The Oxygen Pressure Leaching of Pyrite in Sulfuric Acid Has Been Studied at Pressures up to 976 Psi G"2 and Temperatures From
THE OXYGEN PRESSURE LEACHING OF PYRITE
by
LEONARD KEITH BAILEY B.S. University of Utah, 1973
A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF APPLIED SCIENCE
in the Department of Metallurgy
We accept this thesis as conforming to the required standard
THE UNIVERSITY OF BRITISH COLUMBIA
December, 1974 In presenting this thesis in partial fulfilment of the requirements for
an advanced degree at the University of British Columbia, I agree that
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Department of Metallurgy
The University of British Columbia Vancouver 8, Canada
Date January 27, 1975
fi ABSTRACT
The oxygen pressure leaching of pyrite in sulfuric acid has been studied at pressures up to 976 psi G"2 and temperatures from
85 to 130°C. The dissolution has been found to follow linear shrinking 1 /3 core kinetics (1 - (1 - 2 m hr. has been obtained for the conditions tested. Factors studied for the reaction include: geologic differences in pyrite, acid concentration, pulp density, the presence of neutral salts and other sulfide minerals, and particle size, along with the dependence on temperature and pressure. An overall equation for pyrite dissolution has been tested and found to correlate well with the experimental data. The distribution of reaction products between sulfate and elemental sulfur, and ferric or ferrous iron has been examined. Higher yields of elemental sulfur were found with lower acid concentration and reduced oxygen pressure. The ferric/ferrous ratio was found to be most affected by the extent of pyrite dissolution. Increases in the ratio were also found when neutral salts were added to the system and when oxygen pressure was increased. The dissolution reaction was found to be an acid-consumer for initial acid concentrations above O.4MH2SO4. For concentrations below this Level, the reaction produces acid. An activation energy of 12. 7 t 0. 7 kcal/mole was observed for the dissolution reaction. iii TABLE OF CONTENTS Page ABSTRACT ii TABLE OF CONTENTS iv LIST OF FIGURES vi LIST OF TABLES vii ACKNOWLEDGEMENTS viii A. INTRODUCTION 1 1. General 1 Z. Occurrence 1 3. Structure 1 4. Thermodynamics 3 5. Use 3 6. Previous work 6 7. Scope of this work 8 B. EXPERIMENTAL 10 1. Materials 10 Z. Apparatus 11 3. Procedure 15 4. Analysis 16 (1) Oxygen Consumption 16 (2) Iron Analysis 16 (3) Sulfur Species 18 C. RESULTS AND DISCUSSION 20 1. Reproducibility 20 Z. Variation between Pyrites 22 3. Particle size 25 4. Pulp density 30 5. Acid concentration 33 6. Effect of sulfide minerals 36 7. Effect of neutral salts 38 8. Reaction model . 38 9. Oxygen pressure 41 iv TABLE OF CONTENTS (Continued) Page C. RESULTS AND DISCUSSION (Continued) 10. Temperature 48 CONCLUSIONS 62 REFERENCES 65 APPENDIX: Actual Experimental Data 67 v LIST OF FIGURES Figure Page 1 Pyrite structure 2 2 Potential - pH diagram for the iron-water- sulfur system 5 3 Oxygen removed from system vs recorder ..... 17 4 Reproducibility .. 21 5 Comparison of pyrites 23 6 Comparison of unleached pyrites (photographs). . 24 7 Effect of particle size for Wards pyrite 26 8 Effect of particle size for Sullivan pyrite 27 9 Rate of oxygen consumption vs surface area .... 29 10 Effect of pulp density 31 11 Iron extraction vs pulp density 32 12 Acid effect 34 12A Acid consumption vs log initial acid concen• tration 35 13 Effect of sulfide minerals 37 14 Effect of neutral salts 39 15 Comparison of reaction models ...... 40 16 Particles of Sullivan pyrite after leaching (photographs) 42 17 Effect of pressure, oxygen consumption vs time * . • 43 18 Effect of pressure, reaction model vs time .... 44 19 Rate of dissolution vs oxygen pressure . i 45 20 Fit of data to Langmuir Adsorption Model 47 21 Fit of data to reversible adsorption with dissociation of oxygen model . 49 22 Iron extraction vs oxygen partial pressure 50 23 Effect of temperature, oxygen consumed vs time . , 51 24 Effect of elemental sulfur 53 25 Temperature data in terms of reaction model ... 54 26 Arrhenius plot of temperature data 55 27 Solubility of oxygen in water 59 28 Effect of temperature on the ferric/ferrous ratio . . . 60 vi LIST OF TABLES Table Page I Thermodynamic data for pyrite 4 II Analysis of iron and sulfur in pyrites 12 III Spectroscopic analysis of pyrites 13 IV Comparison of X-ray diffraction patterns for pyrites with a standard pattern . . 14 vii ACKNOWLEDGEMENTS Sincere thanks are extended to Dr. Ernest Peters for his help in carrying out this research and in bringing it to its present form. The help received from other members of the faculty, graduate students and technicians in the department of Metallurgy is also appreciated. The financial support of the U.B.C. Research Committee and a U.B.C. Summer Session Research Scholarship is gratefully acknowledged. viii A. INTRODUCTION 1. General Pyrite is the most common and wide-spread of the sulfide minerals. * Its name stems from the Greek word "pyr", meaning fire. In the days of the ancient Greeks, the mineral was known for its ability to produce sparks when struck by iron. It is found most often with deposits of chalcopyrite, sphalerite, and galena. 2. Occurrence A partial list of the pyrite-producing nations includes: Spain, which at Rio Tinto and other mines accounts for over 2.2 million tons annually; the U.S.S.R., with an output of over 3.4 million tons; and Japanese sources of over 4.4 million tons per year. North American production is led by the United States where deposits in Tennessee, Pennsylvania and several western states produce 0.8 million tons yearly. Canadian sources include: Noranda, Normetal Mines, and the Anaconda (Canada) Co. Ltd., producing about 0.3 million tons per year. 3. Structure Pyrite has a structure similar to that of NaCl, (see Fig. 1). The iron atoms take the sodium positions and S2 groups replace the chlorine. The S2 molecules are oriented so as to leave no net distortion in any direction of the crystal. The face centered cubic lattice has an FIGURE I. PYRITE STRUCTURE o 3 aQ distance of 5.417A. Pyrite is distinguished from the other FeS£ form, marcasite, by the orthorhombic marcasite structure. The groups in marcasite have a net orientation which produces the orthor• hombic crystal structure and makes the mineral more subject to chemical attack than pyrite. 4. Thermodynamics The thermodynamic properties of pyrite are shown in Table I. The reported values are in good agreement. The potential -pH 9 diagram for the iron-water-sulfur system is shown in Figure 2. 5. Use Since pyrite is often found in massive deposits, many processes have been designed to extract the iron and sulfur values for commercial use. Although the mineral is nearly 47% iron, its-use as a feedstock for iron production is limited by the complexities of removing the sulfur. Some use is made, however, of the iron oxides remaining after the production of sulfuric acid. Such processes require the roast• ing of pyrite to liberate sulfur as SO2 and yield an iron oxide calcine. The SG"2 is converted to sulfuric acid by passing the gas over a catalyst bed. While this process is effective, its use in North America is being reduced by competition from other sulfuric acid-producing methods. Environmental pollution controls, particularly on the petroleum and non- ferrous metal industries, have resulted in new processes for recovering sulfur from stack gasses. ^ Increased production of acid by these methods TABLE I THERMODYNAMIC DATA FOR PYRITE ° o Source: A Gf AS Latimer4 -39.84kcal 12.7cal/Deg. Toulmin & Barton5 -38.3 KelLey and King6 12.7 7 Pemsler -38.1 Note: The difference in the free energy terms is actually not as great as reported. The mea• sured value for the free energy of the ferrous ion has been revised since Latimer's report. If the free energy for the ferrous ion is sub• tracted from that of pyrite, a better idea of the accuracy is given. Latimer -39. 84 - (-20. 3) = -19. 54 kcal Toulmin -38.30 - (-18. 85)18 = -19.45 Pemsler-38. 10 - (-18,85) = -19.25 The net difference in free energies measured is Less than -0.15 kcal. 5 FIGURE 2. POTENTIAL/PH DIAGRAM FOR THE IRON- WATER-SULFUR SYSTEM AT IOO°C. 6 has had a negative effect on the pyritic production. In the 1950's, processes for extracting uranium from ores containing pyrite were designed so that the acid for leaching was provi- 11,12 ded by pyrite oxidation. Leaching the ore at temperatures over 100°C under a slight overpressure of oxygen (lOpsi) resulted in uranium extraction of 90 to 95% in from four to six hours. The investigators found that sulfur values in the pyrite were oxidized to sulfuric acid which then dissolved the uranium values. 6. Previous Work The behavior of pyrite under oxygen pressure leaching was 13 studied by McKay and Halpern. Variables in the research included: acid concentration (0 to 0.15 MH-^SO^.), surface area (275 to 540 cm^/g), pulp density (2-8% solids), 02 pressure (0-4 atrn), temp. (100 to 130°C), and effects of ferrous sulfate, ferric sulfate, and cupric sulfate on the system. It was concluded that the major reaction in the dissolution of pyrite is: 1. FeS2 +• 202 — FeS04 +• S°. Formation of ferric ion is accounted for by the reaction: 2. 4FeS04 + Oz +• 2H2S04—4Fe(S04)1 5 + 2HzO. Several other reactions were found to play only a minor role in the dis• solution. 7 The overall rate of reaction was found to be proportional to the pyrite surface area, and to the oxygen pressure. It was shown to be independent of the solution concentration over the range studied. The distribution of reaction products between sulfate and elemental sulfur depended on the solution composition and temperature. High temperatures and low acidities were found to favor higher yields of sulfuric acid and the converse was also shown to hold true. The reaction mechanism postulated involves chemisorption of oxygen on the pyrite surface, followed by a slower attack of another O2 molecule ori the O2 covered mineral site. 3. FeS2 •• Oz -l- Oo(aq) > [FeS2 * 202] - —> FeS04 +• S° 1 slow J An activation energy of 13. 3 £ 2 kcal/mole was observed for the pyrite dissolution. Further work on the system was done by Pawlek and co-work- 14 ers. The experimental pressure was raised to 16 atm, and the temperature effect was studied from 60 to 160°C. The activation energy for pyrite oxidation was observed to be 13. 1 kcal/mole. In his work, Pawlek raised acid concentrations up to 0.4MH2SO4 with the result that he was able to see a slight increase in the ferric to ferrous ratio as the concentration increased. This effect was not reported by McKay, who studied lower acid levels. Pawlek also proposed a reaction by which acid might be generated: 8 +3 f2 f 4. FeS2 +• 8Fe f 4HzO 9Fe f SOj + S° f 8H . Such a process would explain the uranium-pyrite leach system, and yet the effect would become small in comparison with the case of high initial acid concentrations. 15 Peters has studied the effect of pyrite on the dissolution of other minerals. Lead, copper, and zinc sulfides were shown to have increased dissolution rates when mixed with pyrite, while the attack on the pyrite itself was slowed. This phenomenon is explained by a galvanic control on the reaction. Pyrite is assumed to be anodically controlled and is thus more active for oxygen reduction, while the other minerals are cathodically controlled, allowing more active mineral dissolution. Since the sulfide minerals studied are conductors, electrons can be ex• changed and the kinetics of dissolution altered from the case where no other mineral is present 7. Scope of this work The research reported in this thesis is an attempt to clarify the factors effecting the dissolution of pyrite. Equations 1 and 2 were combined to give an overall reaction of the form: +f 5. FeS2 f ( | f | y H- -| x) 02 + (2 + x - 2y) H*" >(l - x) Fe +H + xFe f (2 - y)S° +• ySOj > (1 - yV| ) H20. The effects of temperature, acid concentration, pressure, pulp density, particle size, neutral salt additions, presence of other minerals,, and possible variations among several pyrites were studied, with respect to rate of dissolution and to the values of the constants x and y. It was hoped that such information would allow a better understanding of the reaction mechanism. 10 B. EXPERIMENTAL 1. Materials Five different pyrites were obtained and ground by hand - using a mortar and pestle - to the desired screen sizes. Enough mineral was ground initially for all the tests conducted. The pyrites originated from: (a) Cominco Mines - Sullivan Mine, (b) Kimberley (also a Sullivan Mine sample), (c) Noranda Mines - Quebec, (d) Ward's Scientific Co.-, - Colorado, and (e) Japan, through Professor Nagai. The Sullivan and Kimberley samples were provided as powders (-65 + ZOOmesh), while those from Ward's and Noranda were massive specimens. The Nagai sample consisted of four or five large single crystals which were crushed and sized. All the screened samples were washed to remove any fines which might have adhered to the surface, then air-dried. Four size fractions were used in the experimental work: (a) -150 +- ZOO mesh (Tyler screen sizes), (b) -ZOO +• Z70, (c) -Z70 f 3Z5, (d) -325' + 400 11 Samples of the ground minerals (rl 50 +• 200 mesh fractions) were sent to Can Test Ltd. for chemical analysis. The results are shown in Table II. Spectroscopic analysis (Table III) showed only slight impurities, the largest of which was silicon with a maximum content of 0.5%., Copper in all samples was below 0.05%. , Identical patterns for all five pyrites were obtained by X-ray diffraction. All of the lines on the patterns were attributable to pyrite. The d values obtained are shown in Table IV and were compared with a standard by Swanson, Gilfrich, and Ugrinic. ^ All chemicals used were reagent grade and demineralized water was used throughout the tests. 2. Apparatus All tests were made in a 106.5 ml zirconium shaking autoclave with a 15.0 ml gas reservoir. Zirconium tubing and fittings were used wherever the hot leach solution would be,in contact with a metal surface. A teflon gasket provided the seal at the top. Shaking was provided hori• zontally at a rate of 288-1.5 inch-strokes per minute. The cylindrical autoclave was tilted at a 45° angle to give better agitation and to. facilitate sample removal. Temperature was controlled within - 1/2°C by a thermistemp temperature controller (Model 71, Yellow Springs Industrial Co.) linked to a resistance wire heater on the autoclave. The thermistor sensing 12 TABLE II ANALYSIS OF IRON AND SULFUR IN PYRITES Pyrite % Fe % S S/Fe Ratio Kimberley 45.86 51.65 1. 97 Wards 45. 76 51.55 I.97 Sullivan 45.76 52.25 I.99 Noranda 46.43 52.60 1. 98 Nagai 46.48 52.45 1. 97 13 TABLE III SPECTROSCOPIC ANALYSIS17 OF PYRITES Kimberley Sullivan Wards Noranda Nagai Al . 0. 003 0. 001 0. 003 0. 003 0.001 As 0. 1 0. 1 ND ND ND Ca 0.01 0. 005 0.005 0. 005 0. 01 Cr 0. 007 0.001 0. 001 0.003 0.003 Co 0. 005 0. 005 ND 0. 03 0. 005 Cu 0. 03 0.05 0.01 0. 05 0. 003 Au TRACE TRACE TRACE TRACE TRACE Fe MATRIX MATRIX MATRIX MATRIX MATRIX Pb 0. 1 0. 07 ND ND 0. 01 Mg 0. 2 0. 1 0. 1 0. 1 0. 1 Mn ; 0.03 0. 03 0. 01 0. 05 0. 05 Mo 0.005 0. 003 0. 001 0.007 0. 005 Ni .. 0.001 0.001 ND ND 0.001 Si 0. 2 0. 2 0.5 0. 2 0. 05 Ag 0.001 0. 001 TRACE TRACE TRACE Sn 0. 1 0. 1 ND TRACE TRACE Ti 0.001 ND ND ND ND Zn TRACE TRACE ND TRACE TRACE Not Detected: Antimony, Barium, Be ryllium, Bismuth, Boron, Cadmium, Gallium, Niobium, Potassium, Sodium, Strontium, Tantalum, Thorium, Tungsten, Uranium, and Vanadium. TABLE IV COMPARISON OF X-RAY DIFFRACTION PATTERN FOR PYRITES WITH A STANDARD PATTERN Line Observed d Value Reported by Swanson 1 3.1057 3.128 2 2.6913 2. 709 3 2.4114 2.423 4 2.2000 2.2118 5 1.9093 1.9155 6 1.6289 1.6332 7 1.5602 1.5640 8 1.4993 1.5025 9 1.4453 1.4448 10 1.3130 - 11 1.2752 - 12 1.2431 1.2427 13. 1.2113 1.2113 1.2112 14 1.1822 1,1823 1.1822 15 1.1551 1.1548 1.1551 Lines 10 and 11 were very weak and have apparently not been reported before. They coincide with possible crystal planes in pyrite (410 and 411, respectively). The double values for lines 13, 14 and 15 stem from re• solved doublets which, in turn, result from the x-ray source emitting two wavelengths, k »j<. and k/x. . 15 unit was located in a well in the zirconium autoclave top. Calibration with an oil bath and a precision thermometer was performed before leaching experiments were conducted. Oxygen consumption throughout the reaction was measured by a pressure transducer (Consolidated Electrodynamics Corp., Model 4-311, 0-1000 psi range) coupled to the gas reservoir. The system was pressurized and sealed at the start of each run and the drop in pressure, as oxygen was consumed, recorded via the pressure transducer by a Sar- geant strip chart recorder. 3. Procedure A solution volume of 50 ml was used in all experimental runs. To initiate the run, the powdered mineral sample was placed in the auto• clave, the solution added, and the autoclave sealed. Shaking and heating were then started. When the temperature reached the desired level (10-15 min), the unit was pressurized with oxygen. At the end of the run, the autoclave was sealed by closing the oxygen valve, the heat turned off, and the system cooled by water through a copper coil wrapped around the top of the autoclave. Experiments were allowed to run for a specified time (3 or 6 hours) or to a specified oxygen consumption. Cooling to approximately 60°C took about 10 min., at which time the autoclave was opened and the contents removed by a suction flask. The material was then filtered, the residue dried,and the analysis work performed. 16 4. Analysis (1) Oxygen Consumption:- As stated, the oxygen pressure drop throughout a run was monitored by a transducer coupled with a recorder. 18 Calibration of this system using steam tables showed the recorder to vary linearly with pressure. Rather than calculate the number of moles of oxygen consumed by using a gas law-pressure relation, a series of tests were conducted in which a specific volume of gas (measured by a gas burette) was removed from the autoclave and the resulting pressure drop recorded. These measurements show the relation to be linear though different at each temperature (see Fig. 3). In each case a slope, moles/chart division, was obtained and these values were used in interpreting the results of the leach tests. All tests were run under constant volume conditions, allowing the pressure to decrease over the course of a run. 19 (Z) Iron Analysis:- The standard ceric sulfate titration was used to determine the ferrous iron present in solutions. A value for total iron was obtained by reducing the ferric ions in the solution with stan• nous chloride then repeating the ferrous determination. Ferric ion concentration was determined by difference. Ferroin indicator (1,10 Ortho phenanthroline Ferrous Sulphate . 0Z5M) was used as an end point in the titratrations. Iron in the leach residues was dissolved with aqua regia. It should be noted that hydrolysis of the iron was not a problem. Iron oxides were formed in only two runs in which initial acid concentra• tions were very low. The resulting solution was then boiled to remove 17 FIGURE 3. OXYGEN REMOVED FROM SYSTEM VS. RECORDER 18 the nitric acid and analyzed for total iron. Checking the accuracy of the method on iron standards showed a variation of . 3% for ferrous analy• sis. Mass balances on the iron going into and coming out of the autoclave showed the techniques to be accurate. (3) Sulfur Species:-Sulfate and elemental sulfur were the only two species considered in this work. The assumption that no other species were present was justified by the rapid dissociation of such species in the presence of oxidizing agents such as the ferric ion. As a check on the assumption, a sample of one of the most stable intermediate sulfur 20 species, dithionate was put into a hot ferric sulfate solution. An in• crease in the ferrous concentration of the solution was noted showing that ferric had been reduced and the dithionate oxidized. The quantities of sulfate and elemental sulfur produced were calculated from the dissolution reaction. Four samples were analyzed chemically for elemental sulfur as a check with accuracy of - 2 percent• age points. The overall reaction in general form for the dissolution is equation 5. 5. FeS2 -y + - x )Oz 4- (2 f x-2y)H > (l-x)Fe + xFe +- (2-y)S° + ySO~ f (1-y + jx) HzO Hence, knowing the iron analyses, the sulfur species concentrations are calculated using the equations: 19 2 1 6. moles SO^ = j (moles O-, - (moles FeS2 consumed) _ (moles Fe^^~ produced) 7. moles S° = 2 (moles FeS2 consumed) - moles S04 produced 20 C. RESULTS AND DISCUSSION . 1. Reproduc ibility Three factors affecting the reproducibility of the results were examined. They are: (a) complete removal of material from the autoclave and loss during filtering and drying, (b) the effect of the reaction vessel material (Zr), and (c) the consistency of oxygen consumption findings. The material loss possibility was checked by adding pre- weighed samples of pyrite to an autoclave full of water, then going through the removal and drying procedure and re-weighing the samples. Losses were shown to be less than 0. 5% and were therefore neglected. The possibility of the zirconium autoclave affecting the reaction was remote, but a check was made using zirconium foil to more than double the area of metal available for reaction. No effect was noted on the pyrite leaching. Oxygen consumption consistency was checked by repeated runs with the same conditions. Figure 4 shows that the reproducibility is excellent. With good agreement in all tests made, the reproducibility of the experimental runs is taken to be quite adequate. There is, how- Zl 0.12 r47 0.10 0.08 OXYGEN CONSUMED 0.06 (MOLE) %Ext 0.04 Time % elem S ferric/ferrous 3hr 33.33 542 55.73 6 33.09 8.78 83.72 10 32.54 2 1.17 98.08 0.02 0.00 4 • 6 8 10 TIME (HR.) FIGURE 4. REPRODUCIBILITY, OXYGEN CONSUMPTION VS. TIME. IM H2S04, 1.10°, 976 PSI 02. 22 ever, a certain amount of caution which must be exercised in interpreting the data. This comes about as a result of changing product concentra• tions through the course of a reaction. This variation is shown in the figure for the same reproducibility curves mentioned above. It is noted that as the reaction proceeds further to completion the ferric/ferrous ratio increases sharply. 2. Variation between Pyrites Five different pyrites were used in this study to determine if geologic differences in the pyrites had an effect on the leaching para• meters. As seen from the analysis of the samples (Table II), there are slight differences in composition. However, all the values fall within the analytical accuracy and may or may not be significant. Results of experi• ments at 976 psi O2, IMH2SO4 and 110°C are shown in Figure 5. It is noted that the pyrites can be classified with two groups, one more reactive than the other. The reason for this difference was not immediately apparent. The five samples were examined using the scanning electron microscope to check the possibility of a surface area difference rather than a structural variation such as grain boundary composition. Photo• graphs of two of the unleached, pyrite s (Figure 6) show that the surfaces are indeed different. The other three pyrites show the same structure in each group. The Sullivan and Kimberley samples have a rougher texture which explains the more rapid attack. A quantitative measurement of the surface areas by BET techniques was not attempted due to the large 23 0.10 TINE (HR.) FIGURE 5. COMPARISON OF PYRITES, OXYGEN CONSUMPTION. IM H2SC^, 110°, 976 PSI 02. 24 SULLIVAN 150 X 200 MESH WARDS 150 X 200 MESH FIGURE 6. COMPARISON OF UNLEACHED PYRITES. particle size. The fact that the more rapidly dissolved Sullivan and Kimberley pyrites are from the same mine in an ore dominated by pyrrhotite explains the similarity in leaching behavior and may account for the deviation from the other three samples. Whether the stoichio- metry of the pyrites is a factor in the explanation is not known. The reproducibility of the sulfur analysis does not permit such an exact calculation. The distribution of products between sulfate and elemental sulfur and ferric and ferrous iron is quite scattered. This can be partially explained by the extent to which the reaction was allowed to continue. All five samples were run for six hours at the same conditions. This gives a different extraction in each case. It is found, however, that in general the more reactive pyrites produced slightly higher ferric/ ferrous ratios. 3. Particle Size Four samples of each of the two types of pyrite were tested, (Sullivan pyrite from the rapidly attacked group and Wards pyrite from the slower group). The size fractions ranging from -150 f 200 to -325 +• 400 mesh were all run at 976 psi O2, IMH7SO4, and 110°C. As shown in Figures 7 and 8, the smaller particle sizes leached more rapidly. It is noted that even at the smallest particle size, the rates of the two types of pyrites are still different due to the different breakage patterns. A com• parison of the ferric/ferrous ratios and the percentage of elemental 26 0.10 325/400 r24 270/325 r23 20r^|70 0.08 150/200 r 14 0.06 OXYGEN CONSUMED (MOLE) 0.04 %elem S % Ext 0.02 Mesh Size ferric/ferrous 325/400 32.58 8.54 86.15 270/325 33.62 9.03 86.43 200/270 32.67 8.54 79.56 150/200 34.03 7.57 70.24 0.00 i i ••1 • i 3 4 5 6 TIME (HR.) FIGURE?, EFFECT OF PARTICLE SIZE ON OXYGEN CONSUMPTION I WARDS PYRITE. I M H^QjJIOt 976 PSI 0^ 27 TIME (HR.) FIGURES. EFFECT OF PARTICLE SIZE ON OXYGEN CONSUMPTION FOR SULLIVAN PYRITE. IM H2SQ|, 110°, 976 PSI 02. 28 sulfur produced shows little change with variation of particle size. The two types of pyrites still show a very slight trend toward higher ferric production and aLso greater production of elemental sulfur in the case of the more rapid leaching Sullivan structure. Again, the effect of extraction time is apparent on the ratios. Figure 9 shows a plot of initial rate (moles G^/hr) vs surface area for the various size fractions and pyrite types. The area term was calculated simplyby assuming a spherical model for the particles. Obviously, this model is not representative of the particular particle shapes involved as seen in the photographs, but it does allow an estimate of the increase in surface area with decreasing particle size. As the figure shows, the rate of oxygen consumption increases with the surface areafor both types of pyrite. This is in agr eement with the findings of McKay and Halpern at lower acid concentrations. The non-linearity is thought to re suit from two oxygen consuming reactions operating at the same time. One, the heterogeneous dissolution at the pyrite surface and the other, the homogeneous oxidation of ferrous ion to ferric. The effect is also noted in the subsequent pulp density section. The heterogeneous dissolution reaction shouldbe affected linearly by changes in surface arjea; the ]!i: This concentration increase raises the homogeneous reaction rate, causing the non-linearity of oxygen consumption with surface area. 29 0.04h AREA (CM2/G) FIGURE 9. RATE OF OXYGEN CONSUMPTION VS. SURFACE AREA. \M H2S04o 110°, 97S PSI 02. 30 4. Pulp Density To reduce the number of experimental runs, the Sullivan pyrite was chosen for use throughout the rest of the tests. As noted earlier, the majority of experiments in this study were run at a pulp density of 10% solids. Two additional runs were made at 5 and 20% solids. Runs were made at 976 psi oxygen pressure, 110°C and lMH^SO^ for 3 hours. Figure 10 compares the oxygen consumption curves obtained and tabulates the other data. The most notable effect is the decrease in extraction with increased pulp density. The relationship is linear as shown in Fig. 11. The acid concentration in the tests is much greater than the reaction could consume in all cases and can be ruled out as a cause of the reduced extraction. Mass transfer of oxygen to the surface was also ruled out as the determining factor by making a run with chalcocite, C^S. Oxygen consumption for the copper mineral was ten times that of the pyrite for the same surface area and leach conditions. A possible explanation for the results stems from the fact that two reactions are proceeding simultaneously; one, the heterogeneous disso• lution of pyrite, and the other, the homogeneous oxidation of ferrous ion to the ferric state. The first of these reactions has been shown to have a linear dependence on surface area in work reported by McKay. The second, homogeneous reaction is, therefore, thought to account for the deviation from linearity. This is presumably due to increases in solution concentrations which affect the reaction rate. 31 FIGURE 10. EFFECT OF PULP DENSITY, OXYGEN CONSUMPTION VS. TIME. IM H2S04, IIO* 976 PS I 02. FIGURE IL IRON EXTRACTION VS. PULP DENSITY. No real trend was seen in the formation of elemental sulfur but a slight increase in the ferric/ferrous ratio is noted on increasing the pulp density. 5. Acid Concentration Samples of the -150 +• 200 mesh Sullivan pyrite were leached at 110°C and 976 psi oxygen in 0. 01, 0.1, 1.0, and 3. 0M sulfuric acid. Oxygen consumption for the reactions is shown in Figure 12, along with the other data obtained. The oxygen consumption decreases with increasing acid concentration above 0.1 MH^SO^.. This translates into a similar decrease in the mineral decomposition rate due to changes in elemental sulfur formation, but the trend is in the same direction. Increased yields of elemental sulfur are found with increased acid concentrations. , This is in agreement with the trend reported by 21 McKay and also by investigators working on other sulfide minerals. The ferric/ferrous ratio stays essentially the same for all but the lowest acid concentration studied. The higher ratio may be related to the presence of hydrolyzed species at this low acid level. Stoichiometric calculations using equation 5 show that sulfuric acid is produced from pyritic sulfur when the initial concentra• tion is 0. 01 and 0. lMH^SO.^. At higher acid levels, however, the disso• lution reaction is an acid consumer. Figure 12A compares acid consump- 34 0.10 1 Mr 10 0.08 3M r 38 0.06| OXYGEN CONSUMED (MOLE) 0.04 0.02 H2S04 %elem S ferric/ferrous % Est O.OIM 11.09 12.35 66.16 0.10 20.06 8.64 77.88 1.00 33.09 8.78 83.72 3.00 36.36 8.40 69.77 • • « I • 2 3 4 5 6 Tl ME (HR.) 1 FIGURE 12. ACID EFFECT, OXYGEN CONSUMPTION VS. TIME, 110°, 976 PSI 02. 35 FIGURE 12 A. ACID CONSUMPTION VS. LOG INITIAL ACID CONCENTRATION. 36 tiori (calculated by the stoichiometry of equation 5) with the initial acid concentration'and shows the transition point to be about O.4MH2SO4. This value'is the maximum acid concentration which would be produced by pyrite in an operation such as 'dump' leaching. This, behavior . agrees with the conclusions of the other investigators who studied the lower acid concentrations. The fact that the stoichiometry of the reactions, complies with equation 5 indicates that the.reaction mechanism does not change from one concentration to another. 6. Effect of sulfide minerals Peters has noted a galvanic effect on the dissolution of other minerals by pyrite (see section on previous work). To check possible effects of these minerals on pyrite, 0. 5g of covellite, sphalerite, and pentlandite were added, respectively, with the pyrite in three successive runs. The oxygen consumption for the tests with the extra minerals is compared with a standard run (pyrite;only) in Fig. 13. The galvanic effect is shown by the higher initial slope, corresponding to the con• sumption of the added mineral, followed by a return to the slope of normal pyrite dissolution. In all three cases, the oxygen consumption parallels the standard pyrite run after the other mineral is consumed. Rather scattered values for elemental sulfur are in the same range as those obtained for pure pyrite. The ferric/ferrous ratio shows an increased production of ferric ion with the other mineral additions. 0.08 Nir50 .Znr48 'Cur49 • None 0.06 r43 OXYGEN CONSUMED (MOLE) 0.641 0.02 Sulfide %elem S ferric/ferrous % Ext Cu 31.17 8.68 56.29 Ni 27.94 6.00 55.55 Zn 3800 6.75 60.29 OnlyFe 33.33 5.42 55.73 0.00 I 2 TIME (HR.) FIGURE 13. EFFECT OF SULFIDE MINERALS. IM H2S04, 110°, 976 PSI 02. 38 7. Effect of neutral salts The effect of neutral salts on pyrite dissolution was studied as a continuation of the work on other minerals to see if the metal ions themselves had an effect on the system. Solutions 1M in the neutral salts CuSO^., NiSO^. and ZnSO^ were made. Acid concentration was kept at IMH2SO4. All three salts seemed to slow the dissolution of pyrite (see Figure 14). In the case of copper sulfate, the most effective salt, reduction of iron-extraction was about 25%. A trend toward higher elemental sulfur yields is indicated but all results are within the analyti• cal accuracy. Larger ferric/ferrous ratios are also evident with the salt additions. The catalytic effect of the cupric ion on the production of •1-2 ferric iron was studied by Pawlek for concentrations less than 0. 05MCu His findings agree with the results obtained in this study. McKay also ran tests at very low copper concentrations which demonstrate the higher ferric/ferrous ratios but dismissed the accompanying reduction in pyrite dissolution as an experimental error. 8. Reaction model In testing reproducibility of oxygen consumption data, a run was made in which over 98% of the pyrite was dissolved. Taking the oxygen consumed to this point as 100%, consumption during the run was used to calculate c< . the fraction of the material reacted for several times. These values of c*. were then inserted into models for reaction mechanisms. Figure 15 shows the correlation of the data with the simple 39 i FIGURE 14. EFFECT OF NEUTRAL SALTS. OXYGEN CONSUMPTION VS. TIME. IM SALT, I M H2S04, 110? 976 PSI Og. 0.9 TIME (HR.) FIGURE 15. COMPARISON OF REACTION MODELS. MODEL VS. TIME. 1-(1- o< ) vs time, shrinking core model. As a check on the sensiti• vity of the model, a plot is also shown using the same equation, but changing from a one-third to one-half power relationship on the (1- ©<- ) term. The sensitivity is seen to be good, particularly at the higher fraction reacted values. This model agreement is the basis for assuming a topochemical reaction in pyrite dissolution. The model also shows that the build-up of elemental sulfur on a dissolving particle is not a major factor in controlling the rate of reaction. This assumption seems reasonable since only about 35% of the pyritic sulfur is converted to the elemental form (the rest going to sulfate). Photographs showing pyrite particles at various stages of dissolution are shown in Figure 16. The formation of elemental sulfur coatings on the particles is shown quite clearly. Oxygen consumption and the other data for these runs was presented in the section on reproducibility. 9. Oxygen pressure The effect of oxygen pressure was studied by leaching samples at five different pressures; .976, 676, 476, 326 and 176 psiOz. The resulting oxygen consumption data are plotted in Figure 17. Figure 1 /3 18 shows the data plotted in terms of the 1-(1- 'X ) shrinking core model. The rate of pyrite dissolution in terms of particle penetration taken from the slope of the lines was plotted versus time in Figure 19. This plot shows that the linear relationship between pressure and rate observed by McKay and Pawlek does.not hold at higher pressures. To FIGURE 16. PARTICLES OF SULLIVAN PYRITE AFTER LEACHING. 150X200 MESH, 110°, IMH2S04, 976 PSI 02. 43 Pressure % elem S ferric/ferrous % Ext 976 psi 02 33.33 5.42 55.73 0.08 676 39.09 3.85 4a07 476 38.82 2.82 39.26 326 41.28 1.26 32.30 176 43.18 0.89 21.44 0.06 OXYGEN CONSUMED (MOLE) 0.04 0.02 O.'OQ TIME (HR.) FIGURE 17. EFFECT OF PRESSURE. OXYGEN CONSUMPTION VS. TIME. IM H2S04, 110°. 44 TIME (HR.) FIGURE 18. EFFECT OF PRESSURE. REACTION MODEL VS. TIME. 45 OXYGEN PARTIAL PRESSURE (PSD FIGURE 19. RATE OF DISSOLUTION VS. OXYGEN PRESSURE. INITIAL PARTICLE SIZE -150 • 200 MESH. 46 explain the experimental results and be consistent with the work of others, an adsorption mechanism was postulated. Figure 20 shows a plot of (l/rate) vs I/PQ^. The linearity indicates a good correlation 22 with the Langmuir Adsorption Isotherm. The equation for the Iso• therm is derived from the assumption that the mineral surface consists of a certain number of sites C of which C02 are occupied by oxygen. The rate of desorption is taken to be proportional to C02> or equal to K^CQ^. The rate of adsorption is proportional to the number of empty sites, C-CQ^; and also to the gas pressure. The rate of adsorption is therefore K2P02 (C-CQ^). At equilibrium: K C l o2 K2Po2 £2 If a •= K. '02 a Po2 (C-CQ^) which works out to a P02C (l+-aP02) The rate of dissolution of pyrite"is proportional to the number of sites or occupied by oxygen, Co2I rate equals KCQ^. Substituting the value of Co- from equation 10 gives: K a P02C lf p 9. Rate = KC = a o2 . °2 Y A plot of l/Rate vs l/Po2 is then of the linear formli = mx4-b; with the slope equal to 1/KaC, and the intercept equal to l/KC. The intercept value is useful since it represents the rate at infinite pressure and 47 FIGURE 20. FIT OF DATA T© LANGMUIR ADSORPTION MODEL. 48 hence complete coverage of the surface by oxygen. This is the maximum rate at which the dissolution reaction can proceed. Under the condit ions studied, this rate is 0.0419 molesyhr. which can be normalized for area to give 0.624 moles^m^ hr. using a spherical particle model. Figure 21 shows the rate data plotted versus the square root of the oxygen pressure. The plot was made to test the reaction model of reversible adsorption with dissociation of oxygen. The result is linear but shows a zero rate of reaction at 25 psi oxygen. The fact that work by other investigators shows pyrite dissolution at pressures lower than 25 psi rules out the use of the model. The production of elemental sulfur is increased with decreased oxygen pressure. The ferric/ferrous ratio is also greatly reduced. Figure 22 shows the effect of pressure on iron extraction. Following the Langmuir model, the slope of the curve will become zero as the rate for total coverage of the pyrite surface by oxygen is approached. 10. Temperature Runs were made at 976 psi and lMH^SO^ at five degree intervals from 85° to 130°C. The results are shown in Figure 23. It will be noted that for temperatures over 110°C, oxygen consumption tapers off after about 2 hours. This is explained by the formation of a layer of liquid sulfur on the surface of the particles which reduces the amount of oxygen reaching the pyrite surface. The liquid sulfur layer 4 6 U*W m/es Fe5zAr > ^zMr. W y fibres, 49 FIGURE 21. FIT OF DATA TO REVERSABLE ADSORPTION WITH DISSOCIATION OF OXYGEN MODEL. 50 60 0 200 400 600 800 1000 OXYGEN PARTIAL PRESSURE (PSI) FIGURE 22. IRON EXTRACTION VS. OXYGEN PARTIAL PRESSURE 51 0.10 Temp %elem S ferric/ferrous % Ext I30°C 25.48 81.78 63.22 125 30.57 79.56 64.76 120 23.08 58.39 63.36 11 o° no 115 25.75 24.28 64.97 110 33.09 8.78 83.72 105 33.1 2 12.92 75.17 100 33.76 1 1.39 66.19 105° r26 95 32.97 6.30 56.52 0.08 90 35.80 5.73 43.01 20° r28 65 28.46 5.95 31.73 115° r25 0130° r30 OI25° r29 100° r27 95° r3l 0.06 OXYGEN CONSUMED (MOLE) 90° r!6 0.04 85° r32 0.02 0.00 3 4 TIME (HR.) FIGURE 23. EFFECT OF TEMPERATURE. OXYGEN CONSUMED VS. TIME. IM H2S04, 976 PSI 0^ • 52 forms at temperatures as low as 115°C which indicates that monoclinic 23 sulfur (mp 119°C) is not formed immediately. The presence of the sulfur layer was evident on the residue from the leaches as the particles agglomerated into pellets at the higher temperatures. To test the validity of this sulfur coating hypothesis, a run was made in which a small amount of elemental sulfur was added with the mineral sample and the autoclave heated to 130°C under an inert atmosphere to prevent oxidation and subsequent formation of sulfur from the pyrite. The temperature was then decreased to 110°C and oxygen introduced to the system. ' Figure 24 shows that oxidation was slowed 'cons.i'def-a.bly.- After two hours, the temperature was raised to 130° and the system repre ssurized to 976 psiO^. The figure shows that oxygen consumption increases slightly as would be expected. The parabolic curve at 110° can be explained by pyrite dissolving after the liquid sulfur coating had solidified. Since only 30 to 35% of the pyritic sulfur is in the elemental form after dissolution, the sulfur layer is porous, allowing further attack on the surface. This behavior is contrasted with the liquid sulfur film at the higher temperature which completely coats the particles. The oxygen consumption data were plotted in terms of the geometric rate law for a topochemical dissolution in Figure 25 with good agreement. Rates taken from this plot were then used to construct the Arrhenius plot shown in Figure 26. The slope of the line was used to FIGURE 24. EFFECT OF ELEMENTAL SULFUR. OXYGEN CONSUMPTION VS. TIME. IM H2S04, 976 PSI 0^. 0.5 TIME (HR.) FIGURE 25. TEMPERATURE DATA IN TERMS OF REACTION MODEL. MODEL VS. TIME. FIGURE 26. ARRHENIUS PLOT OF TEMPERATURE DATA. 56 calculate an activation energy of 12. 7 - 0. 7 kcal/mole. This compares favorably with a value of 13. 3 (t 2) kcal obtained by McKay and 13. 1 kcal by Pawlek at lower pressures and acid concentrations. A more precise determination of the activation energy could be obtained by repeating the pressure series in the preceeding section at each temperature. This would allow the calculation of the limiting dissolution rate for complete surface coverage at each temperature. Such a calculation would eliminate any effect due to variation in coverage by adsorbed oxygen and would permit the energy of adsorption to be estimated by subtracting the activation energy obtained in the limiting case from that obtained at lower pressures where the surface coverage is small. The number of experimental runs necessary for such a determination made it impractical for this study but such work could be performed if the need arose. An estimate of the experimental accuracy for the activation energy was made by performing successive linear regressions on the Arrhenius plot data. One data point was left out of the calculation each time. This technique tends to sort out the effect of one poor run. The results showed a value of 12.7, plus 0. 2 and minus 0. 7 kcal. All but one of the calculations fell within a 12. 7^0.2 kcal range. The report• ing of the value as 12. 7 - 0. 7 kcal, therefore, seems well within the possible error. The effect of gas dissolution enthalpy on the activation energy was checked and found to be negligible. This factor is introduced by the experimental design which controlled the concentration of dissolved oxygen in the leach solution by regulating the oxygen pressure. For the reaction: 10. 02(g) >• 02(aq) the dissolved oxygen concentration is related to the partial pressure by the solubility constant, c<- . P L°2] = <* o 2 The solubility term is composed of two parts; a standard solubility, <^ ; — A Hs /RT and an enthalpy of dissolution term e , such that: _ T - &Hs/RT L°2] = *oe Po2 If the rate of the pyrite dissolution reaction is of the form: EA RT Rate = Kr [02] = K° e" / [Q2] . then the overall equation is of the form: o -Ea/RT _ AHs/RT Rate = K e e P r o o2 ° . „ - (Eat A Hs)/RT K P e = r °^o o2 1 -(Eat A Hs) 11. Rate = K e r 58 The effect of the enthalpy term was checked by reviewing solubility 24 data for oxygen at high pressures. As shown in Fig. 27, the solu• bility is nearly constant for the range 85 to 130°C and the enthalpy correction term, therefore, goes to zero. Increasing the temperature yields little change in ferric/ ferrous ratios and although the elemental sulfur data are scattered, no consistent trend is found. This is in contrast to Pawlek's trend toward increased elemental sulfur at lower temperatures. The trend toward increasing ferric/ferrous ratios shown by the data is misleading; the extent of extraction is the major cause of the higher ratios and was mentioned in the reproducability section. Two short runs were made, one at 120° and the other at 130°C to get a more accurate picture of the effect. By leaching samples for one hour at 1 30?C, approximately the same extraction is obtained as for leaching 6 hours at 90°C. The ferric/ ferrous ratios are then seen to be nearly constant when they are com• pared using similar extractions. The results are shown in Fig. 28. The elemental sulfur production in these short runs is seen to be increased at the shorter times. The ferric/ferrous ratio is probably the best indicator of the mixed electrode potential of the system during leaching. A ratio of 25 about 5 corresponds to an EMF of +0.874 volts, well above the 26 observed rest potential for pyrite (0.620) and in a region of the Eh-pH diagram (Fig. 2) where elemental sulfur is unstable. 59 FIGURE 27. SOLUBILITY OF OXYGEN IN WATER. 60 80 6 0 < to ID S 40 a: LU Apparent Trend a: LU 20 Real Trend —o , r%5 r 36 _J 80 90 100 110 120 130 TEMPERATURE (°C.) FIGURE 28. EFFECT OF TEMPERATURE ON THE FERRIC/FERROUS RATIO. The experimental data which show consistent 30-35% pro• duction of elemental sulfur are not consistent with thermodynamic equilibrium at this potential. This fact shows that the sulfate-elemental sulfur redox system is highly irreversible.. 62 • CONCLUSIONS 1. The dissolution of pyrite in sulfuric acid can be accurately 1/3 modelled using the l-(l-c<) shrinking core approach. The sensitivity of this model is excellent in correlating data up to oc values of 0. 977. 2. The dependence of the pyrite dissolution rate on oxygen con• sumption is a function of the surface covered by oxygen. The coverage can be modelled by the Langmuir Adsorption Isotherm. The limiting dissolution rate at extreme pressure can be cal• culated for a given temperature and set of variables. A rate of 0.624 moles/m^ hr. is the limiting value obtained at 110°C for the conditions tested, assuming a spherical particle geometry. 3. An Arrhenius plot, using temperature data from 85 to 130°C, gives the activation energy for pyrite dissolution as 12.7 - 0. 7 kcal/mole. The initial dissolution rate increases with increased temperature. Subsequent formation of a liquid sulfur film on the particles at temperatures over 110°C reduced the rate. 4. ~ The proposed overall reaction is effective in explaining the formation of dissolution products. It is found that for initial acid concentrations less than 0. 1 MH2SO4, sulfuric acid is formed from pyritic suLfur. At acid concentrations greater than 1 molar, dissolution consumes acid. The transition point is about'd^M acid. The pyrite dissolution rate decreases with increased acid concentration above 0. 1 Mt^SO^. The decrease is on the order of 25% in going from 0.1 to 3 .MFL^SO^. Neutral salts of copper, zinc and nickel inhibit pyrite dissolu• tion and at, the same time catalyze the oxidation of ferrous iroi to ferric. The inhibition is greatest with the copper salts amounting to a 25% reduction in iron extraction. The ferric/ ferrous ratio is increased by a factor of two by the copper addition. Addition of copper, zinc, and nickel sulfides has little effect on pyrite dissolution. The dissolution rate of the sulfides themselves is increased as seen by increased oxygen con• sumption which returns to the typical pyrite reaction curve when the additional minerals have been dissolved. Pyrite dissolution is affected by geologic factors. Variation in dissolution rate of about 30% is found between the two major pyrite groupings studied. One factor in this effect is a variation in breakage patterns resulting in increased surface area on the more rapidLy dissolving minerals. The distribution of reaction products between ferric and ferrous ions is greatly effected by the extent of reaction. Varying iron extraction from 40 to 60% results in up to a 15-fold change in the ferric/ferrous ratio. Other variables which affect the ratio are acid concentration (0. 01 to 3M, 1.5 fold decrease) neutral salts (1 MCuSO^, 2-fold increase), and oxygen partial pressure (976 to 176 psi, 5-fold decrease). The percentage of elemental sulfur formed in the dissolution is found to increase with: increased acid concentration (from 11.09 to 36.3% in going from 0.01M to 3M acid); reduced oxygen partial pressure (from 33.3 to 43.2% in going from 976 to 176 psi O^)', and addition of neutral salts (33 to 36% with addition of copper sulfate). Increases in pulp density and (/t«creases in particle size reduce extraction rates slightly. The non-linearity of these effects is possibly attributed to a concentration effect on the homogeneous ferrous to ferric reaction. REFERENCES 1. Dana, Manual of Minerology, 18th Edition, p. 202, (1971). 2. "Sulfur and Pyrites", Minerals Yearbook, U.S. Bureau of Mines, (1967). 3. Powder Diffraction File, American Society for Testing Materials, (1960). 4. W. Latimer, Oxidation Potentials, 2nd Edition, p. 221, (1952). 5. P. Toulmin, III and J.P. Barton, Geochim. cosmochim. Acta 28, 641, (1964). 6. K. Kelley, and King Bulletin 592, Bureau of Mine s U.S.A., Washington, D.C, (1961). 7. J. Pemsler, Technical Report 260, Ledgemont Publication, (1971). 8. D.D. Wagman, et al. , N.B.S. Technical Note 270-3 and 270-4, Washington, D.C. , (1968, 1969). 9. R.J. Biernat and R.G. Robins, Electrochimica Acta. , 1972, vol. 17, p. 1279. 10. K. Semrau, "Sulfur Oxides Control and Metallurgical Technology", Journal of Metals, Vol. 23, p. 41, (1971). 11. F.A. Forward, and J. Halpern, AIME Trans. , 1955, Vol. 203, p. 463, Journal of Metals, March, 1955. 12. J. Halpern, and F. A. Forward, Trans. Inst. Min. Met. (London), 1957, Vol. 66. p. 181. 13. D.R. McKay, and J. Halpern, Trans. Met. Soc. AIME 212, p. 301, 1958. 14. J. Gerlach, H. Hahne, and F. Pawlek, Z. Erzbergbau Metallhuttenw, 19( 2), p. 66, ( 1966). 15. E, Peters, AIME International Symposium of Hydrometallurgy 1973, P- 224. 16. Swanson, et al. , U. S. Department of Commerce, NBS Circular 539, Vol. V~. 66 17. Analysis, Can. Test Ltd. , File No. 7647A. 18. J.H. Perry, Chemical Engineers Handbook, Third Edition (1950), p. 278. 19. Standard Methods of Chemical Analysis, 6th Edition, Vol. 1, p. 544. 20. G; Nickless, Inorganic Sulfur Chemistry Elsevier Co., 1968, p. 524. 21. F. Lowen, M.A. Sc. Thesis "Pressure Leaching of Copper Sulfides in Perchloric Acid Solutions", U.B.C. 1967. 22. A. Adamson, The Physical Chemistry of Surfaces, 2nd Edition, 1967, p. 570. 23. W.N. Tuller, The Sulfur Data Book, p. 9(1954). 24. H.A. Pray, C.E. Schweikert, and B.H. Minnich, Industrial and Engineering Chemistry, Vol. 44, No. 5, P. 1146 (1952). 25. CRC Handbook of Chemistry and Physics, 52nd Edition, P. Dill, (1972). 26. E. Peters and H. Majima, Canadian Metallurgical Quarterly, Vol. 7, No. 3, p. 111. APPENDIX ACTUAL EXPERIMENTAL DATA Initial Residue 1 ff 1 Tot. Soln. Vol. JtFe J Fe Moles C>2 Run Pyrite Size Acid PQ2 Temp.°C Weight Weight Res.%Fe 11-09 InitiaSul. l run150/20s on 0whic 1hM techniqu976 e wa11s 0perfected - 5.000. No0 result1.461s apply1 .24 . 82 .153 1. 28 12. 52 . 0937 it 11 it it 11 Nor. 5.0010 2.1332 39. 08 .139 1. 28 11.09 . 0753 12 Leak in system. it II ti 11 13 . Nag.. 5.0000 t- 1. 7237 32. 04 . 138 1. 23 11.80 .0719 tt II ti it 14 War. 5.0026 2.0629 32. 85 . 135 1. 39 11.91 . 0777 n it ti 15 Kim. ti 5.0000 . 1.4400 23.67 .130 1. 28 15.09. . 0920 it ti tt 16 Sul. 90 5.0006 3.1422 41. 0 1 . 142 1. 03 6.93 . 0465 it ti it 17 Nor. 11 5.0000 3.6543 43. 10 .135 0. 77 5. 34 . 0333 18 Sul. 325/400 " " 110 5.0012 1.2456 17. 86 . 147 1. 44 13. 70 . 0953 it 19 200/270 " " it 5.0010 1.1483 15. 36 . 130 ' 1.44 15.96 .0979 11 11 20 270/325 " " 5.0005 1.2873 19.05 . 142 1.44 14. 22 . 0953 11 11 21 Nor. 200/270 " 5.0000 1.5751 30. 11 . 136 1.49 13. 19 . 0895 11 22 War. 200/270 " " 5.0000 1.6120 26.84 . 120 1. 59 15.17 . 0895 tt ti 23 270/325 " " 5.0000 2.4060 23.11 .128 1. 54 15.45 . 0961 it II 24 375/400 " " 5.0012 1.4048 21.49 . 126 1. 64 15.65 . 0970 25 Sul. 150/200 " " ' 115 5.0012 2.1747 34. 89 . 105 0. 56 14. 16 . 0790 ON ACTUAL EXPERIMENTAL DATA (Continued) Initial Residue .£ ++ Tot Run Pyrite Size Acid P02 Temp. °C Weight Weight Res. %Fe Soln. Vol. JiFe fFe V Moles O 2 26 Sul. 150/200 IM 976 105 5.0000 1.8875 32. 84 . 126 0. 98 13.65 . 0843 27 " n II 100 5.0000 2.2670 34. 00 . 130 0. 94 11.65 .0736 ii 28 II II 120 • 5.0016 2.2105 36.93 . 074 0. 33 19.60 . 0793 II II it 29 125 5.0017 2.2889. 36. 82 . 080 0. 23 18. 53 .0750 ti it ti 30 130 5.0010 2.2841 36. 72 . 076 • 0. 23 19.04 . 0773 II ti it 31 95 5.0000 2.7183 35.69 . 144 1. 23 8.98 .0632 32 n n II 85 5.0004 3.6679 42. 57 . 122 0. 82 5.95 .0372 II 33 0.1M " 110 5.0005 1.7052 31. 21 . 124 1.49 14.37 . .0996 II 11 34 0.01M " 5.0000 2.1704 31. 99 .119 1. 03 13. 75 .0995 II 35 IM •* 130 5.0007 3.2456 36. 51 . 110 1. 33 9. 03 . 0478 II 11 II 36 120 5.0000 3.7106 43. 64 . 106 1. 39 6.72 . 0351 • it II 37 " 110 10.0000 6.0568 40.49 .118 2.82 18. 53 . 1087 38 it n 3M " 5.0016 2.2654 31.92 . 114 1.49 14.01 . 0753 ti II 39 IM " 2.5000 1.2496 40.00 . 094 1.23 7.24 .0342 40 ii II it " 110/130 7.0069* 6.4565 28. 95 . 066 0. 26 3. 90 . 0147 41 II lM+Zn " 110 5.0034 2.8372 36. 80 .124 0. 92 9. 70 . 0586 42 Experimental error. *' 5. 0026g pyrite + 2. 0043g sulfur. ON oo i i ACTUAL EXPERIMENTAL DATA (Continued) Initial Residue f Tot Run Pyrite Size Acid P Temp. °C Weight Weight Res. %Fe Soln. Vol. k Fe Fe - Moles O-, °2 ** 43 Sul. fZr 150/200 1M . 976 110 5.0000 2.6357 38. 27 . 207 0. 96 6. 16 . 0618 44 II 11 1M4-Ni it II 5.0038 2.8711 4.0. 33 .110 0.96 10.98 . 0577 45 .1! 11 lMt-Cu 11 tt 5.0005 3.4000 44.47 .120 0. 67 . 7. 75 . 0443 46 Expe r ime nta 1 error. 47 it it lM it it 5. 0000 .9063 7.77 . 110 0. 92 20.4 . 1108 48 it rt ti II tt 5.55*1*3* 2. 8852 36.93 .121 1.49 . 11.55 . 0644 **** it 11 49 it ii it 5.5013 2.7212 37. 90 . 100 1. 33 12. 88 . 0644 ***** 50 it it it it it .5.5083 2.7617 37. 30 . 101 1. 95 13. 65 .0711 51 it ti it 676 I'• 5.0004 3.0347 39.78 .094 2.41 11.70 .0497 52 II it .- ti 476 ll 5.0000 •3.4064 42.68 .113 2.08 7.95 . 0406 53 tt ti tt 326 n 5. 0000 3.6712 44.75 . 120 2. 72 6. 16 .0317 54 tt tt it 176 II 5.0000 4. 1131 43.84 . 110 2. 36 4.46 . 0205 ** Zr foil added, did not affect reaction. *** 5. OOOOg pyrite ± 0. 5513g ZnS. **** 5. 0000g pyrite -l- 0.5013g CuS. ***** 5.0075g pyrite +• 0. 5008g NiSFeS.