Environ. Sci. Technol. 1996, 30, 2604-2608
Fundamental Study of an Ambient Temperature Ferrite Process in the Treatment of Acid Mine Drainage WEIXING WANG,† ZHENGHE XU,* AND J. FINCH Department of Mining & Metallurgical Engineering, McGill University, 3450 University Street, Montreal, Quebec, Canada H3A 2A7
The feasibility of making ferrite at ambient temperature from solutions of simulated acid mine drainage is explored. The X-ray diffraction and saturation magnetization of the precipitates obtained clearly show that the spinel ferrite can be made at ambient conditions and pH 9-11. The optimum molar ratio of Fe3+ to Fe2+, a critical parameter for ferrite formation, is found to be 1.75. The effect of [Zn2+] and [Ca2+] in solution on ferrite formation was investigated. The former has no effect on the magnetic properties of ferrite, but the saturation magnetization of the ferrite decreases with an increase of [Ca2+]. The results suggest that an ambient temperature ferrite process for acid mine drainage is worthy of further investigation.
Introduction Acid mine drainage (AMD) is one of the most severe environmental problems facing the mining industry today. In Canada, for example, many of the waterways around mining sites are contaminated, often from mines that have been closed for decades. Two provinces with the most extensive metal mining base, Ontario and Quebec, have over 2000 abandoned mine sites, and many of them do or will generate an acid drainage (1). AMD can be loosely defined as the ground surface water draining from a mining site where sulfide minerals are (or were) processed. When sulfide minerals, particularly pyrite and pyrrhotite, are exposed to oxygen in water, they oxidize to produce sulfuric acid. The acid, in turn, leaches other minerals. As a result, AMD contains a variety of metal contaminants depending on the site (2). The common practice to control AMD involves metal precipitation as hydroxides and/or sulfide. These processes often generate secondary hazardous wastes (such as metal hydroxide sludge) that require highly regulated and costly disposal. Several attempts have been made to precipitate the heavy metals as ferrites (3-6). The term ferrite often refers to * Corresponding author fax: 514-398-4492; e-mail address:
[email protected]. † On leave from Wuhan Yejin University of Science & Technology, People’s Republic of China.
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FIGURE 1. Phase diagram of iron-water system indicating oxidation conditions for the formation of FeFe2O4 (13): [FeSO4] is kept constant at 0.24 (mol/L); solid squares represent a co-existing region of Fe3O4/ r-FeOOH/γ-FeOOH phases.
magnetic oxides containing other metals in addition to iron. In the present paper, ferrites mean the spinel ferrite that can show ferro- or ferrimagnetic behavior. The general chemical formula of ferrites possessing the structure of the mineral spinel (MgAl2O4) is MFe2O4, where M is any divalent ion with an unhydrated ion radius between approximately 0.6 and 1 Å (7). (A well-known example is iron ferrite, or magnetite, where iron is the only metallic constitute.) Ferrite formation in systems with relatively simple solution chemistry (compared to AMD) at temperatures above 60 °C over a given pH range is well documented (8, 9). The major advantages of using the ferrite process to treat AMD are its ability to “scavenge” most divalent metal ions; and the precipitates formed are stable and readily recovered by magnetic filtration (10, 11). An additional incentive to pursue the process is the potential commercial use of ferrite as recording media, magnetic markers, and steel making feeds. Recording media, for example, are made from a variety of magnetic materials, including iron oxides (γ-Fe2O3, Fe3O4), cobalt-doped iron oxides, barium ferrite, and chromium dioxide (12). Ferrite obtained from wastewater treatment could become a source of these materials. In spite of these attractive features, little has been published on the ferrite process applied to the treatment of large volumes of wastewater such as AMD. This may be related to the apparent need for temperatures above ambient (13). Temperature is a key variable in the formation of ferrite. Based on the phase diagram (Figure 1), R-FeOOH and γ-FeOOH are the major precipitates at reaction temperatures below 40 °C. This figure has been used to explain why ferrite does not form at room temperature. It is, however, important to note that this phase diagram was constructed based on air oxidation of a Fe(OH)2 suspension in the absence of ferric species. In practice, most AMDs contain both ferrous and ferric ions. Therefore, this phase diagram may not be suitable to predict the dominant products of precipitation from AMD. Recently, a scheme of ferric/ferrous co-precipitation followed by phase conversion at elevated temperature was
S0013-936X(96)00006-5 CCC: $12.00
1996 American Chemical Society
TABLE 1
Compositions of Metallic Elements in Original AMD from Les Mines Gallen, Leachate of Magnetic Precipitates, and Treated Effluent of Simulated AMD element (ppm)
Fe
Zn
Al
Mg
Cu Mn
704 558 38 original 5880a 2760 leaching 1.8 1.0 0.5 - residual 0.2 25(ppb) 0.9 - a Of 5880 ppm, 3295 ppm are ferric ions. examined.
b
29 -
Ca
Pb
1.0 0.4 NDb ND ND 0.1
ND, not detected. c -, not
proposed to treat AMD (14). In this process, only the precipitates (not the whole AMD) require heating. The challenge posed with this process is the inherently poor solid/liquid separation associated with fine precipitates. The use of co-precipitation was also proposed to treat AMDs from Quebec, Canada, and from Butte, MT, but the need for ferrous ion addition with a controlled oxidation environment and a long reaction time (2 days) raised economic concerns.
FIGURE 2. Effect of Fe3+/Fe2+ molar ratio on the recovery of iron as ferrite (reaction time, 15 min) from iron sulfate solution and the stability measured by weight loss of ferrite during leaching with HCl solution: the initial [Fe2+] is kept constant at 0.01 (mol/L); pH is 2.5, leaching time is 24 h.
The objective of the present work is to examine the possibility of forming ferrite at ambient temperature within a reasonable time frame by controlling the chemistry of AMD. A fundamental study has been conducted using iron sulfate solutions and simulated AMD.
Experimental Section Materials. The chemicals used [Fe2(SO4)3, Al2(SO4)3‚18H2O, FeSO4‚7H2O, ZnSO4, MgSO4, CaCl2‚2H2O, CuSO4‚5H2O, MnCl2, PbSO4, H2SO4, NaOH] were analytical grade. The simulated AMD was prepared to give a similar chemical composition to the drainage from Les Mines Gallen, Quebec, shown in Table 1. Double-distilled water was used unless otherwise stated. Procedure. The experiment was carried out in a 500mL beaker in which the solution was stirred using a mechanical stirrer. A digital pH meter was used to monitor solution pH. The initial pH was set at 2.5. After transferring 500 mL of ferric ion solution or simulated AMD to the beaker, a desired amount of FeSO4‚7H2O was added. The pH was then raised to between 9 and 11 within 15 min by adding NaOH (5 N or 1 N); a fine precipitate was observed to form immediately. The suspension was then transferred to a 500-mL stoppered measuring cylinder, and the precipitates were allowed to settle for 70 min. The magnetic fraction of the precipitates was separated from the nonmagnetic portion using a hand magnet (0.1 T). The magnetic and nonmagnetic portions were filtered separately and weighed after drying in a vacuum oven at room temperature. The recovery of iron as a magnetic phase was calculated using
RFe )
WMP × 100% WMT
where WMP and WMT are, respectively, the weight of the magnetic portion and the theoretical weight of magnetite that could form. The magnetic portion of the precipitates produced from simulated AMD was exposed to hydrochloric acid solution (pH 2.3) for 24 h to examine stability, judged by weight loss in leachate.
FIGURE 3. Effect of ratio Fe3+/Fe2+ on the saturation magnetization of magnetic precipitates produced from iron sulfate solution.
The metal ion concentration in solution was determined using atomic absorption spectroscopy (Perkin Elmer 3110). Representative products were examined by X-ray powder diffractometer (Philips BW 1710) using Cu-KR radiation. The magnetic properties of the ferrites were determined using a vibrating-sample magnetometer (15). The particle size distribution was measured using a laser-particle sizer (Fritsch Analysette 22). All the experiments were carried out at room temperature unless otherwise stated.
Results Ferrite Formation. Figure 2 shows that (a) magnetic precipitates form at ambient temperature within the reaction time of 15 min and (b) the recovery of iron as magnetic precipitates depends largely on the [Fe3+]/[Fe2+] molar ratio. In the absence of ferric ions, the recovery of iron as magnetic precipitates is low (26.2%). With a molar ratio of 0.5, the recovery increased to ca. 46%. The highest recovery, ca. 96%, was obtained at a ratio of 2, which agrees with the stoichiometry of iron ferrite (FeFe2O4). The saturation magnetization as a function of the molar ratio is given in Figure 3. This figure shows that, up to a ratio of 1.5, saturation magnetization increased to a value similar to that for pure iron ferrite, suggesting that the precipitates produced under these conditions are mainly FeFe2O4. Further increase in molar ratio above 2 resulted in a reduction in both recovery of iron as magnetic precipitates (see Figure 2) and the saturation magnetization. This may be attributed to the partial entrainment of Fe-
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FIGURE 4. Powder X-ray diffraction patterns of magnetic precipitates formed in iron (Fe3+/Fe2+ ) 1.5) solution and simulated acid mine drainage: vertical lines represent the diffraction pattern of standard spinel ferrite.
(OH)3 precipitates formed from excess ferric ions. For a ratio below 1.5, some Fe(OH)2 may be incorporated into ferrite as only a limited amount of ferrous ions can be accommodated in FeFe2O4 unless the appropriate oxidation environment is provided. Theoretically, the maximum saturation magnetization should occur at a ratio of 2.0 (as indicated in the formula of iron ferrite, FeFe2O4) not at 1.5-1.75 as determined experimentally. The discrepancy suggests that partial oxidation of Fe2+ to Fe 3+ may be necessary during the conversion of the intermediate product to ferrite. To confirm the formation of ferrites, the magnetic precipitates were characterized using X-ray powder diffraction. The diffraction pattern obtained at a [Fe3+]/[Fe2+] of 1.5 is shown in Figure 4 along with that for standard spinel ferrite (vertical lines). The excellent agreement between the two confirms that the major component of the magnetic precipitate is ferrite. The leach test on the magnetic precipitates showed (Figure 2) that the weight loss is relatively high and constant at [Fe3+]/[Fe2+] below 0.75 and then decreased with ratios above 0.75. The low weight losses at ratios between 1.5 and 1.75 (ca. 2%) confirm that the ferrite formed under these conditions is stable. The above results show that spinel ferrite was formed at ambient temperature when [Fe3+]/[Fe2+] was controlled. The iron recovery as ferrite was optimized at a ratio of 2 and pH above 9.5. Effect of Zinc and Calcium Ions on Ferrite Formation. Zinc and calcium ions are often present in AMD. Their effect on ferrite formation, therefore, needs to be examined. The precipitation in the presence of zinc ions showed that at a [Fe3+]/[Fe2+ ]/[Zn2+] molar ratio of 2/0.8/0.2 (equivalent to [M3+]/[M2+] of 2/1) total metal recovery as magnetic precipitates was ca. 92%, and the saturation magnetization was 60 emu/g. Compared to the case in the absence of zinc, these values are smaller than for the precipitates formed at [Fe3+]/[Fe2+] of 2 but are similar to those obtained with a ratio of 2.25. This finding suggests that zinc ions replace ferrous ions in the lattice, producing some zinc ferrite, (Zn, Fe)Fe2O4. The leach test showed that the weight loss of the zinc-containing ferrite was 6.3%, which is comparable to that in the absence of zinc. Overall, it is demonstrated that zinc ions do not prevent ferrite formation. The presence of calcium, on the other hand, was found to have a significant effect on both metal recovery as magnetic precipitates and the magnetic property of the
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FIGURE 5. Effect of calcium concentration in effluent on the metal recovery as and the stability measured by weight loss of ferrite.
ferrite. The experiments conducted at [Fe3+]/[Fe2+] of 2 with increasing concentration of calcium ions from 100 to 400 ppm (equivalent to molar ratio increasing from 2/1/ 0.13 to 2/1/0.53) showed that the metal recovery (calculated based on the total metal ions, including calcium) decreased (Figure 5). The saturation magnetization decreased with increasing [Ca2+ ], from 78 emu/g in the absence of calcium to 50 emu/g at 100 ppm and down to 30 emu/g at 400 ppm. The residual calcium concentration was ca. 25 and 106 ppm for the initial [Ca2+] of 100 and 400 ppm, respectively. About 75% of the added calcium ions are therefore incorporated in the precipitates. The reduced magnetization in the presence of calcium may be attributed to the inclusion of nonmagnetic CaCO3 or CaSO4 precipitates in the ferrite. Interestingly, only a fraction of calcium in the precipitates was leachable although the weight loss increased with [Ca2+], as shown in Figure 5, to values greater than in the absence of calcium. The results clearly demonstrate that the presence of calcium has an adverse effect on the ferrite process, reducing metal recovery, magnetization, and stability of the ferrite. This finding may be one reason why ferrite is not reported in conventional AMD treatment sludge produced using calcium hydroxide (lime). A detailed study of the calcium effect on ferrite formation is in progress. Simulated AMD. The chemical composition given in Table 1 shows a ∑M3+/∑M2+ ratio of 0.75, which is lower than required for the optimum ferrite formation. Indeed, the precipitates formed upon raising the pH above 10 were nonmagnetic, as expected. If ferric ions were added to adjust the ratio to 2, the [Fe3+]/[Fe2+] will be greater than 4, which is not suited to ferrite formation. An oxidant (H2O2) was therefore used to convert some Fe2+ to Fe3+. The amount of H2O2 was determined based on reaction 1 in order to obtain a ∑M3+/∑M2+ ratio of 2.0:
2Fe(OH)2 + H2O2 ) 2Fe(OH)3
(1)
The precipitates formed under these conditions were also nonmagnetic. Consequently, an additional amount of ferrous salt was added to obtain a [Fe3+]/[Fe2+ ] of 1.75, which gave a (∑Fe3+ + Fe2+)/(∑M3+ + M2+) molar ratio of 1.79. After these adjustments, a metal recovery of 87% as magnetic precipitates was obtained. The XRD pattern of the magnetic precipitates (Figure 4) confirmed that the magnetic phase had a spinel ferrite structure, and no new
phase was identified. It is interesting to note that the incorporation of other metal ions such as zinc, aluminum, and magnesium that are present in the simulated AMD did not alter the crystal structure of the ferrites substantially. The ferrites showed excellent stability (Table 1): only 1.8 ppm of iron and 1 ppm of zinc were returned to solution in the leach test. The mean particle size of the magnetic precipitates was ca. 2 µm, which was finer than ca. 15 µm for the ferrite formed from a solution containing Fe3+ and Fe2+ only. It should be noted that a saturation magnetization of 34 emu/g was measured. This value of saturation magnetization is considerably lower than for iron ferrite, which can be attributed to the incorporation of a significant amount of zinc and aluminium ions in the crystal lattice. However, this magnetic property is still sufficient for magnetic filtration. The residual concentrations of major metal ions in the effluent water are given in Table 1. It is evident that the removal of metal ions from AMD using the room temperature ferrite process is effective, with iron being reduced from 5880 to 0.2 ppm, zinc from 2760 ppm to 25 ppb, and aluminum from 704 to 0.9 ppm. These experimental findings suggest that an ambient temperature ferrite process can be used to treat AMD if appropriate Fe3+/Fe2+ chemistry can be obtained and the calcium effect is controlled.
Discussion Ferrite formation in Fe3+ and Fe2+ (and/or M2+) solution is considered to be a two-step process. At a Fe3+/Fe2+ ratio corresponding to FeFe2O4, sodium hydroxide will coprecipitate ferric and ferrous ions as hydroxides, forming an intermediate complex hydro sol often referred to as “green rust”(8, 9, 16), which is followed by dehydration of the sol as shown below: Fe2+ + 2Fe3+ + 8OH–
–OH
OH > Fe3+
Fe2+
initial sol OH HO– ] < > Fe3+ < OH HO–
n
(2)
nFeFe2O4 + (4n)H2O spinel ferrite
This reaction can be generalized by replacing Fe2+ by M2+, and the ferrite formed as such is MFe2O4. A reasonable value for n in eq 2 appears to be 8 as the unit cell of a spinel lattice that has cubic symmetry contains 8 “units” of MFe2O4 (7). The ferric ion in MFe2O4 can be completely or partially replaced by other trivalent ions such as Al3+ or Cr3+, giving rise to mixed crystals with aluminate and chromite. This accounts for the observation here with simulated AMD, where no significant deviation of crystal structure was observed although the zinc and aluminum ions were incorporated in the magnetic precipitates. The role of ionic radius also helps to explain the adverse effect of calcium ions on the formation of ferrite. Unlike zinc ions, which have a radius (0.82 Å) similar to that of ferrous ions (0.83 Å), calcium has a significantly larger radius (1.14 Å) and cannot be easily incorporated into the ferrite lattice. As a result, calcium ions co-precipitate as inclusions of CaCO3 and/or CaSO4, which interfere with the growth of crystalline ferrites. When the ratio of Fe3+ to Fe2+ is less than 2, spinel ferrite formation may follow the path:
Fe(OH)2 + 2Fe(OH)3 + [O2]
–OH
> Fe3+
[ –OH
initial sol OH OH HO– < ] > Fe2+ < > Fe3+ < OH OH HO–
n
(3)
intermediate sol
nFeFe2O4 + (4n)H2O spinel ferrite
in which oxygen is required for the formation of the intermediate sol. The two schemes represented by eqs 2 and 3 are supported by the fact that when the ratio of Fe3+ to Fe2+ is 2, ferrite forms readily once the pH reaches about 10, while ferrite forms only after a period of oxidation (induced by stirring the suspension for 15-20 min) if the ratio is lower than 2. The function of oxygen here is to regulate the proportion of Fe2+ to Fe3+. In the case of excess ferric ions in the original AMD, the addition of ferrous ions appears to be necessary. Judicious mixing of effluent streams may permit some control of the ferric to ferrous balance. Dehydration is the rate-limiting step in the ferrite process (reactions 2 and 3) as indicated by a long aging time. Heating accelerates the dehydration stage and shortens the aging time. One option to speed up the process for treating large volumes of AMD is to separate the intermediate sol, followed by conversion of the precipitates to ferrite by heating or by natural aging. The heating cost will be much less than that incurred heating the original AMD. A key for this option is to increase solid-liquid separation efficiency by increasing the density and size of the precipitates. It was found that the size of precipitates (as inferred from measurement of settling rate) is largely dependent of the rate of neutralization and the [Fe3+]/ [Fe2+] ratio. The maximum settling rate was obtained at a ratio of 1.75, with a mean particle size of 16 µm. The finer the precipitates, the more difficult it is to separate the solid from liquid even magnetically. Other options have to be considered, and research is in progress for exploiting the use of catalysts to enhance the rate of dehydration. The above experimental results showed that the spinelstructure ferrite can be made from Fe3+/Fe2+ and synthetic AMD solutions at ambient temperature and pH 9-11 adjusted with NaOH. The molar ratio of Fe3+ to Fe2+ is critical to the formation of ferrite. A ratio of 1.75 gave the highest metal recovery as ferrites and the greatest saturation magnetization, lowest leachability, and coarsest particle size. The presence of zinc ions had relatively little effect on these properties compared to calcium ions, which decreased the metal recovery, magnetic property, and stability. Nearly all metal ions could be removed as ferrites from a simulated acid mine drainage.
Acknowledgments The authors wish to acknowledge the financial support of the Natural Sciences and Engineering Research Council of Canada (Z.X., J.F.) and National Educational Committee of People’s Republic of China (W.W.). Useful discussions with Q. Liu are gratefully acknowledged.
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(3) Uchino, K.; Ogasawara, T. Kawasaki Seitetsu Giho 1980, 12, 665675. (4) Okuda, T.; Sugano, I.; Tsuzi, T. Filtr. Sep. 1975, Sep/Oct, 472478. (5) Tamaura, Y.; Katsura, T.; Rojarayanont, S.; Yoshida, T.; Abe, H. Water Sci. Techol. 1991, 23, 1893-1900. (6) Mandaokar, S. S.; Dharmadhikari, D. M.; Dara, S. S. Environ. Pollut. 1994, 83, 277-282. (7) Smit, J.; Wijin, H. P. J. Ferrites: Physical Properties of Ferrimagnetic Oxides in Relation to their Technical Applications; Wiley: New York, 1959; pp 136-142. (8) Feitknecht, W. Z. Elektrochem. 1959, 63, 34-40. (9) Bernal, J. D.; Dasgupta, E. R.; Mackay, A. L. Clay Miner. Bull. 1959, 4, 15-30. (10) Okuda, T.; Sugano, I.; Tsuzi, T. Kawasaki Jpn. Filter 1975, 12, 472-478. (11) Takada, T. Inst. Chem. Res. 1977, 13, 37-41.
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(12) Bate, G. Ferromagnetic Materials; Wohlfarth, E. P., Ed.; NorthHolland Publishing Company: Dordrecht, 1980; pp 381-507. (13) Kiyama, M. Bull. Chem. Soc. Jpn. 1974, 47, 1646-1650. (14) Yang, K.; Misra, M.; Mehta, R. Proceedings of Waste Processing and Recycling in Mineral and Metallurgical Industries; Rao, S. R., Ed.; CIM: Montreal, 1995; pp 425-438. (15) Foner, S. Rev. Sci. Instrum. 1959, 30, 548-557. (16) Ding, M.; Zeng, H. Environ. Sci. 1992, 13, 59-67.
Received for review January 3, 1996. Revised manuscript received March 22, 1996. Accepted March 25, 1996.X ES960006H X
Abstract published in Advance ACS Abstracts, May 15, 1996.
