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The oxygen pressure leaching of pyrite in sulfuric acid Bailey, Leonard Keith 1974

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T H E O X Y G E N PRESSURE L E A C H I N G O F P Y R I T E IN SULFURIC ACID  by  LEONARD KEITH BAILEY B . S . University of Utah, 1973  A THESIS SUBMITTED IN P A R T I A L F U L F I L M E N T O F T H E REQUIREMENTS FOR T H E D E G R E E O F M A S T E R O F A P P L I E D SCIENCE in the Department of Metallurgy  We accept this thesis as conforming to the required standard  T H E UNIVERSITY O F BRITISH COLUMBIA December,  1974  In p r e s e n t i n g t h i s  thesis  an advanced degree at  further  agree  fulfilment  of  the  requirements  the U n i v e r s i t y of B r i t i s h Columbia, I agree  the L i b r a r y s h a l l make i t I  in p a r t i a l  freely  available  for  this  thesis  f o r s c h o l a r l y purposes may be granted by the Head of my Department  of  this thesis for  It  of  financial  gain s h a l l not  Metallurgy  The U n i v e r s i t y o f B r i t i s h Columbia Vancouver 8, Canada  Date  fi  J a n u a r y 27, 1975  or  i s understood that copying or p u b l i c a t i o n  written permission.  Department  that  r e f e r e n c e and study.  t h a t p e r m i s s i o n f o r e x t e n s i v e copying o f  by h i s r e p r e s e n t a t i v e s .  for  be allowed without my  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 -<X)  ). The dependence of the reaction rate on  oxygen pressure has been modelled using a Langmuir Adsorption Isotherm. A limiting dissolution rate for total adsorption of G. 624 moles/ 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 f e r r i c 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 i n i t i a l a c i d concentrations above O . 4 M H 2 S O 4 .  F o r concentrations  below this Level, the r e a c t i o n produces a c i d .  A n a c t i v a t i o n energy of 12. 7 t 0. 7 k c a l / m o l e was o b s e r v e d for the d i s s o l u t i o n r e a c t i o n .  iii  T A B L E OF CONTENTS Page ABSTRACT  ii  T A B L E OF CONTENTS  iv  LIST O F FIGURES  vi  LIST OF T A B L E S  vii  ACKNOWLEDGEMENTS  viii  A. INTRODUCTION 1. General Z. Occurrence 3. Structure 4. Thermodynamics 5. Use 6. Previous work 7. Scope of this work  1 1 1 1 3 3 6 8  B. E X P E R I M E N T A L 1. Materials Z. Apparatus 3. Procedure 4. Analysis (1) Oxygen Consumption (2) Iron Analysis (3) Sulfur Species C. R E S U L T S AND DISCUSSION 1. Reproducibility Z. Variation between P y r i t e s 3. Particle size 4. Pulp density 5. A c i d concentration 6. Effect of sulfide minerals 7. Effect of neutral salts 8. Reaction model 9. Oxygen pressure  10 10 11 15 16 16 16 18 20 20 22 25 30 33 36 38 . 38 41  iv  T A B L E O F C O N T E N T S (Continued) Page C.  R E S U L T S AND DISCUSSION (Continued) 10.  Temperature  48  CONCLUSIONS  62  REFERENCES  65  APPENDIX: Actual Experimental Data  67  v  LIST OF  FIGURES  Figure 1 2 3 4 5 6 7 8 9 10 11 12 12A 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28  Page P y r i t e structure Potential - pH diagram for the iron-watersulfur system Oxygen removed from system vs recorder ..... Reproducibility .. Comparison of pyrites Comparison of unleached pyrites (photographs). . Effect of particle size for Wards pyrite Effect of particle size for Sullivan pyrite Rate of oxygen consumption vs surface area .... Effect of pulp density Iron extraction vs pulp density Acid effect Acid consumption vs log initial acid concentration Effect of sulfide minerals Effect of neutral salts Comparison of reaction models ............... P a r t i c l e s of Sullivan pyrite after leaching (photographs) Effect of pressure, oxygen consumption vs time *. • Effect of pressure, reaction model vs time .... Rate of dissolution vs oxygen pressure . i F i t of data to Langmuir Adsorption Model F i t of data to reversible adsorption with dissociation of oxygen model . Iron extraction vs oxygen partial pressure Effect of temperature, oxygen consumed vs time . , Effect of elemental sulfur Temperature data in terms of reaction model ... Arrhenius plot of temperature data Solubility of oxygen in water Effect of temperature on the f e r r i c / f e r r o u s ratio .. .  vi  2 5 17 21 23 24 26 27 29 31 32 34 35 37 39 40 42 43 44 45 47 49 50 51 53 54 55 59 60  LIST O F T A B L E S  Table  I  II  Page  Thermodynamic data for pyrite  4  Analysis of iron and sulfur in pyrites  12  III  Spectroscopic analysis of pyrites  13  IV  Comparison of X - r a y diffraction patterns for pyrites with a standard pattern . .  vii  14  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 f i r e . In the days of the ancient Greeks, the m i n e r a l 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 m i l l i o n tons annually; the U.S.S.R., with an output of over 3.4 m i l l i o n tons; and Japanese sources of over 4.4 m i l l i o n tons per year. North American production is led by the United States where deposits in Tennessee, Pennsylvania and several western states produce 0.8 m i l l i o n tons yearly. Canadian sources include: Noranda, Normetal Mines, and the Anaconda (Canada) Co. Ltd., producing about 0.3 m i l l i o n tons per year. 3. Structure Pyrite has a structure similar to that of N a C l , (see F i g . 1 ) . The iron atoms take the sodium positions and S 2 groups replace the chlorine.  The S 2 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  a  Q  o 3 distance of 5.417A.  P y r i t e is distinguished f r o m the other F e S £  form, m a r c a s i t e , by the orthorhombic m a r c a s i t e structure.  The  groups in m a r c a s i t e have a net orientation which produces the o r t h o r hombic c r y s t a l structure and makes the m i n e r a l m o r e subject to c h e m i c a l attack than pyrite.  4.  Thermodynamics The thermodynamic p r o p e r t i e s of pyrite are shown in T a b l e  I.  The r e p o r t e d values a r e i n good agreement.  The potential -pH  9 d i a g r a m for the i r o n - w a t e r - s u l f u r system  is shown in F i g u r e 2.  5. Use Since pyrite is often found in m a s s i v e deposits, many p r o c e s s e s have been designed to extract the i r o n and sulfur values for c o m m e r c i a l use.  Although the m i n e r a l i s n e a r l y 4 7 % iron, its-use as a  feedstock for iron production is limited by the complexities of r e m o v i n g the sulfur.  Some use is made, however, of the i r o n oxides r e m a i n i n g  after the production of sulfuric acid.  Such p r o c e s s e s r e q u i r e the r o a s t -  ing of pyrite to liberate sulfur as S O 2 and yield an i r o n oxide calcine. The SG"2 is converted to sulfuric acid by passing the gas over a catalyst bed.  While this p r o c e s s is effective, its use in North A m e r i c a is being  reduced by competition f r o m other sulfuric acid-producing methods. E n v i r o n m e n t a l pollution controls, p a r t i c u l a r l y on the petroleum and nonferrous metal industries, have resulted in new p r o c e s s e s for r e c o v e r i n g sulfur f r o m stack gasses. ^  Increased production of a c i d by these methods  TABLE I THERMODYNAMIC D A T A FOR P Y R I T E  A G°  Source: Latimer  A So  f  -39.84kcal  4  Toulmin & Barton KelLey and K i n g  5  12.7cal/Deg.  -38.3 12.7  6  7 Pemsler  -38.1  Note: The difference in the free energy terms is actually not as great as reported. The measured 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 subtracted from that of pyrite, a better idea of the accuracy is given. Latimer -39. 84 - (-20. 3) = -19. 54 kcal Toulmin -38.30 - (-18. 8 5 ) = -19.45 Pemsler-38. 10 - (-18,85) = -19.25 18  The net difference in free energies measured is Less than - 0 . 1 5 kcal.  5  F I G U R E 2.  P O T E N T I A L / P H DIAGRAM FOR THE I R O N WATER-SULFUR S Y S T E M 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 provi11,12 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), 0  2  pressure (0-4 atrn), temp. (100 to 1 3 0 ° 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.  FeS  +•  2  20  — F e S 0 +• S ° .  2  4  Formation of ferric ion is accounted for by the reaction: 2.  4 F e S 0 + O +• 2 H S 0 — 4 F e ( S 0 ) 4  z  2  4  4  1  5  + 2H O. z  Several other reactions were found to play only a minor role in the dissolution.  7  The o v e r a l l rate of reaction was found to be p r o p o r t i o n a l to the pyrite surface area, and to the oxygen p r e s s u r e .  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 m e c h a n i s m postulated involves c h e m i s o r p t i o n of oxygen on the pyrite surface, followed by a slower attack of another O2 molecule ori the O2 c o v e r e d m i n e r a l site. 3.  FeS 1  2  •• O  z  -l- O o ( a q )  >  [FeS  2  * 20 ]  - —>  2  slow  FeS0  4  +• S°  J  An activation energy of 13. 3 £ 2 kcal/mole was o b s e r v e d for the pyrite dissolution.  F u r t h e r work on the system was done by Pawlek and co-work14 ers.  The experimental p r e s s u r e was r a i s e d to 16 atm,  temperature effect was  studied f r o m 60 to 160°C.  The activation energy  for pyrite oxidation was observed to be 13. 1 k c a l / m o l e . Pawlek r a i s e d a c i d concentrations up to 0 . 4 M H S O 2  and the  4  In his work,  with the r e s u l t that  he was able to see a slight i n c r e a s e in the f e r r i c to f e r r o u s ratio as the concentration i n c r e a s e d . studied lower a c i d levels. acid might be generated:  T h i s effect was not r e p o r t e d by McKay,  who  Pawlek also proposed a r e a c t i o n by which  8 4.  FeS  2  +• 8 F e  + 3  f 4H O z  9Fe  f 2  f SOj + S° f 8 H . f  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 other minerals.  has studied the effect of pyrite on the dissolution of 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. control on the reaction.  This phenomenon is explained by a galvanic  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 exchanged 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:  5.  FeS f ( | f | 2  +xFe  + H  y H- - | ) 0 x  2  + (2 + x - 2y) H*"  >(l  - x) F e  + f  f (2 - y ) S ° +• y S O j > (1 - y V | ) H 0 . 2  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.  1.  EXPERIMENTAL  Materials F i v e different p y r i t e s were obtained and ground b y hand -  u s i n g a m o r t a r a n d p e s t l e - to the d e s i r e d s c r e e n s i z e s . m i n e r a l w a s g r o u n d i n i t i a l l y f o r a l l the t e s t s c o n d u c t e d .  Enough The pyrites  originated from:  The  (a)  Cominco Mines - Sullivan Mine,  (b)  K i m b e r l e y (also a S u l l i v a n M i n e  (c)  Noranda Mines - Quebec,  (d)  Ward's Scientific Co.-, - Colorado,  (e)  Japan,  sample),  and  through P r o f e s s o r Nagai.  S u l l i v a n and K i m b e r l e y s a m p l e s w e r e p r o v i d e d as  (-65 + ZOOmesh), w h i l e t h o s e f r o m W a r d ' s a n d N o r a n d a specimens.  powders were  massive  T h e N a g a i s a m p l e c o n s i s t e d of four or five l a r g e  single  crystals which were crushed and sized.  A l l the s c r e e n e d s a m p l e s w e r e w a s h e d to r e m o v e a n y f i n e s w h i c h m i g h t h a v e a d h e r e d to the s u r f a c e ,  then a i r - d r i e d .  f r a c t i o n s w e r e u s e d i n the e x p e r i m e n t a l w o r k : (a)  - 1 5 0 +- ZOO m e s h ( T y l e r s c r e e n  (b)  -ZOO +• Z 7 0 ,  (c)  - Z 7 0 f 3Z5,  (d)  -325' + 400  sizes),  Four  size  11 Samples of the ground minerals (rl 50 +• 200 mesh fractions) were sent to Can Test Ltd. for chemical analysis. shown in Table II.  The results are  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 - r a y diffraction.  A l l 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. ^ A l l chemicals used were reagent grade and demineralized water was used throughout the tests.  2.  Apparatus A l l tests were made in a 106.5 ml zirconium shaking autoclave  with a 15.0 m l 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 4 5 ° 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  T A B L E II ANALYSIS O F 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  T A B L E III SPECTROSCOPIC A N A L Y S I S  Kimberley  Sullivan  17  O F PYRITES  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  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  Mn  ;  Not Detected: Antimony, Barium, Be ryllium, Bismuth, Boron, Cadmium, Gallium, Niobium, Potassium, Sodium, Strontium, Tantalum, Thorium, Tungsten, Uranium, and Vanadium.  T A B L E IV COMPARISON O F X - R A Y DIFFRACTION P A T T E R N FOR PYRITES WITH A STANDARD P A T T E R N  Line  Observed d Value  1 2 3 4 5 6 7 8 9 10 11 12 13.  3.1057 2.6913 2.4114 2.2000 1.9093 1.6289 1.5602 1.4993 1.4453 1.3130 1.2752 1.2431 1.2113 1.2112 1.1822 1.1822 1.1551 1.1551  14 15  Reported by Swanson 3.128 2. 709 2.423 2.2118 1.9155 1.6332 1.5640 1.5025 1.4448  -  -  1.2427 1.2113 1,1823 1.1548  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 r e 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 C o r p . , 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 Sargeant 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 autoclave, the solution added, and the autoclave sealed. were then started.  Shaking and heating  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 6 0 ° C took about 10 m i n . , 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:- A s stated, the oxygen p r e s s u r e drop  throughout a r u n was m o n i t o r e d by a t r a n s d u c e r coupled w i t h a r e c o r d e r . 18 C a l i b r a t i o n of this s y s t e m u s i n g steam tables v a r y l i n e a r l y with p r e s s u r e .  showed the r e c o r d e r to  R a t h e r than c a l c u l a t e the n u m b e r of m o l e s  of oxygen c o n s u m e d by u s i n g a gas l a w - p r e s s u r e r e l a t i o n , a s e r i e s of tests w e r e conducted i n w h i c h a s p e c i f i c volume of gas ( m e a s u r e d by a gas burette) was r e m o v e d f r o m the autoclave and the r e s u l t i n g p r e s s u r e drop r e c o r d e d .  These m e a s u r e m e n t s show the r e l a t i o n to be l i n e a r  though different at each t e m p e r a t u r e  (see F i g . 3). In e a c h case a slope,  m o l e s / c h a r t d i v i s i o n , was obtained and these values w e r e u s e d i n i n t e r p r e t i n g the r e s u l t s of the l e a c h tests.  A l l tests w e r e r u n under  constant volume conditions, a l l o w i n g the p r e s s u r e to d e c r e a s e o v e r the course of a run. 19 (Z) I r o n A n a l y s i s : - The to determine  s t a n d a r d c e r i c sulfate t i t r a t i o n  the f e r r o u s i r o n present i n solutions.  was  used  A value for total  i r o n was obtained by r e d u c i n g the f e r r i c ions i n the solution with stannous c h l o r i d e then repeating the f e r r o u s d e t e r m i n a t i o n . concentration was  determined  F e r r i c ion  by d i f f e r e n c e . F e r r o i n i n d i c a t o r (1,10  Ortho phenanthroline F e r r o u s Sulphate . 0Z5M) was used as an end point in the titratrations. Iron i n the l e a c h r e s i d u e s was d i s s o l v e d with aqua regia.  It should be noted that h y d r o l y s i s of the i r o n was not a p r o b l e m .  Iron oxides w e r e f o r m e d i n only two runs i n w h i c h i n i t i a l a c i d c o n c e n t r a tions w e r e v e r y low.  The r e s u l t i n g solution was then b o i l e d to r e m o v e  17  FIGURE  3.  OXYGEN  REMOVED  FROM S Y S T E M  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 analysis. 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 f e r r i c 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 f e r r i c sulfate solution. An i n -  crease in the ferrous concentration of the solution was noted showing that f e r r i c 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 percentage points.  The overall reaction in general  form for the dissolution is  equation 5. 5. F e S  2  - y + - x )O 4- (2 f x-2y)H  > (l-x)Fe  z  +- (2-y)S° + ySO~  f (1-y + jx)  + xFe  HO z  Hence, knowing the iron analyses, the sulfur species concentrations are calculated using the equations:  19  6.  2 1 moles SO^ = j (moles O-, (moles F e S consumed) 2  _  7.  (moles Fe^^~ produced)  moles S° = 2 (moles FeS2 consumed) - moles S 0  4  produced  20 C.  R E S U L T S 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 preweighed 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 a l l tests made, the reproducibility of the experimental runs is taken to be quite adequate. There i s , how-  Zl  0.12 r47  0.10  0.08 OXYGEN CONSUMED 0 . 0 6 (MOLE) Time  0.04  % elem S ferric/ferrous  3hr 6 10  0.02  0.00  33.33 33.09 32.54  4 • 6 TIME (HR.)  542 8.78 2 1.17  8  %Ext  55.73 83.72 98.08  10  FIGURE 4. REPRODUCIBILITY, OXYGEN CONSUMPTION VS. TIME. IM H S 0 , 1.10°, 9 7 6 PSI 0 . 2  4  2  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 i s noted that as the reaction proceeds further to completion the f e r r i c / f e r r o u s ratio increases sharply.  2.  Variation between P y r i t e s Five different pyrites were used in this study to determine  if geologic differences in the pyrites had an effect on the leaching parameters. 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 ments at 976 psi  O2,  or may  I M H 2 S O 4  and  not be significant. Results of experi110°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. Photographs 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 H SC^, 110°, 976 PSI 0 . 2  2  24  FIGURE  6.  SULLIVAN  150 X 2 0 0  MESH  WARDS  150 X 200  MESH  COMPARISON  OF U N L E A C H E D P Y R I T E S .  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 s i m i l a r i t y in leaching behavior and may account for the deviation from the other three samples. Whether the stoichiometry 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 f e r r i c and ferrous iron is quite scattered.  This can be  partially explained by the extent to which the reaction was allowed to continue. A l l 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 f e r r i c / 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 a l l run at 976 psi O2,  I M H 7 S O 4 , 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 s t i l l different due to the different breakage patterns. A comparison 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  Mesh Size %elem S  0.02  0.00  E F F E C T  I  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  i  i  ••1 •  3  4  O F P A R T I C L E  W A R D S  % Ext  325/400  TIME F I G U R E ? ,  ferric/ferrous  P Y R I T E .  5  i  6  (HR.) SIZE ON OXYGEN CONSUMPTION I  M H^QjJIOt 976 PSI 0^  27  TIME  FIGURES.  (HR.)  E F F E C T OF PARTICLE SIZE ON OXYGEN CONSUMPTION FOR SULLIVAN PYRITE. IM H SQ|, 1 1 0 ° , 9 7 6 PSI 0 2 . 2  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. is also noted in the subsequent pulp density section.  The effect  The heterogeneous  dissolution reaction shouldbe affected linearly by changes in surface arjea; the ]!i:<J<*M:ogeneous reaction however, shows no such effect since it will be dependent only on solution concentration. Reducing the particle size allows more surface for attack which results in increased solution concentrations. This concentration increase raises the homogeneous reaction rate, causing the non-linearity of oxygen consumption with surface area.  29  0.04h  AREA  FIGURE 9.  RATE OF OXYGEN  \M  H S0 2  4 o  1 1 0 ° , 97S  (CM /G) 2  CONSUMPTION VS. SURFACE AREA. PSI  0 . 2  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  solids. Runs were made at 976 psi oxygen pressure, 110°C  20%  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 F i g . 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 dissolution of pyrite, and the other, the homogeneous oxidation of ferrous ion to the f e r r i c 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 H S 0 , IIO* 9 7 6 PS I 0 . 2  4  2  FIGURE I L  IRON  EXTRACTION VS. PULP DENSITY.  No real trend was seen in the formation of elemental sulfur but a slight increase in the f e r r i c / f e r r o u s ratio is noted on increasing the pulp density.  5. A c i d 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 i n 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 concentration is 0. 01 and 0. lMH^SO.^. At higher acid levels, however, the dissolution reaction is an acid consumer. Figure 12A compares acid consump-  34  0.10  1 M r 10  0.08 3M  r  38  0.06| OXYGEN  CONSUMED (MOLE)  0.04  0.02  H2S04  %elem S  ferric/ferrous  % Est  O.OIM 0.10 1.00 3.00  11.09 20.06 33.09 36.36  12.35 8.64 8.78 8.40  66.16 77.88 83.72 69.77  •  • «  I  2  3  4  •  5  6  ME (HR.) FIGURE 12. ACID EFFECT, OXYGEN CONSUMPTION VS. TIME, 110°, 9 7 6 PSI 0 . Tl  2  1  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 lower acid concentrations.  .  studied the  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 i s  compared with a standard run (pyrite;only) in F i g . 13.  The galvanic  effect is shown by the higher initial slope, corresponding to the consumption 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 f e r r i c / f e r r o u s ratio shows an increased production of f e r r i c ion with the other mineral additions.  0.08  Nir50 .Znr48 'Cur49 • None r43  0.06  OXYGEN CONSUMED (MOLE) 0.641  0.02  Sulfide Cu  0.00  F I G U R E 13.  %elem S 31.17  Ni Zn  27.94  OnlyFe  33.33  TIME  2 (HR.)  I  3800  ferric/ferrous 8.68 6.00 6.75 5.42  E F F E C T OF S U L F I D E M I N E R A L S . IM H S 0 , 110°, 976 PSI 0 . 2  4  2  % Ext 56.29 55.55 60.29 55.73  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 i f the metal ions themselves had an effect on the system. Solutions 1M in the neutral salts CuSO^., NiSO^. and ZnSO^ were made. A c i d concentration was kept at  A l l three salts seemed to slow the dissolution of  I M H 2 S O 4 .  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 analytical accuracy. salt additions.  Larger f e r r i c / f e r r o u s ratios are also evident with the The catalytic effect of the cupric ion on the production of •1-2  f e r r i c 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 e r r o r . 8. Reaction model In testing reproducibility of oxygen consumption data, a run was made in which over 9 8 % of the pyrite was dissolved. oxygen consumed to this point as 100%, used to calculate  Taking the  consumption during the run was  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 S A L T S . OXYGEN CONSUMPTION V S . TIME. IM SALT, I M H S 0 , 110? 976 PSI Og. 2  4  0.9  TIME  FIGURE  15.  (HR.)  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 ) model.  shrinking core  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.  P A R T I C L E S OF SULLIVAN PYRITE A F T E R L E A C H I N G . 150X200 MESH, 110°, I M H S 0 , 976 PSI 0 . 2  2  4  43  Pressure  0.08  % elem S 33.33  976 psi 0 2 676 476 326 176  39.09 38.82 41.28 43.18  ferric/ferrous  % Ext  5.42 3.85 2.82 1.26 0.89  55.73 4a07 39.26 32.30 21.44  0.06 OXYGEN CONSUMED (MOLE) 0.04  0.02  O.'OQ T I M E (HR.)  FIGURE  17. EFFECT OF PRESSURE. OXYGEN VS. TIME. IM H S 0 , 110°. 2  4  CONSUMPTION  44  TIME (HR.)  FIGURE 18.  EFFECT OF P R E S S U R E .  REACTION MODEL VS. TIME.  45  OXYGEN PARTIAL PRESSURE (PSD  FIGURE  19.  R A T E OF DISSOLUTION VS. OXYGEN PRESSURE. INITIAL PARTICLE SIZE -150 • 2 0 0 MESH.  46 explain the experimental r e s u l t s and be consistent with the work of others, an adsorption m e c h a n i s m was postulated. plot of ( l / r a t e ) vs I / P Q ^ .  F i g u r e 20 shows a  The linearity indicates a good c o r r e l a t i o n 22  with the L a n g m u i r Adsorption Isotherm.  The equation for the Iso-  therm i s derived f r o m the assumption that the m i n e r a l surface consists of a certain number of sites  C  of which  C  are occupied by oxygen.  0 2  The rate of desorption is taken to be proportional to C K^CQ^.  or equal to  The rate of adsorption is proportional to the number of empty  sites, C - C Q ^ ; and also to the gas p r e s s u r e . therefore K P 2  0 2  (C-CQ^).  The rate of adsorption i s  At equilibrium:  l  K  If a •=  0 2 >  C  o  2  K Po 2  <C-C )  2  02  £ 2  K.  '02  a Po a P  2  0 2  (C-CQ^)  which works out to  C  (l+-aP ) 02  The rate of dissolution of pyrite"is proportional to the number of sites occupied by oxygen, C o I  o  r  2  rate equals K C Q ^ .  Substituting the value  of Co- f r o m equation 10 gives: K a P 9.  Rate  A plot of l/Rate vs l / P o  = KC = . °2  l  f  a  p  0 2  o  C  2  is then of the l i n e a r f o r m l i = mx4-b; with the Y  2  slope equal to 1/KaC, and the intercept equal to l / K C .  The intercept  value is useful since it represents the rate at infinite p r e s s u r e and  47  FIGURE 2 0 .  FIT  OF DATA T© LANGMUIR  MODEL.  ADSORPTION  48  hence complete coverage of the surface by oxygen.  T h i s i s the maximum  rate at which the dissolution r e a c t i o n can proceed.  Under the condit  ions  studied, this rate i s 0.0419 m o l e s y h r . which can be n o r m a l i z e d for a r e a to give 0.624 m o l e s ^ m ^ hr. using a s p h e r i c a l p a r t i c l e model.  F i g u r e 21 shows the rate data plotted versus the square root of the oxygen p r e s s u r e .  The plot was made to test the r e a c t i o n m o d e l  of r e v e r s i b l e adsorption with dissociation of oxygen. but shows a zero rate of r e a c t i o n at 25 p s i oxygen.  The r e s u l t is linear  T h e fact that work  by other investigators shows pyrite dissolution at p r e s s u r e s lower than 25 psi r u l e s out the use of the model.  The production of elemental sulfur i s i n c r e a s e d with decreased oxygen p r e s s u r e . reduced.  The f e r r i c / f e r r o u s ratio i s also greatly  F i g u r e 22 shows the effect of p r e s s u r e on i r o n extraction.  Following the L a n g m u i r model, the slope of the curve w i l l become zero as the rate for total coverage of the pyrite surface by oxygen is approached.  10.  Temperature Runs were made at 976 p s i  intervals f r o m 85° to 130°C.  and l M H ^ S O ^ at five degree  The r e s u l t s are shown in F i g u r e 23. It  w i l l be noted that for temperatures over 110°C, oxygen consumption tapers off after about 2 hours.  T h i s i s explained by the formation of a  layer of liquid sulfur on the surface of the p a r t i c l e s which reduces the amount of oxygen reaching the pyrite surface.  4 6 U*W  m/es  Fe5z  Ar  >  ^ zMr.  W y  The liquid sulfur layer  fibres,  49  FIGURE  21.  FIT WITH  OF  DATA  TO  REVERSABLE  DISSOCIATION  ADSORPTION  O F 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  0.08  Temp %elem S I30°C 25.48 125 30.57 120 23.08 115 25.75 110 33.09 105 33.1 2 100 33.76 95 32.97 90 35.80 65 28.46  ferric/ferrous 81.78 79.56 58.39 24.28 8.78 12.92 1 1.39 6.30 5.73 5.95  % Ext 63.22 64.76 63.36 64.97 83.72 75.17 66.19 56.52 43.01 31.73  11  o° no  105° r26 20° 115° 0130° OI25° 100°  r28 r25 r30 r29 r27  9 5 ° r3l  0.06 OXYGEN CONSUMED (MOLE)  9 0 ° r!6  0.04 8 5 ° r32  0.02  0.00  FIGURE  3 4 TIME (HR.)  23. EFFECT OF TEMPERATURE. OXYGEN CONSUMED VS. TIME. IM H S 0 , 976 PSI 0^ 2  4  • 52  forms at temperatures as low as 115°C which indicates that monoclinic 23 sulfur (mp 119°C)  i s 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 m i n e r a l 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 3 5 % of the pyritic sulfur i s i n the elemental form after dissolution, the sulfur layer i s porous, allowing further attack on the surface. This behavior i s 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 i n 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 H S 0 , 976 PSI 0^. 2  4  0.5  TIME  (HR.)  FIGURE 25. TEMPERATURE DATA IN TERMS OF REACTION MODEL. MODEL VS. T I M E .  FIGURE 26.  ARRHENIUS P L O T 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 M c K a y 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. A l l 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 e r r o r .  The was  effect of gas dissolution enthalpy on the activation energy  checked and found to be negligible.  T h i s factor is introduced by  the  experimental design which controlled the concentration of d i s s o l v e d oxygen in the leach solution by regulating the oxygen p r e s s u r e .  For  10.  the  reaction:  0 (g)  >• 0 ( a q )  2  2  the dissolved oxygen concentration is related to the p a r t i a l p r e s s u r e by the solubility constant,  c<- .  L°]  =  2  The  <* o2 P  solubility t e r m i s composed of two parts; a standard solubility, <^ ; — A Hs  /RT  and an enthalpy of dissolution t e r m e _  T  L°2]  =  , such that: - &Hs/RT  *oe  Po  2  If the rate of the pyrite dissolution reaction is of the f o r m :  Rate  =  K  [0 ]  r  =  2  K°  e  "  E  A  /  R  T  [Q ] . 2  then the o v e r a l l equation is of the f o r m : o Rate  =  K  -Ea/RT  _ AHs/RT  e  e  r  =  11.  Rate  =  K  r  K  °  1 r  P  o  °^o  e  . „ o  - (Eat A H s ) / R T  P  e 2  -(Eat  A  Hs)  o  2  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 f e r r i c / 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 f e r r i c / ferrous ratios are then seen to be nearly constant when they are compared using similar extractions. The results are shown in F i g . 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 E M F of +0.874 volts, 26 observed rest potential for pyrite (0.620)  well above the  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:  Apparent Trend  LU  a: LU  20  Real Trend  r%5  —o , r 36  _J  80  FIGURE 28.  90  100 110 TEMPERATURE (°C.)  EFFECT OF TEMPERATURE FERRIC/FERROUS RATIO.  120  ON THE  130  The experimental data which show consistent 30-35% production of elemental sulfur are not consistent with thermodynamic equilibrium at this potential.  This fact shows that the  sulfur redox system is highly irreversible..  sulfate-elemental  62 •  1.  CONCLUSIONS  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 o c values of 0. 977. 2.  The dependence of the pyrite dissolution rate on oxygen consumption 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 f i l m 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 M H 2 S O 4 , sulfuric acid is  formed from pyritic suLfur. At acid concentrations greater than 1 molar, dissolution consumes acid.  The transition  point i s 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 2 5 % in going from 0.1 to 3 .MFL^SO^. Neutral salts of copper, zinc and nickel inhibit pyrite dissolution and at, the same time catalyze the oxidation of ferrous iroi to f e r r i c .  The inhibition is greatest with the copper salts  amounting to a 2 5 % reduction in iron extraction.  The f e r r i c /  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 consumption 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 3 0 % 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 f e r r i c and ferrous ions i s greatly effected by the extent of reaction. Varying iron extraction from 40 to 6 0 % results in up to a 15-fold change in the f e r r i c / f e r r o u s 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 i s  possibly attributed to a concentration effect on the homogeneous ferrous to f e r r i c 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). K. Kelley, and King Bulletin 592, Bureau of Mine s U . S . A . ,  6.  Washington, D . C , (1961). 7.  J . Pemsler,  8.  D . D . Wagman, et al. , N . B . S . Technical Note 270-3 and 270-4, Washington, D . C . , (1968, 1969). R . J . Biernat and R . G . Robins, Electrochimica Acta. , 1972, vol. 17, p. 1279.  9.  Technical Report 260, Ledgemont Publication, (1971).  10.  K. Semrau, "Sulfur Oxides Control and Metallurgical Technology", Journal of Metals, V o l . 23, p. 41, (1971).  11.  F . A . Forward, and J . Halpern, AIME Trans. , 1955, V o l . 203, p. 463, Journal of Metals, March, 1955.  12.  J . Halpern, and F . A . Forward, Trans. Inst. M i n . 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, 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~.  Z . Erzbergbau Metallhuttenw,  66  17. 18.  Analysis, Can. Test Ltd. , File No. 7647A. J . H . Perry, Chemical Engineers Handbook, Third Edition (1950), p. 278.  19.  Standard Methods of Chemical Analysis, 6th Edition, V o l . 1, p. 544.  20.  G; Nickless,  21.  F . Lowen, M . A . Sc. Thesis "Pressure Leaching of Copper Sulfides in Perchloric Acid Solutions", U . B . C . 1967. A . Adamson, The Physical Chemistry of Surfaces, 2nd Edition,  22.  Inorganic Sulfur Chemistry Elsevier C o . , 1968, p. 524.  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, V o l . 44, No. 5, P . 1146 (1952). CRC Handbook of Chemistry and Physics, 52nd Edition, P . D i l l , (1972).  25.  26.  E . Peters and H . Majima, Canadian Metallurgical Quarterly, V o l . 7, No. 3, p. 111.  APPENDIX ACTUAL EXPERIMENTAL Initial  DATA  Residue Soln. V o l .  Size  P Q  Run  Pyrite  Acid  1-9 10  Initial runs on which was No results apply.24. 82 150/200 1M technique 976 Sul. 110perfected. - 5.0000 1.4611 it  11  11  Nor.  12  L e a k in system.  13  . Nag..  it  it  .153  1. 28  12. 52  . 0937  5.0010  2.1332  39. 08  .139  1. 28  11.09  . 0753  11  5.0000  t- 1. 7237  32. 04  . 138  1. 23  11.80  .0719  II  ti  it  5.0026  2.0629  32. 85  . 135  1. 39  11.91  . 0777  ti  5.0000  . 1.4400  23.67  .130  1. 28  15.09.  . 0920  5.0006  3.1422  41. 0 1  . 142  1. 03  6.93  . 0465  5.0000  3.6543  43. 10  .135  0. 77  5. 34  . 0333  5.0012  1.2456  17. 86  . 147  1. 44  13. 70  . 0953  ' 1.44  15.96  .0979  15  Kim.  n  it  ti  16  Sul.  it  ti  tt  17  Nor.  it  ti  it  18  Sul.  20  2  Res.%Fe  ti  War.  11  Weight  II  14  it  Weight  M o l e s C>  it tt  19  2  Temp.°C  1 ff 1 Tot. JtFe J Fe  90 11  325/400  "  "  200/270  "  "  it  5.0010  1.1483  15. 36  . 130  "  11  5.0005  1.2873  19.05  . 142  1.44  14. 22  . 0953  11  5.0000  1.5751  30. 11  . 136  1.49  13. 19  . 0895  270/325  "  110  21  Nor.  200/270  "  22  War.  200/270  "  "  11  5.0000  1.6120  26.84  . 120  1. 59  15.17  . 0895  23  tt  270/325  "  "  ti  5.0000  2.4060  23.11  .128  1. 54  15.45  . 0961  24  it  375/400  "  "  II  5.0012  1.4048  21.49  . 126  1. 64  15.65  . 0970  150/200  "  "'  115  5.0012  2.1747  34. 89  . 105  0. 56  14. 16  . 0790  25  Sul.  11  ON  A C T U A L E X P E R I M E N T A L D A T A (Continued) Run  Pyrite  26  Sul.  27  "  Size  Acid  150/200 I M n  28  P  0 2  Temp. °C  Initial Weight  Residue Weight  Res. % F e  .£  Soln. Vol. JiFe  ++  fFe  T o t  V  Moles O 2  976  105  5.0000  1.8875  32. 84  . 126  0. 98  13.65  . 0843  I I  100  5.0000  2.2670  34. 00  . 130  0. 94  11.65  .0736  120  • 5.0016  2.2105  36.93  . 074  0. 33  19.60  . 0793  0. 23  18. 53  .0750  0. 23  19.04  . 0773  ii  II  I I  29  II  II  it  125  5.0017  2.2889.  36. 82  . 080  30  ti  it  ti  130  5.0010  2.2841  36. 72  . 076  31  II  ti  it  95  5.0000  2.7183  35.69  . 144  1. 23  8.98  .0632  32  n  n  I I  85  5.0004  3.6679  42. 57  . 122  0. 82  5.95  .0372  33  II  0.1M  "  110  5.0005  1.7052  31. 21  . 124  1.49  14.37  34  . .0996  II  0.01M  "  11  5.0000  2.1704  31. 99  .119  1. 03  13. 75  .0995  35  II  IM  •*  130  5.0007  3.2456  36. 51  . 110  1. 33  9. 03  . 0478  36  II  120  5.0000  3.7106  43. 64  . 106  1. 39  6.72  . 0351  10.0000  6.0568  40.49  .118  2.82  18. 53  . 1087  5.0016  2.2654  31.92  . 114  1.49  14.01  . 0753  2.5000  1.2496  40.00  . 094  1.23  7.24  .0342  11  I I  it  I I  110  38  it  3M  "  n  39  ti  IM  "  II  37  40 41 42  "  "  •  ii II  •  II  it  110/130  7.0069*  6.4565  28. 95  . 066  0. 26  3. 90  . 0147  lM+Zn  "  110  5.0034  2.8372  36. 80  .124  0. 92  9. 70  . 0586  Experimental e r r o r . *' 5. 0026g pyrite + 2. 0043g sulfur. ON  oo i i  ACTUAL EXPERIMENTAL  Run 43  Pyrite  Acid  Size  ** Sul. f Z r 150/200  1M .  44  II  45  .1!  46  Expe r ime nta 1 e r r o r .  11  11  D A T A (Continued)  1M4-Ni lMt-Cu  P  Temp. ° C  °2  Initial Weight  Residue Weight  Res. % F e  Soln. V o l .  k Fe  f e  T o t  F  -  6. 16  M o l e s O-,  . 0618  976  110  5.0000  2.6357  38. 27  . 207  0. 96  it  II  5.0038  2.8711  4.0. 3 3  .110  0.96  10.98  . 0577  11  tt  5.0005  3.4000  44.47  .120  0. 67  . 7. 75  . 0443  .9063  . 110  0. 92  20.4  . 1108  . 11.55  . 0644  47  it  it  lM  it  it  5. 0000  48  it  rt  ti  II  tt  2. 8852  36.93  .121  1.49  49  it  ii  it  it  11  2.7212  37. 90  . 100  1. 33  12. 88  . 0644  50  it  it  it  it  it  5.55*1*3* **** 5.5013 ***** .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  ** *** **** *****  Z r foil added, 5. OOOOg pyrite 5. 0000g pyrite 5.0075g pyrite  did not affect reaction. ± 0. 5513g ZnS. -l- 0.5013g CuS. +• 0. 5008g N i S F e S .  7.77  


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