Ind. Eng. Chem. Res. 1987,26,824-830
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Larson, M. A.; White, E. T.; Ramanarayanan, K. A.; Berglund, K. A. AIChE J. 1985,31, 90. Matuchovi, M.; Nfilt, J. Kristall. Technik 1976,11, 149. Mullin, J. W. Crystallization, 2nd ed.; Butterworths: London, 1972. Mullin, J. W.; Gaska, C. Can. J. Chem. Eng. 1969,47,483. Mullin, J. W.; Gaska, C. J . Chem. Eng. Data 1973,18,217. Njvlt, Y.;Matuchovh, M. J . Chem. Eng. Data 1976,11, 245. Randolph, A. D.;Rajagopal, K. Ind. Eng. Chem. Fundam. 1970,9, 165. Randolph, A. D.; Sikdar, S. K. Znd. Eng. Chem. Fundam. 1976,15, 64.
Rosen, H. N.; Hulburt, H. M. Chem. Eng. Prog. Symp. Ser. 1971, 67(110),18. Sikdar, S. K. Ind. Eng. Chem. Fundam. 1977,16,390. Wey, I. J.; Terwilliger, J. P. AIChE J. 1974,20, 1219. White, E.T.; Bendig, L. L.; Larson, M. A. AZChE Symp. Ser. 1976, 17(153),41. Zheng, Dao-hong; Budz, J.; Jones, A. G.; Mullin, J. W. J . Crystal Growth 1986,79,658.
Received for review November 14, 1985 Accepted December 5, 1986
Kinetic Study of Pyrite Oxidation in Basic Carbonate Solutions Terry R. Guilinger Sandia National Laboratories, Albuquerque, N e w Mexico 87185
Robert S. Schechter' and Larry W. Lake*$ Department of Chemical Engineering and Department of Petroleum Engineering, T h e University of Texas at Austin, Austin, Texas 78712
The general goal of this experimental study was to find ways to control the unwanted oxidation of pyrite during the in situ leaching of uranium ores. We investigated the effect of particle size, leaching pH, flow rate, total carbonate concentration, and cation type of column leaching rates. Our work appears to be the first dealing with pyrite leaching in test columns at high p H in which the identity and percentages of the various eluted sulfur compounds was measured. The most interesting of the experimental observations was the ultimate ceasing of the pyrite reaction when using an ammonia leachant. That we observed no such ceasing with other leachants suggests the formation of a reaction-inhibiting intermediate. However, no such compound was isolated in microprobe and X-ray diffraction analyses. The existence of the reaction shut-off indicated, nevertheless, that there are mechanisms whereby pyrite reactions can be inhibited. Solution mining is a process which extracts minerals by direct dissolution from an ore body without bringing the solid ore to the surface. Generally, it involves pumping a leachant (or lixiviant) into the ore zone to dissolve the ore minerals, pumping the leachate (pregnant lixiviant) to the surface, recovering the desired constituents, regenerating the leachant, and recycling it. The technique is now being either applied or considered for the recovery of copper, uranium, gold, nickel, silver, cobalt, zinc, vanadium, and manganese as well as a number of nonmetallic minerals. The state of art in solution mining is manifest in papers presented at two international symposiums (Aplan et al., 1974; Schlitt and Hiskey, 1982). Most often the target mineral must first be oxidized before it is soluble. The leachant must therefore contain an oxidizing agent which passes through that part of the ore body which has been denuded of the desired mineral to reach the unleached mineral. The distance that the oxidant travels to reach the unleached mineral increases as the leaching process proceeds. If the ore zone contains one or more oxidant consumers (competitors) which oxidize slower than the desired mineral and if these are present in large quantities, the competitor will limit the distance between injection and production wells. This limit occurs because as the oxidant-rich leachant flows through the stripped ore body, it will be consumed in reacting with the unwanted mineral. Thus, it is of considerable importance to at least limit the amount
* Author to whom
*
correspondence should be addressed. Departments of Chemical and Petroleum Engineering. Department of Petroleum Engineering. 0888-5885/87/2626-0824$01.50/0
of oxidant reacting with the competing mineral. One of the most ubiquitous competitors for oxidant and one which oxidizes rather slowly in both acidic and basic solutions is pyrite (FeS2). This mineral is potentially a highly efficient oxygen consumer since, if the sulfur is oxidized to its hexavalent state, 3.75 mol of oxygen is required per mole of FeS2. Thus, pyrite is often a limiting factor in solution mining. This paper presents a study of pyrite oxidation in basic solutions which are characteristic of leachants used to recover uranium. The goal of this research is to defiie the parameters which affect both the pyrite reaction rate and the distribution of sulfur products. Since oxidation is an undesirable reaction, we will seek factors which can suppress the consumption.
Review of Aqueous Pyrite Oxidation Studies The most commonly proposed reaction for the oxidation of pyrite in aqueous solutions is 4FeS2 + 1502 + 14H20 4Fe(OH), +88SOd2- + 16H'
-
(1)
The rate of this reaction is first-order in oxygen at low concentrations (Goddard and Brosnahan, 1982),decreases to half-order at higher concentrations (Stenhouse and Armstrong, 1952; Kosikov et al., 1974), and becomes zero-order at the highest concentrations (Smith and Shumate, 1970). The reaction is insensitive to pH (Smith and Shumate, 1970; Goddard and Brosnahan, 1982) and does not seem to be affected by different crystalline forms of FeS2 (Goddard and Brosnahan, 1982). Once the bond between iron and sulfur has been severed, the fates of species containing these elements are deter0 1987 American Chemical Society
Ind. Eng. Chem. Res., Vol. 26, No. 4, 1987 825 mined independently (Stenhouse and Armstrong, 1952; Sato, 1960; Mishra, 1973). The nature of the ultimate reaction products for iron and sulfur is not generally understood. If sulfur forms in an intermediate oxidation state, it can be oxidized to sulfate through a reaction similar to S2032-+ 2 0 2 + HzO 2SOd2-+ 2H+ (2)
I
-
However, the extent of this progression is unclear. Stenhouse and Armstrong (1952) identified only sulfate in the reaction products of their experiments. Others have found these also and varying amounts of thiosulfate (S203), trithionate (S,06),tetrathionate (S,O,), and sulfite (SO,) (Goddard and Brosnahan, 1982; Steger and Desjardins, 1978). The nature of the iron product depends on the solution pH. At intermediate pH, reaction 1 forms lepidocrocite (y-FeOOH); as the pH increases, an amorphous iron hydroxide forms at the expense of the lepidocrocite (Nelson, 1978). Other works (Marshall, 1964; Weiser and Milligan, 1939) indicate that goethite (a-FeOOH) forms at high pH. Geothite is thermodynamically favored at high pH, but many times it forms very slowly (Liddell, 1979).
Experimental Procedures and Analytical Techniques Flow Reactor. The flow reactor was capable of maintaining a compressive force on a packed bed at any desired absolute pressure. The pressure level was 3.06 MPa which was enough to ensure that oxygen, added as hydrogen peroxide, remained in solution through the length of the reactor. The details of the reactor design are reported elsewhere (Guilinger, 1983). The reactor was 25 cm long and 2.5 cm in diameter. Ore Preparation. The pyrite was received from Combustion Engineering Minerals Company in three size fractions: through 325-mesh US standard sieve (particle diameters less than 0.044 mm), through 60 mesh onto 200 mesh (0.074-0.25 mm), and through 20 mesh (less than 0.841 mm). Each sample was first blended and stored in an air-tight container at -5 "C. For the leaching experiments, the pyrite was mixed with sea sand to create a 200-g mixture of 10 wt % pyrite which was then packed in the reactor. Operating Procedure. Each experiment consisted of a preflush of deoxygenated, distilled water followed by leachant and then a postleach chloride tracer. The postleach chloride tracer was used to determine the pore volume of the pyrite sample. The leach and postleach solutions were made up by using deoxygenated, distilled water and reagent-grade chemicals. The pressure drop and oxygen concentration were recorded continuously with calibrated probes. Effluent samples were analyzed for sulfate (SO,-), total soluble sulfur (TSS), and pH. After each experiment, the contents of the packed bed were sectioned into four parts and analyzed by X-ray diffraction for crystalline phases and by scanning electron microscope (SEM) for iron oxide/hydroxide solids. We also analyzed the spent pyrite for precipitated elemental sulfur to determine if all the oxidized pyritic sulfur was collected in the effluent stream. All analytical procedures used to make these measurements were obtained from the literature (Guilinger, 1983). Experimental Results Effect of Oxygen Concentration on the Pyrite Reaction Rate. The base case leachant was a potassium carbonate solution at a pH of 10.1 with 2.0 kg/m3 of total carbonate and 100 g/m3 of equivalent oxygen flowing at
; t O0
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Figure 1. Effect of injected oxygen concentration on total soluble sulfur history.
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Figure 2. Effect of injected oxygen concentration on sulfate history.
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Figure 3. Effect of injected oxygen concentration on effluent oxygen history. IO 40,
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Figure 4. Effect of injected oxygen concentration on pH history.
an interstitial velocity of about 0.118 mm/s. Table I shows the formulations for each experiment and the plateau effluent values of pH, oxygen concentration, total soluble sulfur (TSS),sulfate concentration, and the calculated pyrite reaction rate. The base case is experiment 1. To evaluate the effect of oxygen concentration on the pyrite reaction rate, we performed experiments on the large-grain pyrite using the base case leachant in one experiment and halving the injected oxygen concentration to 50 g/m3 of O2 in another, experiment 2. Figures 1-4 show comparisons of TSS, sulfate, oxygen, and hydrogen ion effluent concentrations for these two experiments. The
826 Ind. Eng. Chem. Res., Vol. 26, No. 4, 1987 Table I. Summary of Pyrite Experiments
expt 1 2
3 4 5
6 7 8 9 10 11 12 13 a
ore mesh size, mm
