Environmental Geochemistry of Sulfide Oxidation - American

Chapter 34

Suppression of Pyrite Oxidation Rate by Phosphate Addition 1,3

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Downloaded by UNIV OF CALIFORNIA SAN DIEGO on June 9, 2015 | http://pubs.acs.org Publication Date: December 20, 1993 | doi: 10.1021/bk-1994-0550.ch034

Xiao Huang and V. P. Evangelou 1

2

Department of Agronomy and Agricultural Experiment Station, University of Kentucky, N-122 Agricultural Science Center North, Lexington, KY 40546-0091 Pittsburgh Research Center, U.S. Bureau of Mines, Pittsburgh, PA 15236 3

Pyrite, commonly found in various ore deposits, produces highly acidic drainage water when exposed to the atmosphere. Current acidic drainage remediation technologies are not long lasting or cost effective. In this study we demonstrate that a ferric phosphate coating can form on pyrite surfaces when contacted with a solution of KH PO and H O . This ferric phosphate coating was shown to inhibit pyrite oxidation. 2

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Pyrite is a mineral commonly found in coals or other ore deposits formed in chemically reduced environments (1). Mining operations expose the overburden or ore containing pyrite to the atmosphere. Pyrite is also a waste product of ore processing plants. As a consequence of pyrite oxidation, the drainage from mining sites becomes highly acidic and enriched with sulfate, iron, manganese, and sometimes many other heavy metals (2). This acidic drainage finds its way into streams and lakes, causing a severe environmental pollution problem. In view of these facts, pyrite oxidation mechanisms and possible controls are of great interest to mining engineers, chemical engineers, and environmental scientists. Research on pyrite oxidation and its control has focused mainly on preventing oxidizing components from coming in contact with pyrite (3). For instance, Fe-oxidizing bacteria are considered to be responsible for the rapid oxidation rate of pyrite (4). This leads to the development and use of bactericides and slow-release bactericidal formulations (5). Other methods include sealing and insulation. The effectiveness of these techniques is site-specific and high cost precludes their use. Inspired by the phosphating technology widely used to treat the surface of steel for rust-proofing purposes (6,7), we conceived that a FeP0 coating could also be established on pyritic surfaces by treating them with a solution of phosphate and 4

0097-6156/94/0550-0562$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.

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Suppression of Pyrite Oxidation Rate

hydrogen peroxide. We hypothesized that H 0 can be used to oxidize pyritic surfaces, generating Fe *, so that insoluble FeP0 would form directly on the pyrite surfaces. Thus, at the expense of a certain fraction of pyrite, pyrite oxidation can be prevented. In this study, we examined the potential of forming a FeP0 coating on pyritic surfaces and then evaluated the influence of this coating on the kinetics of pyrite oxidation, employing a solution of H 0 . 2

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Methods Iron sulfide was separated from a shale by density separation using 97% tetrabromoethane (density of 2.97 g mL" ). The sulfide separate was washed with 4 M hydrofluoric acid (HF) and distilled water. The iron sulfide was then dried in a vacuumed desiccator and several properties relevant to this study were determined. X-ray diffraction was employed to establish that iron sulfide obtained from the shale was pyrite. It had a specific surface area of 7.15 m g* and contained 76% pyrite, as determined by dissolution of the sample with 30% H 0 . The impurity was expected to be iron oxyhydroxide, which was produced due to the exposure to air, and hydrolysis of Fe during washing with distilled water. We removed the iron oxyhydroxide impurity by leaching the pyrite sample with 0.1 M H Q before conducting oxidation experiments. Leaching-oxidation experiments were conducted at 40°C, employing a porous bed-reactor system. This porous bed-reactor system consisted of a chromatographic column (threaded chromaflex borosilicate glass column with acrylic water jacket and 20 micrometer polyethylene bed support). Fifty milligrams of pyrite were suspended in 500 milligrams of sand that had passed through a 140 mesh sieve. The mixture was placed on top of a styrene filter that was located on the bottom of the chromatographic column. An oxidizing solution was passed through the porous bedreactor, at a constant flow rate of 0.5 mL min' , employing a peristaltic pump. A water jacket was used to maintain a constant temperature. Aliquots were collected at certain time intervals using a Buchler Alpha 200 fraction collector. Sulfate (S0 ) was determined turbidometrically and iron was determined with atomic adsorption spectrophotometry.


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Results and Discussion We conducted three leaching experiments with the following solutions: 1) 0.5% H 0 in 0.2 M NaCl at pH 4; 2) 0.5% H 0 and 0.013 M EDTA in 0.2 M NaCl at pH 4; 3) 0.5% H 0 and 0.02 M KH P0 in 0.1 M NaCl at pH 4. Experiment 1 was designed to test maximum pyrite oxidation potential. Experiment 2 was designed to represent the situation where the influence of Fe * on kinetics of pyrite oxidation was eliminated through complexation of Fe * by EDTA; and experiment 3 was aimed at examining the potential of creating phosphate coatings and observing the influence of these coatings on the kinetics of pyrite oxidation. The oxidation of pyrite by hydrogen peroxide can be represented schematically 2

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

ENVIRONMENTAL GEOCHEMISTRY OF SULFIDE OXIDATION

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3

2

FeS + H 0 - » Fe * + S0

4

FeS \

2

2

2

2

J H 0

2

2

2+

+H

+

(1) 2

Fe + S0 " + H

+

4

Based on the mechanism proposed in equation 1, the pyrite oxidation rate is the sum of the direct oxidation of S " by H 0 and the oxidation of S ' by Fe *. Note that because the oxidation of Fe by H 0 is a rapid reaction, it is expected that at any time t, all the Fe released from pyrite oxidation is in the form of Fe *. Based also on the above, the rate law of oxidation can be expressed as: 2

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2+

3

dM - - ^ = OqfHA] + k^Fe *]) S 3

(2)

where M represents the number of moles of pyrite remaining in the system and S represents the surface area of pyrite attimet; k and k denote rate constants; [H OJ and [Fe *] represent concentrations of H 0 and Fe * at time t Examination of the pyrite sample by a scanning electron microscope revealed that the particles were relatively homogeneous in size. Thus, we assumed that surface area (S) at any time t was proportional to the number of moles (M) of pyrite remaining in the system during oxidation (8) and can be described by: t

2

3

2

3

2

2

S = Κ [M]

(3)

where Κ is a constant. Substituting equation 3 into equation 2 gives: dM 3

-

= fcfHA] + k^Fe *]) Κ [M]

(4)

By moving M to the left-hand side and integrating with respect to M, equation 4 can be rearranged to: d (In M) ~ ~j = ( K k J H A ] + Kk^Fe *]) 3

(5)

t

According to equation 5, if kinetic data of pyrite oxidation (S0 release) obtained by employing a constant concentration of H 0 were plotted as a first-order reaction, In (M/M ) vs. t (M=M at t=0) will give a curvilinear function. In leaching oxidation experiments, the concentration of Fe * was expected to decrease with time. Thus, a plot of In (M/M ) vs. t was expected to concave up (progressively less negative values of the slopes with increasing time). When oxidation of S " by Fe * 4

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was inhibited, a In (M/M ) vs. t plot was expected to be a straight line as kjtFe *] approaches zero. We hypothesized the following: if phosphate, introduced into the leachingoxidation system, reacted with all Fe * to form a discrete phase of FeP0 , the In (M/MJ vs. t plot would produce a straight line. Any further suppression of pyrite oxidation by phosphate beyond what was described by precipitation of Fe * as a discrete phase of FeP0 , would be attributed to the formation of a FeP0 coating established on the surface of pyrite. In this study, we used EDTA to complex Fe * by forming Fe-EDTA complexes, thereby preventing oxidation of pyrite by Fe * and the formation of Fe(OH) , which might also coat pyrite particles. Peck (9) indicated that the standard redox potential of Fe^/Fe * could be lowered from +0.77 to 0 by EDTA. Thus, we expected that pyrite oxidation by H20 in the presence of EDTA represented a situation where the observed rate of pyrite oxidation was solely due to the direct oxidation of S " by H 0 . The kinetic oxidation data from this treatment were expected to give a In (M/M ) vs. t plot with a single slope. Figure 1 shows that pyrite oxidation by H 0 was very rapid in the first 500 minutes and became slow from 500 to 1000 minutes. The first-order plot (In (M/M ) vs. t) shown in Figure 2 is a concave curve, demonstrating the influence of Fe * on the oxidation of pyrite. In the presence of EDTA, the oxidation of pyrite was suppressed. Thefirst-orderplot of the data representing the EDTA treatment shown in Figure 2 exhibits a straight line. This straight line indicates that the suppression was due to the prevention of the direct oxidation of S " by Fe * (Figure 3). As shown in Figure 1, phosphate suppressed pyrite oxidation to a much greater extent than EDTA. The slope of the In (M/M ) vs. t plot, representing the phosphate treatment is almost parallel to that representing the EDTA treatment in the initial 300 minutes of the oxidation process. This observation indicates that, over this period, phosphate played the same role as EDTA, i.e. to remove Fe * and to inhibit the direct oxidation of S " by Fe *. After 300 minutes,the rate of oxidation of pyrite in the presence of phosphate dropped rapidly, as evidenced by the change in slope of the plot around 300 minutes (Figures 1 and 2). The plot representing the phosphate treatment deviated from the straight line obtained with the data representing the EDTA treatment (Figure 2). This deviation resulted from the faster decrease in active surface area of pyrite or from an increase in surface coating coverage of pyrite. As shown in Figure 3, almost all Fe * produced during oxidation was precipitated by phosphate. These results strongly suggest that FeP0 was coating the pyrite surfaces, rather than simply precipitating as a discrete phase. To obtain direct evidence of a FeP0 surface coating, we separated residual pyrite particles from the sand-pyrite mixture, after oxidation and phosphatation were terminated, and examined the surfaces of the pyrite particles by scanning electron microscopy (SEM) and element-specific X-ray energy-dispersive analysis from the SEM (e.g. Figure 5). Figure 4A shows the morphology of residual pyrite particles after oxidation with 0.5% H 0 . Pyrite particles were coated with a thin layer of presumably amorphous iron (HI) hydroxide. As shown in Figure 3, slightly less than 0

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Time (minutes)

Figure 1. Release of S0 during pyrite oxidation under the following conditions: (1) 0.5% H 0 ; (2) 0.5% H 0 in the presence of 0.013 M EDTA; and (3) 0.5% H 0 in the presence of 0.02 M ΚΗ^Ο . 4

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Λ

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g> c

Phosphate

-0.1

ε Φ £ ο c ο

-0.2

Β

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œ ο ε •δ ε

Χ.

EDTA

-0.4

8>

Control

-0.5 0

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Time (minutes)

Figure 2. First-order plots of pyrite oxidation under the following conditions: (1) 0.5% H 0 ; (2) 0.5% H 0 in the presence of 0.013 M EDTA; and (3) 0.5% H 0 in the presence of 0.02 M KH P0 . 2

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Suppression of Pyrite Oxidation Rate 567

HUANG & EVANGELOU

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Δ Actual values


CO

ω ο ε ε

.o.o-0-o-o-o-o-o

ο Theoretical values 0.2

Control Φ Φ

+

^ ^ Α - Δ - Δ - Δ - Δ - Δ - Δ - Δ - Δ -- Δ - Δ - Δ - Δ - Δ - Δ - Δ - Δ

t/ 0.1

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Δ Actual values EDTA

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Phosphate \

fi*

D

Δ-Δ-Δ-Δ-Δ^4 100 200 300

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Time (minutes)

Figure 3. Release of Fe during pyrite oxidation under the following conditions: (1) 0.5% H 0 ; (2) 0.5% H 0 in the presence of 0.013 M EDTA; and (3) 0.5% H 0 in the presence of 0.02 M K H ^ Q i (actual values are those determined by measuring Fe in the leachate; theoretical values are those calculated with S0 " data according to the stoichiometry). 2

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half of the Fe produced during oxidation was not released to solution and presumably ended up as Fe(OH) . Comparison of the kinetic data from experiments 1 and 2 (Figure 1) suggests that the iron(HI) hydroxide formed during oxidation did not inhibit pyrite oxidation. Actually, it might have accelerated pyrite oxidation due to its function as a reservoir of Fe *, because the localized high acidity on the surface of pyrite during oxidation would prevent the deposition of Fe(OH) . Thus, most of the Fe(OH) would have formed as a discrete phase. Figure 4B shows the residual pyrite particles displayed a morphology typical of framboidal pyrite (15). These particles consist of small pyrite crystals with easily identified octahedrons. The surfaces of these particles were free of coatings. The holes observed on the octahedrons indicate the locations where oxidation took place. The absence of any coating on the surfaces of pyrite shown in Figure 4B is probably due to removal of all Fe released by pyrite oxidation as Fe-EDTA complexes (Figure 3), which prevents the formation of iron hydroxide precipitates. Figure 4C shows the morphology of pyrite particles oxidized in the presence of phosphate. In this photograph the surfaces of the framboidal pyrite particles are heavily coated. We conducted X-ray scanning of these coated pyrite particles to examine the distribution of Fe, S, and P. The distribution of Ρ was similar to the distribution of Fe and S but the intensity of Fe was much higher than that of S and Ρ (Figure 5). The higher density of Fe is due to the presence of FeP0 and FeS , both of which contain Fe. The above results further suggest that the differences in oxidation rates between the EDTA and phosphate treatments resulted from the formation of an FeP0 coating. In order to further understand the chemical properties of the iron phosphate coating, we repeated the leaching experiment with a set of columns of pure pyrite and solutions containing 0.147 M H 0 and 0.01 M KH P0 . At the end of this experiment, each column was leached with 50 mL of 2 M HC1. The leachate was analyzed for iron and phosphate. We found that the mole ratio of iron to phosphate was 1.0, indicating that the coating was most likely amorphous FeP0 . 3

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Conclusions Pyrite particles can be coated with a protective coating by treatment with a mixed solution of KH P0 and H 0 , at the expense of a certain fraction of pyrite. The data in Figure 6 show that phosphate coatings established with a mixed solution of ΙΟ M KH P0 and 0.5% H 0 consumed 20% of the pyrite sample. We believed that this consumed fraction can be decreased by decreasing the concentration of H 0 . To test the stability of the coatings, we exposed pyrite with coatings to 0.5% H 0 in the absence and presence of phosphate. As shown in Figure 6 (curve B), the pyrite sample with phosphate coating gradually decomposed, suggesting the partial collapse of the coating with time. The collapse of the coating resulted from the dissolution of FeP0 by the strong acid produced by oxidation of exposed pyritic surfaces. However, when compared with the oxidation of pyrite with no coating (curve A), the suppression of oxidation due to the coating was significant. Moreover, FeP0 dissolution can be inhibited in the presence of phosphate 2

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HUANG & EVANGELOU



34.

Figure 4. Scanning electron microscope (SEM) photos of residual framboidal pyrite particles after oxidation experiments: (A) pyrite particle oxidized with 0.5% H 0 ; (B) and (C) pyrite particles oxidized with 0.5% H 0 in the 2

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presence of 0.013 M EDTA. Continued on next page.


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Figure 4. Continued. Scanning electron microscope (SEM) photos of residual framboidal pyrite particles after oxidation experiments: (D) and (E) pyrite particles oxidized with 0.5% H 0 in the presence of 0.02 M KH P0 . 2

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Figure 5. Element-specific X-ray energy-dispersive analysis (from the SEM) of pyrite particles oxidized in the presence of phosphate, showing the distribution of Fe, S, and P: (A) SEM microphoto of the particle examined; B) distribution of Fe; (C) distribution of S; and (D) distribution of P.


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Time (minutes) Figure 6. Pyrite oxidation as a function of time. Before 1000 min represents coating with a mixed solution of ΙΟ M KH P0 and 0.1 % H 0 . After 1000 min, curve A represents pyrite oxidation with coating removed by leaching with 50 mL of 2 M HC1 and exposed to 0.25% H 0 ; curve Β represents the oxidation of FeP0 -coated pyrite exposed to 0.25% H 0 in the absence of phosphate solution; and curve C represents oxidation of FeP0 -coated pyrite exposed to 0.25% H 0 in the presence of ΙΟ M KH P0 . 3

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concentration as low as ΙΟ" M (curve C in Figure 6), which decreased the solubility of iron phosphate (10). Thus, when pyritic ores are to be exposed to strong oxidizers, the stability of phosphate coating will need to be strengthened by a small concentration of phosphate in solution. In summary, pyritic surfaces can be coated with FeP0 , which can prevent pyrite oxidation. However, when coated pyrite is exposed to strong oxidizers, the coating can be stabilized by maintaining 10" M KH P0 in solution. 4

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Acknowledgments The investigation in the paper (91-3-179) is in connection with a project of the Agric. Exp. Stan, of the University of Kentucky and is published with the approval of the director. Also, the authors wish to thank Mr. R.W. Hammack, and Dr. R. Kleinmann of the U.S. Department of the Interior, Bureau of Mines, Pittsburgh, PA, and Ms. Patricia M. Erikson of the U.S. EPA, Cincinnati, OH, for the enlightening discussions on pyrite oxidation, and Ms. Libby Reed for typing and proofing the manuscript Literature Cited 1. 2. 3. 4. 5.

7. 8. 9. 10.

Lowson, R.T. Chem. Rev. 1982, 82, 461-493. Krothe, N.C.; Edkins, J.E.; Schubert, J.P. In Proceedings of 1980 Symposium on Surface Hydrology, Sedimentology, and Reclamation; University of Kentucky: Lexington, KY, 1980; pp. 455-564. Singer, P.C.; Stumm, W. Science 1970, 167, 1121-1123. Temple, K.L.; Delchamps, E.W. Appl. Microbio. 1953, 1, 255-258. Kleinmann, R.L.P. In Proceedings of 1980 Symposium on Surface Hydrology. Sedimentology. and Reclamation: University of Kentucky: Lexington, KY, 1980; pp. 333-337. Philips, D. Plating and Surface Finishing 1990, 77, 31-35. Gorecki, G. Metal Finishing 1988, 86, 15-16. Peck, H.P.J. An. Rev. Microbiol. 1968, 22, 489-518. Lindsay, W. Chemical Equilibrium in Soils: J. Wiley and Sons: New York, NY, 1979.

RECEIVED October 13,

1993


Environmental Geochemistry of Sulfide Oxidation - American

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