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Leaching of metal oxides Devuyst, Eric August Pierre 1973

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c-l LEACHING OF METAL OXIDES  by  ERIC AUGUST PIERRE DEVUYST Ing. Civ. Mines, U.L.B. Brussels, 1968 M.A.Sc., U.B.C., 1970 A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHY in the Department of METALLURGY  We accept this thesis as conforming to the required standard  THE UNIVERSITY OF BRITISH COLUMBIA September, 1973  In presenting this thesis in partial  fulfilment of the requirements for  an advanced degree at the University of B r i t i s h Columbia, I agree that the Library shall make it freely available for reference and study. I further agree that permission for extensive copying of this thesis for scholarly purposes may be granted by the Head of my Department or by his representatives.  It  is understood that copying or publication  of this thesis for financial gain shall not be allowed without my written permission.  Department of  . Metallurgy  The University of B r i t i s h Columbia Vancouver 8, Canada  Date  September 26, 1973  f  ABSTRACT  The leaching of metal oxides in acids has been investigated. The experiments were focussed on the leaching of ferric oxides in perchloric, hydrochloric, sulphuric, oxalic and malonic acids. Additional studies were made of the leaching of aluminum, cuprous, cupric and manganous oxides in the above acids. In dilute solutions of the acids (<0.2M), the rates of leaching of the oxides showed a dependence on the mean activity of the acids ( )> a  +  varying between f i r s t and second order.  In acids which do not  form strong complexes with the metal ions considered, e.g. HCIO^, the order was approximately one; in acids with strong complexing power the order was close to two.  In more concentrated acids (>0.2M), the  order decreased progressively by one unit (i.e. from 1 to 0 or 2 to 1), with increasing acidity.  The addition of ferrous salts to oxalic and  malonic acids greatly enhanced the rates of leaching of ferric oxides. A general mechanism for the direct leaching of metal oxides in acids has been proposed.  It is postulated that the oxide surface be-  comes rapidly hydroxylated,  followed by the successive adsorptions of  hydrogen ions, anions of the acid and again hydrogen ions at hydroxylated sites.  With the exception of dehydrated aluminum oxides, the  kinetics of the leaching of the oxides in acids were consistent with the assumption that the rate determining step was the desorption of  metal species which are formed at the oxide surface during the above adsorption reactions.  For the dissolution of dehydrated aluminum  oxides, i t appeared that the rate of surface hydroxylation was ratecontrolling under some conditions.  The adsorption pre-equilibria  could be correlated to the pH of zero-point of charge (Z.P.C.) of the oxide surface.  Oxide surfaces exhibiting a high pH of Z.P.C, e.g.  C^O, CuO and MnO, are suggested to become more rapidly saturated by ionic species from solution with increasing acidity than oxide surfaces having a lower pH of Z.P.C, 'e.g. a-Fe20.j. This saturation of the oxide surface i s postulated to be the reason for the observed decreasing dependence of the rates of leaching on a with increasing acidity. +  It is also concluded that the complexing ability of the acids for metal ions i s essentially correlated with the rates of desorption of the metal-anion species formed at the oxide surface. The catalytic effect of ferrous ion additions to oxalic and malonic acids i s explained by an electrochemical mechanism involving the formation and rapid desorption of ferrous and ferric species at the ferric oxide surface.  . iv  TABLE OF CONTENTS Page 1.  REVIEW OF PREVIOUS INVESTIGATIONS 1.1  Introduction  1.2  The Chemistry of Oxide Surfaces  1  .........  1  ;  ..  .  ;  2  1.2.1  Hydration-Dehydration of Oxide Surfaces  2  1.2.2  Zero-Point of Charge  6  ..............  (a) Variables Affecting the Zero-Point of Charge  6  (b) Selectivity of Adsorption 1.3  The Direct Leaching of Metal Oxides..: 1.3.1  1.4  1.5  10  Kinetics of Leaching  15 .............  15  1.3.2 Mechanisms of Leaching of Metal Oxides The Leaching of Metal Oxides Involving ElectronTransfer .  26  1.4.1  Kinetics of Electron-Transfer Reactions ......  26  1.4.2  Mechanism of the Leaching involving Charge Transfer at the Oxide-Electrolyte Interface ..  33  C r i t i c a l Summary  2.  SCOPE OF THE PRESENT INVESTIGATION  3.  EXPERIMENTAL  21  •  ..  44 '  45  3.1 . Minerals and Reagents 3.1.1  Natural Minerals  3.1.2  Synthetic Minerals (a) Hematite (a-Fe.O.)  37  45 45 .... .  45 45  V  Page (b) Cuprous Oxide (Cu 0)  .. ...  2  ...... •  47  (c) Cupric Oxide (DuO) ............  47  (d) Manganous Oxide (MnO) .....................  47  (e) Aluminum Oxides . . . . 3.1.3  4.  Reagents  3.2  Apparatus Design  3.3  Experimental Procedure  3.4  Analytical Methods  4.2 4.3 4.4 4.5  ... •'.  . ..  50  ........  Iron  3.4.2  Aluminum  3.4.3  Copper  3.4.4  Manganese ....................... v  3.4.5  Determination of the Ferrous Content of Hematite Specimens . ... '  50  52  3.4.1  RESULTS 4.1  50  53  .  53  ;  54  ••  ..  54  .......  54 54  •  .. .  55  The Leaching of Metal Oxides in Aqueous Perchloric Acid Solutions .........  55  The Leaching of Metal Oxides in Aqueous Hydrochloric Acid Solutions  59  The Leaching of Metal Oxides in Aqueous Sulphuric Acid Solutions  '  .  The Leaching of Ferric Oxide in Oxalic Acid in the Absence of Added Ferrous'Salt in Solution  70 71-  The Leaching of Ferric Oxide in Oxalic Acid in the Presence of Added Ferrous Oxalate in Solution..  76  4.5.1  Preliminary Experiments  76  4.5.2  The Effect of Sample Weight  •  ..  79  vi  Page 4.5.3 The Effect of Added Ferrous Oxalate Concentration 4.5.4 4.5.5  79  The Effect of Adding Ferrous Ion Complexing Agents to the Oxalate Electrolyte  82  The Effect of Adding Various Cations in Solution  82  4.5.6  The-Effect of the Concentration of Oxalic Acid..  83  4.5.7  The Effect of the T i Content of Synthetic Ferric Oxide  4.5.8  y  The Effect of Temperature  V  83 ...  ..  85  4.5.9 The Effect of pH  4.6 4.7  4.5.10 Distribution of Ferrous Species i n 0.2 M Oxalic Acid as a Function of pH The Leaching of Ferric Oxide in Malonic Acid in the Presence of Added Ferrous Ion ................... The Leaching of Ferric Oxide in Various Other Acids ... 4.7.1 4.7.2  5.  88  In the Absence of Added Ferrous Salts in Solution  :  ...  91 95  ... .95  In the Presence of Added Ferrous Salts in Solution  DISCUSSION . .  88  95  '  97  5.1 The Direct Leaching of Metal Oxides in Acids  97  5.1.1 Model for the Mechanism of Leaching ............  97  5.1.2 Leaching of Metal Oxides in HCIO^ Solutions .. . .' 104 5.1.3 Leaching of Metal Oxides in HCI Solutions  •. 108  5.1.4  Leaching of Metal Oxides in ^SO^ Solutions .... 116  5.1.5  The Leaching of Ferric Oxide in H^O^  5.1.6  The Leaching of Ferric Oxide in Various Other Acids ' •  Solutions 119 120  vii  Page 5.2  The Acid Leaching of Ferric Oxides i n the Presence of Added Ferrous Salts in Solution  ..... 122  5.2.1 The Leaching of Ferric Oxide in H„C 0. Solutions 2 24 5.2.2 The Leaching of Ferric Oxide in Malonic Acid  122  o  129  5.'2.3 The Leaching of Ferric Oxide in HCl 6. CONCLUSIONS  •'  •  .. . .  7. SUGGESTIONS FOR FUTURE WORK 8. APPENDIX A  '  •  REFERENCES  • '  '  !  ........ :  . ...  ....  130 .131  ... 133 135  9. APPENDIX B 10.  . ..  :  . ..."  140 167  viii  LIST OF FIGURES  Figure 1  2  Page Rate of leaching of Ck^O i n H „ S 0 , and HC10> -at- various concentrations of H+ '(TH31*C)T. . ..  . 24  Rate of leaching of goethite i n HCl -versus the mean a c t i v i t y of HCl (T=85°C) ........  24  3  Electron microprobe pictures for T i or Mg of synthetic a-Fe 2 0 3 samples (Table 5) (xl,000) . .v . 48 (a) 0.1% T i ; (b) 0.2% T i ; (c) 0.5% T i ; (d-)- 1.3% T i . . . 48 (e) 3.0% T i ; (f) 0.5% Mg •.• ' • . . . . / 49  4  Apparatus design . . . . . . . . . . . . . . . . . . . . . .  5  Relatives rates of leaching of versus the concentration of  6  ,. .  CU2O and CuO HCIO4 at 12°G  in  HOIO4  Relative rate of leaching of goethite (Surana and Hay,' T=110°C) and hematite (T=90°C) versus the concentrat i o n of HC10. .  51  56  57  4 7 8  9  10  11  12  13  14  Relative rate of leaching of f e r r i c oxide i n d i l u t e HCl as a function of the mean a c t i v i t y of HCl . . . . . . . .  60  Relative rate of leaching of f e r r i c oxide i n HCl as a function of the mean a c t i v i t y of HCl . . . . . . . . . . . . . . .  61  Ratio of the r e l a t i v e rate of leaching oxide and a+ as a function of a+  62  of f e r r i c  E f f e c t of adding L i C l on the r e l a t i v e rate of leaching of f e r r i c oxide (Michigan) i n 2.4 M HCl at 80°C  64-  HCIO4  The e f f e c t of adding NaOH or on the r e l a t i v e rate of leaching of a-Fe203 (Michigan) i n 2.4 M HCl at 80°C  65  Relative rates of leaching of aluminum oxides i n HCl versus the mean a c t i v i t y of HCl (T=80°C)  67  Relative rates of leaching of Cu20 and CuO i n d i l u t e HCl versus the mean a c t i v i t y of HCl (T=12°C) ... .•  68  Relative rates of leaching of CuO i n HCl versus the mean a c t i v i t y of HCl (T=12°C) . . . .  v  69  ix  Figure 15 16  Page Relative rate bf'leaching of ferric oxide in H2SO4 versus the concentration of H2SO4 >.'..•• . . ..... Relative rate of leaching of CU2O, CuO.and MnO i n H SO^ versus the concentration of H2SO4 (T=12°C) .. 2  17 .  Distribution of species i n oxalic acid at 80°C versus pH  . ...  72 .73 74  18  Rate.of leaching of ferric oxide.(Michigan) in 0.3 M oxalic acid versus pH (T=80°C) ..  19  Leaching of goethite (Minnesota) i n 0.2 M oxalic acid in the presence of air, 0„ and He versus time (T=80°C,, pH=2.8) .... .V  77  Rate of leaching of ferric oxide i n 0.2 M oxalic acid versus the concentration of added ferrous ion (T=80°C, pH=2.8) '  80  Log-log plot of the rate of leaching of ferric oxide in 0.2 M oxalic acid versus the concentration of added ferrous ion {T=80°C, H=2.8) ' p  81  Rate of leaching of ferric oxide i n oxalic acid versus the concentration of oxalic acid (T=80°C, pH=2.8, added ferrous=6 mg/liter)  84  Relative rate of leaching of ferric oxide i n 0,2 M,' oxalic acid (pH=2.8, added ferrous=6 mg/liter) and 2.4 M HCI at 80°C versus the titanium content of the oxide '  86  Arrhenius plots for the leaching of ferric oxide i n 0.2 M oxalic acid (pH=2.8)  87  25  Normalized rate of leaching of ferric oxide i n 0.2 M oxalic acid versus pH (T=80°C, Fe(II)=6 mg/liter)  89  26.  Log-log plot o f the total solubility of ferrous species in 0.2 M oxalic acid versus the concentration of oxalate ion at 80°C  92  Distribution of ferrous species in 0.2 M oxalic acid versus pH at 80°C  93  20  21  22  23  24  27  75  x Page  Figure 28  Rate of leaching of ferric oxide (Michigan) i n 0.3 M malonic acid versus pH (T=80°C, Fe(II)=9 mg/liter) 94  29  Rate of leaching of ferric oxide (Michigan) in HCI versus the concentration of added ferrous ion (T=80°C)  30  Morphology of the acid attack on the basal plane of a a-Fe203 single crystal ................ 128 (a) 9 M HC10 , 80°C, 10 days, x 2,000 (b) 6 M HCI, 60°C, 10 min, x 2,000 (c) 6 M H S0,, 60°C, 20 min, x 2,000 (d) 0.2 M oxalic acid, 6 m g / l i t e r Fe(II),, 80°C, 20 min, x 1,000 4  2  96  xi  LIST OF TABLES Table.  Page  1  Zero Points of Charge for Various Oxides ...............  2  Experimental Values of n in Rate=k [Acid]  3  Activation Energies (kcal/mole) for the Leaching of Various Oxides i n Various Acids  22  4  Rate Constants and Enthalpies.and Entropies of Activation for -Various Homogeneous Ferrous-Ferric Electron Transfer Reactions •  29  5  Synthetic Hematite Specimens  46  6  Leaching of Metal Oxides in HCIO^. Calculated Constants in Rate Expression (5.6) 106  7  Leaching of Metal Oxides in HCl Calculated Constants in Rate Expression (5.6) '  110  8  Leaching of Cu„0 in HCl. (5.14) and (5.15)  115  9 10  .  9 18  Constants i n Rate Expressions  Leaching of Metal Oxides in H2SO4. Calculated Constants in Rate Expression (5.17)  •  Leaching of a-Fe20.3 (Michigan) in...Oxalic Acid. Calculated Constants in Rate Expressions (5.18) and , (5.19)  "118  121  A.l  Analysis of Goethite (Minnesota) and Hematite (Michigan) 135  A.2  X-Ray Diffraction Patterns of.Synthetic Hematite (Table 5)  136  A.3  X-Ray Diffraction Patterns for Synthetic Cu 0 and CuO .. 137,  A.4  Chemical Analysis of Pyrolusite  A. 5  X-Ray Diffraction Patterns of A1(0H) , - A 1 0 „ and.  B. l -  a-Al 0 ......7. ... PH of a 0.2 M Oxalic Acid Aqueous Solution at 80°C as a Function of Added HC10. and NaOH  2  138 Y  2  3  9  139 140  xii  Table B.2 B.3  Page Calculated Distribution of Oxalate Species i n 0.2 M. Oxalic Acid at 80°C  141  Total Solubility of Ferrous Species i n 0.2 M Oxalic Acid as a Function of pH at 80°C (Figure 18)  142  B.4  Calculated Mean Activities of HCIO4 Solutions at 25°C ...  143  B.5  Experimental and Calculated Rates of Leaching of Metal Oxides in HCIO^ Solutions (Table 6, Figures 5 and 6) Calculated Mean Activities of HCI Solutions .  .  144 146  Experimental and Calculated Rates, of Leaching, of Metal Oxides in HCI Solutions (Table 7, Figures 7, 8, 9, 12, 13 and 14) •  147  Calculated Relative Rates of Leaching of Ferric Oxide Using Simplified Rate Expressions  151  Ratios of the Relative Rates of Leaching of Ferric Oxides and a+ as a Function of a+  152  Experimental and Calculated Rates of Leaching of Ferric Oxide (Michigan) in: HCl-LiCl, HCl-NaOH and HCI-HCIO4 Solutions (Table 7, Figures 10 and 11) ..  153  B.10  Calculated Mean Activities of H2SO4 Solutions  154  B.ll  Experimental and Calculated Rates of Leaching of Metal Oxides in ^SO^ Solutions (Table 9, Figures 15 and 16) ... 155  B.12  Calculated and Experimental Rates of Leaching of a-Fe20^ (Michigan) i n 0.3 M Oxalic Acid at 90°C versus pH (Figure 18) .........  B.6 B..7  B.7,a. B.8 B.9  157  B.13  Experimental Relative Rates of Leaching of a-Fe203 i n 0.2 M Oxalic Acid at 80°C versus the T i Content (Figure 23) 158  B.14  The Effect of Added Ferrous Oxalate on the Leaching of ct-Fe 0 i n 0.2 M Oxalic Acid at 80°C and pH 2.8 Figures 20 and 21).:... . .......  159  Effect of Sample Weight (Sample Q, Table 5). Leaching of ci-Fe 0 i n 0.2 M Oxalic Acid at 80°C and pH 2.8, with 6 mg FeTII)/liter  160  B.15  xiii  Table B.16 B.17 B.18  Page Effect of Temperature on the Rate of Leaching of a-Fe_0 in 0.2 M Oxalic Acid at pH 2.8 (Figure* 24) .... . .... Calculated Distribution of Ferrous Oxalate Species in 0.2 M Oxalic Acid versus pH, at 80°C (Figure 27)  162  Experimental and Calculated Rates of Leaching of a-Fe 0 (Sample 0, Table 5) at 80°C versus pH (Figure 25)  163  Effect of Oxalic Acid Concentration on the Rate of Leaching of a-Fe 0 (Sample Q, Table 5) at 80°C and at pH 2.8 (Figure 22) .........  164  Rate of Leaching of a-Fe 0 (Sample H, Table 5) in 0.5 M Malonic Acid at 80°C versus pH in the Presence of 9 mg/liter of added Ferrous IOn (Figure 28)  165  Effect of Adding Ferrous Ion on the Leaching of a-Fe 03 (Michigan) in HCl Solutions at 80°C (Figure 29).........  166  2  B.19  3  2  B.20  B.21  161  3  2  3  2  ACKNOWLEDGEMENT  The author wishes to express his gratitude for the advice and aid of Dr. I.H. Warren during the course of the work and desires to thank members of faculty, fellow graduate students and the technical staff for their helpful collaboration. Financial support from the National Research Council of Canada in the form of a Research Assistantship i s gratefully acknowledged.  1  1.  1.1  REVIEW OF PREVIOUS INVESTIGATIONS  Introduction Metals can be leached from their oxides by direct reaction with  an aqueous solution of an acid or an alkali, or by reaction with either of these reagents in the presence of an oxidizing or reducing agent. Reactions of the f i r s t type are represented by the historic Bayer process for alumina production (1) and the recent Jarosite process for zinc recovery from zinc-ferrite (2), whilst the leaching of uranium oxide with sulphuric acid in the presence of oxygen (3) and the leaching of manganese dioxide with sulphurous acid (4) are examples of the second type.  Although much has been published in recent years on the  thermodynamics of metals in oxide-water systems (5), the kinetics and mechanisms of oxide leaching reactions have been only sparsely studied and this despite the potential significance of such studies in the field of corrosion in addition to extractive metallurgy. Burkin (6), in a 1966 review of the chemistry of hydrometallurgical processes, commented briefly on the kinetics of dissolution of ferric oxide and cuprous oxide in acids and of uranium dioxide in oxygenated carbonate solutions.  More recently Habashi (7) assembled an extensive  bibliography on the leaching of oxides,but did not attempt a comprehensive discussion of the kinetics and mechanism of their dissolution. The present study, whilst concerned principally with the leaching  2  of iron oxides, was undertaken with the ultimate objective of attempting to develop a general mechanism to explain the dissolution of oxides. Warren et al (8,9) have proposed a simple model for the dissolution of goethite and hematite in perchloric and hydrochloric acids which could account for the leaching of these oxides in dilute solutions of the acids.  In this model i t was postulated that the oxide surface immersed  in the aqueous solution of the acid is subjected to rapid hydroxylation followed by rapid equilibration with the ionic species in solution. The much more rapid dissolution of ferric oxide in hydrochloric acid than in perchloric acid was explained in terms of the activation of positively charged surface sites by adsorbed chloride ions. As a preliminary in attempting to expand this model to a variety of acids and oxides, a detailed review of the factors affecting the oxides surface hydroxylation, the charging of the oxides-electrolytes interfaces and the selectivity of anion and/or cation adsorption at . these interfaces w i l l be reviewed.  Other, mechanisms of oxides dissolu-  tion under a variety of conditions which have been proposed in earlier work w i l l also be considered.  1.2 1.2.1  The Chemistry of Oxides Surfaces Hydration-Dehydration of Oxide Surfaces If the f i r s t step in the overall leaching mechanism of oxides is  hydroxylation of the surfaces, as proposed by Mackay and Wadsworth (10) for leaching UC^ in oxygenated dilute sulphuric acid and by Warren and Monhemius (8) and Warren and Surana (9) for leaching goethite, then  3  clearly the kinetics of hydration of oxides are of considerable interest. A very complete survey to 1967 of the studies of hydration and dehydration of oxide surfaces has been made by Hair (11). For aluminum oxide Peri and Hannan (12) have concluded from infrared studies that the surfaces of the oxide produced by heating y-AlJd^ above 800°C s t i l l retain some hydroxyl groups but that no increase in their number occurs on exposure of the surfaces to water vapour at room temperature and pressure.  The surface of the oxide calcined at 800°C  revealed the presence of at least five types of isolated hydroxyl groups. Peri (13) was able to propose a computer model for the dehydration process of the y-AlJO^  surface in which the possible remaining isolated  hydroxyl groups are indeed on five types of sites on which they have from zero to four nearest oxide neighbours.  These hydroxyl groups  apparently show a similar behaviour to those isolated hydroxyl groups produced on the surface of s i l i c a calcined above 400°G.  It has been  observed (14,15) that water molecules cluster around these isolated hydroxyles without reacting with adjacent oxide groups to rehydroxylate them. Bielanski and Sedzimir (16) in a study of the adsorption of water vapour on boehmite calcined at various temperatures between 500°C and 1300°C showed that the rate of water adsorption decreased with increasing calcination temperature until at between 1100°C and 1300°C oxide (a-Al^O^) with an essentially hydrophobic surface was produced. Unfortunately, conflicting views on the kinetics of hydrationdehydration have been advanced.  Wade and Hackerman (17) and Hendriksen  et a l (18) concluded from studies of the heats of immersion of a-Al 0  4  in water that the rehydroxylation of ot-Al^O^ was rapid and independent of the temperature of dehydration,but Morimoto et a l (19) had observed earlier that a maximum occurred in the heat of immersion of Y and cx  -  aluminas with increasing temperature of dehydration pretreatment of the oxides, suggesting that irreversible dehydration of the oxide surface had taken place. Hendriksen et a l (18) suggested that the aluminas used by Morimoto possibly had annealed upon heat-pretreatment resulting i n a decreased surface area of the samples. Very recently, Baker et a l (20) identified six different mechanisms for the sorption of water by oxides, namely: (a)  Hydrogen bonding between adsorbed water molecules and surface  hydroxyl groups. (b)  Hydrogen bonding between sorbed water molecules and hydroxyl  groups in micropores. (c)  Hydration of exposed surface cations by adsorbed water mole-  cules . (d)  Dissociative chemisorption of water with the formation of  hydroxyl groups. (e)  Hydration i n depth of poorly ordered cations.  (f)  Hydroxide or oxide-hydroxide formation in depth.  According to Baker et a l (20) the slowness of processes (c) and (d) are at the origin of the irreversible rehydration of dehydroxylated s i l i c a and chromia.  Apparently process (c) contributed to a large  extent to the rehydration of the dehydroxylated a-alumina surface; Baker also concludes that process (d) i s rapid for a-A^O^, but he agrees however that slow adsorption of water vapour on a-A^O^ continued over a period of months. Moreover, their water vapour uptake measurements for  5  a-Al^O^ were made after outgassing this oxide at 500°C, and thus i t appears possible that many hydroxyl groups are s t i l l present on the alumina surface at this  temperature.  Titanium dioxide i s a.typical tetravalent metal oxide.  However,  in contrast with silica,the surface of dehydrated TiC^ is at least partly rapidly rehydroxylated upon rehydration (21,22).  Primet et  al (22) showed that dehydroxylation of crystallized TiO^ i s only partially reversible, as the decrease in surface area of TiC^ during the dehydration-rehydration cycles was not sufficient to account for the observed decrease in rehydration. Primet et a l (22) postulate the formation of three types of sites at the TiC^ surface upon dehydration.  The f i r s t type of sites are basic in character and appear  by the condensation of adjacent hydroxyl groups.  The second and third  types of sites are acidic (Lewis); the strongest Lewis sites are created by the removal of isolated hydroxyl groups and the weakest Lewis sites are due to the removal of molecular water (around 15.0°C) .  Rehydration  of dehydrated TiO^ apparently proceeds by dissociative adsorption of water on Ti-0 pairs (basic sites) until 50% of the surface i s hydroxylated and by molecular adsorption on isolated T i ions (strong) and on isolated oxygen ions (weak) (23). In contrast with the behaviour of alumina, s i l i c a and titanium dioxide, ferric oxide which has been dehydrated by calcination appears , to react readily with water in a process that has been suggested (24) as involving interaction of one surface Fe-0 species with a HO mole-  6  cule to produce two surface OH groups.  Recent observations by  McCafferty and Zettlemoyer (25) suggest that the f i r s t layer of physically adsorbed water on a-FeJO^ i s immobile and doubly hydrogen bonded to the underlying hydroxyl layer, but that succeeding layers are mobile. Infra-red studies on the surface hydration of divalent metal oxides are.rendered d i f f i c u l t by the presence of a high background adsorption in the,spectral regions of interest.  Anderson et a l (26)  observed an irreversible modification of the surface of MgO following/  complete dehydration.  crystals  As with the s i l i c a surface,  the species formed during the readsorption process are dependent upon the prior thermal history of the oxide sample*, but contrary to s i l i c a , MgO rehydroxylates rapidly. In contrast with the behaviour of the oxides mentioned above,"isolated hydroxyl groups are apparently not formed at the surface of BeO upon dehydration (27).  It also seems that the hydration-dehydration  cycle of BeO i s reversible on material heated to temperatures of at least 550°C.  L.2.2. (a)  Zero-Point of Charge Variables Affecting the Zero-Point of Charge  Oxides , especially the hydrous oxides, exhibit ion exchange proper1  ties (28).  The ion exchange capacity of oxides arises from the existence  of a pH-dependent surface charge.  Charge 'can develop on a hydroxylated  7  surface through amphoteric d i s s o c i a t i o n of the surface hydroxide groups. Dissociation reactions can be written as follows,  -MOH + OH,  ,  |-MOH+HJa s  U  q  == ;  -MO  + H.O  (1.1)  J-MOH+ 1  s  (1.2)  (| i s a symbol r e f e r r i n g to the surface of the oxide), s The Z.P.C. (zero-point of charge) of an oxide refers to the p  i n any  system, however complex, at which there i s no net charge on the surface of the oxide.  +  I f the charge i s established by H , OH , and species +  capable of interacting with H , OH  or H^O to. form species present i n  the s o l i d l a t t i c e (called p o t e n t i a l determining i o n s , P.D.I.) , then the Z.P.C. may be given the special name I.E.P.(s) (29) ( i s o e l e c t r i c point of the surface, as compared to the I.E.P. of species i n s o l u t i o n ) . Adsorption of species (molecules or ions) under the combined influence of i o n i c and non-ionic bonding i s c a l l e d " s p e c i f i c adsorption".  The  following r e l a t i o n s h i p among the I.E.P.(s)- of an oxide or hydroxide, the charge or oxidation state of the cation and i t s radius was derived by Parks (29):  I.E.P.(s)- = A - B • [(§•) + 0.0029C + a]  where  Z = cationic charge R = r , + 2r + o  (1.3)  8  = oxygen ion radius A,B = constants f o r a l l oxides C. = correction f o r c r y s t a l f i e l d s t a b i l i z a t i o n of M-OH bond a = combined corrections f o r coordination number and state of hydration. Table I shows the range of values of Z.P.C. extracted from a 1964 review by Parks (29). The role of such variables as c r y s t a l structure and e l e c t r o l y t e composition i n determining the solution pH at which there w i l l be no net charge on the oxide surface have been the subject of extensive study.  Increasing c r y s t a l l i n i t y as observed i n aging p r e c i p i t a t e s ,  for example, s h i f t s the Z.P.C. i n the basic d i r e c t i o n .  Healy et a l  (30) have interpreted the wide range of Z.P.C.'s (pH 1.5 to 7.3) they observed for various polymorphs of MnO^ i n terms of v a r i a t i o n i n crystallinity.  They conclude that as the atomic packing i n the MnO^  l a t t i c e increases, the e l e c t r o s t a t i c f i e l d within the l a t t i c e increases and the pH of the Z.P.C. increases; the pH of Z.P.C. can be approximated by the following r e l a t i o n , based on the Huckel equation for the electros t a t i c f i e l d strength of s o l i d s :  pH (Z.P.C.) = A^- + B ¥  c where A and B are positive constants f o r an oxide series and r the c shortest M-0 i n t e r i o n i c distance.  9  TABLE 1 Zero Points of Charge for Various Oxides  Oxide  Coordination . » t \ Z ; £ B i & . M-0 ..... . . (pH) '.' . r  1  M0  2-4.  >11". 5  MO  6-6  ' 8.5-12.5  2  Structure  Example:with Z.P.C* . r-"-(at ^25°C) ...  Octahedral  Cu 0.  Cubic  MgO(12.4)  '  2  Cd0(10.4) NiO(10.3) Cu0(9.5) 4-4 M0 2  3  6T4  6.5-10.4  Hexagonal  Zn0(9-10)  Hexagonal  &-Al 0 (6.5-9.5)  Rhombohedral  o-Fe 0 (8.5)  2  3  2  3  a-Cr 0 (7.0) 2  M0  2  8-4  0-7.5  Cubic  3  U0 (3.5-6.5) 2  Th0 (8.5-ll) 2  6-3  Monoclinic  Z 0 (4-6.7)  Tetragonal  Ti0 (4.7)  2  2  Sn0 (5.5) 2  B-Mn0 (7.0) 2  M0  3  6r2  <0.5  Hexagonal  a-Si0 (2.2)  Rhombohedral  W0 (0.5)  *Selected values of Z.P.C after Parks (29).  2  3  10  Partial oxidation or reduction leading to nonstoichiometry i n solids such as TiO^, Fe^'O^ or  may be expected to shift the Z.P.C.  toward that characteristic of the oxidation (or reduction) state produced (29). The effect of temperature of the electrolyte i n contact with an oxide can be roughly anticipated; the decrease bf the dissociation constant of water, Kw, with increasing temperature, would result in a shift of the pH of Z.P.C. Specific adsorption of ions other than OH  and H at the oxide +  surface w i l l undoubtedly influence the Z.P.C. To achieve zero charge in the presence of a specifically adsorbed ionic species, the pH must be shifted away from the I.E.P.(s) to increase or decrease hydrogen ion adsorption, whichever i s appropriate.  In general, anionic impur-  ities shift the Z.P.C. to a more acid value; cationic impurities shift i t to more basic values . (29). (b) .Selectivity of Adsorption . Selectivity of adsorption and the relative tendency toward adsorption are fundamental factors i n discussing the kinetics of the leaching of oxides and may be inferred from ion exchange behaviour.  O'Connor et  al (31) investigated the behaviour of natural Al^O^ and AlO'OH, and of the former after ignition to temperatures up to 1100°C in acid and alkaline solutions.  On ignition to high temperatures, the original  disordered surface of A1(0H) crystallized successively to a.  '.'\J  11  layer of g.-A10'OH (^300°C), y A l ^ (300°C-900°C) and a-Al 0 ' (>900°C). According to O'Connor et a l (31), a-AlJd^  only re-hydrates to a limited  extent when exposed to a solution, to give a layer approximating to AlO'OH. O'Connor et. a l (31) pointed out that AlO'OH is likely to be weakly acidic in comparison to basic A1(0H) , resulting, in a net nega3  tive charge on the surface of the solid in water.  Indeed, Robinson  et a l (32) observed the Z.P.C. of a-Al^O^ having a fully hydroxylated surface to occur at pH = 9.0 - 9.4; the latter when calcined at temperatures above 1000°C exhibited i t s Z.P.C. at a pH  6.7.  The effects of  HCl and B.JS0^ on the zeta-potential, z* , of alumina samples was explained by O'Connor et al (31) in terms of physical adsorption of anions and anion exchange processes.  In dilute HCl solutions, hydrated aluminas  are subjected to increasing OH  - Cl  anion exchange with increasing  HCl concentration, but heat-treated aluminas show preferential physical adsorption of Cl  in very dilute HCl ( 0.001N) and anion exchange in  more concentrated solutions.  In dilute H S,Q4  solutions, both hydrated  and calcined alumina's showed preferential SO^ - OH  anion exchange.  Earlier, Graham and Crawford (33) had studied the adsorption of oxalate ^2^4^ ^  hydrous alumina.  The. adsorption of oxalate by hydrous alumina  from either an acid solution or a neutral salt solution was greater than that of chloride; Graham and Crawford (33) suggested that the favourable C^O^ - OH  anion exchange can be related to the much greater tendency of  oxalate anions to complex with aluminum cations than do chloride ions. It should be noted that firing the hydrous alumina to 1300°C for three  *£ is "the potential difference at the interface between the oxide and the electrolyte; C i s chosen to be zero at the Z.P.C. of the oxide.  12  hours, producing A^O-j.at the surf ace, lowered the adsorption of oxalate by two orders of magnitude. Ions which can form insoluble compounds or undissociated complexes with a component of the s o l i d c r y s t a l l a t t i c e appear to adsorb more strongly than those which cannot (34).  The observed order of adsorption  of organic e l e c t r o l y t e s onto a-Al^O^ i s (35):  RCOOH > RCONH2 > ROH > RNH^ > RCOOCH3 > RN(CH 3 ) 2 > RN02 > ROCH3 > RH , and the organic e l e c t r o l y t e s with larger hydrocarbon chains form indeed . less soluble compounds. Selective i o n i c adsorption at oxide surfaces can also be i n t e r preted by considering ion-ion i n t e r a c t i o n s .  Ions having a high electro-  s t a t i c f i e l d are structure-promotors f o r surrounding water as opposed to large ions with a r e l a t i v e l y weak f i e l d strength which are structurebreakers and are weakly hydrated.  Berube and De Bruyn (36) based their  model of the TiCV-water interface on ion-ion i n t e r a c t i o n s .  The OH  s u p e r f i c i a l groups firmly anchor the neighbouring water molecules by hydrogen bonding, this phenomenon being strengthened by the large crysta4+ l l i n e f i e l d of the small T i  i o n . This results i n the presence of  "frozen" water near, the surface, the l a t t e r behaving as a structurepromotor macro-ion.  Thus,strong s p e c i f i c adsorption i s to be expected  by those ions which also favour structure-promotion. The observed order of s p e c i f i c adsorption of a l k a l i - c a t i o n s ,  13  Li  +  > -Na > C s +  +  on a negatively charged TiCV surface i s in accord with this prediction. This concept may also be applied to the anions of acids for their adsorption on a positively charged TiC^ surface, but no clearly defined order of adsorption can be.obtained as i n the case of cations. • Specific adsorption of some inorganic anions i s in the order-: Cl  - CIO. - NO > I 4 3  compared to the order, of structure-prompting  effect  NO'I > Cl > CIO, > B > I 3 4 r  For a-Ye^O^y the f i e l d strength exerted by the surface upon the electrolyte i s somewhat weaker than that of TiO^.  Dumont and Watillon (37)  developed the following series of adsorption selectivity i n acidic media, I0~ > F~ > CH C00~ > CH C1C00~ > BrO~ > SCN~ > CHC1 C00~ 3  2  2  > B" > N0~ > C10~ > C l ~ > ClOT - I~ r 3 3 4 The observed adsorption sequence on a-¥e^0^ can also be compared to the order of decreasing.mean activity coefficients of the corresponding acids and bases which reflect ion-ion interaction properties; these are (38), HI > HBr > HC10. > HCI > HN0„ > H S0, 4 3 2 4 o  and  14  CsOH > KOH > NaOH > LiOH Finally, the increasing order of Baume coefficients of viscosity (39) also appear to reflect a similar sequence, namely, I~ < ClOT < N0~ < Br~ < C10~ < .Cl" < F~ < 10~ and 4 3 3 3 Ca  + ••+ + +. < K < Na < L i  A reasonable parallelism between the various sequences i s observed. Nevertheless, some discrepancies arise; e.g., SCN., which can undergo 3+ chemical binding with the Fe oxide.  ion, is more strongly adsorbed on ferric  Moreover, CH^CICOO and CHC1 COO which are structure-breakers  as a whole, are specifically adsorbed on a-Fe20.j i n acid media; the COO  group which can organize water around i t s e l f i s therefore probably,  turned toward the surface. The problem of competitive adsorption at oxide surfaces w i l l arise in solutions which contain more than one type of ionic species, and this i s almost always the case when an oxide i s dissolving in an electrolyte.  Recently, Hingston et a l (40) investigated the competitive,  adsorption of phosphate + arsenate and phosphate + selenite ions on goethite and gibbsite. It appears from their results that the oxide surfaces contain sites common to both anions on which adsorption takes place and sites on which only one or the other anion is able to adsorb. The maximum amount of anions adsorbing from a mixture i s approximately equal to the sum of the maximum adsorption for each anion i n the absence  15  of a competitor.  In mixed systems i t is thus possible to occupy more  sites with anions than when either ion is present alone.  Hingstoh et  al (40) suggest that possibly one type of anion is shared between two Fe atoms on the crystal surface through a bridging link, whereas the other type of ion has two bridging ligand links to each Fe atom.  1.3  The Direct Leaching of Metal Oxides  1.3.1  Kinetics of Leaching In studies of the leaching of goethite (FeO'OH) and hematite  (FeJdy)  i - perchloric, sulphuric and hydrochloric acids various workers n  (9,41,42,43,44) have shown rates of attack which increase for both oxides in the order HCIO^, IL^SO^, HCl of equal normality (>IN). Because of the drastic change in surface area which occurs with HCl attack, due to pitting, i t is impossible to quote rates for the different acids on an equivalent surface area basis. For the leaching of hydrated aluminum oxide (Al^O^ • 2.7 - 2.9  H^O),  Clay and Thomas (45) and Graham and Thomas (46) have observed in their studies that the rates of leaching of the oxide in several organic and inorganic acids (0.2N) are in the following sequence: HF > H P0, > Oxalic > H„S0, > HCl > HBr ^ HN0„ '= HC10, > 3 4 2 4 3 4 o  Maleic > Tartaric > Formic > Citric > Acetic Parts.of these results were recently confirmed for the dissolution of gibbsite (A1(0H) ) in HC10., HCl and H SO  solutions by Packter and  (1.4)  16  and Dhillon (47).  Gibbsite dissolves about .five times more rapidly in  HCI solutions than in HCIO^ solutions of equal strength, and tl^SO^ solutions react about five times faster than HCI solutions of equal mean activity. Azuma and Kametani (41,48) correlated the increasing absolute rates of leaching in the different acids with the increasing complexity constants of the anions of the various acids for ferric iron. A similar correlation appears to be applicable to the leaching of the alumina hydrates because the order (1.4) is in the order of complexing power of the anions for the aluminum ion, provided corrections are made for the differences in dissociation constants of the acids.  In addition,  Wadsworth and Wadia (49) observed a more rapid rate for the leaching of cuprite in sulphuric than perchloric acid, consistent with the sulphate complex for cupric ion being relatively strong, whilst the perchlorate ion i s a non complexer (50)..  Finally, the observations  by  Koch (51) on the leaching of beryllia (BeO) are also consistent with the above pattern since the order of complexing powers for the beryllium ion by the anions (50), namely C^O^ of absolute leaching rates,  > SO^ > Cl , i s in the order  whilst apparent uniformity exists in the  properties required of an anion of an acid to achieve rapid leaching, there appears at present to be none in the observed effects of the variation of concentration of acids on the rates of leaching of the different oxides. The rate dependence on the acid concentration in dilute solutions (<IM)  can be qualitatively expressed by the following relation: Rate  =  k • [Acid]  n  (1.5)  17 A plot of log (Rate) versus log [Acid] should give a slope of n; the values of n obtained by various workers for several metal oxides are listed in Table 2.  It can be concluded that in dilute solutions:  (a) monobasic acids: - i f the anion of the acid is a strong complexer for the metal, n - 2 - i f the anion of the acid is a weak complexer for the metal, n - 1 or smaller. (b) dibasic acids: n is always smaller than 1, and close to 0.5. The value of n = 1 obtained for HF does not contradict the observed sequence as twice the amount of acid is needed to obtain an equivalent concentration of HF^ ions in solution as in the case of the other monobasic acids. The values of n reported in Table 2 only hold for dilute solutions. In more concentrated solutions n becomes equal to 1 for sulphuric acid and increases (sometimes up to 2.5) for strong monobasic acids.  This  apparent complex behaviour of the acids has not abeen explained. Very recently Kabai (53) showed that the rate ,of leaching of any oxide could be described by an 'empirical' differential equation of the form: v  dC a — = K • -JZ^.d-C) dt t where K is a constant depending on the nature and temperature of the electrolyte and type of oxide, C is the weight fraction of the total, mineral which has dissolved (total weight is equal to one), t is the  (1.6)  TABLE 2 Experimental Values of n in Rate = k • [Acid]  Oxide  Acid " (<IM)  Fe.O^-xH 0 2 3 2 (0ixs3)  HF  Ions i n Solution +  1.06  HCl  Strong  1.92-2.2  HBr  H ,Br~  Strong  1.94  41,48  HN0  H,NO~ + H ,C10 H+ ,HSO, 4  Weak  0.93  41,48  Weak  •0.93-1.0  Weak  0.56  41,48  .0.59  41,48  H  A1 0 '3H 0 2  3  2  +  +  2  BeO  ZnO  2 °4 S  3 °4 P  . HCl  so;  41,42,48  41,42,48  Strong .  H ,H P0  Weak  HPOT,PO^  Strong  2  41,48  4  Weak  *1  47  2 °4  H ,C1 + - = H ,HSO,,S0, 4 4  Weak  <1  47  HCIO.  H ,C10^  Weak  <1  47  H ,HSOT,SO7  Strong  <1  49  Weak.  <1  49  H  Cu 0  Reference  Strong  HC10. 4  H  Slope n  H ,HF7 + H ,C1  3  -  Complexing Ability  n  S  4  2 °4 HCIO. 4 HCl H  S  H  2 °4  H  2 2°4  S  C  H S0 2  4  ;  +  +  4 4 H ,C10 + H ,C1 + H ,HSO ,S0 + = H HC 0 ,C 0 2  4  2  H ,HSO.,SO. 4 4  4  Weak  0.53  51  Strong  0.70  51  Strong  0.42  51'  Strong  <1  52  19  time and a i s a dimensionless number depending on the chemical compos i t i o n and structure of the oxide.  Expression (1.6) d i f f e r s from the  Nerns.t equation (54) e s s e n t i a l l y i n the constant a and i s i d e n t i c a l to the Nernst equation when a = 1.  Kabai obtained the values of con-  stants a and K from plots of the log |log (jZ^oj versus log t which were l i n e a r according to equation (1.7)  log [log : ( ^ ) ] = log K + a. • log t  This equation has no meaning when t = 0 or when G = 1.  (1.7)  Changes i n the  nature and concentration of the e l e c t r o l y t e only influenced K according to the 'empirical' equation (1.8), namely n  K = B •e * '  a  (1.8)  —ct  where B i s a constant [t ] , n i s the concentration of the acid [gm.eq/liter] and y i s a constant, for a given acid ' [liter/gm.eq].  Kabai  obtained the a c t i v a t i o n energies f o r the dissolution of the various oxides i n acids from Arrhenius plots of log K versus ^  ant  ^  w a s  able to .  derive equation (1.9). AH^ =  6  •a  (1.9)  where a i s the structure factor as defined i n rate equation (1.6) and 6 = 21.2 keal/mole i s the energy required f o r the dissolution of any oxide independent of i t s composition and structure and of the properties of the e l e c t r o l y t e .  It -eari be shown however that expression (1.9) f o r ,  20  the a c t i v a t i o n energy obtained.by Kabai depends on the mathematical form of h i s rate, expression (1.6) and that (1.9) does not give the true activation energy.  Indeed, the true a c t i v a t i o n energies should  take into account the v a r i a t i o n of rate of leaching of the oxides with increasing temperature for a constant amount of dissolved mineral, C, i . e . constant surface area, and this condition leads to equation (1.10)  5  (1.10)  K„  where the subscripts 1 and 2 refer to the absolute temperatures. and T^.  The corresponding r a t i o of the rates of leaching at the two  temperatures i s then given by  Rate 1  K  Rate 2  K  l  .  V  1-a  _K  K  2  1  1 a (1.11)  2  and hence r e l a t i o n (1.11) y i e l d s 1 •S  K  a  - J L  = k  RT  (1.12)  where k i s a rate constant independent of temperature and E i s the true a c t i v a t i o n energy for the d i s s o l u t i o n of the oxide.  Kabai, however,  postulated that K = A  AH' RT  From the comparison of (1.12) and (1.13) the following r e l a t i o n i s derived:  (1.13)  21  E =  a  =  (1.14)  6  It follows that the true activation energies E for the dissolution of the oxides in acids are equal to the <5 defined by Kabai..  This leads  to a remarkable suggestion, namely that the activation energy for the dissolution of any oxide in whatever acid does not vary by much more than 6 kcal/mole, i.e. E = 20.00 kcal/mole ± 3 kcal/mole  (1.15)  and this can be seen in Table 3. Cuprous oxide seems to dissolve with a much lower activation energy but this oxide also shows some particular behaviour during dissolution as w i l l be discussed in the present investigations.  1.3.2  Mechanisms of Leaching of Metal Oxides The hypotheses developed to explain the observed kinetics of  leaching individual oxides are basically of two types.  In the f i r s t ,  developed by Wadsworth and Wadia (49) for the leachingvof C^O no hydroxylation or charging of the oxide surface i s assumed and the following sequence of steps i s envisaged: |Cu20 + H2S0  cti'o  (1.16)  (aq)  --'H2SO4--  (1.17)  (aq)  22  TABLE 3 Activation Energies (kcal/mole) for the Leaching of Various .Oxides in Various Acids  Oxide  Fe(OH)  3  Acid  E* (kcal/mole)  IN HCl  References.  20.17-22.18  (53)  a-FeO-OH  HC1,H S0 ,HC10  17.8-22.5  (9,42)  a-Fe 0  HC1,H SO ,HC10  19.2-22.9  (41,42,43,53)  HC1,H S0 ,HC10  14.7-22.18  (47,53)  10.5  (49)  0.5N C^COOH  18.1  (53)  Mg(0H)  0.75N H B0  17.28  (53)  Cr(OH)  0.7N H S0. 2 4  21.3-23.12  (53)  0.5N HCl  22.86  (53)  2  3  A1(0H)  3  Cu 0 2  2>  2  3  Mn(OH)  2  4  4  2  H  Cu(0H)  *E =  2  4  4  4  2 °4 S  3  o  3  for the results reported by Kabai (53).  23  |Cu 0 • H SO  + H  2  +  (aq)  k — ^  Cu " + Cu° + H 0 + HSO~ (1. (aq) (aq) 44  2  Equation (1.16) represents the hydrolytic adsorption of ^2^4  o n  t  '  i e  Cu 0 surface and the f i r s t leaching reaction (1.17) the thermal decom2  position of occupied surface sites. indicates the influence of H^0  +  on which H^SO^  i s adsorbed.  The second leaching reaction (1.18)  ion and i t s ability to react with sites  A rate equation (1.20) can be developed  which includes a Langmuir type equation. (1.19) for the fraction of sites, 9 , covered byJ H„S0,: x 2 4 ••  K  3  • [H S0 ] 2  4  = 9' • [k . • k' - * ( ) '+• k  and  • k. ]  H+  1  ,  X  .  1 9 )  [H 2 SO 4 I  i + KX  x  ( 1  O  N  ;  _ .  2.  (1.20)  1  O  (k includes the surface roughness factor). Q  When  • [H S0 ] i s much greater than one the value of 0^ approaches 2  4  one and equation (1.20) becomes the linear portion of the rate versus [H ] plot as shown i n Figure 1. +  be equal to 1.59 x 10  Note that Wadsworth calculated  liter/M, and . .  Cu 0 would be saturated by H S0. at 9 I 2 4 x o  o  to  thus,; the active surface of  = 0.9 for [H„S0.] = 5 x 10 2 4 / . (aq)  —6  M/liter. Other workers who have used the above hypothesis of an uncharged :  surface (or have not taken into account the variation of charge at the oxide surface) are Koch (51) for the dissolution of BeO i n HCl, ^2^2^*4 and H S0 , Pearson and Wadsworth (55) for the dissolution of U0 2  4  2  in  24  1.2 h  0  I  1  1  0  0.4  0.8  Figure I.  1  '  1.4 1.6 ( H ") ( M/liter) 4  Rate of leaching of Ci^O in H£0 and HCI0 at various concentrations of H ( T = 3 l ° C ) ( After Wadsworth (49).) 4  4  +  Rate of leaching of goethite in HCI activity  of  HCh ( T =  85  °C ) ( After  versus the  SuranaO)  )  mean  25  carbonate solution, Takeuchi et a l (56) for the dissolution of ThC^ in hydrofluoric acid and n i t r i c acid mixtures, and Judge (57) for the leaching of SiC^ in hydrofluoric acid solutions. The second hypothesis assumes that the surface of the oxide becomes hydroxylated and then charged by protonation or ionization according to equations (1.21) and (1.22) for goethite in dilute HCI (9) +  K  +  1  |0 - Fe - OH + H 0 =^= s (aq)  |0 - Fe s  3  ,  |0 - Fe  s  K  + 2^0  .:  + C l(,a q ). = ^ r ' " | 0 - Fe - Cl  s  l 0 - Fe - Cl - i —  (1.21)  ;  (1.22)  k  1  s  Fe 0 C l , . (aq)  (1.23)  (Rate determining step)  This leads to a simple rate equation of the form: R = K • [(0 - Fe - OH] • a (K = h±  • K±  H +  • aQ1 _  (1.24)  • K) 2  In equation (1.24) [|0 - Fe - OH] i s assumed to be large in comparison s with [|0 - Fe ]. For perchloric acid no specific adsorption of the s anion i s expected (58) and the rate determining step then becomes desorption from a simple protonated site.  Equation (1.24) however,  cannot describe the reactions of "-solutions containing high concentrations of HCI with goethite (Figure 2), nor can i t account for the 'two  26  part' type leaching curves observed .in the sulphuric acid leaching of cuprite (Figure 1). Several workers have included the charging of the oxide surface in acids into their studies of the kinetics of leaching.of the oxides. Biermann and Heinrichs (59) proposed a qualitative mechanism for the dissolution of chromite in sulphuric acid based on an i n i t i a l protonic attack, followed by formation of various sulphate complexes of chromium.  A mechanism for the dissolution of gibbsite in perchloric, hydro-  chloric and sulphuric acids based on the protonation of the hydrated gibbsite surface has been advanced by Packter and Dhillon (47).  They  proposed a common rate expression (1.25) for the three acids. R = k • .a„. • a — •w  2  (1.25)  with k a rate constant typical for each acid, a^ the mean activity of the acids and a^ the activity of water i n the corresponding acids.  1.4 The Leaching of Metal Oxides Involving Electron-Transfer 1.4.1  Kinetics of Electron Transfer Reactions Heterogeneous electron transfer reactions at the oxide-electro-  lyte interface are similar to homogeneous electron transfer reactions in solution for which there is ample information i n literature.  Major  theoretical treatments of electron transfer have been given by Libby (60), Weiss (61), Halpern and Orgel (62) , Hush (63) , Sacher and Laidler (64), Marcus (65) and Ruff (66) . The b r i l l i a n t experimental work of Henry Taube and his associates (67, 68) forms a most important chapter  in the recent studies of electron transfer reactions.  Electron-trans-  fer is restricted by the "Franck-Condon" principle, i.e. the "electronjump" process involving a net transfer of an electron from an orbital belonging essentially to one metal to an orbital belonging essentially to the other metal occurs in a time short (^10  sec) compared to -13  that required for nuclear position change (^10  sec).  There are  two major consequences of the Franck-Condon^ principle for electron transfer reactions.  The f i r s t is that the total energy of the react-  ants' activated complex must be identical with the energy of the products' activated complex.  That i s , the energy of the'activated complex  as described by nuclear coordinates must be two fold degenerate, and degenerate in a special way that places the migrating electron on one reactant before transfer and on the other after transfer. When two complex ion reactants share one or more ligands of their f i r s t coordination, spheres in tbe activated complex, i t is termed an innersphere activated complex and the mechanism an inner-sphere mechanism. Outer-sphere activated complexes are formed when the inner coordination shells of the reactant complex ions are left intact as to the number and kind of ligands present. A generalized pathway for inner-sphere electron transfer has been given by Sutin (69) which is represented by the following sheme; =r=  A+ X + B AX + B =•= AXB  =•=  ~AXB  +  AX + B  (1.26)  AXB  (1.27)  ~AXB  (1.28)  +  =s=  A~ + BX  +  A~ + BX  (1.29)  +  A" + X + B  +  (1.30)  28  This i s an example of the oxidation.of cation B by cation A in the presence of anion X. Possible rate determining steps are the formation of a reactant complex (1.26) or a precursor complex (1.27), the electron transfer step (1.28) or the dissociation of the successor complex (1.29). The bridging group X i n an inner-sphere perform several functions.  activated complex can  Libby (60) stressed the importance of  reducing coulombic repulsions between two cations with an intervening negative ion.  But, additionally, the negative ion might complex the  reducing agent as i t is oxidized, generally stabilizing i t in the higher valence state.  However, in general the available data on redox reac-  tions do not show that coulombic attractions and repulsions play a dominant role.  Perhaps the most important factor in bridging is that an  easier pathway for an electron transfer is made. The electron-transfer reactions between FeCjI) - Fe(.III) complexes may be of special interest in the present work. Many anions catalyze the Fe(.II-) + Fe(III) electron exchange. . Exchange paths involving F . Gl , ^O^, SO^, EDTA, phenanthroline,  and CN  are known. Rate constants  arid enthalpies and entropies of activation, when known, are listed in Table 4. Attempts to interpret the kinetics and establish the mechanism of the Fe(:II) - Fe(.|III) exchange have tended to f a l l into two principal categories - anion bridging theories (67, 68) and water bridging theories (79).  If the electron is transferred across an anion bridge, one might  reasonably, expect that the activation energy of the exchange process  TABLE 4 Rate Constants and Enthalpies and Entropies of Activation for Various Homogeneous Ferrous-Ferric Electron Transfer Reactions  Reaction  y (g.eq/1)  Fe Fe 2+ 2+ Fe + FeOH 2 +  3 +  +  •  F e + FeF 2+ + Fe 2+ + FeF Fe + FeF 2+ 2+ Fe . + FeCl 2+ + Fe + FeCl 2+ + Fe + FeC 0. 2+ Fe^ + Fe(C 0 ) 2 +  2+  T° (°C)  0.55  0  0.55  k ( 1 ) V mole sec'  2  4  2  -25  (70)  0  io  6.9  -18  (70)  0.50  0  9.7  8.6  -21  (72.)  0.50  0  2.5  9.0  -22  (72)  _  (72)  3  0.50  0  0.5  —  0.55  20  29.0  8.3  -24  (7Q)  0.55  20  53.0  9.5  -20  (70)  0.55  0  7xl0  9.2  -14  (73)  0.55  0  3.6xl0  -  (73)  -  (74)  Fe  + FeS0+  0.25  2.5  Fe  2 +  + Fe(S0 )~  0.25  25  Fe  2 +  + Fe(EDTA)"  Fe  2 +  + Fe(ph)  Fe(ph) + Fe(ph) 2+  25  3+  3+  Fe(CN)^ + Fe(ph) -  Fe(CN)4- + Fe(CN) 6 " J  0  -  1.94xl0  -  <4xl0~  0  >10  25  >10  3.7xl0  4  0.2  5  8  355  o  (74) (75)  4  —  0.1  -  3  4  25  -  2  692  3+  J  References  9.3  2 +  4  AS (e.v.)  0.87  2  o  AH /kcal\ \ mole '  -37  (76)  -  (71)  _  4.1  (77) -32  (78)  30  should change as the complexing anion i s changed (80) . However', i f the exchange involves a water bridge a marked heavy water isotope effect, even for the anion catalyzed processes, and l i t t l e dependence on the nature of the complexing anion is anticipated (81) . Home (73) calculated an activation energy of 9.2 kcal/mole for the electron-trans++ fer process of Fe  + + FeC^O^ which lies well within the range of values  reported for catalysis by OH , F , Cl (Table .4), This, according to Home (73) , provides further evidence in favour of water-bridging  rather  than anion-bridging in the Fe(II) - Fe(jill) electron-exchange reactions. The mechanism of the oxalate catalyzed Fe(li) - Fe(;III) electron-transfer proposed aby Home (73) is based on the rapid formation of an activated complex by reaction of the ferric oxalate ion and the ferrous ion: ( H  +  2 4 or 5*2°t < 2°C ^ F e  0 )  C  Fe  i  H  \Z..-(1.31)  [( 2 H  0 )  4or5  F e  * 2°4^¥V  W  C  The function of the centrally located complexing anion is to overcome coulombic repulsion, form a stabilized activated complex, and draw the reactants into sufficient proximity so that their solvation spheres overlap.  The actual effective electron transfer then proceeds via  waters of solvation adjacent to the complexing anion:  [ ( H  2 3 or 4 * — 0 )  F e  C  2°4  F e  <  H  0 ) 2  5  ]  31  and.the final step i s the dissolution of the activated complex and any rearrangements of the waters of solvation:  [ ( H  2 4 or 5 * 2 ° 4 0 )  Fe  C  <  F e  H  0 ) 2  6  =  ]  (1.33)  +  Fe*(H 0) 2  6  + (H 0) ^ 2  4  5  FeC^  Conflicting evidence was brought up later by Sheppard and Brown (82) in their study of the catalyzed electron-transfer reactions of Fe(T.I) Fe(III) by acid phosphate, oxalate and sulphate anions.  The large  energies of activation, 15, 13.5 and 21.0 kcal/mole for H^PO^, 0,^0^ and SO^ respectively, suggest that the process of electron transfer for these oxyanions may be different from that for the halide paths. The transfer of electrons between a metal or semiconductor and a dissolved or surface-bound reactant is not different in kind from homogeneous solution processes described above. Laxen (83) has compared the rate of dissolution of U0  2  in the presence of Fe  with the rates  of electron transfer between Fe(.'II) - Fe(III) complexes in solution. In a perchlorate medium both the Fe(II) - Fe(III) exchange and 3+ the dissolution of U0  2  by Fe  were strongly catalyzed by the presence  of small concentrations, of sulphate, while both reactions are also affected in a similar manner by the concentration of H  +  in solution.  Of the anions tested, NO^ did not improve the dissolution rate and Cl had only a slight effect.  In perchlorate solutions, when the H  +  addi-  tion was increased the dissolution rate increased up to a pH value of 2 and decreased at lower pH values.  The increase in dissolution rate  32  was  2+ ascribed by Laxen to the increase i n concentration of Fe(OH) ,  the most e f f e c t i v e electron-transfer species.  I t should be noted  2+ that the maximum of Fe(OH)  concentration does not occur at pH 2  according to Needes.and F i n k e l s t e i n (84) and this p a r t i c u l a r aspect suggests that other factors may be involved i n the leach. In sulphuric 3+ acid s o l u t i o n s , the rate of leaching of UO^ by Fe with pH also reached a maximum at pH - 2 and could be attributed to the combined e f f e c t of  A.  I|  Fe(OH)  —  , FeSO^ and YeiSO^)^ species i n s o l u t i o n . At constant pH,  however, the rate of d i s s o l u t i o n of -UO^  showed a square root dependence  on the concentration of f e r r i c ion i n s o l u t i o n . The very high rates of d i s s o l u t i o n of UO^ reported by Hunt and 3+ 3+ Taube (85) with Fe (dipy)^ and Fe(0-phen)^ i n IM HCI serve to confirm the c o r r e l a t i o n between rate of d i s s o l u t i o n of UC^ and the very fast 3+ homogeneous electron transfer of these two complexes with Fe  in  solution. Recent work published by Needes and N i c o l (86) on the oxidative d i s s o l u t i o n of UO^ i n d i l u t e perchloric acid showed that the order of leaching rates of UO^ with various oxidants was T1(:III) >, VO^ > Fe(III) H g ( l l ) , whereas the equivalent order of electron-exchange rates was VO~2 >' H g C t ' l )  > Fe(III) > Tl(.III).  The conclusion to be drawn from  this information i s that the rate of d i s s o l u t i o n of UX^ i s a function of both the p o t e n t i a l and the electron-exchange rate of the redox couple used. 2+ The reductive d i s s o l u t i o n of Mn02 i n Fe has been investigated by Koch (87).  containing acid solutions  The rate of leaching of Mn02 i n  33  Fe  containing sulphuric acid solutions was two orders of magnitude  larger than in perchloric acid solutions of equal strength. . Koch (87), however, excluded the possibility of an electron-transfer rate-controll2+ ing.step because.the rate of dissolution of MhO^ by Fe was independ2+ + ent of the concentration of Fe , H and H^SO^. It should be empha2+ sized here that Koch used large concentrations of Fe  (0.05-0.075  M/liter) and that the possibility of surface, saturation by the active ferrous species should be 1.4.2  considered.  Mechanisms of the Leaching involving Charge Transfer at the Oxide-Electrolyte Interface To date,essentially two types of mechanisms have been developed  to explain the observed kinetics of leaching individual oxides involving an oxidation - reduction step. In the f i r s t , developed .by. Mackay and Wadsworth (10) for the oxidative leaching of UO^ in dilute acid, the formation of an activated complex of uranium at oxide active surface sites is postulated, followed by charge transfer through the activated state to form a U(VI) inter2+ mediate and desorption of UO^  in solution; the following sequence of  steps i s envisaged: [UO s  K i  + HO  OH |0U s *"0H  K =  OH  2  |0U" +0, s -OH (aq) 2  OH lou'" s ^ OH 0~ |0U'" _ + 2H s ^0 k  | s  A c t i  C O m  P  ^  (1.34) (1.35) k a t e d  l e X  2 HO: + (aq)  +  +  u c  2  2  OH-  34  Equations (1.34) represents the formation of a hydroxylated surface and equation (1.36)  the reaction of these hydroxylated sites  with dissolved oxygen producing a surface activated complex of U(VI) which is then readily soluble in the electrolyte.  The surface hydro-  xylated sites are in equilibrium with the solution according to the deprotonation equilibrium equation (1.35).  A rate equation (1.38)  can be developed which includes a Langmuir type equation (1.37) for the fraction of sites, 0, covered by hydroxyl ions: +  [H ] 0=  and  K  2  [H ] . +  2 +  =9 •  2  ( 1  • pQ  '  3 7 )  (1.38)  2  The important feature of the f i r s t type of mechanism i s that no attempt i s made to subdivide the overall reaction into anodic and cathodic reactions. A mechanism involving the formation of a surface activated complex has also been considered by Warren and Devuyst (88) r.in^an^.attempt: to explain the kinetics of the reductive dissolution of pyrolusite by hydrazine i n ammonium carbamate solutions, and the reductive dissolution of manganese dioxide in the presence of S0 was 2  approached in a similar manner by Herring and Ravitz(89). The second hypothesis which was proposed by Habashi and Thurston (90) for the mechanism of the oxidative dissolution U0 assumes that 2  the)dissolution of this oxide proceeds by an electrochemical mechanism in a similar way to the corrosion of metals.. Habashi and Thurston  35  propose that the following two electrochemical reactions proceed simultaneously: 0  + 2H 0 .+ 4e — 4 0 H 2  cathodic reaction '  (1.39)  anodic reaction  (1.40)  2+ U0  2  — —  U0  2  + 2e  In general, the rate of the cathodic reaction can be given by: V = k -A • [D] c c c n  where k  c  (1.41)  i s a rate constant, A the cathodic surface, [D] the concenc  tration of the depolarizer, i.e. 0  and n i s the order of the reaction  2>  with respect to the depolarizer. The rate of the anodic reaction can similarly be given by the equation: V = k • A • [C] a a a  (1.42)  m  where k i s a rate constant, A the anodic surface fraction, [C] the a a .+ concentration of a complexing agent, i.e. H , and m* the order of the ;  reaction with respect to the complexing agent.  At the steady state,  V = V , but, since A + A = A, total surface area of the oxide, suba c a c stituting the value of A^ in the rate equation giving V ,  V  C  = V a  k =—  •k  • A-[D] [C] = k-[D] + k.[C] c a n  n  m  At high concentration of C, or i f k  m  (1.43)  i s large, the velocity equation  3.  (1.43) simplifies to : V = V = k • A • [D] c a c  n  (1.44)  36  and at high concentration of [D] , or i f k c is large, the rate equation (1.43) becomes: V  = V a  = k c  - A- [C]  m  (1.45)  a  An alternative model for the electrochemical dissolution of oxides was recently proposed by Needes and Nicol (86).  In this model  i t is assumed, in agreement with Habashi and Thurston (90), that the overall reaction corresponding to the dissolution of an oxide can be subdivided into an anodic and a cathodic part;.  A fundamental mathe-  matical expression for the relation between the current density and the overpotential n is given by the Butler-Volmer equation: i =  i^j^exp ((1-a) |fr'n)-  exp  (-o^|r  n)J  (1.46)  where i , the exchange-current density, represents the speed of the o forward and backward reactions at equilibrium, n'represents the difference between the applied potential and the equilibrium potential of the reaction, and F is the Faraday constant.. The transfer coefficient a is defined as the fraction of the overpotential contributing to the increase in the rate of the reaction. Experimental values of a are often found to be close to 0.5.  The exchange - current density, i ,  is directly proportional to k Q , the potential independent rate-constant of the reaction at the surface.  Thus, the larger the value of k Q , the  faster w i l l be the rate of electron transfer at the oxide surface. The potential, E M , at which the anodic and cathodic currents are equal is termed the "mixed" or "open circuit" potential, i.e. the  37  potential at which no external current is flowing.  In the case of an  oxide dissolving by an electrochemical mechanism in which there is no barrier to the dissolution or complexing of the species at the surface of the oxide once charge transfer occurs, the dissolution current density is a direct measure of the rate of dissolution of the oxide. Under these conditions, i t can be seen from equation (1.45) that the rate of leaching of an oxide w i l l also depend on the equilibrium potentials of the oxide and the redox couple, since n a  1  c  1  in  where n  and n  =  E° - E a M  (1.47)  =  E°- E c M  (1.48)  are the anodic and cathodic overpotentials, E ° and E °  are the anodic and cathodic equilibrium potentials, and E ^ is the mixed potential defined above.  1.5  C r i t i c a l Summary A major difficulty of considering previous studies of the leaching  of oxides is that few extensive studies of single oxides have been made. Additionally, the range of conditions used by various workers to study individual oxides varies from one to the other.  Although  the effect of acid anions upon the rate of leaching of several metal oxides in different acids can be correlated with the complexing affinity of the anions for oxide cations, there appears to be no explanation of the observed effects of the variation of concentration of acids on  38  the rates of leaching of metal oxides. A satisfactory general hypothesis of the mechanism of the direct dissolution of oxides must be able to explain in addition to the effects of anions on the relative rates, at least the following observations : An apparent dependency of the rates of leaching of a - F e 2 0 3  (a)  and a-FeO'OH in perchloric acid over the range of acid concentration studied (0-1.5N) on either [H ] or [HCIO4] added, whilst the +  rate of leaching of C^O i n the same acid appears to show some type of 'saturation' dependence followed by a rate which appears to be proportional to ([H ] + C) or ([HCIO4] + C) (where C i s +  some constant) (Figure 1 ) . A dependency of the rates of leaching of a-¥e20^  (b)  and a-Fe0*OH  2 in low concentrations of hydrochloric acid (<2.5M) on either a ., H+ 2 a or a , • a ,and at high concentrations an apparent, dependClH+ Glency on a , or a (where a = activity of various ionic species) H+ Cl" (Figure 2). (c)  An apparent 'saturation dependency' of leaching rate of  C^O i n dilute sulphuric acid and a possible similar behaviour by goethite, which causes both oxides to leach by a rate 'law,- in stronger sulphuric acid which shows an apparent dependency on ([H ] + C), ([HS0~] + C) or even possibly .([S0'~*I + C) . +  (d)  The widely differing rates of leaching observed for Cu^O and  for BeO i n sulphuric acid under the same conditions (C^O leaching, about lO^x faster than BeO).  39  For the leaching of metal oxides involving a change in oxidation state during dissolution, a mechanistic model must be able to explain at least the following observations: (a)  A maximum of the rate of leaching of UC^, in the presence  of Fe HC10.  , . occurring at a pH value of approximately 2, in both and H„SO, solutions.  (b) The large differences in rates of leaching of MnO^ with I I Fe-,'  in  HCIO^ and H^SO^  solutions of equal normality and indeed  the same for the leaching of UO^ in the presence of Fe (c) The square root dependency of the rate of leaching of UO^ with Fe on the concentration of Fe ,whilst the rate of leaching of this oxide with 0^ shows a f i r s t order dependency on pC^.'. Whether or not hydroxylation of the oxide surface has to be considered in a general mechanism, for the leaching of oxides is open to question.  If oxides adsorb water dissociatiVely very rapidly to the  extent of one hydroxyl group per metal atom, no distinct behaviour difference might be observed between a totally hydroxylated or bare oxide surface.  This may justify Wadsworth's and Wadia's (A.g.) choice of  a bare cuprous oxide surface i f this surface becomes rapidly hydroxylated in comparison to the overall rate'of dissolution of this oxide. However, Peri(12) has shown that y~  an  a  d ~^2®3  ^  e a t - t r e a t e d a t  temp-  eratures above 800°C do not rehydroxylate rapidly. If hydroxylation of the oxide surfaces is indeed a prerequisite for dissolution, and i f this under some conditions, in the case of a-klJO^ is the slow step in the overall leaching of this oxide, the rate of leaching of a-AlJd^  40  .  v  ,  ,.  would be expected to depend only on the activity of water. As already discussed, a net positive charge develops on the hydroxylated oxide. surf ace in solutions of pH below the pH of Z--P.C of the oxide.  +  This charge is established by H , OH  and anions of the acid  present in solution and may arise in one of the. following ways: (a)  +  Simultaneous or consecutive adsorption of H , OH  and anions  of the acid at the oxide-electrolyte interface. (b)  Adsorption of undissociated acid at neutral oxide surface  sites and of H  +  at the same oxide surface sites.  Warren et al(\8,9) suggest that process (a) occurs during the leaching of ferric oxide in perchloric and hydrochloric acids, whereas Wadsworth and Wadia(49) postulated process (b) to explain the leaching of cuprous oxide in sulphuric acid. It thus appears important to study the leaching kinetics of cuprous oxide in both perchloric and hydrochloric acids, as neither process (a) nor process (b) seem to be sufficient to describe the kinetics of the leaching of oxides in general. Leaching studies in which the concentration of anionic species are varied independently or in a controlled manner should.be able to indicate whether anions or undissociated acids of the anions are taking part in the leaching of oxides.  Although i t is observed that the  complexing power of anions in solution for the oxide cation has a large effect on the leaching of metal oxides, i t is not clear i f this effect is due either to the preferential adsorption of the anion or to the enhanced desorption rate of metal-anion complexes from the oxide surface.  In-  deed, Berube and De Bruyn(36) and Dumont and Watillon(37) have corre^  41  lated the driving force for adsorption of anions to t h e i r action upon surrounding water molecules, and thus of water adjacent to the oxide surface,and found l i t t l e correlation between the adsorption sequence and the complexing a f f i n i t y of the anions for the oxide cation. Kabai (53) proposed an empirical equation correlating the rate of lea_ching of oxides- to • the^concentration of the acid as exponent of an exponential, but l i t t l e fundamental information i s obtained from this relation.  F i n a l l y , several workers (8,9,41,45,46) have suggested that  the adsorption a f f i n i t y of an anion may be associated with the complexing  power of the anion for the oxide cation.  C l e a r l y , study of a se-  l e c t i o n of various acids which produce anions having d i f f e r e n t complexing  power f o r an oxide cation might provide more insight into the role  of anions i n the leaching of metal oxides. The pH of Z.P.C. of an oxide may be a very important characteri s t i c for the leaching of oxides i n a c i d s , mainly for two reasons: (a)  I t gives an indication of how  favourable production of a  net p o s i t i v e charge at the oxide surface i s with decreasing pH. (b)  I t may be related to the anion-exchange capacity of the  oxide surface. The concept of Z.P.C. has been i n t u i t i v e l y used by various investigators ,(8,9,47,59). Warren et al(8,9) for example represented the rapid formation of an excess positive charge by an equilibrium equation involving the adsorption of H  +  ions at the oxide surface. This  equation of course, would only be acceptable i f the pH of the solution i f far enough away from the pH of Z.P.C. of the oxide.  The equilibrium  42 + constant for H adsorption is then a measure of the relative tendency of oxides to adsorb H  +  of Z.P.C. of the oxide.  ions and thus might also be associated with the pH The question may now arise regarding the possi-  b i l i t y of saturating the oxide surface in H  +  and eventually in the  anion(s) of the acid. According to Wadsworth and Wadia (49) the cuprous oxide surface is suggested to already become saturated by undissociated sulphuric acid in dilute solutions, but as mentioned earlier, by a process which involves the direct adsorption of the acid at neutral oxide sites.  One could equally suggest that the oxide surface becomes  saturated in hydrogen ions from solution, followed by increasing adsorption of HSO^  ions at these sites.  Studies using oxides of different  pH's of Z.P.C. might bring a solution to the problem of the species involved in the leaching of oxides and to the significance of the Z.P.C. of oxides. So far,no explanation has been given for the large differences observed between the absolute rates of leaching of some oxides, i.e. C^O  leaching 10 x faster than BeO.  Thermodynamically the leaching of  BeO in water at a given pH is more favourable than the leaching of Cu„0 as the change in standard free energies for the reactions: (1.49) (1.50) are respectively -10.7  and -6.17 kcal/mole (91).  This clearly shows  that kinetic factors can overrule drastically the expected driving forces from equilibrium considerations.  The observed energies of acti-  43  vation for the leaching of most metal oxides are nearly constant from one oxide to the other, i r r e s p e c t i v e of the a c i d , as they vary between 17 and 23 kcal/mole (Table 3); this may  suggest that a similar rate-  determining step i s operative during the leaching of metal oxides i n acids, possibly the desorption of metal species into s o l u t i o n . The acid leaching of metal oxides involving an  oxidation-reduc-  t i o n step i n the presence of a redox couple i n solution has been studied, for the most p a r t , under r e l a t i v e l y r e s t r i c t e d conditions., It i s l o g i c a l to expect that surface hydroxylation and charging.may also be involved i n the o v e r a l l k i n e t i c s of the oxidative or reductive leaching of the oxides.  Mackay and Wadsworth (10) have proposed that  oxygen adsorbs at uncharged hydroxylated  uranium dioxide surface s i t e s  and that the concentration of these neutral s i t e s i s increased by the reaction of hydrogen ions and the negatively charged portion of the UO^  surface.  This approach i s consistent with the  properties of UO^  i n acids.  hydration-cnarging  Laxen (83) and Needes and N i c o l (86) how-  ever, did not consider the UO^  surface properties i n their model of  the oxidative d i s s o l u t i o n of this oxide and were indeed unable to explain the v a r i a t i o n of leaching rate of JJO^ with pH i n the presence of F e ( I I I ) . It appears l i k e l y from t h e i r studies that both the oxide-electrolyte ' double layer properties and the type of f e r r i c species present i n solution at each pH have to be considered . i n order^ to explain the observed k i n e t i c s of dissolution of  U0_.  44  2.  SCOPE OF THE PRESENT INVESTIGATION  The present work had as one objective the resolution of the d i f f erences which are apparent i n the proposed mechanisms of leaching of metal oxides. F e r r i c , aluminum, cuprous, cupric and manganous oxides were selected f o r the present i n v e s t i g a t i o n s .  Extensive studies on the  leaching of f e r r i c oxides were planned i n the hope that they might provide a basis f o r comparing the behaviour of oxides i n general. Leaching experiments on aluminum oxides were undertaken to attempt to provide some understanding of the role of surface hydroxylation i n the k i n e t i c s of the leaching of oxides. Studies on the leaching behaviour of cuprous oxide were included i n the present investigations because ..the. kinetics..of •• leaching of this oxide i n sulphuric acid have been explained i n terms of a unique mechanism, involving as a f i r s t step the adsorption of undissociated acid at the oxide surface.  Cupric and manganous oxides were chosen, i n  addition to the other oxides, i n an attempt to correlate;the pH of the zero point of charge (Z.P.C.) of an oxide to i t s leaching characteri s t i c s i n acids. P e r c h l o r i c , hydrochloric, sulphuric and oxalic acids were selected as reagents i n order to study the e f f e c t of anions on the rate of leaching of the oxides.  Amongst other p r o p e r t i e s , these acids d i f f e r  i n the complexing power of their anions f o r metal ions i n s o l u t i o n .  45  3.  3.1  EXPERIMENTAL  Minerals and Reagents  3.1.1  Natural Minerals Massive specimens of Micaceous hematite and goethite were  obtained from Ward's Natural Science Establishment Inc., New York. The hematite originated from Ishpeming, Michigan and the goethite came from Minnesota.  The quantitative chemical analyses and semi-  quantitative spectrographic analyses for both minerals are given in Table A.l, Appendix A.  For a l l the experiments the specimens were  ground in a porcelain mortar and then wet screened to the 65-150 mesh Tyler sizes.  3.1.2  Synthetic Minerals (a)  Hematite  Synthetic a-Fe^O^ was prepared from pigment grade goethite powder obtained from Harrison and Grosfield Ltd., Canada; i t s purity was 99.95% a-FeO'OH with 0.05 of insoluble matter and traces of copper. The goethite powder was calcined at 800°C for 10 minutes producing pure hematite powder. This powder was cold pressed in discs (1 cm in diameter) followed by sintering at elevated temperature for various times and under various conditions as given in Table 5. Titanium doped specimens were obtained by the method used by Morin (92)<Reagent grade.  HO2 pigment powder from Matheson Co., U.S.A.,  was wet mixed with the hematite powder, followed by" the sintering  TABLE 5 Synthetic Hematite Specimens  Sample No  Pressure (psi)  Atmosphere  air  Sintering T°(F) Time/days  Fe  Ti (%wt)  Uwt).  0"  0.12  2 +  2,100  4  1,650  3  0  0.24  2,100  2  1.3.  1.52  air  2,100  2  0  0.24  15,000  air  2,100  2  3.0  2.94  F  15,000  air  2,100  2  0.5  .0.65;';  G  15,000  air  2,100  1  0  0.14  H  15,000  air  2,100  8  0  0.10  I  15,000  air  2,100  1  0.1  0.16  J  15,000  air  2,100  1  0.2  0.24  K  15,000  air  2,400  1  0.8  1.03  L  15,000  air  2,400  1  0.4  0.77  P  20,000  air  2,400  2  0  0.2  N  20,000  air  2,400  2  0  A  15,000  B  15,000  C  15,000  °2 air  D  15,000  E  Q=A+D+G+H  ON  99.999 Fe 0.5% Ca  0  0=A+D+G+H X  Remarks  0 15,000  air  2,400  1  0  Mixtures 0.5% Mg  47  operations under the conditions indicated i n Table 5.  Calcium doped  specimens were prepared by a method described by Geiger and Wagner (93) and magnesium doped samples were obtained i n the fashion proposed by Gardner et a l (94). X-ray diffraction patterns of the samples were consistent with the ASTM card for hematite and are reported i n Table A.2, Appendix A. Electron-Microprobe pictures of the T i and Ca doped a-YeJd^ surfaces are shown in Figure 3a to 3f. Clustering of T i i s observed for Ti contents of 1.5.and 3.0%. No homogeneous Ca doping could be obtained. (b)  Cuprous Oxide (Cu20)  Cuprite was obtained by the thermal oxidation of pure copper wire, at 900°C under air for 24 hours.  The C^O powders obtained after  crushing the samples in a porcelain mortar presented the X-ray diffraction pattern characteristics of CU2O as given by the ASTM card (Table A.3, Appendix A). •(c)  Cupric Oxide (CuO)  Cupric oxide was obtained by the further oxidation of CuJO at 700°C, under air, for 48 hours.  The CuO powder obtained after crushing  the sample presented a l l the X-ray diffraction pattern characteristics of CuO as given by the ASTM card (Table A.3, Appendix A).  Traces of  CuJO probably contaminated these samples, approximately 1% by wet analysis. (d)  Manganous Oxide  Manganous oxide was obtained by the reductive roasting of natural pyrolusite (analysis given i n Table A.4, Appendix A) at 900°C under  48  (c)  (d)  Figure 3 . Electron microprobe pictures for Ti or Mg of synthetic a-Fe2C>3 samples (Table 5) (x 1,000) (a) 0.1% T i ; (b) 0.2% T i ; (c) 0.5% T i ; (d) 1.3% T i .  (e)  (f) Figure 3.  (e) 3.0% Ti (f) 0.5% Mg  50  cracked ammonia for 24 hours.  The weight loss indicated that a l l  MnO^ present i n the ore had been converted to MnO. (e) Aluminum Oxides Pure gibbsite (a-Al^O^• 3^0) was obtained from Alcan, Canada. Calcination of the gibbsite powder at 300°C for 24 hours under air produced boehmite (a-A^O^-H^O) , and calcination at 600°C for 24 hours under air transformed the gibbsite into yAl^O^.  Calcination of the  gibbsite powder at 1400°C for 24 hours under air resulted i n the formation of pure a-Al^O^. The X-ray diffraction patterns of the synthetic aluminum oxides mentioned above were consistent with the data given by the ASTM cards and are reported in Table A.5, Appendix A.  3.1.3  Reagents Perchloric, hydrochloric and sulphuric acids were obtained from  Allied Chemical, Canada. Co., U.S.A.  Oxalic acid was provided by J.T. Baker Chem.  Ferrous oxalate was from Griffin and. George, England.  A l l other'chemicals which were used were reagent grade.  Helium and  oxygen came from-Canadian Liquid Air Ltd.  3.2 Apparatus Design Leaching experiments were carried out i n a glass reaction vessel maintained at constant temperature i n a heat controlled water bath ' and open to the atmosphere through a reflux condenser.  The main fea-  tures of the apparatus are schematized on Figure 4. The 1500 mis V . capacity, cylindrical  glass reaction flask was fitted with a gas inlet  1. Water Bath 2. Immersion 3. Contact  Heater  Thermometer  4. Reaction Flask 5. Fritted Glass Filter 6. Sampling T u b e 7. Gas Inlet Tube 8. Stirrer  Motor  9. Reflux Condenser 10. Spin B a r I I.  Magnetic  F i g u r e 4.  Stirrer  Apparatus  Desi  52  tube and a sample tube terminating with a fritted glass f i l t e r .  The  solution in the flask was stirred by means of a Teflon-covered magnet rotated by a magnetic stirrer unit below the water bath vessel.  Heat  was supplied by a 100 watt immersion heater, which controlled the temperature within 0.2°C ,.through connection with a mercury relay which was i t s e l f connected to the contact thermometer. The water bath was stirred continuously by,a variable speed stirrer.  3.3 Experimental Procedure . The experimental procedure consisted of the following steps: (a) The temperature controller was set at the required temperature. (b)  The reaction flask, containing 1000 mis of solution of the  required concentration, was immersed i n the water bath, and the various connections made. (c)  The system was allowed to come to thermal equilibrium.  Flushing with He or 0^, i f desired, was carried out simultaneously. (d)  The powder specimen (usually 1 gm) was added to the solution,  and the flask was closed. (e) ' Stirring of the solution was started and a f i r s t sample taken (usually 5 mis) by applying an 0^, air or He overpressure above the  solution.  The analysis of this f i r s t sample was considered  as a blank for the successive samplings at regular intervals. The f i r s t 10 mis of solution removed i n a l l samplings were .immediately returned to the. flask via the reflux condenser. • (f) The. samples were analysed for the desired metal content.  53  In the leaching experiments of ferric oxides with oxalic acid i  i t was necessary to prevent the photo-catalyzed reduction of ferric ion in solution.  The reaction flask used in these experiments was  covered by black masking tape which prevented light from reaching the solution. The pH's of the 0.2M oxalic acid solutions which were used were adjusted with NaOH and HCIO^ additions and were measured at 80°C against standard buffer solutions of pH 2 and 4 using an expanded pH meter.  The measured and calculated pH's are reported in Table B.l,  Appendix B.  3.4  Analytical Methods  3.4.1  Iron The iron content of the solutions was determined by measuring the  absorbance at 510y of the orthophenanthrbline complex of ferrous ions after reduction with excess hydroxylamine hydrochloride (95). In the presence of oxalic acid the solutions had to be heated up to 60°C in order to ensure complete conversion of ferric to ferrous ions.  However, when the samples contained over 0.2fi(l/liter of oxalic  acid, ferrous oxalate precipitated upon adding the o-phenanthroline reagent buffer solution. acid before analysis.  It was then necessary to destroy the oxalic  This was accomplished by adding an excess of  sodium persulfate (Na S 0 ) to the sample and boiling i t to eliminate o  o  o  2 2 o  the excess of oxidizing agent.  54  Aluminum  3.4.2  The aluminum contents of the solutions were determined by measuring the absorbance at 580y of the pyrocatechol v i o l e t complex of aluminum ions i n an ammonium acetate buffer solution (pH 6.1-6.2)(96).  Copper  3.4.3  The copper contents of the solutions were obtained by measuring the absorbance at 640u of the tetraethylene pentamine complex of copper(Il)(97).  Manganese  3.4.4  The manganese contents of the solutions were obtained by measuring the absorbance at 524u of the permanganate: ion obtained by heating the sample i n the presence of excess potassium periodate (98).  Determination of the Ferrous Content of Hematite Specimens  3.4.5  A f i v e gram powder sample of the hematite was dissolved i n 200 mis of 20% sulphuric acid at 80°C, under an He atmosphere.  The solution  was then cooled and an excess of phosphoric acid was added to eliminate the colour of f e r r i c sulphate.  This solution was t i t r a t e d with a  standardized O.lN eerie sulphate solution i n the presence of i n d i c a t o r . The red colour of the o-phenanthroline ferrous complex changed to the green colour of the f e r r i c complex upon completion of the reaction  „ 4+ . _ 2+ Le  + te  3 + 3 + — •  Ce  + Fe  .  This method was found to be sensitive  to as l i t t l e 0.05% ferrous ion content i n the f i v e gram hematite sample (absolute error of ± 0.02%).  55  4. RESULTS  A l l the rates of leaching were obtained from measurement of the i n i t i a l slopes of the plots representing the amount of metal dissolved versus time.  It was usually observed that the rate of leaching of an  oxide did not vary, with time up to 10% dissolution of the contained metal.  If the rate of leaching of the oxide was indeed varying stead-  i l y with time i n the early stages of the leach, the experiment was repeated i n order to obtain the best approximation of the i n i t i a l rate.  4.1  The Leaching of Metal Oxides in Aqueous Perchloric Acid Solutions The rates of leaching of cuprous oxide (C^O), cupric oxide (CuO)  and ferric oxide (a-Fe^O^, Michigan) were measured at constant temperature as a function of the concentration of the acid (Figures 5 and 6). Hay's (99) and Surana's; (100) results on the rates of dissolution of goethite (a — FeO'OH) are also included on Figure 6. The absolute rates of leaching of the oxides vary widely from one to the other and do not serve as a convenient basis for comparison. Hence only relative rates of leaching, that i s the ratios of the actual rates of dissolution of the oxides over their rates of leaching in 0.9M HCIO^, were plotted against the concentration of HC10, (Figures 5 and 6).  0  0.4  0.2 Figure 5.  ( HCI0 )  ( M/liter)  4  Relative  rotes  of  leaching  concentration  of  HCI0  4  0.8  0.6  of  Cu 0 2  and  CuO in  at 12 ° C . ( T a b l e  B.5  HCIO^ ,  versus  the  Appendix B ) •  Calculated  b  Measured  o o  o:  3  4  5  ( HCI0 ) 4  Relative (T=  rate af  9 0 ° C)  leaching of goethite  versus the  8  6  (M/liter) (Surana  concentration  of  and HCI0  H a y , T= 110° C ) and hematite 4  ( Table B 5  , Appendix B ) •  58  The rates of leaching of the oxides show a f i r s t order dependence on the concentration of added perchloric acid in dilute solutions, but a lower order in more concentrated solutions; this divergence of the rates to a lower order dependence on the concentration of HCIO^ begins at different acid concentrations for the various oxides, approximately in the order O.IM  (Cu 0) < 0.3M 2  (CuO) < 1.5M  ( a - F e ^ and a-FeO-OH) .  Very large differences are observed between the absolute rates of leaching of the oxides at a given concentration; in 0.9M HCIO^ at 12°C, for example, the absolute rates of leaching are approximately correlated in the following way: rate (Cu„0> = 9x rate (CuO), - -8xl'0x rate (a-FeO-OH) Z '4. - 2.7xl0v rate ( a - F e ^ ) . . 7  x  The relatively high energies of activation obtained for the leaching of the oxides (42,100) suggests that the dissolution was not controlled by diffusion under the conditions of the experiments.  The  rates of leaching of the oxides were not dependent on solution agitation and this i s in support of the.statement above.  Finally, experiments  with varying amounts of ore samples showed, that the slow step in the leaching reactions for a l l the oxides was heterogeneous in nature. In the present investigations, the maximum concentration of HCIO^ had to be limited to IM HCIO^ for the leaching of Cu^O and CuO powders, and 6M HC10.  for the dissolution of a-FeO-OH and a-Fe„0 , because in the  59  former case the rates of leaching became too large to measure accurately and in the latter diffusion control of the dissolution appeared to be unavoidable due to the increase of solution viscosity.  4.2 The Leaching of Metal Oxides in Aqueous Hydrochloric \• Acid Solutions The rates of leaching of ferric oxides (goethite, hematite), cuprous and cupric oxides (C^O and CuO) and aluminum oxides (gibbsite, y-Al^O^ and a-Al^O^) were investigated at constant temperature as a function of the concentration of the acid and were plotted against the calculated mean activities of HCI (Table B.6, Appendix B). The absolute rates of leaching vary greatly from one oxide to the other and only relative rates were plotted, every observed rate being divided by the rate of leaching of the oxide i n 1.2N HCI. The rates^of dissolution of ferric oxides show a second order dependence on the mean activity of HCI (^ a .) in dilute solutions +  (a+ < 1) (Figures 7 and 9), slowly decreasing to an apparent f i r s t order dependence on a+  i n concentrated  solutions (a+ > 1.0) (Figures  8 and 9). Earlier results on the rates of leaching of various natural ferric oxide powders obtained by Bath (101), Surana (100) and Roach' (42) are included in Figures 7,8 and 9. The transition from the second order to an apparent f i r s t order dependence of the rate on +. i s :a  shown on Figure 9 i n a plot of the ratios of the relative rates of leaching and the mean activity-of HCI versus the mean activity of HCI. Such a plot yields a f i r s t order relation of the ratios with • + ^ for a  rates which show a second order dependence on a+ . and a zero order  60  4  3 QHRelative rate HCl  of  ( Mo Ial)  leaching of  as a function of the  ( Tables  f e r r i c oxide  mean activity of  B . 7 and B 7a , A p p e n d i x B ) .  in HCl .  dilute  Relative r a t e mean  of  activity  leaching of of  HCI  ferric  (Table  B . 7,  oxide  in HCI  as  A p pendix B ) .  a function  of  the  Calculated  A"  '6  I  work  80 • c  A  This  O  Bath  85  • c  •  Bath  85  • c  A  Roach  80 • c  •  Surana  85 • c  JL  0  10  20  Figure 9 .  30  ( Molal) Ratio  of the r e l a t i v e  function  of  a  +  •  rate of l e a c h i n g of ( Table  ferric  B-8 , Appendix B )  oxide  a n d a+ v  as  a  63  dependence on a + .for rates which exhibit a f i r s t order dependence on a+.  Figure 9 clearly shows that the rates do not become proportional  to a + up to a+=30. Most of the earlier studies on the leaching of ferr i c oxides are in good agreement with the present work, with the exception of some of Bath's results (Figures 7, 8 and 9).  Bath obtained  rates of leaching of synthetic a-Fe^O^ powders which appear to yield much higher relative rates of dissolution in dilute HCI solutions when his results are compared with the present work in concentrated solutions.  Moreover, his results suggest that the relative rates of leach-  ing of ferric oxide are truly second order with respect to a + up to a+—1.8 and then become rapidly proportional to a + at higher activities (Figure 9).  It should be mentioned however, that Bath used 0.1 gm/liter  powder samples in his experiments whilst a l l other workers used at least 1 gm/liter powder specimens. Additionally, Bath himself obtained contradictory results for his experimental rates of leaching of the basal plane of a a-Fe20^ single crystal.  It transpires that these lat-  ter results are in agreement with the present work (Figures 7, 8 and 9). The rates of leaching of ferric oxide appears to depend on both the activities of the hydrogen and the chloride ions as can be seen in Figures 10 and 11. and 2.4M  Figure 10 shows the effect of adding 0.6,  of LiCl to a 2.4N HCI solution on the relative rates of dis-  solution of a-Fe203 and Figure 11 shows the effects of adding 0.9, and 1.8M  1.2  HC10.. and 1.2 and 1.8M 4  1.2  NaOH to the same HCI solution on the  relative rate of leaching of a-Fe^O^ at 80°C. The effect of HCI concentration on the relative rates of leaching  64  Effect of HCl  leaching at  of of  8 0 ° C . (  odding ferric Table  LiCI oxide B.9,  on the  relative  (Michigan)  in  Appendix  B ) .  rate  24  M  65  i  1  L_  1.8  1  1  r  I  I  I 0  1.2  0.6  ( NaOH) ( M / l i t e r ) Figure  1  r  1  I  I  I  0.6  1.2  1.8  (HCI0 ) ( M/liter) 4  II. The  effect  rate of at  80°C  of  adding  leaching  NaOH  or HCIO^  on the  of Qt - FegO^ ( M i c h i g o n )  . ( Table  B . 9 , Appendix  B) -  in  relative  2 . 4 M HCl  66  of various aluminum oxides at constant temperature i s shown in Figure 12.  Gibbsite (AlCOH)^) dissolves i n HCI solutions with a rate which  shows approximately a f i r s t order dependence on - a + . i n dilute solutions, slowly becoming zero order with respect to a , i n concentrated +  HCI solutions.  The y-form of A^O^ also appears to leach with a rate  which exhibits a f i r s t order correlation with .'"a--, in dilute solutions, +  but a sudden switch to a zero order function of . for  can be observed  a+ > 2. Finally, a-kl^O^ does not leach to any measurable extent  in HCI even after 4 hours in 7.2N HCI at 80°C. The kinetics of the constant temperature leaching of cuprous and cupric oxides i n HCI are represented i n Figure 13. The oxides both appear,; to dissolve with a f i r s t order dependence on ' a  +  i n dilute  solutions, slowly varying towards a zero order dependence on  a +  in  concentrated solutions (Figure 14). It i s concluded from the above experimental results that a l l the oxides show a decreasing dependence of their rates of leaching on,a + with increasing acid concentration. Additionally, a l l the oxides leach more quickly i n HCI than i n HCIO^ at equal acidity and temperature, but not by the same magnitude.  Ferric oxide, for example, leaches  about 10 times more quickly i n l . 2 N HCI than i n 1.2N HCIO^, whilst cuprous oxide dissolves around 5 times more rapidly and gibbsite about 2.5 times faster.  This enhancement effect exhibited by HCI over HCIO^  increases with increasing acidity; ferric oxide, for example, dissolves about 20 times more quickly i n 2.4N HCI than in 2.4N HCIO^. As i s the case with the leaching of the oxides i n HC10. solutions,  1  1—I  1  A AI(OH)  1  1  r  1  3  A  Figure  r - Ay>3  Measured  12.  a  Relative mean  rotes  activity  of of  +  leaching  ( Molal ) of  oluminum  oxides  HCI . ( T = 8 0 ° C ) ( T a b l e  B 7 ,  in  HCI  versus  Appendix  the  B) •  Relative the  rates of  mean  leaching of  activity  CtigO and  CuO in  dilute  HCl  versus  of H C l . ( T = 12 ° C ) ( T a b l e B . 7 , Appendix B ) •  Relative activity  rate of  of l e a c h i n g  HCI  • (  of  CuO  T= 1 2 ° C ) (  in  HCI  versus the  mean  Table B . 7 , Appendix B ) •  70  the magnitude of the absolute rates of leaching of the oxides in HCl solutions varies widely from one oxide to the other, approximately in the sequence: Rate (Cu 0) = 6x Rate (CuO) = 10 x Rate A1(0H) = 10 -10 5  5  7  2  x Rate (a-Fe 0 ) 2  3  The energies of activation for the dissolution of the oxides in HCl are 21-23 kcal/mole for a-Fe 0  (42,101), 17-18 kcal/mole for  a-FeO-OH (42,100), 13 ± 0.2 kcal/mole for Y - A l ^ 22.2 kcal/mole for A1(0H> (47,53). 3  (this work),and  14.7-  It is to be noted that the acti-  vation energy for the leaching of y-Al"j0^ f a l l s out of the range of activation energies which are generally observed for the dissolution of oxides in acids (Table 3).  4.3  The Leaching of Metal Oxides in Aqueous Sulphuric Acid Solutions The rates of leaching of ferric, cuprous, cupric and manganous  oxides were investigated at constant temperature as a function of the concentration of the acid (Figures 15 and 16).  The average results  obtained earlier by Surana (100) and Roach (42) on the leaching of various natural goethite and hematite minerals are included on Figure 15.  The absolute rates of dissolution of the various oxides cannot  be compared.conveniently and only relative rates of leaching were considered; that is the ratios of the observed rates over the corresponding rates in 0.9M  H„S0. for each oxide.  71  The rate of leaching of ferric oxide shows a decreasing dependence on the concentration of ^SO^  in dilute solutions (0-LM)., becoming pro-  portional to the concentration of H^SO^ (Figure 15).  in stronger solutions (1-7.2M)  The present work also appears to correlate very well  with earlier work. Cuprous, cupric and manganous oxides a l l dissolve with a rate versus ^SO^  concentration which shows a similar dependence on the  concentration of ^SO^  as does ferric oxide (Figure 16).  A priori,  no fundamental difference between the behaviour of the oxides can be detected from the results. Again, as is observed for the leaching of the oxides in HCIO^ and HCI solutions, the absolute rates of leaching of the various oxides in H^SO^  solutions are quite different; in IN H^SO^ at 12°C, for example,  the rates are approximately in the sequence: Rate (MnO)  6  8  = 3.5'xRate (Cu20) = 20.xRate (CuO) = 2.5 x 10 -10  x Rate (a-Fe 2 0 3 ) The energies of activation for the leaching of the oxides in  (49), and 21 ~t 1 kcal/mole  are respectively, .10.5 kcal/mole for C^O for a-Fe 2 0 3  4.4  H^SO^  (42)  The Leaching of Ferric Oxide in Oxalic Acid in the Absence of Added Ferrous.Salt in Solution The rate of leaching of o&-Fe,-,03 (Michigan) was investigated as a  function of pH in 0.3M  9  oxalic acid at 85 C (Figure 18).  Three distinct  ZL  73  1  T -  Calculated Measured Cu 0  A  2  • O  1.0  CuO MnO  a> a or 0.5 0>  > a> or  0.5 Figure 16. Relative  rate  versus the  (  1.0 ) ( M / liter)  H2 S04  of leaching: of concentration  ( Table B.I I , Appendix  CUgO , CuO  of B )  HgSO^  and MnO in H g S O ^  ( T=I2°C )  lOOh  50 h  Distribution (  Table  B.2  of ,  species  Appendix  in  oxalic  B ) •  acid  at  8 0 ° C . v e r s u s pH.  1*J  o Calculated  £ \ c  A  Measured  e i£  E w  10  o  or  0 Figure  PH  18. Rate of leaching of ( T= 8 0 ° C )  ferric  ( table  oxide ( Michigan)  B.I2  , Appendix  in 0 . 3 M  B ) •  oxalic  acid  versus  pH.  76  pH regions can be observed in Figure 18; between pH 0.3 and 1 an exponential decrease in rate of leaching of a-Fe^O^ was observed; from pH 1 to 3.5 a f a i r l y constant rate of dissolution was obtained, and finally above pH 3.5 a steady decrease of the rate towards zero was measured.  No ferrous ion was detected in solution during the  leaching, which was performed i n the absence of light.  The distribution  of H.2C20, .HC^O^ and C^O^ present in oxalic acid versus pH at 80°C 4  has been calculated using the dissociation constants of 1^2020^ which were extrapolated from the data of Kurz and Farran (102) and Pinching and Bates (103) between 25°C and 55°C and is represented i n Figure 17 (Table B.2, Appendix B).  4.5  The Leaching of Ferric Oxide in Oxalic Acid in the Presence of Added Ferrous Oxalate i n Solution  4.5.1  Preliminary Experiments  During preliminary experiments on the leaching of various natural hematite and goethite specimens i n a 0.2M oxalic acid solution at 80°C, i t was observed that the rate of leaching of the oxides was small i n the f i r s t hour of the run, but then began to increase exponentially with time up to complete dissolution of the contained iron (Figure 19). This leaching behaviour suggested that a time dependent change occurred either at the oxide-electrolyte interface or in the electrolyte. It was further observed that the exponential increase in rate, of leaching did not appear when oxygen or air was bubbled through the electrolyte. In contrast, the i n i t i a l period of slow dissolution disappeared when helium was introduced i n solution prior to and during a run (Figure 19).  77  T  Figure  19.  B&aching of  1  Time  (minutes)  goethite ( M i n n e s o t a )  acid in the  presence  of  ( T = 80°C  , pH = 2 . 8  r  1  in  0.2 M  oxalic  oir , Og and He versus time • ) •  78  As any impurity appearing i n solution could be responsible for the observed catalytic effect upon the rate of dissolution of the oxides, i t was necessary to prepare pure synthetic ferric oxide (Table 5).  Indeed, synthetic ferric oxides leached slowly independ-  ently of time and of the presence of either 0^ or He i n solution. -4 However, upon adding as l i t t l e as 10 M of ferrous oxalate to the electrolyte, a high but constant rate of leaching of synthetic ferric oxide with time was obtained, provided 0^ was eliminated from the system.  It was deduced that in the preliminary experiments on natural .  impure ferric oxide samples, ferrous species (or possibly other cations) appeared i n solution. These, were thought, to be due to the presence of ferrous i n the ores and possibly to the small iron contamination of these samples which were uniquely ground in an iron mortar. start of these runs any leached ferrous, was probably  At the  oxidized:to  ferric by oxygen present i n solution, but since the.oxygen was not renewed i n solution during the run, some ferrqus/would  finally persist  in solution and cause the exponential increase i n the rate of leaching with time*  This also explains the absence of the region of slow dis-  solution when 0^ was eliminated by He prior to and during a run (Figure 19), because unoxidized.  any ferrous appearing i n solution would then remain  79  4.5.2  The Effect of Sample Weight The rates of leaching of 1 gm and 2 gm portions of ferric oxide  (sample Q, Table 5) in 0.2M  oxalic acid at pH 2.80 and at 80°C and in  the presence of 6 mg/liter of ferrous, species are given in Table ;  B.15,  -Appendix B. It i s observed that approximately twice the amount of iron is.leached from the 2. gm ore sample than from the 1 gm sample in a given time, suggesting that the leaching of ferric oxide under these conditions is controlled by a reaction at the oxide-electrolyte interface, i.e. by a heterogeneous reaction.  4.5.3  The Effect of Added Ferrous Oxalate Concentration The effect of adding ferrous oxalate to a 0.2M> oxalic acid solu-  tion at pH 2.80 and at 80°C upon the rate of leaching of a-Fe^O^ (sample Q, Table 5) and a-Fe20 (sample C, Table 5) is represented in 3  Figure 20.  The rates, of leaching- of both hematites show only an approx-  imately f i r s t order dependence on the ferrous species concentration in solution up to the solubility of ferrous oxalate in 0.2M at which a constant rate of leaching i s observed. versus log [Fe  1 Total ^  v e i 1  F  i§  u r e  oxalic acid,  Plots of log Rate  21 are straight lines.  The  slopes of these lines indicate that the rate of leaching of a-Fe20  3  free of Ti varies with the 0.66 power of the ferrous concentration in , solution whereas the rate of leaching of a-Fe^O^ containing 1.3% Ti apparently only varies with the 0.60 power of the ferrous concentration. No change in the concentration of ferrous species in solution was detected during ;  the leach.  i  1  0 Figu re  1  1  10 20.  20 ( Fe*)  Rate of  of  leaching  added ferrous  of ion.  t o t a  ferric  r  1  1  30  | ( mg. / liter)  oxide  in  0.2  ( T = 80 °C , pH= 2.8  M oxalic )( Table  acid  versus  the  B.I 4  , Appendix  concentration B )  81  0.4  -  0.2  -  A  A  0  -  A Sample C A  0.4  0.6  Figure  21.  0.8  in  added  ferrous  ix  B )  0.2  1.2  1.0  1.4  1.6  Log, ( F e ( I I ) ) ( mg./liter) 0  L o g - l o g plot oxide  Sample Q  M  oxalic  of acid  the  rate  of  versus the  ion. ( T = 8 0 ° C , pH= 2 . 8  leaching  of  ferric  concentration ) ( Tabje  BI4  of  , Append-  82  4.5.4  The Effect of Adding Ferrous Ion Complexing Agents to the Oxalate Electrolyte The addition of an excess of o-phenanthroline or ferrozine, which  are known ferrous ion complexing agents, to a 0.2M  oxalate solution at  pH 2.8,0 results in the reduction of the rate of leaching of ferric oxide (sample 0, Table 5) to the rate obtained in the absence of added ferrous salts in solution, suggesting that the active ferrous species in solution during the enhanced dissolution of a-Fe^O^.are oxalatoferrous complexes.  4.5.5  The Effect of Adding Various Cations iii Solution The observed catalyzed rate of dissolution of natural ferric  oxides may not be due solely to the appearance of ferrous species in solution but also to such cations as Cr  2+  , Mn  2~f"  , Ni  2+  4"  , Cu , Cu  | |  ,  2 + 2 + Co  and Zn  . Experiments were performed separately in the presence  -4 of 10  M of each of these cations, and no effect on the rate of leach-  ing of a-Fe20 at pH 2.80'in 0.2M -4 + 3  in the presence of 10  M of Cu  oxalic acid was detectable, except  ion.  In this case, an exponential  increase in the rate of dissolution of synthetic a-Fe20 was observed, 3  but indeed the ferrous content in solution was also observed to increase steadily. This is attributed to the possible homogeneous reduction of Fe  in solution by Cu  (99), producing Fe  species  which then in the presence of oxalate can catalyze the rate of leaching of a-Fe 0 . 2  3  83  4.5.6  The Effect of the Concentration of Oxalic Acid The effect of 0.05,  0.10,  0.15,  0.20, 0.30,  0.40  and.0.60 M/liter  oxalic acid on the rate of leaching of a-Ye^O^ (sample Q, Table 5) at pH 2.80 and at 80°C, in the presence of 6 mg/liter of added ferrous is represented in Figure 22-  The. maximum concentration of oxalic,acid  which could be used was limited to 0.7M/liter by the solubility of the acid at 80°C.  The results could suggest a Langmuir type adsorption  isotherm dependence of the rate of leaching of a-Fe20. on the concentration of oxalic acid.  4.5.7  The Effect of the T i Content of Synthetic Ferric Oxide Synthetic a-Fe20 samples were prepared under various sintering 3  conditions and with or without additions of T i (Table 5).  Undoped  a-Fe20^ samples usually contained between 0.1 and 0.2 wt % of ferrous ion and additions of T i to the oxide resulted in an increase of ferrous content of a-Fe20^ due to the replacement of Fe''' atoms by 4+ Ti  atoms.  The effect of sintering conditions and T i doping upon the  absolute rates ,of leaching of a-Fe20,j .in ferrous species containing oxalic acid solutions can only be estimated i f the total surface areas of the samples are known. However, surface area measurements are more convenient when fine powders can be used and such powders are unsuitable for this study because they produce f i l t e r i n g problems and possibly diffusion control of the leaching reactions. Also, surface area measurements do not give information on the number, size and crystallographic orientation of the grains in each particle and these factors  84  T  1  Figure 2 2 .  1  1  ( H^C^) Rote of leaching of  the  concentration  of oxalic  f e r r o u s = 6 m g / l i t e r )(  acid-  Table  1  r  (M/liter) ferric oxide  in  ( T = 8 0 ° C ,  BI9,  oxalic a c i d pH = 2 . 8 ,  Appendix B ) •  versus  added  85  are known to produce variations in the absolute rates of leaching of the oxide (44, 101).  In order to eliminate the effect of unknown  total surface area variations between the various Ti doped synthetic a-Fe^O^ samples, relative rates were considered by taking the ratios of the absolute rates of leaching of each oxide in 0.2M oxalic acid at pH 2.80 and the absolute rates of dissolution of the oxides in 2.4M HCl, both at 80°C.  These relative rates were indeed, found to depend  only on the T i content of a-Fe^O^ (Table B.13, Appendix B). The relative rates of leaching of T i doped a-Fe^O^ versus the Ti content of the oxides are plotted in Figure 23.  The points correspond-  ing to the dissolution of a-IeJO^ containing 0.5 wt % Mg,and pure a-  FeJO^ sintered under 0^ at 900°C, are added on this figure. The addition of T i to a-Ye^O^ can be seen to produce an enhancement of the relative rate of leaching of the oxides up to 0.8 wt % T i , followed by a slight decrease between 0.8 and . 3% T i . The pure a-Fe 0 specimen 2  3  sintered under 0^ exhibits only a slightly lower relative rate than the average relative rate obtained with pure specimens sintered under a i r , whilst the Mg doped a-Fe^O^ sample shows a more pronounced decrease in relative rate of leaching.  4.5.8  The Effect of Temperature Arrhenius plots were obtained between 50 and'?0°C for two a-Fe^O^  samples (samples D and ,E, Table 5), in 0.2M oxalic acid at pH 2.8; (Figure 24).  The apparent activation energies of leaching of the  oxides do not seem to be influenced by the Ti content of the oxides,  Figure 23-  Titanium Relative  rate of  leaching  content  of f e r r i c  added ferrous = 6 mg / liter ) and content  of the  oxide . ( Table  oxide  2 .41*  BI3,  (Wt%)in 0 . 2  M oxalic  HCI at 8 0 ° C  Appendix  B ) •  acid ( pH= 2 .8,  versus the titanium  87  1  1  1  1  1  A Sample  o  1  0  24 mg Fe(ll)/liter o Sample  E  6 mg F e ( l l ) / l i t e r A  Sample  D  6 mg F e ( l l ) / l i t e r  A  O  \  0.5 \-  1  2.7  2.8  Figure  1  2.9  ^  l  1  3.0  3.1  3.2  i  1,000/T(°K)  24.  in 0.2 M oxalic  1  Arrhenius acid.  plots  for  the  ( pH = 2 .8 ) ( Table  leaching B.I6  of  ferric  , Appendix  oxides  B ) •  88  but decrease with increasing additions of ferrous oxalate to the electrolyte, i.e. from 12.9 kcal/mole tolO.5 kcal/mole for additions from 20 mg/liter up to the limit of solubility of ferrous oxalate.  4.5.9  The Effect of pH  The rates of leaching of synthetic ferric oxide (sample 0, Table I |  5) at 80°C i n 0.2M oxalic acid containing 6 mg/liter of added Fe were measured as a function of pH (Figure 25).  The absolute rates  of leaching of the oxide show a maximum around pH 2.8 and become small below pH 1 and above pH 4.5. The rates of leaching plotted i n Figure 25 were normalized to 1 for the maximum rate of leaching.  The decrease  in rate of dissolution of a-Fe20^ above pH 2.8 was not due to the decrease i n solubility of ferric species i n solution, since at pH 4, for example, a constant rate of leaching was obtained up to 20 percent dissolution of a-Fe^O^ and total dissolution of the 1 gm ore was eventually attained. 4.5.10  Distribution of Ferrous Species i n 0.2M Oxalic Acid as a Function of pH  The experimental results suggest that there is a definite relation between the observed rates of leaching of ferric oxide and some ferrous species i n the electrolyte.  Hence, i t is of considerable  interest. t° estimate the distribution of ferrous species i n 0.2M oxalic acid at 80°C versus the pH of the solution. Unfortunately the stability constants corresponding to the formation of the mono-, d i and tri-oxalato ferrous complexes are only known at 25°C and large  1  1  1  1  1 Calculated  a = 0.6  ••--^l 1 Figure 2 5 .  I  l  I  2  I 3  i  i  l  4  pH  Normalised; rate of leaching} of ferric oxide in 0.2 M oxalic ocid versus pH . ( T= 80 °C , Fe(ll) = 6 mg/liter )( Table B-l 8 , Appendix B ) •  5  90  differences are found between the pK's reported in the literature (104,105,106).  These pK'.s can easily be estimated experimentally  at 80°C by using the following relations: Fe  F e C  + C 0 2  (aq)  2°4,  ,  (aq)  +  =  4  (aq) \  C 2  FeC 0 2  Ee(C 0 ) " + C 0 (aq) (aq) [Fe**]  [Fe"""]  2  4  • [C 0 ] 2  (aq)  2  (aq)  [Fe(C0)j'] J  -  Fe(C 0 ) " (aq)  (4.3)  4  2  4  = Ks  4  (aq)  + [FeC 0 ]  K ^=  (4 2)  , (aq)  C  (aq)  4  (aq)  **< 2°A~f  , =  2  2  (4.1)  4  (4.4)  + [Fe^O^ "] 2  4  (aq.)  (aq)  =  [Fe"""]  (aq)  + (4.5)  Total(aq)  where (4.1), (4.2), and (4.3) refer to the. formation of the successive oxalato-ferrous complexes, K  i s the-solubility constant for ferrous  g  oxalate and (4.5) represents the balance of the ferrous content of the solution.  From the combination of these five equations the total j |  solubility of ferrous species, [Fe  ^Total'  c a n  b e  calculated as a  function of the concentration of C~0, and becomes: 2 4  = T~2^ [C 0 ]  t^Total  2  +  K  4  l s K  +  W s  ' 2°4 [ C  _ ] (  , (aq)  (aq) + K  i 2 3 s K  K  K  ' 2°4 ' [ C  _ ]  , (aq)  ( 4  '  6 )  91  2The concentration of C„0, in 0.2M oxalic acid at 80°C can be I j 24 calculated as a function of pH (Table B.2, Appendix B) and [Fe 1Total in these solutions can be obtained by adding an excess of ferrous oxalate to the solution at the pH of interest. The experimental curve ++ = of log[Fe ] m i versus log[C„0,] in 0.2M oxalic acid i s given i n Total 24 Figure 26. The values of K , K^, K g  and  can be reasonably well  estimated by comparing expression . (4.6) to the experimental results (Table B.3, Appendix B).  The distribution of the various ferrous ,  species in 0.2M oxalic acid at 80°C versus pH was then calculated by using relations (4.1), (4.2) and (4.3) and i s represented i n Figure 27 (Table B.17, Appendix B).  2-  and FeiCJO^)2  The maximum concentrations of FeCJO^  occur at pH 2.1 and 3.8 respectively and  ^  aq)  42FeXC^O^)^ reaches a. maximum of 55% above pH 5 as [CJO^ ] is (aq) _ (aq) limited to 0.2M/liter; the 45% remainder i s then Fe(C 0 ) (aq). ( a q )  2  1  k  2  4.6 The Leaching of Ferric Oxide in Malonic Acid in the Presence.of Added Ferrous Ion. The rates of leaching of ferric oxide (Michigan) were investigated in 0.5M malonic acid at 80°C in the presence of 9 mg/liter of added ferrous ion versus the pH of the solution (Figure 28).  The rate of  leaching shows a maximum around pH 5.0 and becomes small below pH 2.7 and above pH 6.4. Unfortunately too l i t t l e data aire, available to enable the concentration of malonato-ferrous complexes to be calculated as a function of pH.  Log-log acid  plot  versus  (Table  B.3  of the  the  total  solubility  concentration  , Appendix  B ) -  of  of  ferrous  oxalate  ion  species in at  0.2  8 0 ° C.  M  oxalic  I  1  1  1  ( Table  ,  1  1  1  A  Fe(C204)|"  o FeC204  A  Fe(C204)*~  of ferrous  B.I7  1  Fe*"  »  Distribution  1  species in 0.2 M oxalic  Appendix  B ) •  acid  versus  pH  at  80 ° C .  r  10 \IQ  o  E c E  VO  5  r  E o  or  Figure  PH  28. Rate  of  ( T = 80  leaching of °C ,  Fe(ll) =  ferric  oxide  ( Michigan)  9 mg / l i t e r ) ( T a b l e  in 0 . 3 M  malonic  B . 2 0 , Appendix B ) •  acid  versus  pH  95  4.7  The Leaching of Ferric Oxide in Various Other Acids  4.7.1  In the Absence of Added Ferrous Salts in Solution Leaching experiments with a-Fe^O^ (Michigan) were carried.out at  80°C in 0.2M  solutions of maleic, tartaric, c i t r i c and sulphamic acids  and with 4 gm/liter of ethylene-diamine tetraacetic acid (E.D.T.A.). None of these acids was found to provide high enough-rates of leaching of a-YeJO^ to be of interest for detailed studies in relation to the present work.  4.7.2  In the Presence.of Added Ferrous Salts in Solution No effect of adding 10 ^M/liter of ferrous salts was observed on  the rates of leaching of a-FeJO^ (Michigan) under the conditions of leaching mentioned above in 4.7.1. Additional experiments were carried out on the leaching of aFe^O^ FeCl  in HCl solutions at 80°C in the presence of Fe(.II) added as 2  (Figure 29).  The addition of 12, 24 and 50 mg/liter of ferrous  did not influence the rate of leaching of a-Fe^O^ in 2.4N HCl solutions, but a pronounced increase in the rate in 6N HCl was observed.  This  increase in rate does not appear to show a f i r s t order dependence on the added ferrous concentration in solution.  Fe(ll) ( m g . / l i t e r ) .  Figure 29-  u Rote versus  the  ( Table B.2I  of  leaching  concentration of ,  Appendix  B  of ferric  oxide ( Michigan)  added ferrous  in HCI  ion. ( T= 8 0 ° C  )  97  5.  5.1  DISCUSSION  The Direct Leaching of Metal Oxides i n Acids  5.1.1  Model for the Mechanism of Leaching The results obtained on the leaching of aluminum oxides in hydro-  chloric acid solutions indicate that the absolute rates of leaching of y-A^O^-remain close to those of a-Al(OH> up to 1.2N HCI (Figure 3  12), but then rapidly diverge from the latter towards a constant rate of leaching at higher HCI concentrations. This may suggest that the rate of hydroxylation of y-A^O^ eventually becomes slower than that of the dissolution of the hydroxylated surface, causing the overall rate of leaching of the oxide to reach a constant value as the activity of water in the electrolytes i s approximately constant.  If, as i t  appears, the rate of hydroxylation of aluminum oxides decreases for oxides which have been heated to increasingly higher temperatures  (16),  this may explain the immeasurably low,rates of leaching obtained with a-Al 0 2  3  calcined at 1200°C.  In general, however, i t appears that the hydroxylation of oxide surfaces i s rapid and the f i r s t step of the leaching of oxides can be represented in the case of a M^O^  oxide, for example, by the following  equilibrium equation:  J |- 2° s M  + 3  7' 2° ===' H  l" s  M 0  *  0 H  (5.1)  The oxide surface species which are formed upon hydroxylation are  98  not known, but infrared studies on oxide surfaces (11) suggest that at least one hydroxyl group per surface oxide cation is present. is also, supported by Orioda and De Bruyn (107) who,  This  in studies on the  hydroxylation of the hematite (a-Fe^O^) surface identified the presence of a hydrated surface layer approximating the goethite (a-Fe0*OH) composition. The species present in an aqueous solution of an acid, HX, are in general H  +  and X  cussed before, H  ions, and to a lesser extent OH +  and HX.  As dis-  is at the origin of the positive charge which deve-  lops at oxide surfaces in acids of pH below the pH of Z.P.C. of the oxide.  This can be represented either by.the adsorption of H  +  ions  or the dissociation of chemisorbed water at the surface of the oxide. If i t may be assumed that charging of the oxide surface is a rapid reaction, the following equilibrium equation can be written: •  |-MO s where  K • OH + H ' zS= (aq)  |-MO s  • 0H  2  =  |-MO s  + HO  (5.2)  is the protonation equilibrium constant for the adsorption  of hydrogen ions at the oxide surface.  As proposed in earlier work  by Warren and coworkers . (8,9), a mechanism of the leaching of ferric oxides in perchloric acid, involving the adsorption of H  +  ions f o l - ,  lowed by the rate controlling desorption in solution of the resulting surface species, was in good agreement with the observed f i r s t order kinetics of the leaching of these oxides, in dilute solutions of the acid. Whether undissociated acid w i l l adsorb at the uncharged oxide  99  surfaces, as postulated by Wadsworth and Wadia (49) for the leaching of Cu„0 in H„SO, solutions i s considered to be doubtful for the follow2 2 4 ing reasons: (a)  The concentration of undissociated acid in solution is often very  low compared to the concentration of ionized species (i.e. ^SO^jHCljHClO^). (b)  If i t is assumed that the leaching of ferric oxides in perchloric  acid involves in a f i r s t step the adsorption of the undissociated acid, the overall rate of dissolution of these oxides would be expected to be at least proportional to the activity of undissociated HCIO^. This in turn is proportional to the product of the activities of the. species into ,  which the acid dissociates, namely a j. a-,1_,. H  +  CIO4  This product, with the 2  assumption that a^-a^Q- is also proportional to a R + or, in dilute solutions of the acid, to the square of the total concentration of perchloric acid. (c)  These findings do not correlate with the experimental results(Figure 6).  The adsorption of undissociated acid, HX, at the oxide surface cannot  be distinguished in a rate expression from the adsorption of X  ions at  +  sites already protonated by H ions as proposed in equation (5.2), simply because aTT =K, • aTT. • a _ where K. is the dissociation constant for HX in water. HX d HT x d The three factors mentioned in (a), (b) and (c) above make i t reasonable to assumed that undissociated acid does not' participate in the leaching of oxides. Anions other than OH , i.e. X , may, under favorable conditions, adsorb at positively charged oxide surface sites produced in equilibrium reaction (5.2).  If this adsorption process is a fast reaction, i t can be  conveniently written under the form of the following equilibrium equation:  100  I-MO • OH* + X s where K  a  2  =§=  ( ) aq  |-MO s  • OH  •X  (5.3)  1  is the equilibrium constant for the adsorption of X  tonated oxide sites.  at pro-  An equilibrium of the form of equation (5.3) can  be written for every anion present in solution, and thus also for multicharged species which are obtained for example in the case of polyacids. The adsorption of OH  ions does not need to be considered, as this  reaction is already taken care of by the surface hydroxylation step (5.1). A model for the leaching of metal oxides involving the steps (5.1), (5.2) and (5.3) proposed above is no different to the one proposed by Warren and coworkers (8,9)* for the leaching of ferric oxides in dilute hydrochloric acid solutions (<2N HCl), and i t is recalled that this model does not appear to provide an explanation for the leaching of ferric oxides in sulphuric acid and in concentrated perchloric and hydrochloric acids (Figure 8, curve A).  However, examination of the present results  on the direct leaching of metal oxides suggests that a l l oxides studied show a dependence of their rates of leaching on the concentration of the acids which decreases up to a maximum of one order with increasing acidity. This observation is thought to correspond to the progressive saturation of the oxide surfaces by a species whose concentration in solution is proportional to the concentration of the acid, i.e. H was suggested that X  +  or X .  Since i t  species adsorb preferentially at sites protonated  by H  +  ions, i t is proposed that the oxides surfaces may become saturated  by H  +  ions f i r s t , followed eventually by saturation by X Saturation of the oxides surfaces by H  +  and X  ions (109).  ions can be  101  accounted for by writing a mass balance equation stating that the t o t a l surface area of the oxide, i . e . u n i t y , i s equal to the sum of the surface portions created i n p r e - e q u i l i b r i a equations (5.1, (5.2) and (5.3), namely:  [|-M0 s  • OH] + [|-M0 s  • OH*] + [|-M0 s  • OH2 • X] = 1  (5.4)  Taking into account this surface balance r e s t r i c t i o n , the proposed model was tentatively compared with the results on the leaching of f e r r i c oxides i n HCl solutions (Figure 8, curve B ) . I t was concluded that although this model could account f o r the rates of leaching up to somewhat higher acid concentrations, an increasingly poorer correl a t i o n with the results at high a c i d i t i e s was s t i l l obtained.  I t may  thus be suggested that at least one more step i s involved i n the leaching of metal oxides i n acids and that the contribution of this step to the o v e r a l l rate of leaching becomes more apparent at high acidities.  Hence, the further reaction of s i t e s created at the oxide  surface i n preceding reactions, i . e . protonated and anion covered s i t e s , i s suggested to occur. (a)  The adsorption of H  i . e . [|-M0 s (b) [|-M0 s  (c)  The possible steps are: +  ions at p o s i t i v e l y protonated s i t e s ,  +  • 0H2 ].  The adsorption of X  ions at anion containing s i t e s , i . e .  • 0H2 • X ] .  + The adsorption of H  ions at anion containing s i t e s .  Steps (a) and (b) can probably be neglected as the former involves the interaction of two p o s i t i v e l y charged species and the l a t t e r  102  suggests that the oxide surface may become negatively charged. If step (c), the adsorption of H  +  ions at anion containing sites,  is assumed to be a quick reaction, the following equilibrium equation can be proposed: K.  |-MO • 0H2 • X + H  s where K  (aq)  iM=r |-MO • 0H2 • XH  (5.5)  s  i s the equilibrium constant for the protonation of anion  covered oxide surface sites.  The formation of positively charged  anion containing sites has been proposed by other workers. Ahmed(58),. suggested that an equilibrium of the form [-M(H 2 0)(OH) 2 ] s u r f a c e  +  +  2H  +  X-  —  [-M(H 0) 2 X]+ urface 2  + H20 is at the origin of the excess positive charge observed at oxide surfaces in the presence of anions which are known to adsorb at these surfaces.  It can easily be seen that this equilibrium equation is  the combination of equilibrium equations (5.1), (5.2), (5.3) and (5.5) proposed in the present model. Wadsworth and Wadia (49) postulated in their model for the leaching of Cu20 in ^SO^ that one of the dis+  solution steps i s the reaction of H .ions with the oxide surface portion which is covered with undissociated acid.  This step i s also con-  sistent with equation (5.5) i f the latter i s written under the form of a* rate equation* because,; although the steps leading to the formation of |-MO • 0H2 * X s  species are assumed to be different in the present  103  model, the same surface species are also produced by the adsorption of undissocxated a c i d , HX, at the hydroxylated oxide surface. With a few exceptions, which w i l l be discussed l a t e r on, the model  may  now,  ;account for a l l the results on the d i r e c t  leaching of metal oxides.which were investigated i n the present work, i f i t i s assumed that the rate determining steps for the leaching of these oxides are the desorptions into solution of the metal surface species produced i n p r e - e q u i l i b r i a (5 .2) , (5.3) and (5.5).  The basis  for this assumption i s . t h a t i n order to correlate the proposed model to the experimental r e s u l t s , i t was necessary to consider together the contribution of a l l three adsorption steps: (5.2), (5.3) and (5.5). However, i t i s clear from the discussion that these steps are successive and i n t e r - r e l a t e d .  As a r e s u l t i t i s not possible to consider  the case in.which these steps.are simultaneously rate c o n t r o l l i n g , since i f i t i s assumed that one of the adsorption reactions i s rate c o n t r o l l i n g i t i s automatically implied that the preceding steps are rapidly achieved e q u i l i b r i a and the succeeding steps are quick reactions.  This i s not the case f o r the desorption steps since these reac-  tions are independent.  The p o s s i b i l i t y remains that the formation of  activated complexes at .the oxide  surfaces, which indeed also are  independent r e a c t i o n s , are rate determining, but i f this was  the case  i t was expected that the energies of activation f o r the d i s s o l u t i o n of oxides would depend on the nature of the acid and, as was discussed i n the review of l i t e r a t u r e (Table 3 ) , the energies of activation show l i t t l e dependence on the type of acid used.  104  A general form of the rate expression f o r the d i r e c t leaching of metal oxides i n acids can be derived using the p r e - e q u i l i b r i a equations (5.1), (5.2), (5.3) and (5.5) and the surface balance equation (5.4) which has the form:  "-Rate =  2 + k -K -a + + k -K -K -a..-a _ + k -K •K *K -a u + -a 1 p H • 2 p a H X 3 p a a,p ET X 1 + K -a + . + K -K -a .-a v _ p H p a X  (5.6)  i n which K , K and K are the equilibrium constants defined above, p a a,p and k^,  and k^ are rate constants respectively corresponding to the  desorption reactions from protonated, anion containing, and protonated anion containing oxide surface s i t e s .  In rate expression (5.6) i t  was also assumed that the concentration of unhydroxylated oxide surface s i t e s i s zero, i . e .  i s l a r g e , and that the concentration of  protonated anion containing s i t e s , i . e . [ |-MO s  1  • OH^XH.""] , i s s u f f i c i -  ently small to be neglected i n the surface balance equation (5.4). In the following sections of the discussion an attempt i s made to correlatec the estimated values of the constants i n rate expression (5.6) f o r each of the systems studied to thermodynamic equilibrium properties exhibited by the oxide-electrolyte interface and species in solution.  5.1.2  Leaching of Metal Oxides i n  HCIO4  Solutions  The results on the k i n e t i c s of leaching of f e r r i c oxides, cuprous and cupric oxides (Figures 5 and 6) were compared with the proposed general rate expression (5.6) with X EC107.  The conversion from concentrations  105  to the corresponding activities of H  and ClO^ in solution was made by  +  using literature values of the mean activity coefficients of the acid at 25°C (108) and with the assumption that where  a  a ±  = a .= a H+ C10 ' 4  is the mean activity of perchloric acid (Table B.4,  +  Appendix B).  The calculated values of the constants in rate expres-  sion (5.6) obtained from the best f i t with the experimental results for each oxide are given in Table 6,and the corresponding rate curves are plotted together with the results in Figures 5 and 6 (Table B.5, Appendix B).  Rate expression (5.6) in the case of HCIO^ can be simpli-  fied and becomes:  , 1  with a  H +  = a  H+  P  Rate =  =  +  K  p ' H+  a  (  k  l  a  +  k  2  * a * C10T 4 K  a  )  ( 5  '  7 )  ±  4 Although the calculated rates appear l i e within 5% of the measured rates, the constants in Table 6 may only be considered as approximate because the activities a + and a . . were estimated for 25°C, 4 in  whereas the experiments were sometimes conducted at quite different temperatures.  The  values in Table 6 are the largest for the oxides  exhibiting the higher pH's of Z.P.C. This observation agrees well with the expected behavior of oxide  surfaces, because oxide surfaces  showing a basic character are anticipated to become more completely protonated by hydrogen ions with increasing acidity of the leach solution than those of oxides exhibiting an acid character. In other words, i t is calculated from rate equation (5.7) that 90% of the active  TABLE 6 Leaching of Metal Oxides in HCIO4 Calculated Constants in Rate Expression (5.6)  OXIDE  Z.P.C.  ^(PH) Gu 0  T° of leach (°C)  k  l  k  2  k  3  / mg Metal \ \ min-gm 7  K P  K a  K a,p  (molal ^)  ^  K  jmgmgeMetal \ min.gm.molal/  1\  >11.5  12  30.0  >200  N  19.0  <io"  1  CuO  9.5  12  5.3  >10  N  5.0  <io"  1  a-FeO-OH*  8.5  110  5.97xl0~  N  N  1.25  -3 <10  N  <io"  5  a-Fe 0  8.5  90  2.00xl0~  _N  N  1.25  <io"  N  <io"  6  2  2  3  * Hay (99) N = negligible  2  3  3  N  E Activation energy (Table 3)  V a  N  (kcal/mole) -  19.4 1.64  ^18.0 17 R_  17.8-22.5 19.2-22.9  107  oxide surf ace .portion becomes ..covered, by hydrogen ions at an activity of H  +  of 0.5 in .the. case.of cuprous.oxide.and  ferric oxides. duct k *K 2 k„ and K I  of 8 in the case of  Equation (5.7) shows that whilst the value of the pro-  can be calculated, the.individual values of the constants a  a  cannot be obtained.  However, limiting .values of these  constants can be estimated by modifying rate equation (5.7) by the following relation:  K  *a * V  p  K  ' C10a  < < K  p * H+ a  + 1  ( 5  '  8 )  This equation states implicitly,that the anion covered oxide surface portion i s negligibly small, and with the assumption that the experimental measured rates may vary by a maximum of 10%, the inequality (5.8) becomes: 0 • 1 • (K_ • a + 1) ^TT+  K  <  a  (5.9)  P  K  P ' V  ' C10T a  Expression (5.9) i s valid up to the highest concentrations of acid which were used, i.e. a,_j_ = a ._.= 1 in the case of Cu.O and CuO and H ClU^ z oir  a^ = +  SQ-^Q-  =  100 in the case of ferric oxides.  mate of the higher limits of K  This allows an esti-  for these oxides, namely K a  CuO and Cu_0 and K I  a  < 10  -3  < 0.1 for a  for the ferric oxides.  These rather small  values.of the equilibrium constants for perchlorate ion adsorption do agree with the thermodynamic properties of this ion. It".should be mentioned that 0^0  dissolves only half i t s copper  in HCIO^ solutions according to the overall reaction:  108  Cu 0 + 2H 2  +  ——  Cu° + Cu""" + H 0  (5.10)  2  This, is because the Cu ion dlsproportionates in.the presence of ClO^, +  and  suggests that the applicability of the model to the case of Cu 0 2  may only be fortuitous. The proposed model appears to describe very well the variation of absolute rates .of leaching of the oxides with the concentration of perchloric acid, but i t fails to account for the wide differences observed between these rates from one. oxide to the other (cf.results). It i s suggested that whole or part of these differences may be associated with one or more of the following factors: (a)  Differences in entropies and enthalpies of activation for  leaching.  However, withtthe exception of Cu 0, the energies of 2  activation of the oxides are not very different (Table 6). (b)  Differences in the types and concentrations of defects pre-  sent, in the oxides, i f i t can be assumed that the active surface for dissolution i s controlled by defects.(110). (c)  Differences in the total surface areas i n relation to grain  and particle sizes, and .shapes and porosity of the particles.  5.1.3  Leaching of Metal Oxides in HCI Solutions With the exception of y- and a - A l ^ ^ , the results on the leaching  of the metal oxides which were investigated correlate well with-the proposed model leading to rate equation (5.6), with X  = Cl . The  activities of the hydrogen and chloride ions as a function of acid concentration were calculated with  .tti6 3.ssumption tti3.t,,a._^' — .3.  ^p-t—  '  109  where .'a  i s the mean a c t i v i t y of HCl, and with the help of l i t e r a -  ture values of the mean a c t i v i t y c o e f f i c i e n t s of HCl for temperatures up to 60°C (111,112) and by extrapolation of these c o e f f i c i e n t s up to 8 5 ° e (101) '(Table B..6, Appendix B) .  The values of the constants  i n the general rate expression (5.6) were calculated for each oxide with the aim of obtaining the best f i t with the experimental results and are reproduced i n Table 7. The calculated rate curves and the experimental points are shown i n Figures 7,8 and 9 (Tables B.7, and B.8, Appendix B).  Rate expression (5.6) can be rewritten for the leaching of  oxides i n HCl solutions arid becomes:  R a t e  % ' *«+  =  1  +K ;  a  p- H+  +  K  K  a  a  p- a- H+' Cl-_  V^p'^^Cl-)  W ± t h  <  •(ki+VVaci-+V =  a  H+  =  a  Cl"  ( 5  -  1 1 }  As can be seen i n Table 7, both K^, the protonation equilibrium constant, and K , the equilibrium constant for chloride i o n adsorption, are small for f e r r i c oxide i n comparison with the corresponding constants obtained for the other oxides, whilst the complexity constant for the formation of the mono-chloro-ferric complex,. K •, i s r e l a t i v e l y large i n comparison with the complexity constants associated with the formation of the mono-chloro complexes of aluminum and cupric ions.  This suggests that there i s no direct correlation between the  a f f i n i t y of an anion to adsorb at an oxide surface and i t s tendency to complex the oxide metal cation.  Indeed, i t appears that  and  vary i n the same way from one oxide to the other and that these cons-  TABLE 7 Leaching of Metal Oxides in HCl Calculated Constants in Rate Expression (5.6)  OXIDE  k -K 3 a,p  Z.P.C  T° of leach  (pHi))  (°C)  >11.5  12  N  55.8  19.0  9.5  12  6.10  52.5  5.0  0.6  <io-  3  9.1  80  3.0  0.333  3  8.5  80  0.525 7.3 _4 4'i39xl0~ 3x10  1.2  0.142  Cu 0 2  k  lX  k  K  2z  /mg Metal j «\ min•gm  K  K  K* c (ref 50) / mg Metal \ (molalu}) \ min- gm. molal/ 0  a,p  «- (molal"'')  3.16  (Table 3) / kcal | (mole/  -  118  IO '  1.26  10  <4xl0"  0.110' '  small small  15-22  <10"  0.016"  300  19-23  -  4  73  1  CuO A1(0H) a-Fe 0 2  *  stability constant for the equilibrium:  M + r = MCI z+  (Z  1)+  C  For Cu 0, K corresponds to Cu + 2C1 z c N = negligible +  o  CuCl" /  3  3  2  ^18.0  Ill  tants thus depend on the acid-base properties of the oxide surface, i.e. the electrostatic field exerted by this surface. Apart from cuprous oxide, the magnitude of the complexity constants, K^., seems to correlate fairly well with those of the rates-of desorption of the metals from chloride containing oxide sites relative to the rates of desorption of the metals from protonated oxide sites. Indeed, the e s t i mated ratio of the rate constants, k^/k^, in Table 7, i s large for ferric oxide and.small for gibbsite and cupric oxide.  Although the  product k^-K^ ^ can.be calculated, only limiting values of the individual constants k„ and K 3  can be estimated, by assuming that the  a,p  experimental results may be affected by a maximum error 10% (Table 7). Additional evidence for the validity of the general rate expression (5.11) i s given by the fact that the calculated rates of leaching of ferric oxide for the experiments in which the activities of H . and +  Cl" ions were controlled by. using LiCl-HCl-H^O, NaOH-HCl-R^O and HCK^HCl-H^O electrolytes are very close to the measured rates of leaching (Figures 10 and 11) (Table B.9, Appendix B).  These calculations were  carried out with the following assumptions: (a) . At constant total ionic strength, the mean activity coefficients are the same for the HCl-H^O and the HCl-LiCl-H^O systems (113) (b)  At constant total ionic strength,' the mean activity coeffi-  cients for the HCIO^-HCI-H^O and NaOH-HCl-H^O systems are the average values of the mean activity coefficients of the individual HCIO^-H^O and HCl-H^O systems in the former case,and of the  112  individual NaOH-H 0 and HC1-H 0 systems i n the latter. The kinetics of the leaching of cuprous oxide i n HCI merits special attention. Indeed, the values of rate constants, k^, and k„ * K for the leaching of this oxide do not seem to show'. a,p any correlation with the known behavior of cuprous species in chloride solutions. For example, the value of rate constant  for the  desorption of the mono-chloro-cuprous complex in solution is large (Table 7), but i t i s known that CuCl is almost insoluble in water _3 (114) (1.1 x 10  M/liter at 25°C).  Contrary to i t s behavior i n  HCIO^ solutions, in which C^O produces cupr^ic ions in .solution with half the copper becoming elemental, in HCI solutions this oxide apparently dissolves only i n the form of cuprous species.  It was also experimentally  observed that, i n contrast to the behavior of ferric oxide, cuprous oxide dissolves more quickly i n dilute HCIO^ than i n dilute HCI at equal acidity (Figures 5 and 13).  Eventually, the rate of leaching  of C^O becomes greater i n HCI than i n HCIO^ with increasing acidity (a , > 1). Moreover, the addition of 0.09M of NaCl to a 0.09M HC10. + 4 H leach solution resulted i n the same rate of dissolution of C^O as in 0.09M HCI.  It i s therefore proposed that due to the relative  insolubility of CuCl, the C^O surface already becomes covered withu.CuCl i n very dilute HCI solutions, and that this surface then profS tonates as proposed i n the last step of the general model and eventually even becomes di-chlorinated by Cl ion adsorption at theseoprotonated sites, according to the following equilibrium: +  -CuClH + Cl (aq) s  |-CuCl • HCI s  (5.12)  113  Equation (5.12) thus represents an additional step i n the proposed general model i n 5.1.1. By adding a term for di-anion adsorption to rate expression (5.6), and with the assumption that the concentration of |-M0  • OH^-X-H"*"  sites is not negligible, equation (5.6)  s for the-Hei-Cu^O system^becomes:. , k, -K -a + + k„-K K -a TT+ -a ri _ + k.-K K K - a „ + 1 p ir 2 p a Cl 3 p a a , p _H+ c i 1 + K -a . + K -K -a+ -a-,_ p IF p a H Cl a  Rate  2  + k.-K -K -K -K •aL.-a ,. _ 4 p a a,p a, a BT Cl  (5.13)  2  + K -K -K -a .-a_._ p a a,p Cl By dividing the numerator and denominator of rate expression (5.13) by K -K «a .'a^,--, the following equation is obtained: p a H Cl ( V ^ a ^ f ^ f k 3 >K a < p .a H + + k 4 -K a < p .a H + -a C I _ Rate ""=' C  1  +  K  P'  a H f ) / K  K  p' a-VaCl)  1 +  +  K  a  a, P ' H*  (5.14) be assumed that the surface of C ^ O becomes rapidly  If i t can  saturated by Cl ions i n dilute HCl, then (1 +  K  'a^p «  K  K  a  a  ' p' H+' C l " a  and (k /Ka'a _ + k ) - 0. Rate expression (5.14) can thus be simplified to:  Rate =  Ka,p H+ a  1 + K -au+ a,p H^ with • a +  = a H + = acl_  (k  ' 3  +  V  ^  a  ^  (5.15)  114  The numerical values of the constants in rate expression (5.15) were calculated by comparing this equation with the experimental results (Figure 13) (Table B.7, Appendix B). individual constants k., K H  and K a  Limiting values of the  were obtained by allowing a a, a  10% error in the experimental results (Table 8).  These new values  of the constants are in agreement with the surface properties of Cu^O .in HCI*  Furthermore,; the chloride covered cuprous oxide surface i s  expected to exhibit a more acid pH of Z.P.C. than the hydroxylated surface, i.e. K < K i n Table 8. The limiting values of k, and a,p p 4 6  K  suggest that chloride ions show l i t t l e affinity for adsorption Si y SL  at protonated chlorinated cuprous oxide sites, but that the desorption of cuprous species from dichlorinated oxide sites i s rapid in comparison with the desorption of species from the other oxide sites. This observation \ seems to- agree : with the fact .that the complexity constant for the formation of CuCl^ i s large (Table 7). As discussed in 5.1.1, the results on the leaching of y-A^O^ and a-A^O^ (Figure 12) are consistent with the proposed model i f i t is assumed that.the rate of leaching of these oxides in solutions of acidity greater than IN HCI i s controlled by the slow hydroxylation reaction at the oxide surfaces. It i s to be noted that the energy of activation of 13.1 kcal/mole obtained for the dissolution of y-Al 0 2  3  in 3; N HCI between 50 and 90°C i s close to the reported  energy of activation of 15.8 kcal/mole for the hydroxylation of the surface of y-Al^O^  (19).  The wide differences.observed between the rates of leaching of the oxides from one to the other have been discussed in 5.1.2.  TABLE 8 Leaching of C^O in HCl Constants in Rate Expressions (5.14) and (5.15) k1  m-Cu_\ min-gm/  k_ 2  k_ . 3  k. 4  K p  /mS_Cu_) /mS_Cu_) ( 5S_Cu_ ] 1 min. gm / lmin-gm/ \mxn.gm/  N  N = negligible  N  38.6  >10  3  ( m o l a  19  K a\:  fl)  K a,p  K . . a, a  (molar ) (molal' ) 1  >2xl0  4  1  3.0  k. • K 4 a, a  (molal" ) ( . S , ) \ mm- gm-molal I 1  m  C u  n  <10  _1  150  116  5.1.4  The Leachingc of Metal Oxides in H„SO, 2 — 4 Solutions— The general model for the leaching of oxides given in 5.1.1  was applied to the leaching of ferric, cuprous, cupric and manganous oxides in H.SO. solutions, with X z 4  = HSO,  4  and X~ = SOT, since ELSO, 4  2  4  produces two types of anions in solution upon i t s dissociation. The mean, activity coefficients of H^SO^  have been estimated from  literature data up to 60°C (115,116) and by extrapolation up to 80°C, with the assumption that . a^-= • . _~... mean activity of R^SO^  = a^gQ  -  where a is the 4 , - .• +  T  (Table B.jJO,Appendix B) . The following  expression for the rate of leaching of oxides in H^SO^  in terms of  the general model was sufficient to describe the experimental results: -K -a .-a„^= + k -K -K -K  k -K - a ^ + k_-K R a t e  1  =  P  2  p  a  H+  S04  3  -a +  p a , a,p + 1 + K -a . +H P H  .. .  +  ' S07 * 4- p 'K ' V " H 4 4 ••• =K -K -a -a = p a S0 a  •••  k  K  a  V j  +  S 0  (5.16)  H  ori  H  4  This expression for the rate of leaching contains terms for the protonation of the oxide surface', SO^ and HSO^ sites^and protonation of sulphated sites.  adsorption at protonated  It is also suggested that  the oxide surface may become saturated successively by adsorbed H , +  and SO., ions. ..- • 4  -  > Since i t is assumed that a + = a„„__ and having H HbU4  a 4.'a„._ = K, „*a. „„- where K, „ is the second dissociation constant H  of  SO4  +  H  S 0 2  4  d,2  i n  w a t e r  T1S04  d,2  T  ( ) > 117  it:  i s  deduced that  a S Q  rate equation can be simplified to yield:  = - &  d  2 <  The above  117  k R a t e s  a)-V  + k  (2)-V  (5  .  17)  where the constants i n equation (5.16) have been grouped in single 2 constants k,. and K for the terms in a,^ and a„+. (l) (l) HtF H x  T  /-N  The numeri-  cal values of the constants i n rate expression (5.17) were calculated by comparing this equation with the experimental results (Figures 15 and 16) (Tables B . l l j Appendix B) and are reported in Table 9.  The numerical values of the individual constants k^,  ,  and K  and K , and of the combination of constants (k «K -K *K p a 3 p a a,p K.... „ + k ,-K -K') and (K 'K - *K, _ + K') in rate equation (5.16) d,'Z 4 p a a a,p d,2. a can be estimated when the same values as obtained for the leaching of the oxides in HC10, are assigned to k., and K . However, the 4 1 p limited range of acidities which could be used for the leaching of some, of the oxides makes this estimation only reasonable in the case of a-Fe^O^ (Table 9).  The large value of  suggests that SO^ ions  are strongly adsorbed at the ferric oxide surface and the much smaller limiting values of K  and K' suggest that the adsorption of H a ,p a  +  ions  at SO^ containing sites and of HSO^ ions at protonated ferric oxide sites i s not as spontaneous.  The desorption of ferric species from  protonated sulphated sites cannot be distinguished from that of the species from bisulphate containing sites and therefore k^ = k^. large value of rate constant k^ compared with k  2  The  suggests that the  desorption of ferric species from a HSO^ containing surface site i s . more rapid than from a SO^ containing site.; This might be due to the double coordinating ability exhibited by SO^ ions, which i s lacking  TABLE 9 Leaching of Metal Oxides in H2SO4 Calculated Constants i n Rate Expression ( 5 .17)  OXIDE  Cu20 CuO  a~ 2°3  C  (°C)  >11.5  12  >9  Fe  *K  (pH)-  9.5  MnO  k  Z.P.C. TP of leach  k  (D  K* (ref 50)  K  (l)  (2)  /. . mg -Metal \ / mg Metal \ (molal "*") (molal "*") z \ min-gm-molal/' [ min .gm-molal ) 4  3  2.31 x 10  2.87 x 10  3  2  12  3.91 x 10  12  1.85 x 10  4  1.80 x 10 2  2  15  io '  20.7  10  1.89  1Q  1.8 x io  14.0 (ref = 18.0  2.26  -2  6.27 x 10"  8.5  15  123  3.14 x 10  3  2  io '  123  E Activation energy(Table 3) *. / kcal | I mole '  -  4.04  19.0 - 23.0  stability constant for the equilibrium Z  M + + SO; = 4  (z  2)+  MSO, 4  Calculated Constants in Rate Expression(5.16)  OXIDE  k  l  / mg Metal \ [ min-gm / a-Feo0o ' 23  -4 3 x 10  k  2  *4.2  -«-  (ref 117) -1 (molal )  ^ 0.09  1.86 x 10  -3  K P -1  (molal )  1.2  K' a  K a  -1  -1  (molal )  (molal )  3.06 x l O  K  2  <3 x 10  -2  V  k  4  a,P / mg Metal \ -1 (molal ) I min-gm / <16  >0.5 .  119  with HSO,. 4  5.1.5  The Leaching of Ferric Oxide in IL^C,^ Solutions The general rate expression (5.6) was applied to the leaching  of ferric oxide i n oxalic acid solutions, with X = C 2 0~ and X = HI^O^, and becomes:  Rate = k - K 1  +k  +  H  p 2  ... *a„+'a  p  p +  «a_  tv  a  n  = + k -K -K -K  C2O4  3  p  a  a,p  = + k -K «K''a„.'a _, + k -K 4 p a tiT HC^O^ 5 p  -a ,^- -  (5.18)  2  a,p  H  ...  u  2 4  ... •Ki.-Kja  -K'Ka-a  2  n  +  ^  2 ° 4  C  In writing equation (5.18) i t was assumed that the concentration of the oxide surface species produced in the corresponding pre-equilibria reactions i s negligibly small, i.e. that the ferric oxide surface did not become saturated by species from solution.  As no  distinction can be made between the desorption of ferric species from protonated oxalate containing sites and UC^O^ containing sites, k^ i s 2 • + . HC2O4 H  equal to k. . Moreover, a .-a,,,, -= K,. ,~'.a+ij+ 'a n  4  •,-K,~, * a. 4.*a-&.„•• d,l  d-2- H  2 H  C -_ and a +'a H ,  -HC204  TTr  2°4  3 , where K. , and K, „ are respectively the f i r s t and  H2C204'  d,l  n  *  .d,2  -  second dissociation constants of H^C^O^ in water. ^Thus", rate expression (5.18) can  ' be simplified and becomes:  :  (1)'V  +  K  (2)' H ' C 2 0^ a  . .k,, *a 4_'a. 1 VJ (4) V H C O 0, " O2"2 4 N  +  a  +  A  k  (3)'V HC 0  2"4  (5.19)  120  where the constants in (5.18) have been grouped in single rate constants k ^  s  . As the experiments were carried out in dilute  solutions of oxalic acid (0.3 M/liter), the concentration and activities .of species in solution were assumed to be equal.  The numeri-  cal values of the constants, .k^^, were obtained by comparing rate expression (5.19) with the experimental results (Figure 19) (Table 10) and limiting values of the individual rate and equilibrium constants in (5.18):'were estimated by allowing a 10% error in the ,  experimental results (Table B..12, Appendix B). It is concluded that the model for the leaching of oxides given in 5.1.1  is applicable to the leaching of ferric oxide in  oxalate solutions. The large values of the rate constants k^,  k^  and k,_ suggest that the desorption of ferric species from oxalate containing sites is rapid and is in agreement with the large complexing affinity of the CJO^ ion for Fe'' '" ions in solution (50).  5.1.6  The Leaching of Ferric Oxide in Various Other Acids According to the previous discussions, a rapid leach rate  should be obtained i f an acid is selected producing anions in solution which are strong metal ion complexers and show affinity.for adsorption at oxide surfaces.  A typical acid which was chosen is  ethylene-diamine-tetra-acetic acid, i.e. E.D.T.A., or H^X. the pK for the formation of FeX  Although  species in this acid is 25.1  (118),  E.D.T.A. solutions did not leach ferric oxide at an appreciable dissolution rate. observation:  At least three factors may be.suggested for this  TABLE 10 Leaching of a-Fe^O^ (Michigan) in Oxalic Acid. Calculated Constants in Rate Expressions (5.18) and (5.19)  V  V  (1)  / mg Fe j [ min•gm-molal/  ao  -3  V  (2)  \c  *(3)  / mg Fe j ( min.gm-molal / 2  4.6 x 10  /  *(4)  mg Fe | / \ min-gm-molal / 2  5.8 x 10  mg Fe I -gm-molal min  -1  1.1 x 10  -1  1 mg Fe min.gm  '3 x 10  / mg Fe |  I min • gm / -4  K• P (molal "*")  1.2  >4.0 x 10  -2  K a (molal "*")  / mg Fe j \ min• gm '  / mg Fe | min .gm /  1  >5.0 x 10  -2  >5 x 10  K' a (molal )  K a,p -1 (molal )  <10  <10  1  -2  K' : a,p -1 (molal )  <1  122  (a)  The solubility of E.D.T.A. at pH 2.7 i s only around  0.01 M/liter. 4(b)  At this pH, the complexing species, X  are only present  at a concentration of approximately 5 X 10 ^ M/liter. (c)  The anionic species present i n solution at pH 2.7 are  H^X  and R^X  and these anions not only are weak ferric ion  complexers, but also may compete for the adsorption sites at the oxide surface. These factors are also suggested to be the cause of the very low rates of leaching of ferric oxide i n maleic, malonic, tartaric, c i t r i c and sulphamic acids. 5.2  The Acid Leaching of Ferric Oxides i n the Presence of Added Ferrous Salts i n Solution  5.2.1  The Leaching of Ferric Oxide in H^C^O^ Solutions  At least two different mechanisms can be suggested for the leaching of ferric oxide in oxalic acid i n the presence of added ferrous oxalate (119)., In one mechanism, adsorbed ferrous oxalate species may lose whole or part.of their oxalate groups to neighboring ferric sites of the oxide surface. The-desorption of the resulting ferric oxalate complexes and the now oxalate-depleted ferrous species w i l l result in the dissolution of the oxide whilst the ferrous content i n solution i s kept constant.  In another mech-  anism adsorbed ferrous oxalate species may lose electrons to the oxide lattice, probably at defects. The desorption of the resulting  123 ferric oxalate complexes and the ferrous ions from the substrate w i l l again lead to the leaching of the oxide and restoration of the ferrous content of the solution. As the leaching of a-Fe^O^ was not affected by the presence of 2+ 2+ 2+ 2+ 2+ Mn , Cu , Co , Ni and Zn oxalate complexes in 0.2 M oxalate solution,it i s concluded that a mechanism by group transfer in"the ferrous oxalate catalyzed dissolution of a-Fe^O^ is not likely to be operative. Hence, a mechanism involving leaching by electron transfer at the oxide-electrolyte interface is proposed.  The increasing rela-  tive rates of leaching of a-Fe^O^ with increasing T i content of the oxide can either be attributed to the corresponding  rise of the elec-  tronic bulk conductivity of the oxide or to the increasing ferrous content of the oxide, i f ferrous sites are active dissolution centres. Although ,the bulk conductivity of a-Fe^O^ increases by more than ten orders of magnitude for T i contents from 0 to 0.8 wt % (92), the corresponding  relative rates of leaching of a-Fe^O^ only increase by  a factor of six (Figure 23), indicating that there is very l i t t l e correlation between ,rate of leaching and bulk conductivity of a-Fe^O^. However-,', the possibility remains that the surface .conductivity of a-Fe20.j in the oxalate electrolyte does not vary much (120), due to the adsorption of the oxalato-ferrous complexes at the oxide surface. This may also be the reason why the addition of 0.5% Mg to a-Fe^O^ /reduces the relative rate of leaching of the oxide by only 25%, despite the fact that the bulk conductivity of the oxide becomes p-type (Figure 23).  124  If, as i t i s assumed, the overall leaching reaction involves electron-transfer, an anodic and a cathodic reaction can be included in the model of the leaching.  It i s suggested that these reactions  are: (a)  anodic Fe(C 0 ) ( 2 4 n  2 n - 2  )"  Fe(C 0,) 2 4n ?  adsorbed  (2n  - 3)  adsorbed  +  6  (5.21) (b)  cathodic -FeO • OH +'H 0 +  l-Fe(OH);  e  (5.22)  It is further suggested that the electron-transfer reaction is also the rate determining step in the overall dissolution reaction of a-Fe20 . 3  The Butler-Volmer equation can be written for the above  anodic and cathodic reactions using the high-field approximation (121) i f ri  a  and n , the anodic and cathodic overpotentials, are assumed to c  be sufficiently large.  Hence, the anodic current per unit area at  the oxide surface can be expressed as: 1 = 1  a  a,o  • exp  ( 1  -°a>  i#  (5.23)  where i i s the exchange-current density, and a , the transfer a,o a' &  coefficient, i s defined as the fraction of the overpotential contributing to the increase in the rate of the reaction. The cathodic current per unit area becomes: l  c  = -l  c,o  exp  -a  RT  (5.24)  125  At a potential E^, i.e. the mixed potential, | i | = | l I- , Using c  the equilibrium potentials E  and E cl  corresponding to the anodic and C  cathodic reactions, and by rearrangement of equations (5.23) and (5.24), the expression of E^ i s given by: 1-a 1-a +a  a c  E  a  +  1-a +a a c  (5.25)  c;  Delahay and Berzins (122) introduced an equation which correlates the exchange-current density, i , to the potential independent Q  rate-constant of the reaction at the surface, k , and to the actio vities of the oxidized and reduced species of the couples, with the assumption that the potential i n the outer plane of closest approach of the redox species i s constant: i  o  =. k o  1-a  F • a.\  a  (5.26)  • a^^  The difference",' E -E , can be expressed i n terms of the equilibrium c a constant K of the overall reaction by use of the Nernst equation as follows: [|-FeO-OH] • [ F e ( C 0 ) ^  l  a  2  -  • [E c -E a ] = In K + in  2 n _ 2 )  4  "]  (2n-3)[|-Fe(OH)-] • C F e ( C 0 ) ^  l  ads.  a  2  4  ^  (5.27) The suitable combination of equations (5.23), (5.25), (5.26) and (5.27) leads to the general expression for the current density involved in the electrochemical reaction at the oxide surface, namely:  126  a  1  c \ l - a +a J = F * (k ) ° • (k o,a o,c a  a  a  (2n-2)-  "«<W»  ' W  n  1 - a  V \  c  a  + a  c/  / 1-a ( a / \\ K|-FeO.OH]f 1  •  a  + a a  N  (5.28) Thus, i f the overall rate of leaching of a-Fe20^ is controlled by the electron transfer step, and i f only negatively charged ferrous oxalate 2species are involved i n the dissolution, i.e. Fe(C20 )2 4  4and Fe^C^O^)^ ,  the rate of dissolution of a-Fe20^ can be expressed by:  (5.29) where k^ and k2 contain the constant terms in expression (5.28) and the conversion factors and the activity of. |-FeO0H i n (5.28) is assumed to s be unity.  The activities of the adsorbed oxalato-ferrous complexes  depend i n turn on the activities of H and the corresponding oxalato+  ferrous complexes in solution, i f the same model as for the direct dissolution of ferric oxide i n the presence of adsorbing anions can be applied.  The rate expression i n i t s complete form, i.e. including the  dependence on H and ferrous species i n solution, becomes: +  / Rate = k  ( 1 )  • (a .a H +  V+ a  Y  Vl-a ^/  F e  2-)  a  + ^  2 )  ' Ca^a  4-A 2 43  / K I l-a'+a A  (5.30)  127  The rates of leaching of a - F e ^  (0% T i ) and a - F e ^  i n 0.2 M oxalic acid at a constant pH of 2.80 and 0.60 3'  (1.3% T i )  show respectively a  power dependence on the concentration of added ferrous oxalate  (Figure 21).  It thus appears that the e f f e c t of the T i content of  a-FeJ0^ i s to modify the values of the transfer c o e f f i c i e n t s a ' and a ' i n rate expression 3.  (5.30).  a ,  For s i m p l i c i t y , i t i s assumed  C  that a  = a cl  of  0.66  = a ' = a ' (Table B.18, Appendix B). C  and  3. i n  The numerical values  O  (5.30) were calculated using the experimental re-  sults on the v a r i a t i o n of leaching of a - F e 2 0 3 versus pH (Table B.18, Appendix B).  (Figure 25)  Rate expression (5.30) i s found to correlate  very w e l l with these r e s u l t s .  F i n a l l y , the experimental rates of leach-  ing of a - F e ^ ^ as a function of the concentration of oxalic acid at a constant pH of 2.80  and i n the presence of 6 mg/liter of added ferrous  ion compare w e l l with the calculated rates using rate equation (5.30) (Figure 22)  (Table B.19, Appendix B).  It should be noted that the morphology of the acid attack of the basal plane of a a - F e 2 0 3 single c r y s t a l i s s i g n i f i c a n t l y d i f f e r e n t i n ferrous species containing oxalic acid than i n HCl, E^SO^ solutions (Figure 30a to 30d)..  and HCIO^  In the l a t t e r acids, uniform  (HCIO^) or evenly d i s t r i b u t e d p i t t i n g attack (HCl, H^O^)  attack  of the basal  plane of a - F e ^ ^ i s observed, whilst i n oxalic acid i n the presence of ferrous species l o c a l i z e d etching can be seen (Figure 30d).  The  etch p i t s appear to be aligned along three crystallographic directions with t h e i r edges, which are p a r a l l e l to these d i r e c t i o n s , forming pseudo hexagons i n the basal plane.  These directions may  correspond  128  (c)  (d)  Morphology of the acid attack on the basal plane of a a-Fe^O^ single c r y s t a l . (a) 9 M H C 1 Q , 80°C, 1 0 days, x 2 , 0 0 0 ; (b) 6 M HCI, 60°C, 1 0 min, x 2 , 0 0 0 (c) 6 M H 2 S O 4 , 60°C, 2 0 min, x 2 , 0 0 0 ; (d) 0 . 2 M oxalic a c i d , 6 mg/liter F e ( I I ) , 80°C, 2 0 min, x 1 , 0 0 0  Figure 30.  A  to the intersection of rhombohedral planes and the basal plane of  a-¥e20y  I t : i s  suggested that the cathodic reaction, i.e. (5.22),  takes place at defects associated with the observed crystallographic directions and that the anodic reaction, i.e. (5.21), proceeds at evenly distributed protonated sites of the oxide surface.  The cathodic  reaction w i l l cause pitting of the oxide surface, since i t i s proposed that ferrous ions which are formed during the reduction of ferric ions in the oxide lattice desorb from cathodic sites.  Conversely, the anodic  reaction w i l l not modify the morphology of the oxide surface, since only species from solution are involved i n this reaction, i.e. the oxidation of oxalato-ferrous to oxalato-ferric species.  5.2.2  The Leaching of Ferric Oxide in Malonic Acid  Due to the lack of data on the stability constants of equilibria reactions associated with the formation of malonato-ferrous species, the results on the leaching of ferric oxide in malonic acid (Figure 28) can only be interpreted qualitatively.  The similar variation of the  rate of leaching of a-Fe20 in malonic and oxalic acids as a function 3  of pH (Figures 25 and 28) suggests that the oxide dissolves by the same mechanism in both.acids. adsorb at protonated  It i s proposed that malonato-ferrous species  a-Fe^O^  sites, followed by the rate determining  electron transfer between these adsorbed species and the oxide lattice. The desorptions of the resulting malonato-ferric species and ferrous ions from the oxide surface w i l l result i n the leaching of the oxide. The results (Figures 25 and 28) show that the pH corresponding to the  130  maximum rate of leaching of a-FeJO^ i n o x a l i c acid i s s h i f t e d towards a more basic pH i n the case of malonic a c i d , i . e . from pH 2.8 5.  to about  It i s suggested that this difference can be associated with the  d i s t r i b u t i o n of the complexing ions i n the two acids, i . e . malonic acid becomes completely dissociated i n water at a higher pH than o x a l i c acid (123).  5.2.3  The Leaching of F e r r i c Oxide i n HCl The r e s u l t s presented i n Figure 29 (Table B.21, Appendix B) i n d i c -  ate that the rate of leaching of a-Fe20.j i n HCl i n the presence of ferrous species i n s o l u t i o n i s only enhanced i n strong HCl s o l u t i o n s . This suggests that only highly complexed ferrous species, i . e . FeCl^ 2and FeCl^ , are active i n producing an increase i n the rate of leaching of the oxide.  Due  to the possible s i m i l a r i t i e s between the ferrous  catalyzed leaching of f e r r i c oxide i n both o x a l i c , malonic and hydroc h l o r i c a c i d s , i t i s proposed that ferrous chloride complexes act as redox couples at the oxide surface.  The necessity of having highly  complexed ferrous species i n solution may  be due to either the increased  adsorption a f f i n i t y of negatively charged complexes at p o s i t i v e l y charged oxide s i t e s and/or to the enhanced rate of electron-transfer between these adsorbed ferrous species and the oxide l a t t i c e (Table 4).  131  6. CONCLUSIONS  Iv" The direct leaching in acids of most of the oxides which were investigated can be described quantitatively by a general model written in terms of the rate controlling desorptions into solution of surface metal complexes formed i n rapid adsorption prequilibria.  These  surface metal complexes are essentially created at three kinds of oxide sites: (a) (b)  Positively protonated sites, |-MO • OH* s Anion containing sites, |-MO • OH X. s 1  (c)  Positively protonated anion containing sites, |-MO • OH s  • XH  l  2.  The order of complexing power of the anions of the acids for the  oxide metal ions in solution is in the. order of the calculated rate constants for the desorption of metal complexes from oxide sites containing these anions and does not correlate with the adsorption a f f i nity of the anions. 3.  The affinity for adsorption of the anions of the acids at a given  oxide surface appears to depend essentially on the negative charge and the relative water structure promoting effect of the anions, whereas the tendency for adsorption of a given anion at different oxide surfaces can be related to the pH's of Z.P.C. of the oxides. 4. The.rates of leaching of the dehydrated forms of aluminum oxides  132  appear to become controlled by the rates of hydroxylation of the oxide surfaces with increasing a c i d i t y of the e l e c t r o l y t e .  This suggests  that the hydroxylation of the oxide surface i s a prerequisite f o r enhanced speed of d i s s o l u t i o n by species i n s o l u t i o n . 5.  The leaching of f e r r i c oxides i n acids may be considerably enhanced  by the presence of small quantities of ferrous species i n s o l u t i o n . It seems that at least two conditions have to be f u l f i l l e d to observe this c a t a l y t i c effect with f e r r i c oxides: (a)  The e l e c t r o l y t e should form highly complexed ferrous species  which are susceptible to fast electron transfer with f e r r i c ions at the oxide surface. (b)  These ferrous complexes should exhibit a f f i n i t y f o r adsorp-  tion at the oxide surface. The experimental results suggest that oxalato-, chloro- and malonato-ferrous complexes may f a l l i n this category of complexes. The mechanism of the ferrous catalyzed leaching of f e r r i c oxides i s thought to involve electrochemical reactions at the oxide surface. It appears that the electron-transfer steps between the adsorbed ferrous complexes and the a-Fe 0  surface are rate c o n t r o l l i n g .  133  7. SUGGESTIONS FOR FUTURE WORK  Although the proposed general mechanism for the leaching of metal oxides in acids can account for the rates of leaching on a relative basis, i t does not provide an explanation for the observed large differences in the absolute rates of leaching. Future studies on pure polycrystalline and single crystals of a variety of metal oxides should be made to elucidate this problem. At the same time this might provide more information on the following aspects of the leaching: (a)  anisotropy, i.e. preferential attack on characteristic  crystal faces of the oxides. (b)  the effect of crystal defects.  (c)  the effect of impurities; this could be substantiated  through controlled oxide doping. The fundamental studies should be extended to applied problems of oxide leaching and should include some of the current problems, for example the separation of mixed nickel and copper oxides and aluminum and iron oxides.  The extraction of metals from pyrometallurg-  i c a l fumes such as lead-zinc-iron oxides which are produced i n the iron blast furnace, and iron-manganese oxides which are obtained i n the ferro-manganese production should also be considered.  •134--  The possible positive or negative catalytic effects of small quantities of complexed cations in solution on the rates of leaching of metal oxides also warrants further research.  135  8. APPENDIX A Chemical Analysis and X-Ray Diffraction Patterns TABLE A.l Analysis of Goethite (Minnesota) and Hematite (Michigan) Hematite Element  V Al  Weight %  Weight as Oxide MO mn  0.84  0.13  5.00  9.55  Weight %  7.20 -  Weight as Oxide MO ,m n 0.11 69.93 FeO'OH 11.81 Fe 0 -  Ca  1.00  Fe  '57.94  Na  0.50  0.85  Mn  0.10  0.14  0.30  0.44  Si  3.00  5.88  8.95  17.52  Ti  0.50  Others  0.05  0.85 -  Total  1.40 76.00 Fe203 7.10 FeO•OH  Goethite  101.90  52.21 -  -  2  -  0.01  -  Total  99.81  3  • TABLE A.2 X-Ray Diffraction Patterns of Synthetic Hematite (Table 5) (Using the k Fe Radiation) a  Reported  Sample B  dA  29  I/I  3.66  30.66  2.69  Sample 0  Sample P  Sample E  29  I/I  29  I/I  29  I/I  29  25  30.6  20  30.7  10  30.5  43  30.6  30  42.18  100  42.2  100  42.3  100  42.0  100  42.1  100  2.51  45.36  50  45.5  50  45.5  60  45.2  75  45.3  75  2.201  53.18  30  52.3  40  52.3  20  52.1  33  52.1  40  1.838  63.56  40  63.6  40  63.5  50  63.5  47  63.5  50  1.690  69.90  60  69.9  60  69.9  60  68.8  56  69.8  50  1.596  74.68  16  74.8  10  74.6  -  74.5  21  74.6  10  1.484  81.42  35  81.4  30  81.3  20  81.3  37  81.4  40  1.452  83.62  35  83.5  30  83.5  20  83.5  33  83.6  40  1  I/I  TABLE A.3 X-Ray Diffraction Pattern for Synthetic Cu 0 and CuO 2  (Using the k^Cu Radiation) :  Reported dA  29  —  CuO  Cu.O  1 ••  I/I  o  26  I/I  o  This Study  Reported  This Study dA  29  I/I  o  29  I/I i  29.70  15  2.751  32.52  12  32.70  15  • 100  36.50  100  2.530  35.44  49  35.40  50  42.30  37  42.40  .45  2.523  35.54  100  35.60  100  1.510  61.30  27  61.30  30  2.323  38.72  96  38.80  90  1.287  73.50  17  73.60  20  2.312  38.92  30  38.90  50  1.233  77.40  4  77.30  14  1.959  46.30  3  46.30  5  1.866 .  48.76  25  48.90  30  1.714  53.41  8  53.60  10  1.581  58.31  14  58.35  15  1.505  61.56  20  61.50  25  1.418  65.80  12  65.90  10  1.410  66.22  15  66.25  12  1.375  68.14  19  68.15  20  1.304  72.42  7  72.50  5  1.265  75.02  6  75.10  55  1.262  75.22  7  75.25  5  3.020  29.50  2.365  36.40  2.135  •9  138  \  TABLE A.4 Chemical Analysis bf Pyrolusite, 3-MnO  Element or . Compound  Weight %  Mn0  75.20  Si0  5.89  2  2  Fe  0.02  CuO  3.84  P  0.045  Others,  15.0  TABLE A.5 X-Ray Diffraction Patterns of A1(0H) y - A l ^ and a - A l ^ , (Using the k Cu Radiation) 3>  a  A1(0H) Reported. -dA  4.85 4.37 4.32 3.306 3.187 3.112 2.454 2.420 2.388 2.285 2.244 2.168 2.043 1.993 1.921 1.799 1.750 1.689  29  Y"A1 0 2  This Work I/I o  18.28 320 20.30 50 20.54 23 26.94 15 12 27.97 28.66 7 36.58 23 27.12 . 20 37.64 27 39.40 5 10 40.15 41.62 7 44.30 17 11 45.50 47.28 11 13 50,70" 16 52.22 54.26 13  . 29  18.3 20.4 20.6 26.95 28.1 28.8 36.8 37.2 37.9 39.4 40.2 41.8 44.2 45.5 47.4 50.7 52.2 54.6  I/I o 200 55 30 15 13 7 17 6 26 6 11 13 20 12 9 15 14 10  Reported dA  4.56 2.80 2.39 2.28 1.977 1.520 1.395  a _ A 1 3  This Work  2°3  Reported  29  I/I o  29  I/I o  dA  19.44 31.92 37.60 39.50 45.84 60.90 67.13  40 20 80 50 100 30 100  37.5 39.4 45.8 67.2  80 10 50 100  3.49 2.554 2.383 2.088 1.741 1.603 1.512 1.4055 1.3746 1.2396 1.2347  29  This Work I/I  25.50 75 35.10 100 37.72 45 43.30 100 50 52.52 90 57.44 61.25 11 • 66.46 38 50 68.16 18 76.84 77.20 5  29 o  I/I o  25.7 35.2 37.9 43.4 52.7 57.6 61.2 66.6 68.2 76.9 77.2  55 100 33 100 36 85 8 30 40 12 6  140 9. APPENDIX B Calculated and Experimental Results TABLE B.l pH of a 0.2 M Oxalic Acid Aqueous Solution at 80°C as a Function of Added HC10. and NaOH.  HC10  4  (M/liter)  NaOH (M/liter)  pH  pH  Measured  Calculated*  0.9  0  -  0.04  0.45  0  -  0.35  0.18  0  -  0.75  0  0  1.05  1.05  0  0.05  1.3  1.30  0  0.10  1.6  1.59  0  0.15  1.9  2.00  0  0.20  2.8  2.85  0  0.25  3.55  3.68  0  0. 30  4.0.  •4.20  0  0.35  4.5  4.65  141  TABLE  B.2  Calculated*Distribution of. Oxalate Species i n 0.2 M Oxalic Acid at 80°C  pH  [C 0"]  [HC 0 ]  [H C 0 ]  100%=1  100%=1  100%=1  2  2  4  2  2  4  (%)  6  3.82  •X  io"  1.14  X  10"  X  2.84  X  10'  5.55  X  IO"*  5.55  X  2.0  2.53  X  7.95  X  2.5  9.25  X  9.25  X  3.0  2.99  X  9.45  X  3.5  9.00  X  9.00  X  4.0  2.40  X  7.60  X  4.5  5.00  X  io"3 io"2 io"2 io" 1 io"1 io" 1 io"  5.00  X  1.21  X  0.5  1.14  X  1.0  9.00  1.5  5.0  7.60  X  3  2.41  +  2  [H ] •• K  where [H_C„0 ] = K  d,2  [H ] +  d,2  [CO]  10"  1  1  8.85  X  10"  1  1  7.15  X  IO"  4.41  X  2.00  X  7.35  X  2.38  X  7.15  X  1.91  X  3.98  X  6.05  X  io"1 io" 1 io"1 io"1 io"1 io"1 io" IO^  X  1  ^  Vl d,2 K  1  1  io"1 io"2 io"2 io"3 io"3 io"4 io"5 io"  ^  [ H  (b.l)  1  +  [H ] : ; — 9  = 1 + K  X  ' K  K [HC„O4]  9.60  1  — ^ 1 +  't  H] + K  +~2  d,l d,] K  [ H  b 2  (->  ]  =  (b.3) 1 +  and  2  io"5 io"5 io" •  0.0  K  d,2  [H ]  + K  d,l  • d, f "J K  H  2  , = 10 ' (102), K, „ = 10 ' (103) are respectively the f i r s t d,l a,I 1  4  4  5  and second dissociation constants of oxalic acid at 80°C  TABLE B.3 Total Solubility of Ferrous. Species in 0.2 M Oxalic Acid as a Function of pH at 80°C (Figure 18). [c o"]  pH  16g [C 0-]  2  10  Measured  2  Total (mgFe/liter) [ F e  (M/liter) 0.70 5.28x10"^ 1.05 2.19x10 1.35 6.66x10 . 1.60 1.54x10 1.90 3.80x10 2.80 3.78x10 3.55 y ~2;tf)xi0 4.00 4.8x10 4.50 . 1.0x10  -5.280 -4.660 -4.177 -3.813 -3.420 -2.423 -1.700 -1.320 -1.000  *Stability Constants -i „ yr 10 (i)  390-406 90-100 51-55 33-36 28-30 35-40 102-110 320-336 665-680  Ref (104) (25°C)  S  l o g ^  -  10 2' l log (K -K -K ) l o g  ( K  10  K s  K  3  )  2  1  5  ]  ,  1  Ref (105) (25°C) 4 , 5 2  6.21 -  Calculated (Equation (4.6))*  lo  • S.10  [Fe  -  Total  Total (mgFe/liter)  [ F e  2.590-2.610 1.950-2.000 1.706-1.741 1.518-1.556 1.447-1.477 1.544-1.600 2.008-2.041 2.505-2.526 2.823-2.833  Ref (106) (25°C)  -  ]  420 117 52.6 34.7 30.5 37.0 116.0 283.0 705.0  This Work (80°C)  9 , 5 7  5.22  ]  4.0 6 , 3  7.1 3.76xl0"  8  log10L~Fe  2.622 2.068 1.722 1.540 1.484 1.568 2.062 2.451 2.847  143  TABLE B.4 Calculated Mean Activities of HC10  HC10, 4 Molarity  HC10. 4 Molality ~  Solutions, at 25°C  a ,-a =a  Y,;(25°C) Mean Activity Coefficient (ref.108)  ±  H+  c  Mean Activity*  (M/liter)  (Molal)  0.045  0.045  0.850  0.38  0.09  0.09  0.805  0.0725  0.15  0.15  0.790  0.12  0.18  0.18  0.780  0.14  0.36  0.36  0.773  0.28  0.45  0.45  0.770  0.346  0.50  0.50  0.769  0.38  0.75  0.76  0.790  0.60  0.90  0.94  0.820  0.77  1.00  1.05  0.823  0.86  1.50  1.58  0.925  1,46  1.80  1.92  1.05  2.02  3.00  3.34  1.65  5 .50  4.50  5.30  3.60  19.6  6.00  7.56  9.50  72.0  (Molal)  *The mean activity i s calculated as a ; - m. • y where m i s the molality and y +  coefficient of HC10  +  +  the mean activity  TABLE B.5 Experimental and Calculated Rates of Leaching of Metal Oxides in HC10, Solutions (Table 6, Figures 5 and .6). 4 Rates of Leaching  Oxide  HCIO4  (origin, , temperature of leach)  Molarity (M/liter)  '  a;  Measured  <.  Mean Activity (Molal)  Absolute /mg Metal \ ' min-gm /  Relative  Calculated [Equation (5.6),Table 6] Absolute Relative /mg Metal \ \ min•gm '  Cu 0 Synthetic 12°C  0.045 0.09 0.18 0.36 0.90  0.038 0.0725 0.141 0.278 0.770  12.8 17.5 24.4 29.5 43.2  0.296 0.405 0.565 0.683 1.000  13.1 18.5 24.3 30.4 42.2  0.304 0.430 0.563 0.705 0.980  CuO Synthetic 12°C  0.045 0.09 0:45 0.90  0.038 0.0725 0.346 0.770  0.90 1.42 3.91 5.00  1.180 0.282 0.782 1.000  0.985 1.475 3.80 5.20  0.197 0.295 0.760 1.040  a-Fe 0  0.45 0.90 1.80 3.00 4.50  2  2  3  Michigan 90°C  . •  0.346 0.770 2.02 2.02 19 .0  0.65xl0" 1.00x10 ~ 1.40xl0~^ 1.70x10 ^ 1.80x10 3  0.65 1.00 1.40 1.70 1.80  0.603x10";? 0.975x10 ^ 1.435x10" 1.77x10 ^ 1.91x10  0.603 0.975 1.435 1.77 1.91 continued  TABLE B.5 continued  Absolute  Relative  Absolute  Relative  2.55xl0_ 3.96xl0_^ 5.20x10 5.74xl0_^. 5.97x10  0.850 1.32 1.73 1.91 1.99  a-FeO-OH Minnesota 110°C (after Hay)  0.75 1.50 3.00 4.50 6.00  0.60 1.46 5.50 19.00 72.00  2.60xl0_ 4.16xl0._f 5.27xl0_^ 5.61xl0_f 6.15x10  0.865 1.39 1.76 1.87 2.05  a-FeO'-OH Natural 1100°C (afterC Surana)  0.15 0.50 0.75 1.00 1.50  0,.12 0.38 0.60 0.86 1.46  4.50xlO_J 1.07xl0_ 1.92xl0_^ 2.22xl0_^ 3.30x10  0.214 0.560 0.915 1.06 1.57  2  )  2  2  )  0.261 0.643 0.850 1.04 1.32  TABLE B.6 Calculated Mean Activities of HCI Solutions  HCI  HCI  Molarity  Molality  (M/liter)  (Molal)  0.06 0.12 0.20 0.24 0.36 0.48 0.50 0.60 0.72 1.00 1.20 1.50 1.80 2.00 2.40 3.00 3.60 4.00 4.80 5.00 5.40 6.00 7.00 7.20  0.06 0.12 0.20 0.24 0.36 0.48 0.51 0.61 0.73 1.02 1.22 1.54 1.85 2.09 2.50 3.20 3.89 4.36 5.35 5.57 6.08 6.84 8.18 8.42  Y (12°C) +  Mean Activity Coefficient  a (12°C)  Y (80°C)  a (80°C)  Y (85°C)  a (85°C)  a +=a _  Mean Activity Coefficient  Mean Activity  Mean Activity Coefficient  Mean Activity  0.73  0.15  0.69  0.35  0.71  0.72  0.78  1.21  0.85  1.78  1.06  3.39  1.35  5.89  1.80  10.0  2.48 3.30  16.9 27.0  +  H  cl  Mean Activity (Molal)  ±  +  (Molal)  0.82 0.79  0.0493 0.095  0.765 0.761 0.757  0.183 0.274 0.364  0.755 0.752  0.453 0.541  0.69  0.42  0.845  1.03  0.75  0.91  0.83  1.54  0.94  2.36  1.24  4.83  12.0  1.70  9.10  28.0  2.07 2.49  12.6 17.0  3.54  29.80  0.970  2.35 4.10  2.42  ±  ±  (Molal)  TABLE B.7 • Experimental and Calculated Rates of Leaching of Metal Oxides in HCl Solutions (Table 7 , Figures 7,8,9,12,13 and 14). Rates of Leaching Oxide (origin, temperature of.leach)  HCl  ' ..v  Molarity  Mean Activity (Molal)  (M/liter)  Measured  >  Absolute /mg Metalj \ min • gm /  Relative  Cu 0 Synthetic 12°C  0.06 0.12 0.24 0.48 0.60 1.20  0.0493 0.095 0.183 0.364 0.453 1.03  5.7 11.3 21.4 53.5 63.0 134.0  0.0425 0.0843 0.16 0.40 0.47 1.00  Cu 0* Synthetic 12°C  0.06 0.12 0.24 0.48 0.60 1.20  0.0493 0.095 0.183 0.364 0.453 1.03  5.7 11.3 21.4 " 53.5 63.0 134.0  0.0425 0.0843 0.16 0.40 0.47 1.00  :'. 2.14" 6.50 12.0 19.5 35.5 55.0 83.5  0.110 0.334 0.615 1.00 1.82 2.82 4.28  2  2  CuO Synthetic 12°C  0.12 0.36 0.72 1.2 2.4 4.8 6.0  -  0.095 0.274 0.541 1.03 2.42 12.0 28.0  Calculated [Equation (5.15) Table 8]";'  Calculated [Equation (5.6),Table 7] Absolute Relative /mg Metal\ I min • gm / 4.31 10.9 24.1 49.3 61.7 134.0 5.7 11.3 22.7 47.6 60.5 144.0 2.85 7.75 13.8 21.6 34.4 59.2 82.6  0.332 0.0815 0.180 0.368 0.460 1.00 0.0425 0.0843 0.169 0.356 0.451 1.07 0.146 0.397 0.707 1.11 1.76 3.04 4.24  continued  TABLE B.7 continued  Oxide  HCI  Calculated  Measured Absolute Rate  Relative Rate  Absolute Rate  Relative Rate  CuO Synthetic 12°C  0.12. 0.36 0.72 1.2 2.4 4.8 6.0  0.095 0.274 0.541 1.03 2.42 12.0 28.0  2.14 6.50 12.0 19.5 35.5 55.0 83.5  0.110 0.334 0.615 1.00 1.82 2.82 4.28  2.85 7.75 13.8 21.6 34.4 59.2 82.6  0.146 0.397 0.707 1.11 1.76 3.04 4.24  A1(0H) Synthetic 80°C  1.2 2.4 3.6 4.8 6.0  0.91 2.36 4.83 9.10 17.0  1.46 3.34 4.75 6.60 7.80  1.00 2.28 3.25 4.51 5.34  1.65 3.36 4.89 6.31 7.86  1.13 2.30 3.35 4.33 5.38  Y-Al 0 Synthetic 80°C  0.6 1.2 2.4 3.6 4.8 6.0  0.42 0.91 2.36 4.83 9.10 17.0  0.957 1.19 1.87 2.05 2.33 2.46  0.81 1.00 1.57 1.73 1.96 2.06  3  continued  TABLE B.7 continued  +  a  Measured Absolute  o-Fe 0 Michigan 80°C  0.6 1.2 1.8 2.4 3.6 4.8 5.4 6.0 7.2  0.42 0.91 1.54 2.36 4.83 9.1 12.6 17.0 29.8  1.00xl0_i? 3.77xl0_^ 7.75xlof 1.58x10 f 4.12x10 f 9.25xl0_^ 1.69xl0_j. 2.15xl0_7 4.50x10  o-Fe 0 Synthetic 85°C (After Bath)  0.2 0.5 1.0 2.0 3.0 4.0 5.0 6.0 7.0  0.15 0.35 0.72 1.78 3.39 5.89 10.0 16.9 27.0  l.lxioj l.lxlO^ 5.5x10 3.3xl0_| 9.6x10 1.85 3.12 5.36 7.83  a-Fe 0 Single Crystal 85°C (After Bath)  3.0 4.0 5.0 6.0  3.39 5.89 10.0 16.9  o  Calculated Relative  0.265 1.00 2.05 4.20 10.9 24.5 44.8 57.0 119.0 0.013 0.130 0.645 3.88 11.3 21.8 36.7 63.0 92.0 6.5 16.2 29.6 55.3  Absolute 1.09x10 3.75xl0_^ 8.53x10 . I,64xl0_ . 4.49x10 1.03x10 1.55x10 2.20x10 4.22x10" 1  '  Relative 0.279 0.995 2.26 4.35 11.9 27^,3 4.1 .1 . 58.4 112.0 0.0476 0.215 0.65 2.78 7.08 15.3 30.8' 58.4 99.5 7.08 15.3 30.8 58.4 continued  TABLE B.7 continued  Oxide  HCl  a  ±  Measured Absolute  3.0 4.0 5.0 6.0  3.39 5 .89 10 .0 16 .9  a-FeO-OH Natural 85°C (After Surana)  1.0 1.2 1.5 2.0 3.0 4.0  0 .72 0 .91 1.21 1.78 3.39 5 .89  Ferric Oxides Natural 80°C (After Roach, average results)  1.2 2.4 3.6 4.8  0 .91 2.36 4.83 9 .10  cr-Fe 0  3  Single crystal 85°C (After Bath)  1,.57x10 ] 2,.40x10 3,.415x10 7 6,.645x10 1..63 3,.56  Calculated Relative  Absolute  Relative  6 .5 16 .2 29 .6 55 .3  7.08 15 .3 30 .8 58 .4  0 .654 1.00 1.42 2.76 6 .80 15 .25  0 .65 0 .995 1.50 2.78 7.08 15 .25  1.0 4 .3 10 .74 19 .38  0 .995 4 .35 11 .9 27 .3  151 TABLE B.7.a Calculated Relative Rates of Leaching of Ferric Oxide Using Simplified Rate Expressions (1) After Surana. and Warren (Curve A, Figure 7): 2 Rate = 'k  •a  •'&+• /w , , >. (Molal) '  H +  •a _ = ^  • (' a"\)  cl  k^ = 1.2  +  Relative Rate of Leaching . . . .  0.91 1.00 2.00 3.00 4.00 5.00  Relative Rate a , 7T7 --1; ± (Molal )  1.0 1.2 4.8 10.8 19.2 30.0  1.1 1.2 2.4 3.6 4.8 6.0  (2) After the following rate expression (Curve B, Figure 7): .2 l * H+ ' C l " l' ± Rate = 1 + K • a^+ 1 + K •a k  a  a  k  ( a  Y  4  with k  a' +  (Molal) 0.91 1.00 2.00 3.00 4.00 5.00 6.00  = 2.85  K = 1.5  Relative Rate °  f L e a c h i n  1.00 1.14 2.85 4.66 6.50 8.40 10.3  §  Relative Rate (Molal" ) 1  1.10 1.14 1.46 1.55 1.62 1.68 1.72  TABLE B.8 Ratios of the Relative Rates of Leaching of F e r r i c Oxides and a, as a Function of a HCI Molarity (M/liter)  Mean Activity (Molal) a-Fe 2 0 3 Michigan  0.2 0.5 0.6 1.0 1.2 1.5 1.8 2.0 2.4 3.0 3.6 4.0 4.8 5.0 5.4 6.0 7.0 7.2  Calculated -1 (Molal )  Relative' Rate  0.15 0.35 0.42 0.72 0.91 1.21 1.54 1.78 2.36 3.39 4.83 5.89 9.10 10.0 12.6 16.9 27.0 29.8  Measured ( i f Table B.7) -1 (Molal ) a-Fe 2 0 3 a-Fe 2 0 3 a-FeO-OH Surana Single Bath Crystal Bath  Ferric Oxides Roach  0.087 0.372 0.631 0.906 1.10 1.17  0.896 1.10  1.10  1.33 1.55  2.18  1.82  1.78 3.32  1.92  2.00  3.70  2.75  2.59  3.67  2.96  3.73 3.41  3.27  2.22  2.26  2.13  2.69 3.56 3.35 4.00  0.318 0.615 0.663 0.902 1.09 1.24 1.47 1.56 1.84 2.08 2.46 2.59 3.00 3.08 3.26 3.43 3.68 3.76  TABLE B.9 Experimental and Calculated Rates of Leaching of Ferric Oxide (Michigan) in HCl-LiCl, HCl-NaOH and HCl-HC-10^ Solutions (Table 7, Figures 10 and 11). HCl Molarity (M/liter)  2.4 2.4 2.4 2.4  HCl Molarity (M/liter)  2.4 2.4 2.4 2.4 2.4 2.4  L'iCl -«^_  Y;(80°C-) +  Mean Activity Coefficient  0 0.61.2 2.4  HC10. <4  <-  0 0 0 0.9 1.2 1.8  m . Molality (Molal)  0.94 1.06 1.28 1.70  NaOH -«-  1.8 1.2 0 0 0 0  2.5 2.5 2.5 2.5  Y (80°C) +  Mean . Activity ' Coefficient' 0.70 0.74 0.94 1.24 1.44 1.90  Relative Rate  Cl . -  H+  Measured  -«-  2.50 3.20 3.89 5.35  . SH+ Molality (Molal)  0.6 1.2 2.5 3.3 3.9 4/6  m  2.36 2.66 3.16 4.25  2.36 3.40 4.83 9.10  a  Measured  -«•<r-  +•  2.5 2.5 2.5 2.5 2.5 2.5 • ' 1  0.42 0.89 2.36 4.10 5.61 8.75  4.44 6.18 8.80 15.50  Relative Rate  H+  ci-  4.17 5.97 8.63 17.0  Calculated Equation(5.6)  1.75 1.85 2.36 3.10 3.60 4.75  1.14 1.93 4.17 7.55 9.30 18.5  Calculated Equation(5. i  1.03 1.85 4.44 7.80 11.0 18.6  TABLE B.10 Calculated Mean Activities of H„SO. Solutions The mean activity of H S0. was calculated as a = y ' m where Y 2 4 o  +  respectively the mean activity coefficient and molality of H S0.. 2 4 n  a  ±  2 °4 Molarity (M/liter) H  iS  0.036 0.09 0.18 0.27 0.36 0.54 0.72 0.90 1.00 1.08 1.8 2.0 3.0 3.8 4.0 5.0 7.4 9.0  =  HS0T 4  H+  a  H S0 2  a  4  Molality (Molal) 0.036 0.09 0.18 0.27 0.36 0.55 0.74 0.93 1.03 1.12 1.93 2.17 3.42 4.45 4.80 6.25 10.7 15.0  and m^ = m^- == m  Y (85°C)  Mean Activity Coefficient  Mean Activity Coefficient  0.735 0.515 0.403 0.366 0.326 0.274 0.259 0.245  0.452 0.307 0.248  0.242  and m  +  It i s assumed that  +  Y (12°C) ±  +  +  ±  +  0.191 0.152 0.151 0.146 0.142 0.132 0.131 0.148 0.164 0.169 0.197 0.296 0.435  a (12°C) +  Mean Activity (Molal) 0.0264 0.0463 0.0726 0.099 0.117 0.151 0.192 0.228 0.271  a (85°C) ±  Mean Activity (Molal) 0.0163 0.0276 0.0446 0.0690 0.0836 0.112 0.136 0.146 0.255 0.284 0.506 0.730 0.810 1.23 3.17 6.55  TABLE B . l l Experimental and Calculated Rates of Leaching of Metal Oxides in H S0 Solutions (Table 9, Figures 15 and 16). 2  4  Rates of Leaching Oxide (origin, temperature of leach)  2 °A Molarity H  S  (M/liter)  a  Measured  ±  Mean Activity (Molal)  Absolute /'mg Metal\ 1, min-gm /  Relative  Calculated (Equation 5.17, Table 9) Absolute Relative |'mg Metal| \i min • gm '  Cu 0 Synthetic 12°C  0.036 0.09 0.18 0.54 1.08  0.0264 0.0463 0.0726 0.151 0.271  21.6 25.0 33.2 49.0 75.0  0.288 0.333 0.442 0.653 1.000  21.6 27.3 33.2 49.0 72.3  0.288 0.364 0.442 0.653 0.965  CuO Synthetic 12°C  0.036 0.09 0.36 0.54 0.72 1.08  0.0264 0.0463 0.117 0.151 0.192 0.271  2.50 3.80 5.46 6.70 7.60 10.2  0.245 0.372 0.535 0.653 0.745 1.000  2.94 3.71 5.76 6.70 7.88 9.85  0.288 0.364 0.565 0.653 0.772 0.965  MnO Synthetic 12°C  0.036 0.09 0.18 0.27 0.36  0.0264 0.0463 0.0726 0.099 0.117  38.2 56.0 89.5 115.0 143.0  0.151 0.221 0.353 0.455 0.565  38.0 63.5 90.0 118.0 135.0  0.150 0.250 0.355 0.466 0.533  0.09 0.18 0.36  0.0276 0.0446 0.069  0.305 0.407 0.543  1.64x10";* 2.58x10"^ 3.84x10  0.259 0.407 0.605  2  a-Fe 0 2  3  Michigan 85°C  1.93x10"^ 2.58x10 ^ 3.44x10  continued  TABLE B . l l continued  Oxide  H  oSO  a  Rates of Leaching Absolute  a-Fe 0 7 Michigan  3  0.90 1.80 ^g7.40 9.00  0.0276 0.255 V 0 > 7 3  3.17 6.55  a-Fe0-0H 1.0 Natural 80°C -=2.0 (After Surana) 3.0 4.0 5.0 '  0.146 0.284 0.506 0.810 1.23  Ferric Oxides Natural 80°C (After Roach) Average Results  0.146 0.284 0.810 1.23  1.0' 2.0 4.0 5.0  1.93xl0_J 1.08xH)J 1.88xl0_^ 2.84xl0_^ 3.64x10 ,  Relative 0.305 1.70 2.96 4.48 5.74  Absolute  Relative  1.64xl0_^ 1.09xl0_ 1.93xl0_ 3.10xl0_^ 3.64x10 3  3  0.259 1.72 3.05 4.89 5.74  1.08 1.88 2.40 3.44 4.15  1.08 1.85 2.00 3.25 3.80  1.08 2.05 3.66 4.60  1.08 1.85 3.25 3.80  TABLE B.12 Calculated and Experimental Rates of Leaching of a-FeJO^ (Michigan) in 0.3 M Oxalic Acid at 90°C versus / pH (Figure 18).  pH  ° [HC 0 ]  [c2o=i  2  4  Rate of Leaching  (M/liter)  Measured  2  (M/liter)  (M/liter)  tH2C 0 ] 4  / ng'Fe ) Imin•gm /  Calculated (Equation (5.19), Table 10) i mg Fe\ \, min-gm/ 1  20.25xl0~  0.35  1.75xlO~  6  2.46xl0"  2  2.76xl0  _1  21.6xl0"  0.75  9.77xl0"  6  5.50xl0"  2  2.45xl0  -1  11.05xl0"  11.28xl0~  1.05  3.28xl0~  5  9.24xl0"  2  2.07xl0  _1  7.50xl0"  8.14xl0"  1.30  8.36xl0~  1.33xl0  _1  1.67xl0  _1  6.80xl0"  3  6.71xl0~  2.80°  5.64xl0~  2.82xl0  _1  . 1.12xl0"  4.40xl0~  3  4.35xl0"  4.00  7.20xl0~  2  2.28xl0  _1  5.73xl0~  4  3.40xl0"  4.50  1.50xl0  _1  1.50xl0  _1  1.89xl0~  4  2.16xl0~  5  3  •  3  3  3  3  2  3  3  3  3  3  3  3.32xl0~  3  2.20xl0~  3  158  TABLE B.13 Experimental Relative Rates of Leaching of a-Fe^C^ in 0.2 M Oxalic Acid at 80°C versus the T i Content (Figure 23)  Sample (cf Table 5) Ti (wt %)  Rate of Leaching i n 2.4 N HCl at 80°C  Rate of Leaching Relative i n 0.2 M Oxalic Rate ' Acid at 80°C  (mg Fe/min.gm.)  (mg Fe/min.gm.)  A  0  1.64 X io"  D  0  1.96 X io"  G  0  1.83 X i o "  H  0  1.71 X io"  P  0  1.35 X io"  B  0  3.31 X io"  I  0.1  1.37 X io"  J  0.2  1.36 X io"  K  0.4  3.-0 X io"  F  0.5  4.3  X  L  0.8  2.0  X  C  1.3  6.6  X  E  3.0  7.8  X  X  0.5  2.37 X  io"  1 1 1 1 1 1 1 1 2 2  6.55 X io" 6.76 X io" 7.12 X io" 4.94 X io"  3.46 3.90  1  1  3.77 3.66 3.26 11.0  2.38  17.5  5.62 X IO"  1  18.7  9.00 X IO"  1  21.0  1.20  io" 1 io"  4.00  1.51  io" . 2  1  1.08  4.44 X i o "  2  1  6.46 X IO"  IO" ' 2  1  1  •>-. 22.2 18.0 15.4  1.15 7.96 X IO"  1  3.00  159 TABLE B.14 The Effect of Added Ferrous Oxalate on the Leaching of a-Fe 0 2  [Fe++]  3  in 0.2 M Oxalic Acid at 80°C and pH 2.8 (Figures 20 and 21)  log^tFe "*"]  Rate of Leaching / mg Fe \ i miri'• gm ' Sample Q (Table 5) Sample C (Table 5)  4  added  / mg Fe | ( liter/  RATE  Rate  2;LO  1 o  Rate  log  1 Q  Rate  3.0  0.477 .  0.330  -0.482  0.772  -0.142  6.0  0.778  0.596  -0.225  1.19  0.076  12.0  1.079  0.950  -0.022  1.19  0.278  18.0  1.255  1.20  0.079  2.16  0.334  24.0  1.380  1.29  0.110  2.45  0.389  36.0  1.556  1.30  0.114  2.50  0.398  In Figure  Rate = k  [Fe ]  n  • . -0.022 + 0.482 _ 0.460 _ „ ,, with: Sample Q: n = _ . " 7J^98 " ° ' A  66  1  n  n Sample C:  n =  -  0  7  9  0  4 7 7  0.278 + 0.142 1 > 0 7 9  _.Q>477  0.420  = Q^98  =  _  °' ° 6  160  TABLE B.15 Effect of Sample Weight (Sample Q, Table 5). Leaching of a-Fe 0  Time (min)  in 0.2 M Oxalic Acid at 80°C and pH 2.8, ++ with 6 mg Fe / l i t e r .  \  / mg Fe \ 'liter/  /-mg -Fe 1 liter/  1 gm sample  2 gm sample  10  5.7  10.1  20  11.6  19.9  30  17.8  28.9  40  23.6  42.4  iRate mg Fe\  (min-gm/ 1 gm sample  0.590  2 gm sample  0.530  TABLE B.16 Effect of Temperature on the Rate of Leaching of a-Ie^O^ in 0.2 M Oxalic Acid at pH 2.8 (Figure 27). Temperature  1000  Rate of Leaching (mg Fe/min.gm.)  T  (0°C)  (°K)  (°K)  Sample D (Table 5) _1  Rate 1  logio Rate  /6mg Fe"^" ] \ liter '  Ra  te (50-65 mesh) 4-1/24mg Fe ] \ liter '  Sample E (Table 5) log10 Rate  Rate 4-1-  log10 Rate  / 6mg Fe ) I liter /  50  323  3.085  0.150  -0.824  0.360  -0.444  0.237  -0.626  60  333  3.000  0.238  -0.624  0.630  -0.200  0.433  -0.364  70  343  2.915  0.409  -0.389  0.891  -0.050  0.736  -0.133  80  353  2.830  0.775  -0.111  1.31  0.117  1.275  0.106  85  358  2.795  -  -  1.72  0.236  -  -  90  363  2.755  1.200  0.079  —  -  2.000  0.300  /liter)  12.2 kcal/mole I 0.5  Activation energies:  Sample D (6mg Fe  Sample D (24mg F e ^ / l i t e r )  10.5 kcal/mole t 0.5  Sample E (6mg F e ^ / l i t e r )  12.9 kcal/mole t Q.5  162  TABLE B.17 Calculated. Distribution of Ferrous Oxalate Species in 0.2 M Oxalic Acid versus.pH, at 80°C (Figure 27)(using the stabilityconstants K , K„, K PH  Fe""" (%)  Fe(C 0 ) (%> 2  in Table B.3). Fe(C 0 ) ~ (%;  Fe(C 0,) (%;  6.23xlO~®  2  4  2  4  Z  4  2  3  1.00  8.47xl0  _1  1.53xl0  _1  5.50xl0~  1.25  6.82xl0  _1  3.19xl0  _1  2.97xl0  1.50  4.70xl0  _1  5.22xl0  _1  1.16xl0~  2  8.10xl0~  1.75  2.81xl0  -1  6.86xl0  _1  3.35xl0  _2  5.15xl0"  2.00  1.53xl0  _1  7.72xl0  _1  7.80xl0"  2.48xl0"  2.25  7.80xl0  _1  7.70xl0  _1  1.51xl0  _1  9.43xlO,"  2.50x  3.80xl0  -1  7.03xl0  _1  2.60xl0  _1  3.03xl0~  2.75  1.76xl0~  2  5.90xl0  -1  3.95xl0  -1  8.35xl0~  3.00  3  4.47xl0  _1  5,36xl0  -1  2.02xl0"  3.25  7.46xl0" -  3.09xl0  _1  6.50xl0.  4.30xl0~  3.50  -  2.00xl0  _1  7.20xl0  -1  8.17xl0  _2  3.75  —  1.22xl0  _1  7.37xl0  _1  1.40xl0  _1  4.00  -  7.45xl0"  2  7.15xl0  _1  2.16xl0  _1  4.25  —  4.57xl0~  2  6.58xl0  _1  2.98xl0  _1  4.50  —  2.97xl0~  2  5.95xl0  _1  3.75xl0  -1  4.75  —  2.lOxlO  -2  5.40xl0  _1  4.36xl0  _1  1.66xl0"  2  5,03xl0  _1  4.82xl0  _1  1.35xl0~  2  4.50xl0  _1  5.25xl0  _1  1.16xl0~  4.40xl0  _1  5.48xl0  _1  5.00  -  5.50  —  6.00  Note:  2  100% = 1  4  _3  2  -1  8.73xl0"  7  5  5  4  4  3  3  2  2  TABLE B.18 Experimental and Calculated Rates of Leaching of a-Fe 0 (Sample 0, Table 5) at 80°C versus pH (Figure 25). pH  rFe(C 0 ) -J 2  2  4  2  (M/liter)  rFe(C 0 ) -] 4  2  4  (M/liter)  Measured Normalized Absolute /.mg Fe ]1  Imfnvgm 1/  Calculation [Equation (5.30)] ..Normalized Normalized a =a =a1=a'=0 .6 a =a =a'=a'=l a c a c a c a c , k, .=1.31 x k .=9.84 x 10^ kj j=5.00 x 10 kj j=1.97 x 10  <  2  • 0.50  9.80xl0 ° -1  -14 1.38x10 1.08xl0 _11  0.113.  1.40xl0  _10  1.05  7,82xl0  1.30  4.00xl0  1.60  1.85xl0~  1.90  . 5.70xl0  2.50  2.60xl0  2.80  4.18xl0~  9.94xl0  3.65  7.37xl0  _5  4.10  6.91xl0  4.70  5.50xl0  5.00  5.03xl0~  _8  _7  2  0.026 •  0.003  0.185  0.170  0.069  0.182  0.298  0.330  0.196  0.256  0.420  0.563  0.465  6  1.80xl0  -6  1.36xl0"  0.370  0.606  0.760  0.740  3.03xl0  0.610  1.000  0.942  1.000  _5  _9  8  _7  0.603  0.988  0.920  0.960  5  1.14xl0~  0.400  0.655  0.676  0.664  _5  2.48xl0  -5  0.220  0.360  0.445  0.440  _5  4.25xl0  -5  0.146  0.239  0.236  0.182  4.82xl0  _5  _  —  0.163  0.100  5  5  _7  164  TABLE B.19 Effect of Oxalic Acid Concentration on the Rate of Leaching of «-Fe 0 2  3  (Sample Q, Table 5) at 80°C and at pH 2.8 (Figure 22).  [H C0 ] • : [Fe(C 0 )2-] [Fe(C 0 ) -] 4  2  2  4  4  2  4  Rates of Leaching  Molarity  Molarity  Molarity  Measured  (M/liter)  „, (M/liter)  (M/liter)  / mg Fej Vmih.gm/  8.65xl0  _8  0.05  1.46xl0~  ;0.10  2.62xl0"  5  3.31xl0~  0.15  3.50xl0~  5  6.25xl0  0.20  4.18xl0~  5  0.30  5.18xl0  -5  0.40  5.85xl0~  0.60  6.61xl0~  * a  a  =a  5  5  5  c  = a ' = a ' = 0.6  a  c  7  0.263 • 0.421 ;..  Calculated [Equation (5.30)]* / mg Fe\ 1 min.gm/ 0.268 0.415  0.560  0.519  9.94xl0~  7  0.605  0.605  1.85xl0~  6  0.723  0.740  2.78xl0~  6  0.780  0.845  4.70xl0"  6  0.870  1.000  _7  165  •TABLE B.2G Rate of Leaching of  a-Ie^^  (  S a m  P l > Table 5) in 0.5 M Malonic e H  Acid at 80°C versus pH i n the presence of 9mg/liter of added Ferrous Ion (Figure 28).  pH  Rate of Leaching  I mg Fe \ \ min•gm /  1.6  9.90  X  IO"  2.7  2.33  X  io"  2  3.2  3.75  X  io"  2  4.3  7.50  X  io"  2  5.0  8.35  X  io"  2  5.9  5.78  X  io"  2  6.4  3.67  X  io"  2  3  166  "TABLE B.21 Effect of Adding Ferrous Ion on the Leaching of a-Fe 0 2  3  (Michigan) in HCl.Solutions at 80°C (Figure 29)  [HCl] Molarity  [Fe  (M/liter).  l  a  d  d  e  d  (mg/liter)  Rates of Leaching ( ^ )  2.4  0.0  1.30 x 10"  3  2.4  50.0  1.31 x 10~  3  6.0  0.0  1.95 x 10~  6.0  12.5  2.44 x IO  -2  50.0  3.32 x 10  _2  6.0  '  2  167  10-. REFERENCES  1.  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