1

1.5

2.0

2.5

3.0

3.5

4

PH

Figure 5. Scatter diagram of dissolved ferric iron concentration as a function of pH in column leaching experiments. Superimposed on the Fe(III) plot are the solubility curves of ferric hydroxide calculated with the following log values: curve 7, -43; curve 2, -42; curve 3, -41; curve 4, -40; curve 5, -39; curve 6, -38;

curve 7, -37. Sulfate concentration was fixed to 0.1 M concentration for these calculations.

pH

Figure 6. Scatter diagram of dissolved ferric iron concentration as a function of pH in column leaching experiments, with superimposed solubility curves of jarosites. The calculations were based on 0.1 M S0 " and either 1 mM K (Kbearing jarosite, curve 7), 0.1 M Na (Na-jarosite, curve 2), 1 mM H 0 (H 0jarosite, curve 3), 1 mM NH (NH -jarosite, curve 4), or 1 mM Na (Na-jarosite, 2

+

4

+

+

3

+

4

+

4

curve 5).


3

87

ENVIRONMENTAL GEOCHEMISTRY OF SULFIDE OXIDATION

88

The concentration of dissolved iron was controlled by secondary precipitates. At pH >2.5, dissolved iron concentrations were low and X-ray amorphous or cryptocrystalline, brown Fe(III)-oxyhydroxide or Fe(IQ)-oxyhydroxysulfate precipitates were formed. At pH

Environmental Geochemistry of Sulfide Oxidation - American

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