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Leaching of metal oxides 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 Brit ish 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 Bri t ish 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 first and second order. In acids which do not form strong complexes with the metal ions considered, e.g. HCIÔ , 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 hydroxyl- ated 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 rate- controlling 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 is postulated to be the reason for the observed decreas- ing 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 is 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 is 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 1.1 Introduction ......... ; 1 1.2 The Chemistry of Oxide Surfaces ..; . 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 10 1.3 The Direct Leaching of Metal Oxides..: 15 1.3.1 Kinetics of Leaching ............. 15 1.3.2 Mechanisms of Leaching of Metal Oxides 21 1.4 The Leaching of Metal Oxides Involving Electron- Transfer . 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 1.5 Critical Summary • .. 37 2. SCOPE OF THE PRESENT INVESTIGATION 44 3. EXPERIMENTAL ' 45 3.1 . Minerals and Reagents 45 3.1.1 Natural Minerals 45 3.1.2 Synthetic Minerals .... 45 (a) Hematite (a-Fe.O.) . 45 V Page (b) Cuprous Oxide (Cu20) .. ... ...... • 47 (c) Cupric Oxide (DuO) ............ 47 (d) Manganous Oxide (MnO) ..................... 47 (e) Aluminum Oxides . . . . 50 3.1.3 Reagents ... 50 3.2 Apparatus Design •'. . .. 50 3.3 Experimental Procedure 52 3.4 Analytical Methods ........ 53 3.4.1 Iron ;. 53 3.4.2 Aluminum 54 3.4.3 Copper . . 54 3.4.4 Manganese ....................... v ....... 54 3.4.5 Determination of the Ferrous Content of Hematite Specimens . ... 54 4. RESULTS ' • • .. . 55 4.1 The Leaching of Metal Oxides in Aqueous Perchloric Acid Solutions ......... 55 4.2 The Leaching of Metal Oxides in Aqueous Hydrochloric Acid Solutions 59 4.3 The Leaching of Metal Oxides in Aqueous Sulphuric Acid Solutions ' . 70 4.4 The Leaching of Ferric Oxide in Oxalic Acid in the Absence of Added Ferrous'Salt in Solution 71- 4.5 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 79 4.5.4 The Effect of Adding Ferrous Ion Complexing Agents to the Oxalate Electrolyte 82 4.5.5 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 Ti Content of Synthetic Ferric Oxide y 83 4.5.8 The Effect of Temperature V ... .. 85 4.5.9 The Effect of pH 88 4.5.10 Distribution of Ferrous Species in 0.2 M Oxalic Acid as a Function of pH 88 4.6 The Leaching of Ferric Oxide in Malonic Acid in the Presence of Added Ferrous Ion ................... 91 4.7 The Leaching of Ferric Oxide in Various Other Acids ... 95 4.7.1 In the Absence of Added Ferrous Salts in Solution : ... .95 4.7.2 In the Presence of Added Ferrous Salts in Solution 95 5. DISCUSSION . . ... ' 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 HCIÔ  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^ Solutions 119 5.1.6 The Leaching of Ferric Oxide in Various Other Acids ' • 120 v i i Page 5.2 The Acid Leaching of Ferric Oxides in the Presence of Added Ferrous Salts in Solution ..... 122 5.2.1 The Leaching of Ferric Oxide in H„Co0. Solutions 122 2 2 4 5.2.2 The Leaching of Ferric Oxide in Malonic Acid 129 5.'2.3 The Leaching of Ferric Oxide in HCl . .. 130 6. CONCLUSIONS •' • .. . . ' • .... .131 7. SUGGESTIONS FOR FUTURE WORK • ' ... 133 8. APPENDIX A ! ' 135 9. APPENDIX B ........ 140 10. REFERENCES : . ... : . ..." 167 v i i i LIST OF FIGURES Figure Page 1 Rate of leaching of Ck^O in H„S0, and HC10> -at- various concentrations of H+ '(TH31*C)T. . .. . 24 2 Rate of leaching of goethite in HCl -versus the mean activity of HCl (T=85°C) ........ 24 3 Electron microprobe pictures for Ti 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 ...................... ,. . 51 5 Relatives rates of leaching of CU2O and CuO in HOIO4 versus the concentration of HCIO4 at 12°G 56 6 Relative rate of leaching of goethite (Surana and Hay,' T=110°C) and hematite (T=90°C) versus the concentra- tion of HC10. . 57 4 7 Relative rate of leaching of ferr i c oxide in dilute HCl as a function of the mean activity of HCl ........ 60 8 Relative rate of leaching of ferr i c oxide in HCl as a function of the mean activity of HCl ............... 61 9 Ratio of the relative rate of leaching of ferr i c oxide and a+ as a function of a+ 62 10 Effect of adding L i C l on the relative rate of leaching of fe r r i c oxide (Michigan) in 2.4 M HCl at 80°C 64- 11 The effect of adding NaOH or HCIO4 on the relative rate of leaching of a-Fe203 (Michigan) in 2.4 M HCl at 80°C 65 12 Relative rates of leaching of aluminum oxides in HCl versus the mean activity of HCl (T=80°C) 67 13 Relative rates of leaching of Cu 2 0 and CuO in dilute HCl versus the mean activity of HCl (T=12°C) ... .• 68 14 Relative rates of leaching of CuO in HCl versus the mean activity of HCl (T=12°C) . . . . v 69 ix Figure Page 15 Relative rate bf'leaching of ferric oxide in H2SO4 versus the concentration of H2SO4 >.'..•• . . ..... 72 16 Relative rate of leaching of CU2O, CuO.and MnO in H2SO^ versus the concentration of H2SO4 (T=12°C) .. .73 17 . Distribution of species in oxalic acid at 80°C versus pH . ... 74 18 Rate.of leaching of ferric oxide.(Michigan) in 0.3 M oxalic acid versus pH (T=80°C) . . 75 19 Leaching of goethite (Minnesota) in 0.2 M oxalic acid in the presence of air, 0„ and He versus time (T=80°C,, pH=2.8) .... .V 77 20 Rate of leaching of ferric oxide in 0.2 M oxalic acid versus the concentration of added ferrous ion (T=80°C, pH=2.8) ' 80 21 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, pH=2.8) ' 81 22 Rate of leaching of ferric oxide in oxalic acid versus the concentration of oxalic acid (T=80°C, pH=2.8, added ferrous=6 mg/liter) 84 23 Relative rate of leaching of ferric oxide in 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 24 Arrhenius plots for the leaching of ferric oxide in 0.2 M oxalic acid (pH=2.8) 87 25 Normalized rate of leaching of ferric oxide in 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 27 Distribution of ferrous species in 0.2 M oxalic acid versus pH at 80°C 93 Figure 28 29 30 x Page Rate of leaching of ferric oxide (Michigan) in 0.3 M malonic acid versus pH (T=80°C, Fe(II)=9 mg/liter) 94 Rate of leaching of ferric oxide (Michigan) in HCI versus the concentration of added ferrous ion (T=80°C) 96 Morphology of the acid attack on the basal plane of a a-Fe203 single crystal ................ 128 (a) 9 M HC104, 80°C, 10 days, x 2,000 (b) 6 M HCI, 60°C, 10 min, x 2,000 (c) 6 M H2S0,, 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 xi LIST OF TABLES Table. Page 1 Zero Points of Charge for Various Oxides ............... 9 2 Experimental Values of n in Rate=k [Acid] . 18 3 Activation Energies (kcal/mole) for the Leaching of Various Oxides in Various Acids 22 4 Rate Constants and Enthalpies.and Entropies of Acti- vation for -Various Homogeneous Ferrous-Ferric Electron Transfer Reactions • 29 5 Synthetic Hematite Specimens 46 6 Leaching of Metal Oxides in HCIÔ . 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. Constants in Rate Expressions (5.14) and (5.15) 115 9 Leaching of Metal Oxides in H2SO4. Calculated Constants in Rate Expression (5.17) • "118 10 Leaching of a-Fe20.3 (Michigan) in...Oxalic Acid. Calculated Constants in Rate Expressions (5.18) and , (5.19) 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 Cu20 and CuO .. 137, A.4 Chemical Analysis of Pyrolusite 138 A. 5 X-Ray Diffraction Patterns of A1(0H) , Y - A 1 9 0 „ and. a-Al 2 0 3 . . . . . . 7 . ... 139 B. l - PH of a 0.2 M Oxalic Acid Aqueous Solution at 80°C as a Function of Added HC10. and NaOH 140 x i i Table Page B.2 Calculated Distribution of Oxalate Species in 0.2 M. Oxalic Acid at 80°C 141 B.3 Total Solubility of Ferrous Species in 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 HCIÔ  Solutions (Table 6, Figures 5 and 6) 144 B.6 Calculated Mean Activities of HCI Solutions . . 146 B..7 Experimental and Calculated Rates, of Leaching, of Metal Oxides in HCI Solutions (Table 7, Figures 7, 8, 9, 12, 13 and 14) • 147 B.7,a. Calculated Relative Rates of Leaching of Ferric Oxide Using Simplified Rate Expressions 151 B.8 Ratios of the Relative Rates of Leaching of Ferric Oxides and a+ as a Function of a+ 152 B.9 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) in 0.3 M Oxalic Acid at 90°C versus pH (Figure 18) ......... 157 B.13 Experimental Relative Rates of Leaching of a-Fe203 in 0.2 M Oxalic Acid at 80°C versus the Ti Content (Figure 23) 158 B.14 The Effect of Added Ferrous Oxalate on the Leaching of ct-Fe 0 in 0.2 M Oxalic Acid at 80°C and pH 2.8 Figures 20 and 21).:... . ....... 159 B.15 Effect of Sample Weight (Sample Q, Table 5). Leaching of ci-Fe 0 in 0.2 M Oxalic Acid at 80°C and pH 2.8, with 6 mg FeTII)/liter 160 x i i i Table Page B.16 Effect of Temperature on the Rate of Leaching of a-Fe_0 in 0.2 M Oxalic Acid at pH 2.8 (Figure* 24) .... . .... 161 B.17 Calculated Distribution of Ferrous Oxalate Species in 0.2 M Oxalic Acid versus pH, at 80°C (Figure 27) 162 B.18 Experimental and Calculated Rates of Leaching of a-Fe203 (Sample 0, Table 5) at 80°C versus pH (Figure 25) 163 B.19 Effect of Oxalic Acid Concentration on the Rate of Leaching of a-Fe203 (Sample Q, Table 5) at 80°C and at pH 2.8 (Figure 22) ......... 164 B.20 Rate of Leaching of a-Fe203 (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 B.21 Effect of Adding Ferrous Ion on the Leaching of a-Fe203 (Michigan) in HCl Solutions at 80°C (Figure 29)......... 166 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 is gratefully acknowledged. 1 1. REVIEW OF PREVIOUS INVESTIGATIONS 1.1 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 first type are represented by the historic Bayer pro- cess 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 leach- ing 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 comprehen- sive 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 attempt- ing 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 will be reviewed. Other, mechanisms of oxides dissolu- tion under a variety of conditions which have been proposed in earlier work will also be considered. 1.2 The Chemistry of Oxides Surfaces 1.2.1 Hydration-Dehydration of Oxide Surfaces If the first step in the overall leaching mechanism of oxides is hydroxylation of the surfaces, as proposed by Mackay and Wadsworth (10) for leaching UĈ  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 dehydra- tion of oxide surfaces has been made by Hair (11). For aluminum oxide Peri and Hannan (12) have concluded from infra- red 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 pro- cess 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 silica 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 hydration- dehydration have been advanced. Wade and Hackerman (17) and Hendriksen et al (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 al (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 al (18) suggested that the aluminas used by Morimoto possibly had annealed upon heat-pretreatment resulting in a decreased surface area of the samples. Very recently, Baker et al (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 in depth of poorly ordered cations. (f) Hydroxide or oxide-hydroxide formation in depth. According to Baker et al (20) the slowness of processes (c) and (d) are at the origin of the irreversible rehydration of dehydroxylated silica 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) is rapid for a-A^O^, but he agrees how- ever 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 is 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^ is 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 al (22) postulate the formation of three types of sites at the TiC^ surface upon dehy- dration. The first 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 is hydro- xylated and by molecular adsorption on isolated Ti ions (strong) and on isolated oxygen ions (weak) (23). In contrast with the behaviour of alumina, silica 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 first layer of physically adsorbed water on a-FeJO^ is 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 difficult by the presence of a high background adsorption in the,spectral regions of interest. Anderson et al (26) observed an irreversible modification of the surface of MgO crystals following/ complete dehydration. As with the silica surface, the species formed during the readsorption process are dependent upon the prior thermal history of the oxide sample*, but contrary to silica, MgO rehydroxylates rapidly. In contrast with the behaviour of the oxides mentioned above,"iso- lated hydroxyl groups are apparently not formed at the surface of BeO upon dehydration (27). It also seems that the hydration-dehydration cycle of BeO is reversible on material heated to temperatures of at least 550°C. L.2.2. Zero-Point of Charge (a) Variables Affecting the Zero-Point of Charge Oxides1, especially the hydrous oxides, exhibit ion exchange proper- 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 dissociation of the surface hydroxide groups. Dissociation reactions can be written as follows, -MOH + OH, , -MO + H.O (1.1) |-MOH+HJ a  == J-MOH+ s  U q ;  s  1  (1.2) (| i s a symbol referring 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 sur- face of the oxide. If the charge i s established by H + , OH , and species capable of interacting with H + , OH or Ĥ O to. form species present in the solid lattice (called potential determining ions, P.D.I.) , then the Z.P.C. may be given the special name I.E.P.(s) (29) (isoelectric point of the surface, as compared to the I.E.P. of species in solution). Adsorption of species (molecules or ions) under the combined influence of ionic and non-ionic bonding i s called "specific adsorption". The following relationship 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] (1.3) where Z = cationic charge R = r, + 2r + o 8 = oxygen ion radius A,B = constants for a l l oxides C. = correction for crystal f i e l d stabilization of M-OH bond a = combined corrections for 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 crystal structure and electrolyte composition in 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 in aging precipitates, for example, shifts the Z.P.C. in the basic direction. Healy et al (30) have interpreted the wide range of Z.P.C.'s (pH 1.5 to 7.3) they observed for various polymorphs of MnÔ  in terms of variation in cry- s t a l l i n i t y . They conclude that as the atomic packing in the MnÔ lattice increases, the electrostatic f i e l d within the lat 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 relation, based on the Huckel equation for the electro- static f i e l d strength of solids: pH (Z.P.C.) = A^- + B ¥ c where A and B are positive constants for an oxide series and r the c shortest M-0 interionic distance. 9 TABLE 1 Zero Points of Charge for Various Oxides Oxide Coordination . » t \ r 1 Z ; £ B i & . Structure Example:with Z.P.C* M-0 ..... . . (pH) '.' . . r-"-(at ^25°C) ... M20 2-4. >11". 5 Octahedral Cu20. MO 6-6 ' 8.5-12.5 Cubic ' MgO(12.4) Cd0(10.4) NiO(10.3) Cu0(9.5) 4-4 Hexagonal Zn0(9-10) M203 6T4 6.5-10.4 Hexagonal &-Al203(6.5-9.5) Rhombohedral o-Fe203(8.5) a-Cr203(7.0) M02 8-4 0-7.5 Cubic U02(3.5-6.5) Th02(8.5-ll) Monoclinic Z 02(4-6.7) 6-3 Tetragonal Ti02(4.7) Sn02(5.5) B-Mn02(7.0) Hexagonal a-Si02(2.2) M03 6r2 <0.5 Rhombohedral W03(0.5) *Selected values of Z.P.C after Parks (29). 10 Partial oxidation or reduction leading to nonstoichiometry in solids such as TiO^, Fê 'Ô  or may be expected to shift the Z.P.C. toward that characteristic of the oxidation (or reduction) state pro- duced (29). The effect of temperature of the electrolyte in 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 will 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 is 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 adsorp- tion are fundamental factors in 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 al (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. al (31) pointed out that AlO'OH is likely to be weakly acidic in comparison to basic A1(0H)3, resulting, in a net nega- tive charge on the surface of the solid in water. Indeed, Robinson et al (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 temper- atures above 1000°C exhibited its 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,Q 4  solutions, both hydrated and calcined alumina's showed preferential SÔ  - 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 Ĉ Ô  - 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 is 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 solid crystal lattice appear to adsorb more strongly than those which cannot (34). The observed order of adsorption of organic electrolytes onto a-Al^O^ i s (35): RCOOH > RCONH 2  > ROH > RNH^ > RCOOCH 3  > RN(CH 3 ) 2  > RN0 2 > ROCH 3  > RH , and the organic electrolytes with larger hydrocarbon chains form indeed . less soluble compounds. Selective ionic adsorption at oxide surfaces can also be inter- preted by considering ion-ion interactions. Ions having a high electro- static f i e l d are structure-promotors for surrounding water as opposed to large ions with a relatively weak f i e l d strength which are structure- breakers and are weakly hydrated. Berube and De Bruyn (36) based their model of the TiCV-water interface on ion-ion interactions. The OH superficial groups firmly anchor the neighbouring water molecules by hydrogen bonding, this phenomenon being strengthened by the large crysta- 4+ l l i n e f i e l d of the small Ti ion. This results in the presence of "frozen" water near, the surface, the latter behaving as a structure- promotor macro-ion. Thus,strong specific adsorption i s to be expected by those ions which also favour structure-promotion. The observed order of specific adsorption of alkali-cations, 13 L i + > -Na+ > Cs + on a negatively charged TiCV surface is in accord with this prediction. This concept may also be applied to the anions of acids for their ad- sorption on a positively charged TiC^ surface, but no clearly defined order of adsorption can be.obtained as in the case of cations. •- Specific adsorption of some inorganic anions is 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 field strength exerted by the surface upon the electro- lyte is somewhat weaker than that of TiO^. Dumont and Watillon (37) developed the following series of adsorption selectivity in acidic media, I0~ > F~ > CH3C00~ > CH2C1C00~ > BrO~ > SCN~ > CHC12C00~ > B" > N0~ > C10~ > Cl~ > 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 corres- ponding acids and bases which reflect ion-ion interaction properties; these are (38), HI > HBr > HC10. > HCI > HN0„ > HoS0, and 4 3 2 4 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 < Li A reasonable parallelism between the various sequences is observed. Nevertheless, some discrepancies arise; e.g., SCN., which can undergo 3+ chemical binding with the Fe ion, is more strongly adsorbed on ferric oxide. Moreover, CĤ CICOO and CHC1 COO which are structure-breakers as a whole, are specifically adsorbed on a-Fe20.j in acid media; the COO group which can organize water around itself is therefore probably, turned toward the surface. The problem of competitive adsorption at oxide surfaces will arise in solutions which contain more than one type of ionic species, and this is almost always the case when an oxide is dissolving in an electro- lyte. Recently, Hingston et al (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 is approximately equal to the sum of the maximum adsorption for each anion in 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 - n perchloric, sulphuric and hydrochloric acids various workers (9,41,42,43,44) have shown rates of attack which increase for both oxides in the order HCIÔ , IL̂ SÔ , 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 differ- ent 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 > HoP0, > Oxalic > H„S0, > HCl > HBr ^ HN0„ '= HC10, > 3 4 2 4 3 4 Maleic > Tartaric > Formic > Citric > Acetic (1.4) Parts.of these results were recently confirmed for the dissolution of gibbsite (A1(0H) ) in HC10., HCl and H SO solutions by Packter and 16 and Dhillon (47). Gibbsite dissolves about .five times more rapidly in HCI solutions than in HCIÔ  solutions of equal strength, and tl^SO^ solu- tions 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 is 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 beryll- ium ion by the anions (50), namely Ĉ Ô  > SÔ  > Cl , is in the order of absolute leaching rates, 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 HF̂  ions in solution as in the case of the other mono- basic 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: dC v  a — = K • -JZ^.d-C) (1.6) 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 TABLE 2 Experimental Values of n in Rate = k • [Acid] n Oxide Acid Ions in Complexing Slope Reference " (<IM) Solution Ability n Fe.Ô -xH 0 2 3 2 HF H + ,HF7 + - Strong 1.06 41,48 (0ixs3) HCl H ,C1 Strong 1.92-2.2 41,42,48 HBr H+,Br~ Strong 1.94 41,48 HN03 H+,NO~ + - Weak 0.93 41,48 HC10. 4 H ,C10 + - Weak •0.93-1.0 41,42,48 H2 S°4 H ,HSO, 4 Weak 0.56 41,48 s o ; Strong . - H3 P°4 H ,H2P04 Weak .0.59 41,48 HPOT,PO^ Strong A1203'3H20 . HCl H ,;C1 + - = Weak *1 47 H2 S°4 H ,HSO,,S0, 4 4 Weak <1 47 HCIO. 4 H+,C10^ Weak <1 47 Cu20 H2 S°4 H + ,HSOT,SO7 4 4 Strong <1 49 HCIO. 4 H ,C10 + Weak. <1 49 BeO HCl H ,C1 Weak 0.53 51 H2 S°4 + -H ,HSO ,S0 + - = Strong 0.70 51 H2 C2°4 H HC204,C204 Strong 0.42 51' ZnO H 2 S 0 4 H ,HSO.,SO. 4 4 Strong <1 52 19 time and a is a dimensionless number depending on the chemical compo- sition and structure of the oxide. Expression (1.6) differs from the Nerns.t equation (54) essentially in the constant a and is identical 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 linear according to equation (1.7) log [log : ( ^ ) ] = log K + a. • log t (1.7) This equation has no meaning when t = 0 or when G = 1. Changes in the nature and concentration of the electrolyte only influenced K according to the 'empirical' equation (1.8), namely K = B • e  n * ' a  (1.8) —ct where B is a constant [t ] , n is the concentration of the acid [gm.eq/liter] and y is a constant, for a given acid ' [liter/gm.eq]. Kabai obtained the activation energies for the dissolution of the various oxides in acids from Arrhenius plots of log K versus ^  a n t ^  w a s  able to . derive equation (1.9). AH^ = 6 • a (1.9) where a is the structure factor as defined in rate equation (1.6) and 6 = 21.2 keal/mole i s the energy required for the dissolution of any oxide independent of i t s composition and structure and of the properties of the electrolyte. It -eari be shown however that expression (1.9) for , 20 the activation energy obtained.by Kabai depends on the mathematical form of his rate, expression (1.6) and that (1.9) does not give the true activation energy. Indeed, the true activation energies should take into account the variation 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 K„ (1.10) where the subscripts 1 and 2 refer to the absolute temperatures. and T^. The corresponding ratio of the rates of leaching at the two temperatures is then given by Rate 1 K l . V 1-a _ K1 Rate 2 K 2 K 2 1 a (1.11) and hence relation (1.11) yields 1 •S a K = k - J L RT (1.12) where k is a rate constant independent of temperature and E is the true activation energy for the dissolution of the oxide. Kabai, however, postulated that K = A AH' RT (1.13) From the comparison of (1.12) and (1.13) the following relation i s derived: 21 E = = 6 (1.14) a 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. 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 will be discussed in the present investigations. leaching individual oxides are basically of two types. In the first, developed by Wadsworth and Wadia (49) for the leachingvof C^O no hydro- xylation or charging of the oxide surface is assumed and the following sequence of steps is envisaged: E = 20.00 kcal/mole ± 3 kcal/mole (1.15) 1.3.2 Mechanisms of Leaching of Metal Oxides The hypotheses developed to explain the observed kinetics of |Cu 2 0 + H 2 S0 (aq) (1.16) cti'o - - 'H 2 SO 4 -- (aq) (1.17) 22 TABLE 3 Activation Energies (kcal/mole) for the Leaching of Various .Oxides in Various Acids Oxide Acid E* References. (kcal/mole) Fe(OH)3 IN HCl 20.17-22.18 (53) a-FeO-OH HC1,H2S0 ,HC104 17.8-22.5 (9,42) a-Fe203 HC1,H SO4,HC104 19.2-22.9 (41,42,43,53) A1(0H)3 HC1,H2S04,HC104 14.7-22.18 (47,53) Cu20 H2 S°4 10.5 (49) Cu(0H)2> 0.5N C^COOH 18.1 (53) Mg(0H)2 0.75N H3B03 17.28 (53) Cr(OH)3 0.7N HoS0. 2 4 21.3-23.12 (53) Mn(OH)2 0.5N HCl 22.86 (53) *E = for the results reported by Kabai (53). 23 k |Cu20 • H SO + H + — ^ Cu44" + Cu° + H20 + HSO~ (1. (aq) (aq) (aq) Equation (1.16) represents the hydrolytic adsorption of ^2^4 o n t ' i e Cu20 surface and the first leaching reaction (1.17) the thermal decom- position of occupied surface sites. The second leaching reaction (1.18) indicates the influence of H^0+ ion and its ability to react with sites on which Ĥ SÔ  is adsorbed. A rate equation (1.20) can be developed which includes a Langmuir type equation. (1.19) for the fraction of sites, 9 , covered by H„S0,: x  J 2 4 • K3 • [H2S04] ( 1 . 1 9 ) x i + K X  [H 2 SO 4 I and = 9' • [k . • k' - * ( H +) '+• k • k. ] (1.20) 1 , X O N ; 2. _ . O 1 (kQ includes the surface roughness factor). When • [H2S04] is much greater than one the value of 0̂  approaches one and equation (1.20) becomes the linear portion of the rate versus [H+] plot as shown in Figure 1. Note that Wadsworth calculated to be equal to 1.59 x 10 liter/M, and . . thus,; the active surface of —6 Cuo0 would be saturated by HoS0. at 9 = 0.9 for [H„S0.] = 5 x 10 I 2 4 x 2 4 / . (aq) 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 in HCl, ̂ 2̂ 2̂ *4 and H2S04, Pearson and Wadsworth (55) for the dissolution of U02 in 24 1.2 h 0 I 1 1 1 ' 0 0.4 0.8 1.4 1.6 Figure I. ( H4") ( M/liter) Rate of leaching of Ci^O in H£04 and HCI0 4at various concentrations of H + ( T = 3 l ° C ) ( After Wadsworth (49).) Rate of leaching of goethite in HCI versus the mean activity of HCh ( T = 8 5 °C ) ( After S u r a n a O ) ) 25 carbonate solution, Takeuchi et al (56) for the dissolution of ThC^ in hydrofluoric acid and nitric 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 + H30 =̂ = |0 - Fe + 2^0 (1.21) s (aq) s , K .: ; |0 - Fe + Cl, . = ^ r ' " | 0 - Fe - Cl (1.22) s ( a q ) s k l 0 - Fe - Cl - i — Fe 0 Cl, . (1.23) 1 (aq) s (Rate determining step) This leads to a simple rate equation of the form: R = K • [(0 - Fe - OH] • a H + • aQ1_ (1.24) (K = h ±  • K ±  • K2) In equation (1.24) [|0 - Fe - OH] is assumed to be large in comparison s with [|0 - Fe ]. For perchloric acid no specific adsorption of the s anion is 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 concentra- tions 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 chrom- ium. 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„. • a2 (1.25) — • w with k a rate constant typical for each acid, a^ the mean activity of the acids and a^ the activity of water in 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 in 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 brilliant 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 "electron- jump" 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 first is that the total energy of the react- ants' activated complex must be identical with the energy of the pro- ducts' activated complex. That is , 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 first coordination, spheres in tbe activated complex, i t is termed an inner- sphere activated complex and the mechanism an inner-sphere mechanism. Outer-sphere activated complexes are formed when the inner coordin- ation 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; A+ X + B =r= AX + B (1.26) AX + B =•= AXB (1.27) AXB =•= ~AXB+ (1.28) ~AXB+ =s= A~ + BX+ (1.29) A~ + BX+ A" + X + B + (1.30) 28 This is 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 elect- ron transfer step (1.28) or the dissociation of the successor complex (1.29). The bridging group X in an inner-sphere activated complex can perform several functions. 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 domi- nant 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^, SÔ , 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 T° k AH A S References (g.eq/1) (°C) ( 1 ) /kcal\ (e.v.) V mole sec' \ mole ' F e 2 + + F e 3 + 0.55 0 0.87 9.3 -25 (70) 2+ 2+ Fe + FeOH 0.55 0 io 3 6.9 -18 (70) Fe 2 + + FeF 2 + • 0.50 0 9.7 8.6 -21 (72.) 2+ + Fe + FeF 0.50 0 2.5 9.0 -22 (72) 2+ 2 Fe + FeF 0.50 0 0.5 — _ (72) 2+ 2+ Fe . + FeCl 0.55 20 29.0 8.3 -24 (7Q) 2+ + Fe + FeCl 0.55 20 53.0 9.5 -20 (70) 2+ + Fe + FeCo0. 0.55 0 7xl0 2 9.2 -14 (73) 2+ -Fe^ + Fe(C 20 4) 2 Fe 2 + + FeS0+ 0.55 0 3.6xl03 - - (73) 0.25 2.5 692 - (74) Fe 2 + + Fe(S04)~ Fe 2 + + Fe(EDTA)" 0.25 25 25 1.94xl04 <4xl0~4 - - o (74) (75) Fe 2 + + Fe(ph) 3 + - 25 3.7xl04 0.2 -37 (76) Fe(ph)2++ Fe(ph) 3 + Fe(CN)^- + Fe(ph) 3 + — 0 25 >105 >108 _ - (71) (77) Fe(CN)4- + Fe(CN)0J 6 " J - 0.1 355 4.1 -32 (78) 30 should change as the complexing anion is 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 + FeĈ Ô  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-trans- fer proposed aby Home (73) is based on the rapid formation of an acti- vated complex by reaction of the ferric oxalate ion and the ferrous ion: ( H2 0 )4 or 5 Fe*C2°t + Fe<H2°C ̂  i \Z..-- (1.31) [ ( H 2 0 ) 4 o r 5 F e * C 2 ° 4 ^ ¥ V W 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: [ ( H2 0 )3 or 4 F e * — C2°4 F e< H 2 0 )5 ] 31 and.the final step is the dissolution of the activated complex and any rearrangements of the waters of solvation: [ ( H2 0 )4 or 5 F e* C2°4 F e< H 2 0 )6 ] = + (1.33) Fe*(H20)6 + (H 20) 4 ^ 5 F e C ^ 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 Ĥ PÔ , 0,̂ 0̂ and SÔ  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 homo- geneous solution processes described above. Laxen (83) has compared the rate of dissolution of U02 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 U02 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, NÔ  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 2+ was ascribed by Laxen to the increase in concentration of Fe(OH) , the most effective electron-transfer species. It should be noted 2+ that the maximum of Fe(OH) concentration does not occur at pH 2 according to Needes.and Finkelstein (84) and this particular aspect suggests that other factors may be involved in the leach. In sulphuric 3+ acid solutions, the rate of leaching of UO^ by Fe with pH also reached a maximum at pH - 2 and could be attributed to the combined effect of I | A. — Fe(OH) , FeSO^ and YeiSO^)^ species in solution. At constant pH, however, the rate of dissolution of -UÔ  showed a square root dependence on the concentration of ferr i c ion in solution. The very high rates of dissolution of UO^ reported by Hunt and 3+ 3+ Taube (85) with Fe (dipy)^ and Fe(0-phen)^ in IM HCI serve to confirm the correlation between rate of dissolution of UC^ and the very fast 3+ homogeneous electron transfer of these two complexes with Fe i n solution. Recent work published by Needes and Nicol (86) on the oxidative dissolution of UO^ in dilute perchloric acid showed that the order of leaching rates of UO^ with various oxidants was T1(:III) >, VO^ > Fe(III) VO Hg(ll ) , whereas the equivalent order of electron-exchange rates was ~2 >' H g C t ' l ) > Fe(III) > Tl(.III). The conclusion to be drawn from this information is that the rate of dissolution of UX^ is a function of both the potential and the electron-exchange rate of the redox couple used. 2+ The reductive dissolution of Mn0 2  in Fe containing acid solutions has been investigated by Koch (87). The rate of leaching of Mn0 2  in 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-controll- 2+ ing.step because.the rate of dissolution of MhÔ  by Fe was independ- 2+ + ent of the concentration of Fe , H and Ĥ SÔ . It should be empha- 2+ 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 considered. 1.4.2 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 fi 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) inter- 2+ mediate and desorption of UÔ  in solution; the following sequence of steps is envisaged: K OH [UO + HO i lou'" (1.34) s s ^ OH OH K2 0~ |0U = |0U'" _ + 2H (1.35) s *"0H s ^0 OH k k |0U" + 0 , | A c t i ^ a t e d u c 2 + +HO: + OH- s -OH 2(aq) s C O m P l e X 2(aq) 2 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 + ] 2 0 = K 2 + [ H + ] 2 . ( 1 ' 3 7 ) and = 9 • • pQ2 (1.38) The important feature of the first type of mechanism is that no attempt is 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.in̂ an̂ .attempt: to explain the kinetics of the reductive dissolution of pyrolusite by hydrazine in ammonium carbamate solutions, and the reductive dissolution of manganese dioxide in the presence of S02 was 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 U02 assumes that 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 + 2H20 .+ 4e — 4 0 H cathodic reaction ' (1.39) 2+ U02 — — U02 + 2e anodic reaction (1.40) In general, the rate of the cathodic reaction can be given by: V = k -A • [D] n (1.41) c c c where k is a rate constant, A the cathodic surface, [D] the concen-c c tration of the depolarizer, i.e. 0 2 > and n is the order of the reaction with respect to the depolarizer. The rate of the anodic reaction can similarly be given by the equation: V = k • A • [C] m (1.42) a a a where k is 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, sub-a c a c stituting the value of A^ in the rate equation giving V , k • k • A-[D]n[C]m V = V = — = (1.43) C a k-[D]n + k.[C] m c a At high concentration of C, or i f k is large, the velocity equation 3. (1.43) simplifies to : V = V = k • A • [D] n (1.44) c a c 36 and at high concentration of [D] , or i f k c  is large, the rate equation (1.43) becomes: V = V = k - A- [C] m (1.45) a c 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 (-ô |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 differ- ence 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 will 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 will also depend on the equilibrium poten- tials of the oxide and the redox couple, since n = E° - E (1.47) a1 a M i n = E ° - E (1.48) c1 c M where n and n 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 Critical 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 affin- ity 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 obser- vations : (a) An apparent dependency of the rates of leaching of a-Fe20 3 and a-FeO'OH in perchloric acid over the range of acid concentra- tion studied (0-1.5N) on either [H+] or [HCIO4] added, whilst the rate of leaching of C^O in 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 is some constant) (Figure 1 ) . (b) A dependency of the rates of leaching of a-¥e20^ 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, depend-Cl- H+ Gl- ency 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 in 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 Cû O and for BeO in sulphuric acid under the same conditions (C^O leach- ing, 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 , . occurring at a pH value of approximately 2, in both HC10. and H„SO, solutions. (b) The large differences in rates of leaching of MnÔ  with I I Fe-,' in HCIO^ and Ĥ SÔ  solutions of equal normality and indeed the same for the leaching of UÔ  in the presence of Fe (c) The square root dependency of the rate of leaching of UÔ with Fe on the concentration of Fe ,whilst the rate of leach- ing of this oxide with 0^ shows a first order dependency on pĈ .'. Whether or not hydroxylation of the oxide surface has to be con- sidered 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 hydroxy- lated in comparison to the overall rate'of dissolution of this oxide. However, Peri(12) has shown that y~  a n d  a ~^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 hydro- xylated 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 leach- ing 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 their 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 is obtained from this relation. Fin 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 complex- ing power of the anion for the oxide cation. Clearly, study of a se- lection of various acids which produce anions having different complex- ing power for an oxide cation might provide more insight into the role of anions in the leaching of metal oxides. The pH of Z.P.C. of an oxide may be a very important character- i s t i c for the leaching of oxides in acids, mainly for two reasons: (a) It gives an indication of how favourable production of a net positive charge at the oxide surface is with decreasing pH. (b) It may be related to the anion-exchange capacity of the oxide surface. The concept of Z.P.C. has been intuitively used by various inves- tigators ,(8,9,47,59). Warren et al(8,9) for example represented the rapid formation of an excess positive charge by an equilibrium equa- tion 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 + ions and thus might also be associated with the pH of Z.P.C. of the oxide. The question may now arise regarding the possi- bility 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 adsorp- tion of HSÔ  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: 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- (1.49) (1.50) 43 vation for the leaching of most metal oxides are nearly constant from one oxide to the other, irrespective of the acid, as they vary between 17 and 23 kcal/mole (Table 3); this may suggest that a similar rate- determining step is operative during the leaching of metal oxides in acids, possibly the desorption of metal species into solution. The acid leaching of metal oxides involving an oxidation-reduc- tion step in the presence of a redox couple in solution has been studied, for the most part, under relatively restricted conditions., It is logical to expect that surface hydroxylation and charging.may also be involved in the overall kinetics of the oxidative or reductive leaching of the oxides. Mackay and Wadsworth (10) have proposed that oxygen adsorbs at uncharged hydroxylated uranium dioxide surface sites and that the concentration of these neutral sites is increased by the reaction of hydrogen ions and the negatively charged portion of the UO^ surface. This approach is consistent with the hydration-cnarging properties of UO^ in acids. Laxen (83) and Needes and Nicol (86) how- ever, did not consider the UÔ  surface properties in their model of the oxidative dissolution of this oxide and were indeed unable to explain the variation of leaching rate of JJO^ with pH in the presence of Fe(III). It appears likely from their studies that both the oxide-electrolyte ' double layer properties and the type of fe r r i c species present in solu- tion at each pH have to be considered . in order^ to explain the observed kinetics 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 in the proposed mechanisms of leaching of metal oxides. Ferric, aluminum, cuprous, cupric and manganous oxides were selected for the present investigations. Extensive studies on the leaching of fer r i c oxides were planned in the hope that they might provide a basis for comparing the behaviour of oxides in general. Leaching experiments on aluminum oxides were undertaken to attempt to provide some understanding of the role of surface hydroxylation in the kinetics of the leaching of oxides. Studies on the leaching behaviour of cuprous oxide were included in the present investigations because ..the. kinetics..of •  leaching of this oxide in sulphuric acid have been explained in terms of a unique mech- anism, involving as a f i r s t step the adsorption of undissociated acid at the oxide surface. Cupric and manganous oxides were chosen, in addition to the other oxides, in an attempt to correlate;the pH of the zero point of charge (Z.P.C.) of an oxide to i t s leaching character- i s t i c s in acids. Perchloric, hydrochloric, sulphuric and oxalic acids were selected as reagents in order to study the effect of anions on the rate of leaching of the oxides. Amongst other properties, these acids differ in the complexing power of their anions for metal ions in solution. 45 3. EXPERIMENTAL 3.1 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; its 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 Atmosphere Sintering Ti (psi) T°(F) Time/days (%wt) Fe 2 + Uwt). Remarks A 15,000 air 2,100 4 0" 0.12 B 15,000 °2 1,650 3 0 0.24 C 15,000 air 2,100 2 1.3. 1.52 D 15,000 air 2,100 2 0 0.24 E 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 Q=A+D+G+H 0 0=A+D+G+H 0 X 15,000 air 2,400 1 0 ON 99.999 Fe 0.5% Ca Mixtures 0.5% Mg 47 operations under the conditions indicated in Table 5. Calcium doped specimens were prepared by a method described by Geiger and Wagner (93) and magnesium doped samples were obtained in the fashion proposed by Gardner et al (94). X-ray diffraction patterns of the samples were consistent with the ASTM card for hematite and are reported in Table A.2, Appendix A. Electron-Microprobe pictures of the Ti and Ca doped a-YeJd^ sur- faces are shown in Figure 3a to 3f. Clustering of Ti is 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 diffrac- tion 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 in 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% Ti; (b) 0.2% Ti; (c) 0.5% Ti; (d) 1.3% Ti. (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 MnÔ  present in 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 in the form- ation 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. Oxalic acid was provided by J.T. Baker Chem. Co., U.S.A. Ferrous oxalate was from Griffin and. George, England. All 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 in a glass reaction vessel maintained at constant temperature in 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 Heater 3. Contact Thermometer 4. Reaction Flask 5. Fritted Glass Filter 6. Sampling Tube 7. Gas Inlet Tube 8. Stirrer Motor 9. Reflux Condenser 10. Spin Bar I I. Magnetic Stirrer Figure 4. 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 itself 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 in 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 first sample taken (usually 5 mis) by applying an 0̂ , air or He overpressure above the solution. The analysis of this first sample was considered as a blank for the successive samplings at regular intervals. The first 10 mis of solution removed in 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 HCIÔ  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. It was then necessary to destroy the oxalic acid before analysis. This was accomplished by adding an excess of sodium persulfate (NaoSo0o) to the sample and boiling i t to eliminate 2 2 o the excess of oxidizing agent. 54 3.4.2 Aluminum The aluminum contents of the solutions were determined by meas- uring the absorbance at 580y of the pyrocatechol violet complex of aluminum ions in an ammonium acetate buffer solution (pH 6.1-6.2)(96). 3.4.3 Copper The copper contents of the solutions were obtained by measuring the absorbance at 640u of the tetraethylene pentamine complex of cop- per(Il)(97). 3.4.4 Manganese The manganese contents of the solutions were obtained by measuring the absorbance at 524u of the permanganate: ion obtained by heating the sample in the presence of excess potassium periodate (98). 3.4.5 Determination of the Ferrous Content of Hematite Specimens A five gram powder sample of the hematite was dissolved in 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 fe r r i c sulphate. This solution was titrated with a standardized O.lN eerie sulphate solution in the presence of indicator. 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+ 3 + 3 + Le + te — • Ce + Fe . This method was found to be sensitive to as l i t t l e 0.05% ferrous ion content in the five gram hematite sample (absolute error of ± 0.02%). 55 4. RESULTS All the rates of leaching were obtained from measurement of the ini 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- ily with time in the early stages of the leach, the experiment was repeated in 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 temper- ature 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 rela- tive rates of leaching, that is the ratios of the actual rates of dis- solution of the oxides over their rates of leaching in 0.9M HCIO^, were plotted against the concentration of HC10, (Figures 5 and 6). 0 0.2 0.4 0.6 0.8 Figure 5. ( HCI0 4 ) ( M/ l i ter) Relative rotes of leaching of C u 2 0 and CuO in HCIO^ versus the concentrat ion of H C I 0 4 at 12 ° C . ( Table B . 5 , Append ix B ) • b o o o: Calculated Measured 3 4 5 6 ( HCI0 4 ) ( M / l i t e r ) 8 Relative rate af leaching of goethite (Surana and H a y , T= 110° C ) and hematite (T= 9 0 ° C ) versus the concentration of HCI0 4 ( Table B 5 , Appendix B ) • 58 The rates of leaching of the oxides show a first 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 HCIÔ  begins at different acid concentrations for the various oxides, approximately in the order O.IM (Cu20) < 0.3M (CuO) < 1.5M (a-Fe^ 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 HCIÔ  at 12°C, for example, the absolute rates of leaching are approximately correlated in the following way: rate (Cu„0> = 9x rate (CuO), - -8xl'07x rate (a-FeO-OH) Z '4. - 2.7xl0vx rate (a-Fe^) . . The relatively high energies of activation obtained for the leach- ing of the oxides (42,100) suggests that the dissolution was not con- trolled by diffusion under the conditions of the experiments. The rates of leaching of the oxides were not dependent on solution agitation and this is 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 HCIÔ had to be limited to IM HCIÔ  for the leaching of Cû 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 in 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 first order dependence on a+ in 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 first order dependence of the rate on : a+. is shown on Figure 9 in 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 first order relation of the ratios with • a+ ̂  for rates which show a second order dependence on a+ . and a zero order 60 3 4 QH- ( Mo Ial) Relative rate of leaching of ferr ic oxide in dilute HCl as a function of the mean activity of HCl . ( Tables B .7 and B 7a , Appendix B ) . Relative rate of leaching of fer r ic oxide in HCI as a function of the mean activity of HCI (Tab le B. 7, Ap pendix B ) . A " '6 I Calculated A This work 80 • c O Bath 85 • c • Bath 85 • c A Roach 80 • c • Surana 85 • c JL 0 Figure 9. 10 2 0 ( Molal) 3 0 Ratio of the re lat ive rate of leaching of ferric oxide and a+ as a function of a + • ( Table B-8 , Appendix B ) v 6 3 dependence on a + .for rates which exhibit a first 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 fer- ric 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 solu- tions. 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. Figure 10 shows the effect of adding 0.6, 1.2 and 2.4M of LiCl to a 2.4N HCI solution on the relative rates of dis- solution of a-Fe20 3  and Figure 11 shows the effects of adding 0.9, 1.2 and 1.8M HC10.. and 1.2 and 1.8M NaOH to the same HCI solution on the 4 relative rate of leaching of a-Fe^O^ at 80°C. The effect of HCI concentration on the relative rates of leaching 6 4 Ef fec t of odding LiCI on the relative rate of leaching of ferric oxide ( M i c h i g a n ) in 2 4 M HCl at 8 0 ° C . ( Table B . 9 , Appendix B ) . 65 i 1 1 r 1 1 r 1 L_ I I I I I I 1.8 1.2 0.6 0 0.6 1.2 1.8 ( NaOH) ( M / l i t e r ) (HCI0 4) ( M/ l i ter) Figure II. The effect of adding NaOH or HCIO^ on the relative rate of leaching of Qt - FegO^ ( Michigon) in 2 . 4 M HCl at 8 0 ° C . ( Table B . 9 , Appendix B) - 66 of various aluminum oxides at constant temperature is shown in Figure 12. Gibbsite (AlCOH)^) dissolves in HCI solutions with a rate which shows approximately a first order dependence on - a +  . in dilute solu- tions, slowly becoming zero order with respect to a +, in concentrated HCI solutions. The y-form of A^O^ also appears to leach with a rate which exhibits a first order correlation with .'"a+--, in dilute solutions, but a sudden switch to a zero order function of . can be observed for 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 in HCI are represented in Figure 13. The oxides both appear,; to dissolve with a first order dependence on ' a+ in dilute solutions, slowly varying towards a zero order dependence on a + in concentrated solutions (Figure 14). It is 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 in HCI than in HCIÔ  at equal acidity and tempera- ture, but not by the same magnitude. Ferric oxide, for example, leaches about 10 times more quickly inl.2N HCI than in 1.2N HCIÔ , whilst cuprous oxide dissolves around 5 times more rapidly and gibbsite about 2.5 times faster. This enhancement effect exhibited by HCI over HCIÔ increases with increasing acidity; ferric oxide, for example, dissolves about 20 times more quickly in 2.4N HCI than in 2.4N HCIÔ . As is the case with the leaching of the oxides in HC10. solutions, 1 1 1—I 1 1 1 r A AI(OH)3 Measured A r - Ay>3 Figure 12. a + ( Molal ) Relative rotes of leaching of oluminum oxides in HCI versus the mean activity of HCI . ( T = 8 0 ° C ) ( Table B 7 , A p p e n d i x B ) • Relative rates of leaching of CtigO and CuO in dilute HCl versus the mean activity of H C l . ( T = 12 ° C ) ( T a b l e B . 7 , Appendix B ) • Relative rate of leaching of CuO in HCI versus the mean activity of HCI • ( T= 1 2 ° C ) ( 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 (Cu20) = 6x Rate (CuO) = 105x Rate A1(0H) = 105-107 x Rate (a-Fe203) 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 ^ (this work),and 14.7- 22.2 kcal/mole for A1(0H>3 (47,53). It is to be noted that the acti- vation energy for the leaching of y-Al"j0^ falls 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 con- sidered; that is the ratios of the observed rates over the correspond- ing 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 Ĥ SÔ  in stronger solutions (1-7.2M) (Figure 15). 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 HCIÔ  and HCI solutions, the absolute rates of leaching of the various oxides in Ĥ SÔ  solutions are quite different; in IN H^SO^ at 12°C, for example, the rates are approximately in the sequence: Rate (MnO) = 3.5'xRate (Cu 2 0) = 20.xRate (CuO) = 2.5 x 106-108 x Rate (a-Fe 2 0 3 ) The energies of activation for the leaching of the oxides in Ĥ SÔ are respectively, .10.5 kcal/mole for C^O (49), and 21 ~t 1 kcal/mole for a-Fe 2 0 3  (42) 4.4 The Leaching of Ferric Oxide in Oxalic Acid in the Absence of Added Ferrous.Salt in Solution The rate of leaching of o&-Fe,-,0 3  (Michigan) was investigated as a function of pH in 0.3M oxalic acid at 859C (Figure 18). Three distinct ZL 7 3 1.0 T - 1 Calculated Measured A Cu 2 0 O CuO • MnO a> a or 0.5 0> > a> or Figure 16. 0.5 ( H 2 S0 4  ) ( M / l iter) 1.0 Relative rate of leaching: of CUgO , CuO and MnO in HgSO^ versus the concentration of HgSO^ ( T = I 2 ° C ) ( Table B.I I , Appendix B ) l O O h 5 0 h Distr ibut ion of species in oxalic ac id at 8 0 ° C . versus pH. ( Table B . 2 , Appendix B ) • 1*J o £ \ c e i£ E w 10 o or 0 Figure 18. Calculated A Measured PH Rate of leaching of ferric oxide ( Michigan) in 0 . 3 M oxalic acid versus pH. ( T= 8 0 ° C ) ( tab le B.I2 , Appendix B ) • 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 fairly 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 in the absence of light. The distribution of H.2C204, .HĈ Ô  and Ĉ Ô  present in oxalic acid versus pH at 80°C has been calculated using the dissociation constants of 1̂ 2020̂  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 in Figure 17 (Table B.2, Appendix B). 4.5 The Leaching of Ferric Oxide in Oxalic Acid in the Presence of Added Ferrous Oxalate in Solution 4.5.1 Preliminary Experiments During preliminary experiments on the leaching of various natural hematite and goethite specimens in a 0.2M oxalic acid solution at 80°C, i t was observed that the rate of leaching of the oxides was small in the first 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 in solution prior to and during a run (Figure 19). 77 T 1 1 r Figure 19. Time (minutes) B&aching of goethite ( Minnesota) in 0.2 M oxalic acid in the presence of oir , Og and He versus time • ( T = 8 0 ° C , pH = 2 .8 ) • 78 As any impurity appearing in 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 in 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 sys- tem. It was deduced that in the preliminary experiments on natural . impure ferric oxide samples, ferrous species (or possibly other cations) appeared in solution. These, were thought, to be due to the presence of ferrous in the ores and possibly to the small iron contamination of these samples which were uniquely ground in an iron mortar. At the start of these runs any leached ferrous, was probably oxidized:to ferric by oxygen present in solution, but since the.oxygen was not renewed in solution during the run, some ferrqus/would finally persist in solution and cause the exponential increase in 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 any ferrous appearing in solution would then remain unoxidized. 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 is 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 inter- face, 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-Fe203 (sample C, Table 5) is represented in Figure 20. The rates, of leaching- of both hematites show only an approx- imately first order dependence on the ferrous species concentration in solution up to the solubility of ferrous oxalate in 0.2M oxalic acid, at which a constant rate of leaching is observed. Plots of log Rate versus log [Fe 1 Total ^ v e i 1 F i § u r e 21 are straight lines. The slopes of these lines indicate that the rate of leaching of a-Fe203 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 1 1 1 1 r 0 10 20 3 0 Figu re 2 0 . ( F e * ) t o t a | ( mg. / liter) Rate of leaching of ferric oxide in 0.2 M o x a l i c ac id versus the concentration of added ferrous ion. ( T = 80 °C , pH= 2.8 )( Table B.I 4 , Appendix B ) 81 0.4 - 0.2 - 0 - A A A Sample C A Sample Q 0.4 0.6 0.8 1.0 1.2 1.4 1.6 Figure 21. Log, 0 ( Fe ( I I ) ) ( mg. / l i ter) L o g - l o g plot of the rate of l each ing of ferric oxide in 0.2 M oxalic ac id versus the concentrat ion of added ferrous ion. ( T = 8 0 ° C , pH= 2 .8 ) ( Tabje BI4 , Append- ix B ) 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 oxalato- ferrous 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 2+ 2~f" 2+ 4" | | in solution but also to such cations as Cr , Mn , Ni , 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-Fe203 at pH 2.80'in 0.2M oxalic acid was detectable, except -4 + in the presence of 10 M of Cu ion. In this case, an exponential increase in the rate of dissolution of synthetic a-Fe203 was observed, 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 leach- ing of a-Fe203. 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 concen- tration of oxalic acid. 4.5.7 The Effect of the Ti Content of Synthetic Ferric Oxide Synthetic a-Fe203 samples were prepared under various sintering conditions and with or without additions of Ti (Table 5). Undoped a-Fe20^ samples usually contained between 0.1 and 0.2 wt % of ferrous ion and additions of Ti 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 Ti 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 unsuit- able for this study because they produce filtering problems and poss- ibly diffusion control of the leaching reactions. Also, surface area measurements do not give information on the number, size and crystallo- graphic orientation of the grains in each particle and these factors 84 T 1 1 1 1 r Figure 22 . ( H ^ C ^ ) (M / l i t e r ) Rote of leaching of ferric oxide in oxalic acid versus the concentration of oxalic acid- ( T = 8 0 ° C , pH = 2 . 8 , added fe r rous = 6 mg/l i ter )( Table B I 9 , Appendix B ) • 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 Ti content of a-Fe^O^ (Table B.13, Appendix B). The relative rates of leaching of Ti 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 addi- tion of Ti 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 % Ti, followed by a slight decrease between 0.8 and . 3% Ti. The pure a-Fe203 specimen sintered under 0^ exhibits only a slightly lower relative rate than the average relative rate obtained with pure specimens sintered under air, 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 content ( W t % ) - Relative rate of leaching of ferr ic oxide in 0 .2 M oxalic acid ( pH= 2 .8, added ferrous = 6 mg / liter ) and 2 .41* HCI at 8 0 ° C versus the titanium content of the oxide . ( Table B I 3 , Appendix B ) • 87 0.5 \- 1 1 1 o 1 1 1 A Sample 0 24 mg Fe(ll)/ l iter o Sample E 6 mg Fe(l l ) / l i ter A Sample D 6 mg Fe(l l ) / l i ter 1 1 1 A O \ l 1 i 2.7 2.8 Figure 24 . 2.9 3.0 3.1 1 , 0 0 0 / T ( ° K ) 3.2 ^ Arrhenius plots for the leaching of f e r r i c oxides in 0.2 M oxalic a c i d . ( pH = 2 .8 ) ( Table B. I6 , Appendix 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 in 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 in 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 de- crease in solubility of ferric species in 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 in 0.2M Oxalic Acid as a Function of pH The experimental results suggest that there is a definite rela- tion between the observed rates of leaching of ferric oxide and some ferrous species in the electrolyte. Hence, i t is of considerable interest. t° estimate the distribution of ferrous species in 0.2M oxalic acid at 80°C versus the pH of the solution. Unfortunately the stability constants corresponding to the formation of the mono-, di- and tri-oxalato ferrous complexes are only known at 25°C and large 1 1 1 1 1 Calculated a = 0.6 ••--^l I l I I i i l 1 2 3 4 5 F igure 2 5 . 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 ) • 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 + C 20 4 = FeC204 (4.1) (aq) (aq) (aq) F e C2°4, , + C 2 \ , = **<C2°A~f ,  (4 - 2) (aq) (aq) (aq) K Ee(C 20 4) 2" + C 20 4 ^= Fe(C 20 4) 4" (4.3) (aq) (aq) (aq) [Fe**] • [C 20 4] = Ks (4.4) (aq) (aq) [Fe"""] + [FeC204] + [Fe^O^ 2"] + (aq) (aq.) (aq) [Fe(C0)j'] = [Fe"""] (4.5) J (aq) Total(aq) where (4.1), (4.2), and (4.3) refer to the. formation of the successive oxalato-ferrous complexes, Kg is the-solubility constant for ferrous 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 ^ T o t a l = T ~ 2 ^ + K l K s + W s ' [ C2°4 _ ] ( , [C 20 4 ] (aq) (aq) + K i K 2 K 3 K s ' [ C2°4 _ ]' , ( 4 ' 6 ) (aq) 91 2- The concentration of C„0, in 0.2M oxalic acid at 80°C can be 2 4 I j 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 oxa- late 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 is given in Total 2 4 Figure 26. The values of Kg, K̂ , K 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 is represented in Figure 27 (Table B.17, Appendix B). The maximum concentrations of FeCJO^ 2- ^ aq) and FeiCJO^)2 occur at pH 2.1 and 3.8 respectively and 4- ( a q ) 2-FeXC^O^)^ reaches a. maximum of 55% above pH 5 as [CJO^ ] is (aq) 2_ (aq) limited to 0.2M/liter; the 45% remainder is then Fe(C 0 ) 1 k 2(aq). 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. L o g - l o g plot of the total solubility of ferrous species in 0.2 M oxalic a c i d versus the concentra t ion of oxa la te ion at 8 0 ° C . ( T a b l e B . 3 , Append ix B ) - I 1 1 1 1 1 1 1 1 r » Fe*" A F e ( C 2 0 4 ) | " o FeC 2 0 4 A F e ( C 2 0 4 ) * ~ Distribution of ferrous species in 0.2 M oxalic acid versus pH at 8 0 ° C . ( Table B.I7 , Appendix B ) • 10 \- IQ o E c E E o or 5 r VO Figure 2 8 . PH Rate of leaching of f e r r i c oxide ( Mich igan) in 0 . 3 M malonic ( T = 8 0 ° C , Fe(ll) = 9 mg /l i ter ) (Table B . 20 , 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, citric 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 a- Fe^O^ in HCl solutions at 80°C in the presence of Fe(.II) added as FeCl 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 first order dependence on the added ferrous concentration in solution. Figure 29- Fe(ll) u Rote of leaching versus the concentration of ( Table B.2I , Appendix B ( mg. / l i ter ) . of ferric oxide ( Michigan) in HCI added ferrous ion. ( T= 8 0 ° C ) 97 5. DISCUSSION 5.1 The Direct Leaching of Metal Oxides in 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>3 up to 1.2N HCI (Figure 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 is 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 20 3 calcined at 1200°C. In general, however, i t appears that the hydroxylation of oxide surfaces is rapid and the first step of the leaching of oxides can be represented in the case of a M̂ Ô  oxide, for example, by the following equilibrium equation: J |- M2° 3 + 7'H2° ===' l " M 0  *  0 H (5.1) s s 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. This is also, supported by Orioda and De Bruyn (107) who, in studies on the hydroxylation of the hematite (a-Fe^O^) surface identified the pres- ence 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 ions, and to a lesser extent OH and HX. As dis- cussed before, H + 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: • K |-MO • OH + H ' zS= |-MO • 0H2 = |-MO + HO (5.2) s (aq) s s where 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 fol- , lowed by the rate controlling desorption in solution of the resulting surface species, was in good agreement with the observed first order kinetics of the leaching of these oxides, in dilute solutions of the acid. Whether undissociated acid will adsorb at the uncharged oxide 99 surfaces, as postulated by Wadsworth and Wadia (49) for the leaching of Cu„0 in H„SO, solutions is considered to be doubtful for the follow- 2 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 first 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 _,. This product, with the H +  CIO4 2 assumption that a^-a^Q- is also proportional to a R +  or, in dilute solu- tions of the acid, to the square of the total concentration of perchloric acid. These findings do not correlate with the experimental results(Figure 6). (c) 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 a TT  =K, • a TT . • 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 reason- able to assumed that undissociated acid does not' participate in the leach- ing 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 =§= |-MO • OH • X (5.3) s 2 ( a q) s 1 where K is the equilibrium constant for the adsorption of X at pro-a tonated oxide sites. An equilibrium of the form of equation (5.3) can be written for every anion present in solution, and thus also for multi- charged 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 + or X . Since i t was suggested that X species adsorb preferentially at sites protonated by H+ ions, i t is proposed that the oxides surfaces may become saturated by H + ions first, followed eventually by saturation by X ions (109). Saturation of the oxides surfaces by H + and X ions can be 101 accounted for by writing a mass balance equation stating that the total surface area of the oxide, i.e. unity, is equal to the sum of the surface portions created in pre-equilibria equations (5.1, (5.2) and (5.3), namely: [|-M0 • OH] + [|-M0 • OH*] + [|-M0 • OH 2  • X] = 1 (5.4) s s s Taking into account this surface balance restriction, the proposed model was tentatively compared with the results on the leaching of ferric oxides in HCl solutions (Figure 8, curve B). It was concluded that although this model could account for the rates of leaching up to somewhat higher acid concentrations, an increasingly poorer corre- lation with the results at high acidities was s t i l l obtained. It may thus be suggested that at least one more step i s involved in the leaching of metal oxides in acids and that the contribution of this step to the overall rate of leaching becomes more apparent at high acidities. Hence, the further reaction of sites created at the oxide surface in preceding reactions, i.e. protonated and anion covered site s , is suggested to occur. The possible steps are: (a) The adsorption of H +  ions at positively protonated s i t e s , i.e. [|-M0 • 0H 2 + ]. s (b) The adsorption of X ions at anion containing s i t e s , i.e. [|-M0 • 0H 2  • X]. s + (c) The adsorption of H ions at anion containing si t e s . Steps (a) and (b) can probably be neglected as the former involves the interaction of two positively charged species and the latter 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 • 0H 2  • X + H iM=r |-MO • 0H 2  • XH (5.5) s (aq) s where K is 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 ] surface +  2H +  +  X- — [-M(H20)2X]+urface + H 2 0 is at the origin of the excess positive charge observed at oxide sur- faces 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 Cu 2 0 in ̂ SO^ that one of the dis- solution steps is the reaction of H + .ions with the oxide surface por- tion which is covered with undissociated acid. This step is also con- sistent with equation (5.5) i f the latter is written under the form of a* rate equation* because,; although the steps leading to the formation of |-MO • 0H 2  * X species are assumed to be different in the present s 103 model, the same surface species are also produced by the adsorption of undissocxated acid, HX, at the hydroxylated oxide surface. With a few exceptions, which w i l l be discussed later on, the model may now, ;account for a l l the results on the direct leaching of metal oxides.which were investigated in the present work, i f i t is assumed that the rate determining steps for the leaching of these oxides are the desorptions into solution of the metal surface species produced inpre-equilibria (5 .2) , (5.3) and (5.5). The basis for this assumption is.that in order to correlate the proposed model to the experimental results, 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 is clear from the discussion that these steps are succes- sive and inter-related. As a result i t is not possible to consider the case in.which these steps.are simultaneously rate controlling, since i f i t is assumed that one of the adsorption reactions is rate controlling i t is automatically implied that the preceding steps are rapidly achieved equilibria and the succeeding steps are quick reac- tions. This is not the case for the desorption steps since these reac- tions are independent. The possibility remains that the formation of activated complexes at .the oxide surfaces, which indeed also are independent reactions, are rate determining, but i f this was the case i t was expected that the energies of activation for the dissolution of oxides would depend on the nature of the acid and, as was discussed in the review of literature (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 for the direct leaching of metal oxides in acids can be derived using the pre-equilibria equations (5.1), (5.2), (5.3) and (5.5) and the surface balance equation (5.4) which has the form: 2 "-- k -K -a  +  + k -K -K -a..-a _ + k -K •K *K -a u+ -a Rate = 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) in 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 si t e s . In rate expression (5.6) i t was also assumed that the concentration of unhydroxylated oxide sur- face sites i s zero, i.e. is large, and that the concentration of protonated anion containing s i t e s , i.e. [ |-MO • OĤ XH." 1 "] , is s u f f i c i - s ently small to be neglected in the surface balance equation (5.4). In the following sections of the discussion an attempt is made to correlatec the estimated values of the constants in rate expression (5.6) for 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 in HCIO4 Solutions The results on the kinetics of leaching of fe 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 a = a . = a ± H+ C104' where a + 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 HCIÔ  can be simpli- fied and becomes: Rate = P H + 1 + Kp ' aH+ , ( k l + k 2 * K a * a C10T ) ( 5 ' 7 ) 4 with a H + = a = a ± 4 Although the calculated rates appear lie within 5% of the measured rates, the constants in Table 6 may only be considered as approximate because the activities a + and a. i n. were estimated for 25°C, 4 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 solu- tion 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. T° of leach k l k2 k3 K P K a K a,p V K a E Activation energy (Table 3) ^(PH) (°C) / mg Metal \ (molal ̂ ) jmgmgeMetal \ (kcal/mole) \ min-gm 7 ^ 1 \ min.gm.molal/ Gu20 >11.5 12 30.0 >200 N 19.0 <io" 1 N 19.4 - CuO 9.5 12 5.3 >10 N 5.0 <io" 1 N 1.64 1̂8.0 a-FeO-OH* 8.5 110 5.97xl0~2 N N 1.25 -3 <10 N <io" 5 17 R_ 17.8-22.5 a-Fe203 8.5 90 2.00xl0~3 _N N 1.25 <io" 3 N <io" 6 19.2-22.9 * Hay (99) N = negligible 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 of 8 in the case of ferric oxides. Equation (5.7) shows that whilst the value of the pro- duct k *K can be calculated, the.individual values of the constants 2 a k„ and K cannot be obtained. However, limiting .values of these I a constants can be estimated by modifying rate equation (5.7) by the following relation: Kp * K a * V ' a C 1 0 - < < K p * a H + + 1 ( 5 ' 8 ) This equation states implicitly,that the anion covered oxide surface portion is negligibly small, and with the assumption that the experi- mental measured rates may vary by a maximum of 10%, the inequality (5.8) becomes: 0 • 1 • (K_ • a T T + + 1) P Ka < P ^ - (5.9) K ' V ' aC10T Expression (5.9) is valid up to the highest concentrations of acid which were used, i.e. a,_j_ = aoir._.= 1 in the case of Cu.O and CuO and H ClU^ z a^ + = S Q - ^ Q - = 100 in the case of ferric oxides. This allows an esti- mate of the higher limits of K for these oxides, namely K < 0.1 for a a -3 CuO and Cu_0 and K < 10 for the ferric oxides. These rather small I a 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 its copper in HCIÔ  solutions according to the overall reaction: 108 Cu20 + 2H+ — — Cu° + Cu""" + H20 (5.10) This, is because the Cu+ ion dlsproportionates in.the presence of ClO^, and suggests that the applicability of the model to the case of Cu20 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 is suggested that whole or part of these differences may be associ- ated with one or more of the following factors: (a) Differences in entropies and enthalpies of activation for leaching. However, withtthe exception of Cu20, the energies of 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 is controlled by defects.(110). (c) Differences in the total surface areas in 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-Al^^, 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. p̂-t— ' 109 where .'a is the mean activity of HCl, and with the help of l i t e r a - ture values of the mean activity coefficients of HCl for temperatures up to 60°C (111,112) and by extrapolation of these coefficients up to 8 5° e (101) '(Table B..6, Appendix B) . The values of the constants in 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 in Table 7. The calculated rate curves and the experimental points are shown in 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 in HCl solutions arid becomes: R a t e = % ' *«+ 1 ; + K p- a H+  + K p- K a- a H+' a Cl-_ V ^ p ' ^ ^ C l - )  W ± t h  <  = a H+  = a C l "  ( 5 - 1 1 } As can be seen in Table 7, both K^, the protonation equilibrium constant, and K , the equilibrium constant for chloride ion adsorption, are small for f e r r i c oxide in comparison with the corresponding cons- tants obtained for the other oxides, whilst the complexity constant for the formation of the mono-chloro-ferric complex,. K •, i s relatively large in 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 aff 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 in the same way from one oxide to the other and that these cons- • ( k i + V V a c i - + V TABLE 7 Leaching of Metal Oxides in HCl Calculated Constants in Rate Expression (5.6) OXIDE Z.P.C T° of k l k 2 leach X z (pHi)) (°C) /mg Metal j «- \ min•gm Cu20 >11.5 12 N 55.8 CuO 9.5 12 6.10 52.5 A1(0H)3 9.1 80 0.525 7.3 a-Fe203 8.5 80 _4 3x10 4'i39xl0~ K «- (molal"'') K K a,p k0-K 3 a,p K* c (ref 50) (Table 3) / mg Metal \ (molalu}) / kcal | \ min- gm. molal/ (mole/ 19.0 3.16 - 118 IO 4' 7 3 - 5.0 0.6 <io- 1 1.26 10 1̂8.0 3.0 0.333 <4xl0"3 0.110' ' small small 15-22 1.2 0.142 <10"3 0.016"2 300 19-23 * stability constant for the equilibrium: Mz++Cr = MCI(Z-1)+ For Cuo0, K corresponds to Cu+ + 2C1 CuCl" z c / N = negligible I l l 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, K̂., 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 esti- mated ratio of the rate constants, k^/k^, in Table 7, is large for ferric oxide and.small for gibbsite and cupric oxide. Although the product k̂ -K̂  ̂  can.be calculated, only limiting values of the indivi- dual constants k„ and K can be estimated, by assuming that the 3 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) is 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-R̂ O and HCK^- HCl-Ĥ 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 coeffi- cients are the same for the HCl-Ĥ O and the HCl-LiCl-H^O systems (113) (b) At constant total ionic strength,' the mean activity coeffi- cients for the HCIÔ -HCI-Ĥ O and NaOH-HCl-Ĥ O systems are the average values of the mean activity coefficients of the indivi- dual HCIÔ -Ĥ O and HCl-Ĥ O systems in the former case,and of the 112 individual NaOH-H 0 and HC1-H 0 systems in the latter. The kinetics of the leaching of cuprous oxide in 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 chlor- ide 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 is known that CuCl is almost insoluble in water _3 (114) (1.1 x 10 M/liter at 25°C). Contrary to its behavior in 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 in the form of cuprous species. It was also experimentally observed that, in contrast to the behavior of ferric oxide, cuprous oxide dissolves more quickly in dilute HCIO^ than in dilute HCI at equal acidity (Figures 5 and 13). Eventually, the rate of leaching of C^O becomes greater in HCI than in HCIÔ  with increasing acidity (a , > 1). Moreover, the addition of 0.09M of NaCl to a 0.09M HC10. H + 4 leach solution resulted in the same rate of dissolution of C^O as in 0.09M HCI. It is therefore proposed that due to the relative insolubility of CuCl, the C^O surface already becomes covered withu.- CuCl in very dilute HCI solutions, and that this surface then profS tonates as proposed in 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 |-CuCl • HCI (5.12) s (aq) s 113 Equation (5.12) thus represents an additional step in the proposed general model in 5.1.1. By adding a term for di-anion adsorption to rate expression (5.6), and with the assumption that the concen- tration of |-M0 • OH^-X-H"*" sites is not negligible, equation (5.6) s for the-Hei-Cu^O system^becomes:. , Rate k, -K -a + + k„-K K -a TT+ -a ri _ + k.-K K K -a„ + - 1 p i r 2 p a C l 3 p a a , p _H+ a c i - 1 + K -a . + K -K -a  + -a-,_ p IF p a H Cl + k.-K -K -K -K •aL.-a 2 ,. _ 4 p a a,p a, a BT Cl + K -K -K -a 2 .-a_._ p a a,p Cl (5.13) 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 CI _ C1  + K P ' a H f ) / K p ' K a - V a C l ) +  1 + K a, P ' a H* Rate ""=' (5.14) If i t can be assumed that the surface of C ^ O becomes rapidly saturated by Cl ions in dilute HCl, then (1 +  K  'a^p «  K a ' K p ' a H+' a C l " and (k /Ka'a _ + k ) - 0. Rate expression (5.14) can thus be simpli- fied to: Rate = K a,p aH+ 1 + K -a u + a,p H^ '  (k 3  +  V ^ a ^ (5.15) with • a +  = a H +  = a c l _ 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). Limiting values of the individual constants k., K and K were obtained by allowing a H 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 Cû O .in HCI* Furthermore,; the chloride covered cuprous oxide surface is expected to exhibit a more acid pH of Z.P.C. than the hydroxylated surface, i.e. K < K in Table 8. The limiting values of k, and a,p p 6 4 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 desorp- tion of cuprous species from dichlorinated oxide sites is 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^ is 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 is controlled by the slow hydroxy- lation reaction at the oxide surfaces. It is to be noted that the energy of activation of 13.1 kcal/mole obtained for the dissolution of y-Al 20 3 in 3; N HCI between 50 and 90°C is 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) k- k_ k_ . k. K K K K . . k. • K 1 2 3 4 p a\: a,p a, a 4 a, a m-Cu_\ /mS_Cu_) /mS_Cu_) ( 5S_Cu_ ] ( m o l a f l ) (molar 1) (molal'1) (molal"1) ( . mS C u n , ) min-gm/ 1 min. gm / lmin-gm/ \mxn.gm/ \ mm- gm-molal I N N 38.6 >103 19 >2xl04 3.0 <10_1 150 N = negligible 116 5.1.4 The Leaching of Metal Oxides in H„SO, Solutions c 2 — 4 — 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 = HSO, and X~ = SOT, since ELSO, z 4 4 4 2 4 produces two types of anions in solution upon its dissociation. The mean, activity coefficients of Ĥ SÔ  have been estimated from literature data up to 60°C (115,116) and by extrapolation up to 80°C, with the assumption that . a^-= = a^gQ- where a + is the • . _~... 4 T, - .• mean activity of R̂ SÔ  (Table B.jJO,Appendix B) . The following expression for the rate of leaching of oxides in Ĥ SÔ  in terms of the general model was sufficient to describe the experimental results: k -K - a ^ + k_-K -K -a .-a„^= + k -K -K -K -a + R a t e = 1 P 2 p a H+ S04 3 p a , a,p H+ .. . 1 + K -a . + P H + 'aS07 * k4- Kp 'K ' V" aH S 04 • • • 4 ••• V j H =- (5.16) K -K -a +-aori= p a H S04 This expression for the rate of leaching contains terms for the pro- tonation of the oxide surface', SÔ  and HSÔ  adsorption at protonated sites^and protonation of sulphated sites. It is also suggested that the oxide surface may become saturated successively by adsorbed H+, and SO., ions. > Since i t is assumed that a + = a„„__ and having ..- • 4 - H HbU4 a 4.'a„._ = K, „*a.T„„- where K, „ is the second dissociation constant H + SO4 d,2 T1S04 d,2 of H 2 S 04 i n w a t e r ( 1 1 7) > i t : i s deduced that a S Q= - &d 2 < The above rate equation can be simplified to yield: 117 R a t e s k a ) - V + k(2 ) - V ( 5 . 1 7 ) where the constants in equation (5.16) have been grouped in single 2 constants k,.x and K / - N for the terms in a,^ and a„+. The numeri-(l) (l) HtF H T cal values of the constants in rate expression (5.17) were calcul- ated by comparing this equation with the experimental results (Fig- ures 15 and 16) (Tables B.llj 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 SÔ  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 + ions a ,p a at SÔ  containing sites and of HSÔ  ions at protonated ferric oxide sites is 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^. The large value of rate constant k^ compared with k 2 suggests that the desorption of ferric species from a HSÔ  containing surface site is . more rapid than from a SÔ  containing site.; This might be due to the double coordinating ability exhibited by SÔ  ions, which is lacking TABLE 9 Leaching of Metal Oxides in H2SO4 Calculated Constants in Rate Expression ( 5 .17) OXIDE Z.P.C. TP of leach k (D k (2) K (l) K* (ref 50) E Activation (pH)- (°C) /. . mg -Metal \ / mg Metal \ (molal "*") (molal "*") energy(Table 3) *. / kcal | \ min-gm-molal/' [ min .gm-molal z  ) I mole ' Cu 2 0 >11.5 12 2.87 x 10 3 2.31 x 10 4 123 io 2 ' 15 14.0 (ref CuO 9.5 12 3.91 x 10 2 3.14 x 10  3 123 io 2 ' 15 = 18.0 MnO >9 12 1.85 x 10 3 1.80 x 10 4 20.7 10 2.26 - a~ Fe 2°3 8.5 6.27 x 10" 2 1.8 x io -2 1.89 1Q 4.04 19.0 - 23.0 *K stability constant for the equilibrium C  M Z + + SO; = MSO, (z - 2) + 4 4 Calculated Constants in Rate Expression(5.16) OXIDE k l  k 2 / mg Metal \ -«- [ min-gm / * 4 . 2 (ref 117) (molal -1 ) K P (molal -1 ) K a (molal -1 ) K' a (molal -1 ) K a,P (molal -1 ) V k 4 / mg Metal \ I min-gm / -4 ^ - -3 2 -2 a-Fe o 0 o  ' 3 x 10 0.09 1.86 x 10 1.2 3.06 xlO <3 x 10 <16 >0.5 . 2 3 119 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 in oxalic acid solutions, with X = C 2 0~ and X = HÎ O^, and becomes: Rate = k - K + + k - K ' K a - a p + « a _ n= + k -K -K -K 1 p H 2 p a tv C 2 O 4 3 p a a,p 2 ... *a„+'a = + k -K «K''a„.'au_, + k -K ... 2 4 4 p a tiT HĈ Ô  5 p ... •Ki.-Kj- - a 2 , ^ - n- (5.18) a a,p H + ^ C 2 ° 4 In writing equation (5.18) i t was assumed that the concentration of the oxide surface species produced in the corresponding pre-equili- bria reactions is negligibly small, i.e. that the ferric oxide sur- face 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^ is 2 • 2 equal to k. . Moreover, a .-a,,,, n-= K,. ,~'.aij+'a -_ and a +'aTTr, 4  H +. HC2O4 d-2- H + -HC204 H H C 2 ° 4 •,-K,~, * a. 4.*a-&.„- n , where K. , and K, „ are respectively the first and •• d,l H 2 C 2 0 4 ' d,l .d,2 * - 3 second dissociation constants of Ĥ Ĉ Ô  in water. T̂hus", rate expression (5.18) can ' be simplified and becomes: : ( 1 ) ' V + K(2)' aH +' aC 2 0^  + k ( 3 ) ' V A H C 0 2"4 (4) V H O C O 0, (5.19) . .k,, N *a 4_'a. " 1 2"2 VJ 4 120 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 activ- ities .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 com- plexing 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 solu- tion 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 Ĥ X. Although the pK for the formation of FeX species in this acid is 25.1 (118), E.D.T.A. solutions did not leach ferric oxide at an appreciable dis- solution rate. At least three factors may be.suggested for this observation: TABLE 10 Leaching of a-Fê Ô  (Michigan) in Oxalic Acid. Calculated Constants in Rate Expressions (5.18) and (5.19) V V V \c (1) (2) *(3) *(4) / mg Fe j / mg Fe 2 j / mg Fe 2| / [ min•gm-molal/ ( min.gm-molal / \ min-gm-molal / I mg Fe min -gm-molal ao -3 4.6 x 10 1 mg Fe min.gm / mg Fe | I min • gm / '3 x 10 -4 >4.0 x 10 -2 5.8 x 10 -1 1.1 x 10 -1 / mg Fe | 1 min .gm / / mg Fe j \ min• gm ' >5.0 x 10 -2 >5 x 10 -2 K • P (molal "*") K a (molal "*") K' a (molal 1) K a,p K' : a,p -1 -1 (molal ) (molal ) 1.2 <10 <10 <1 122 (a) The solubility of E.D.T.A. at pH 2.7 is 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 in solution at pH 2.7 are Ĥ X and R̂ 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 in maleic, malonic, tartaric, citric and sulphamic acids. 5.2 The Acid Leaching of Ferric Oxides in the Presence of Added Ferrous Salts in Solution 5.2.1 The Leaching of Ferric Oxide in Ĥ Ĉ Ô  Solutions At least two different mechanisms can be suggested for the leaching of ferric oxide in oxalic acid in the presence of added ferrous oxalate (119)., In one mechanism, adsorbed ferrous oxalate species may lose whole or part.of their oxalate groups to neigh- boring ferric sites of the oxide surface. The-desorption of the resulting ferric oxalate complexes and the now oxalate-depleted ferrous species will result in the dissolution of the oxide whilst the ferrous content in solution is 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 will again lead to the leaching of the oxide and restoration of the fer- rous 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 is 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 Ti 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 Ti 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 (Fig- ure 23). 124 If, as i t is assumed, the overall leaching reaction involves electron-transfer, an anodic and a cathodic reaction can be included in the model of the leaching. It is suggested that these reactions are: (a) anodic Fe(C 0 ) ( 2 n - 2 ) " 2 4 n adsorbed Fe(C ?0,) ( 2 n- 3 )-2 4 n adsorbed + 6 (5.21) (b) cathodic -FeO • OH +'H 0 + e l-Fe(OH); (5.22) It is further suggested that the electron-transfer reaction is also the rate determining step in the overall dissolution reaction of a-Fe203. The Butler-Volmer equation can be written for the above anodic and cathodic reactions using the high-field approximation (121) if ri and n , the anodic and cathodic overpotentials, are assumed to a c be sufficiently large. Hence, the anodic current per unit area at the oxide surface can be expressed as: 1 = 1 • exp a a,o ( 1 - ° a > i # (5.23) where i is the exchange-current density, and a , the transfer a,o & a' coefficient, is defined as the fraction of the overpotential contri- buting to the increase in the rate of the reaction. The cathodic current per unit area becomes: l = - l c c,o exp -a RT (5.24) 125 At a potential E^, i.e. the mixed potential, | i c | = | l I- , Using the equilibrium potentials E and E corresponding to the anodic and c l C cathodic reactions, and by rearrangement of equations (5.23) and (5.24), the expression of E^ is given by: 1-a 1-a +a a c E + a 1-a +a a c c; (5.25) Delahay and Berzins (122) introduced an equation which correlates the exchange-current density, i Q , to the potential independent rate-constant of the reaction at the surface, k , and to the acti- o vities of the oxidized and reduced species of the couples, with the assumption that the potential in the outer plane of closest approach of the redox species is constant: i =. k o o 1-a a F • a.\ • â ^ (5.26) The difference",' E -E , can be expressed in terms of the equilibrium c a constant K of the overall reaction by use of the Nernst equation as follows: l[|-FeO-OH] • a[Fe(C 20 4)^ 2 n _ 2 )"] - • [E c -E a ] = In K + in ads. (2n-3)-l[|-Fe(OH)-] • aCFe(C 20 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 c \l-a +a J 1 = F * (k ) a ° • (k o,a o,c a V / 1-a c \ ( a a (2n-2)- n 1 - aa + ac/ / \ \ 1 - a a + a N " « < W » ' W • K | - F e O . O H ] f (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 2- 4- species are involved in the dissolution, i.e. Fe(C204)2 and 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 in (5.28) is assumed to s be unity. The activities of the adsorbed oxalato-ferrous complexes depend in turn on the activities of H + and the corresponding oxalato- ferrous complexes in solution, i f the same model as for the direct dis- solution of ferric oxide in the presence of adsorbing anions can be applied. The rate expression in its complete form, i.e. including the dependence on H+ and ferrous species in solution, becomes: / V Y / K Vl-a  + a ^/ I l-a'+a Rate = k ( 1 ) • (a H + . a F e 2-) a + ^ 2 ) ' C a ^ a 4-A  A 2 4 3 (5.30) 127 3' The rates of leaching of a - F e ^ (0% Ti) and a - F e ^ (1.3% Ti) in 0.2 M oxalic acid at a constant pH of 2.80 show respectively a 0.66 and 0.60 power dependence on the concentration of added ferrous oxalate (Figure 21). It thus appears that the effect of the Ti content of a-FeJ0^ is to modify the values of the transfer coefficients a , a ' and a ' i n rate expression (5.30). For simplicity, i t is assumed 3. C that a = a = a ' = a ' (Table B.18, Appendix B). The numerical values c l C 3. O of and  i n  (5.30) were calculated using the experimental re- sults on the variation of leaching of a-Fe 2 0 3  versus pH (Figure 25) (Table B.18, Appendix B). Rate expression (5.30) is found to correlate very well with these results. Finally, the experimental rates of leach- ing of a-Fe^^ as a function of the concentration of oxalic acid at a constant pH of 2.80 and in the presence of 6 mg/liter of added ferrous ion compare well 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-Fe 2 0 3  single crystal i s significantly different in ferrous species containing oxalic acid than in HCl, E^SO^ and HCIO^ solutions (Figure 30a to 30d).. In the latter acids, uniform attack (HCIO^) or evenly distributed pitting attack (HCl, H^O^) of the basal plane of a-Fe^^ is observed, whilst in oxalic acid in the presence of ferrous species localized etching can be seen (Figure 30d). The etch pits appear to be aligned along three crystallographic directions with their edges, which are parallel to these directions, forming pseudo hexagons in the basal plane. These directions may correspond 128 (c) (d) Figure 30. Morphology of the acid attack on the basal plane of a a-Fe^O^ single crystal. (a) 9 M H C 1 Q A , 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 acid, 6 mg/liter Fe(II), 80°C, 2 0 min, x 1 , 0 0 0 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 even- ly distributed protonated sites of the oxide surface. The cathodic reaction will cause pitting of the oxide surface, since i t is 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 will not modify the morphology of the oxide surface, since only species from solution are involved in 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-Fe203 in malonic and oxalic acids as a function of pH (Figures 25 and 28) suggests that the oxide dissolves by the same mechanism in both.acids. It is proposed that malonato-ferrous species adsorb at protonated 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 will result in 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^ in oxalic acid i s shifted towards a more basic pH in the case of malonic acid, i.e. from pH 2.8 to about 5. It is suggested that this difference can be associated with the distribution of the complexing ions in the two acids, i.e. malonic acid becomes completely dissociated i n water at a higher pH than oxalic acid (123). 5.2.3 The Leaching of Ferric Oxide in HCl The results presented in Figure 29 (Table B.21, Appendix B) indic- ate that the rate of leaching of a-Fe20.j in HCl i n the presence of ferrous species i n solution is only enhanced i n strong HCl solutions. This suggests that only highly complexed ferrous species, i.e. FeCl^ 2- and FeCl^ , are active i n producing an increase in the rate of leaching of the oxide. Due to the possible similarities between the ferrous catalyzed leaching of f e r r i c oxide in both oxalic, malonic and hydro- chloric acids, i t is 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 af f i n i t y of negatively charged complexes at positively charged oxide sites and/or to the enhanced rate of electron-transfer between these adsorbed ferrous species and the oxide lattice (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 in rapid adsorption prequilibria. These surface metal complexes are essentially created at three kinds of oxide sites: (a) Positively protonated sites, |-MO • OH* s (b) Anion containing sites, |-MO • OH X. s 1 (c) Positively protonated anion containing sites, |-MO • OH • XH s 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 con- taining 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 acidity of the electrolyte. This suggests that the hydroxylation of the oxide surface i s a prerequisite for enhanced speed of dissolution by species in solution. 5. The leaching of fe r r i c oxides in acids may be considerably enhanced by the presence of small quantities of ferrous species in solution. It seems that at least two conditions have to be f u l f i l l e d to observe this catalytic effect with f e r r i c oxides: (a) The electrolyte should form highly complexed ferrous species which are susceptible to fast electron transfer with fe r r i c ions at the oxide surface. (b) These ferrous complexes should exhibit af f i n i t y for adsorp- tion at the oxide surface. The experimental results suggest that oxalato-, chloro- and malonato-ferrous complexes may f a l l in this category of complexes. The mechanism of the ferrous catalyzed leaching of fer 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 fer- rous complexes and the a-Fe 0 surface are rate controlling. 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- ical fumes such as lead-zinc-iron oxides which are produced in the iron blast furnace, and iron-manganese oxides which are obtained in 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 Weight % Weight as Oxide MO m n Goethite Weight % Weight as Oxide MO ,m n V 0.84 0.13 7.20 0.11 Al 5.00 9.55 - - Ca 1.00 1.40 - - 76.00 Fe203 69.93 FeO'OH Fe '57.94 7.10 FeO•OH 52.21 11.81 Fe 20 3 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 0.85 - - Others 0.05 - 0.01 - Total 101.90 Total 99.81 • TABLE A.2 X-Ray Diffraction Patterns of Synthetic Hematite (Table 5) (Using the k Fe Radiation) a Reported Sample B Sample 0 Sample P Sample E dA 29 I/I 1 29 I/I 29 I/I 29 I/I 29 I/I 3.66 30.66 25 30.6 20 30.7 10 30.5 43 30.6 30 2.69 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 TABLE A.3 X-Ray Diffraction Pattern for Synthetic Cu20 and CuO (Using the k^Cu : Radiation) Cu.O CuO Reported — 1 •• This Study Reported This Study dA 29 I/I 26 I/I dA 29 I/I 29 I/I i o o o 3.020 29.50 •9 29.70 15 2.751 32.52 12 32.70 15 2.365 36.40 • 100 36.50 100 2.530 35.44 49 35.40 50 2.135 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 \ 1 3 8 TABLE A.4 Chemical Analysis bf Pyrolusite, 3-MnO Element or Weight % . Compound Mn02 75.20 Si0 2 5.89 Fe 0.02 CuO 3.84 P 0.045 Others, 15.0 TABLE A.5 X-Ray Diffraction Patterns of A1(0H)3> y - A l ^ and a - A l ^ , (Using the k Cu Radiation) a A1(0H) Y"A1203 a _ A 12°3 Reported. This Work Reported This Work Reported This Work -dA 29 I/I . 29 I/I dA 29 I/I 29 I/I dA 29 I/I 29 I/I o o o o o o 4.85 18.28 320 18.3 200 4.37 20.30 50 20.4 55 4.32 20.54 23 20.6 30 3.306 26.94 15 26.95 15 3.187 27.97 12 28.1 13 3.112 28.66 7 28.8 7 2.454 36.58 23 36.8 17 2.420 27.12 . 20 37.2 6 2.388 37.64 27 37.9 26 2.285 39.40 5 39.4 6 2.244 40.15 10 40.2 11 2.168 41.62 7 41.8 13 2.043 44.30 17 44.2 20 1.993 45.50 11 45.5 12 1.921 47.28 11 47.4 9 1.799 50,70" 13 50.7 15 1.750 52.22 16 52.2 14 1.689 54.26 13 54.6 10 4.56 19.44 40 - - 2.80 31.92 20 - - 2.39 37.60 80 37.5 80 2.28 39.50 50 39.4 10 1.977 45.84 100 45.8 50 1.520 60.90 30 - - 1.395 67.13 100 67.2 100 3.49 25.50 75 25.7 55 2.554 35.10 100 35.2 100 2.383 37.72 45 37.9 33 2.088 43.30 100 43.4 100 1.741 52.52 50 52.7 36 1.603 57.44 90 57.6 85 1.512 61.25 11 61.2 8 1.4055 • 66.46 38 66.6 30 1.3746 68.16 50 68.2 40 1.2396 76.84 18 76.9 12 1.2347 77.20 5 77.2 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. HC104 NaOH pH pH (M/liter) (M/liter) 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 in 0.2 M Oxalic Acid at 80°C pH [C20"] [HC204] [H2C204] ( % ) 100%=1 100%=1 100%=1 0.0 1.21 X io" 6 3.82 •X io" 2 9.60 X 10"1 0.5 1.14 X io" 5 1.14 X 10"1 8.85 X 10"1 1.0 9.00 X io" 5  • 2.84 X 10'1 7.15 X IO"1 1.5 5.55 X IO"* 5.55 X io" 1 4.41 X io" 1 2.0 2.53 X io" 3 7.95 X io" 1 2.00 X io" 1 2.5 9.25 X io" 3 9.25 X io" 1 7.35 X io" 2 3.0 2.99 X io" 2 9.45 X io" 1 2.38 X io" 2 3.5 9.00 X io" 2 9.00 X io" 1 7.15 X io" 3 4.0 2.40 X io" 1 7.60 X io" 1 1.91 X io" 3 4.5 5.00 X io" 1 5.00 X io" 1 3.98 X io" 4 5.0 7.60 X io" 1 2.41 X IO^1 6.05 X io" 5 [H + ] 2 •• K ' K where [H_C„0 ] = — ^ ^ ^ (b.l) 1 + Kd,2 [ H ] + V l K d , 2 [ H 1 K  9 [H + ] [HC„O 4 ] = ; — :  +~2 ( b - 2 > 1 + K d , 2 ' t H ] + K d , l K d , ] [ H ] [CO] = (b.3) 1 + Kd,2 [ H ] + K d , l • Kd, 2f H " J and , = 101'4(102), K, „ = 104'5(103) are respectively the first d,l a,I 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). pH [ c 2 o " ] 16g10[C20-] Measured Calculated (Equation (4.6))* [ F e ]Total •loS.10[Fe ]Total [ F e ]Total log10L~Fe (M/liter) (mgFe/liter) (mgFe/liter) 0.70 5.28x10"^ -5.280 390-406 2.590-2.610 420 2.622 1.05 2.19x10 -4.660 90-100 1.950-2.000 117 2.068 1.35 6.66x10 . -4.177 51-55 1.706-1.741 52.6 1.722 1.60 1.54x10 -3.813 33-36 1.518-1.556 34.7 1.540 1.90 3.80x10 -3.420 28-30 1.447-1.477 30.5 1.484 2.80 3.78x10 -2.423 35-40 1.544-1.600 37.0 1.568 3.55 y ~2;tf)xi0 -1.700 102-110 2.008-2.041 116.0 2.062 4.00 4.8x10 -1.320 320-336 2.505-2.526 283.0 2.451 4.50 . 1.0x10 -1.000 665-680 2.823-2.833 705.0 2.847 *Stability Constants Ref Ref Ref This Work -i „ yr (104) (105) (106) S10 (i) (25°C) (25°C) (25°C) (80°C) l o g ^ - - - 4.0 l o g 1 0 ( K 2 ' K l ) 5 , 1 4 , 5 2 9 , 5 7 6 , 3 log 1 0(K 3-K 2-K 1) 6.21 5.22 - 7.1 K - - - 3.76xl0"8 s 143 TABLE B.4 Calculated Mean Activities of HC10 Solutions, at 25°C HC10, 4 HC10. 4 Y,;(25°C) a±,-aH+=ac Molarity Molality ~ Mean Activity Mean Activity* Coefficient (M/liter) (Molal) (ref.108) (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 *The mean activity is calculated as a+; - m. • y + where m is the molality and y + the mean activity coefficient of HC10 TABLE B.5 Experimental and Calculated Rates of Leaching of Metal Oxides in HC10, 4 Solutions (Table 6, Figures 5 and .6). Rates of Leaching Oxide HCIO4 ' a ; < . Measured Calculated (origin, , Molarity Mean [Equation (5.6),Table 6] temperature Activity Absolute Relative Absolute Relative of leach) (M/liter) (Molal) /mg Metal \ ' min-gm / /mg Metal \ \ min•gm ' Cu20 0.045 0.038 12.8 0.296 13.1 0.304 Synthetic 0.09 0.0725 17.5 0.405 18.5 0.430 12°C 0.18 0.141 24.4 0.565 24.3 0.563 0.36 0.278 29.5 0.683 30.4 0.705 0.90 0.770 43.2 1.000 42.2 0.980 CuO 0.045 0.038 0.90 1.180 0.985 0.197 Synthetic 0.09 0.0725 1.42 0.282 1.475 0.295 12°C 0:45 0.346 3.91 0.782 3.80 0.760 0.90 0.770 5.00 1.000 5.20 1.040 a-Fe 20 3 Michigan . • 90°C 0.45 0.90 1.80 3.00 4.50 0.346 0.770 2.02 2.02 19 .0 0.65xl0"3 1.00x10 ~ 1.40xl0~^ 1.70x10 ̂ 1.80x10 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 a-FeO-OH 0.75 0.60 2.60xl0_2) 0.865 2.55xl0_2) 0.850 Minnesota 1.50 1.46 4.16xl0._f 1.39 3.96xl0_^ 1.32 110°C 3.00 5.50 5.27xl0_^ 1.76 5.20x10 1.73 (after Hay) 4.50 19.00 5.61xl0_f 1.87 5.74xl0_^. 1.91 6.00 72.00 6.15x10 2.05 5.97x10 1.99 a-FeO'-OH 0.15 0,.12 4.50xlO_J 0.214 0.261 Natural 0.50 0.38 1.07xl0_2 0.560 0.643 1100°C 0.75 0.60 1.92xl0_^ 0.915 0.850 (afterC 1.00 0.86 2.22xl0_^ 1.06 1.04 Surana) 1.50 1.46 3.30x10 1.57 1.32 TABLE B.6 Calculated Mean Activities of HCI Solutions HCI Molarity (M/liter) HCI Molality (Molal) Y+(12°C) Mean Activity Coefficient a +(12°C) aH+=acl_ Mean Activity (Molal) Y±(80°C) Mean Activity Coefficient a+(80°C) Mean Activity (Molal) Y±(85°C) Mean Activity Coefficient a±(85°C) Mean Activity (Molal) 0.06 0.06 0.82 0.0493 0.12 0.12 0.79 0.095 0.20 0.20 0.24 0.24 0.765 0.183 0.36 0.36 0.761 0.274 0.48 0.48 0.757 0.364 0.50 0.51 0.60 0.61 0.755 0.453 0.69 0.42 0.72 0.73 0.752 0.541 1.00 1.02 1.20 1.22 0.845 1.03 0.75 0.91 1.50 1.54 1.80 1.85 0.83 1.54 2.00 2.09 2.40 2.50 0.970 2.42 0.94 2.36 3.00 3.20 3.60 3.89 1.24 4.83 4.00 4.36 4.80 5.35 2.35 12.0 1.70 9.10 5.00 5.57 5.40 6.08 2.07 12.6 6.00 6.84 4.10 28.0 2.49 17.0 7.00 8.18 7.20 8.42 3.54 29.80 0.73 0.69 0.71 0.78 0.85 1.06 1.35 1.80 2.48 3.30 0.15 0.35 0.72 1.21 1.78 3.39 5.89 10.0 16.9 27.0 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 HCl ' . .v Measured Calculated (origin, Molarity Mean [Equation> (5.6),Table 7] temperature Activity Absolute Relative Absolute Relative of.leach) (M/liter) (Molal) /mg Metalj /mg Metal\ \ min • gm / I min • gm / Cu20 0.06 0.0493 5.7 0.0425 4.31 0.332 Synthetic 0.12 0.095 11.3 0.0843 10.9 0.0815 12°C 0.24 0.183 21.4 0.16 24.1 0.180 0.48 0.364 53.5 0.40 49.3 0.368 0.60 0.453 63.0 0.47 61.7 0.460 1.20 1.03 134.0 1.00 134.0 1.00 Cu20* 0.06 0.0493 5.7 0.0425 5.7 0.0425 Synthetic 0.12 0.095 11.3 0.0843 11.3 0.0843 12°C 0.24 0.183 21.4 " 0.16 22.7 0.169 0.48 0.364 53.5 0.40 47.6 0.356 0.60 - 0.453 63.0 0.47 60.5 0.451 1.20 1.03 134.0 1.00 144.0 1.07 CuO 0.12 0.095 :'. 2.14" 0.110 2.85 0.146 Synthetic 0.36 0.274 6.50 0.334 7.75 0.397 12°C 0.72 0.541 12.0 0.615 13.8 0.707 1.2 1.03 19.5 1.00 21.6 1.11 2.4 2.42 35.5 1.82 34.4 1.76 4.8 12.0 55.0 2.82 59.2 3.04 6.0 28.0 83.5 4.28 82.6 4.24 Calculated [Equation (5.15) Table 8]";' continued TABLE B.7 continued Oxide HCI Measured Calculated 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 3 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 continued TABLE B.7 continued a+ Measured Calculated Absolute Relative Absolute Relative o-Fe 0 0.6 0.42 1.00xl0_i? 0.265 1.09x10 0.279 Michigan 1.2 0.91 3.77xl0_^ 1.00 3.75xl0_^ 0.995 80°C 1.8 1.54 7.75xlof 2.05 8.53x10 2.26 2.4 2.36 1.58x10 f 4.20 . I,64xl0_ 4.35 3.6 4.83 4.12x10 f 10.9 . 4.49x10 1 11.9 4.8 9.1 9.25xl0_^ 24.5 1.03x10 27̂ ,3 5.4 12.6 1.69xl0_j. 44.8 1.55x10 4.1 .1 . 6.0 17.0 2.15xl0_7 57.0 2.20x10 58.4 7.2 29.8 4.50x10 119.0 4.22x10" 112.0 o-Feo0 0.2 0.15 l . l x i o j 0.013 0.0476 Synthetic 0.5 0.35 l . l x l O ^ 0.130 0.215 85°C 1.0 0.72 5.5x10 0.645 ' 0.65 (After Bath) 2.0 1.78 3.3xl0_| 3.88 2.78 3.0 3.39 9.6x10 11.3 7.08 4.0 5.89 1.85 21.8 15.3 5.0 10.0 3.12 36.7 30.8' 6.0 16.9 5.36 63.0 58.4 7.0 27.0 7.83 92.0 99.5 a-Fe 0 3.0 3.39 6.5 7.08 Single Crystal 4.0 5.89 16.2 15.3 85°C 5.0 10.0 29.6 30.8 (After Bath) 6.0 16.9 55.3 58.4 continued TABLE B.7 continued Oxide HCl a ± Measured Absolute Relative Calculated Absolute Relative cr-Fe 0 3 Single crystal 3. 0 3 .39 6 .5 7 .08 4. 0 5 .89 16 .2 15 .3 85°C 5. 0 10 .0 29 .6 30 .8 (After Bath) 6. 0 16 .9 55 .3 58 .4 a-FeO-OH 1. 0 0 .72 1, .57x10 ] 0 .654 0 .65 Natural 1. 2 0 .91 2, .40x10 1 .00 0 .995 85°C 1. 5 1 .21 3, .415x10 7 1 .42 1 .50 (After Surana) 2. 0 1 .78 6, .645x10 2 .76 2 .78 3. 0 3 .39 1. .63 6 .80 7 .08 4. 0 5 .89 3, .56 15 .25 15 .25 Ferric Oxides 1. 2 0 .91 1 .0 0 .995 Natural 80°C 2. 4 2 .36 4 .3 4 .35 (After Roach, 3. 6 4 .83 10 .74 11 .9 average results) 4. 8 9 .10 19 .38 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 H + • a c l_ = ̂  • (' a+"\) k^ = 1.2 •'&+• Relative Rate Relative Rate /w , , >. of Leaching a , 7T7 --1; (Molal) ' . . . . ± (Molal ) 0.91 1.0 1.1 1.00 1.2 1.2 2.00 4.8 2.4 3.00 10.8 3.6 4.00 19.2 4.8 5.00 30.0 6.0 (2) After the following rate expression (Curve B, Figure 7): .2 Rate = k l * aH+ ' a C l " k l ' ( a ± Y 1 + K • â + 1 + K • a4 with k = 2.85 K = 1.5 a +' Relative Rate Relative Rate (Molal) ° f L e a c h i n § (Molal"1) 0.91 1.00 1.10 1.00 1.14 1.14 2.00 2.85 1.46 3.00 4.66 1.55 4.00 6.50 1.62 5.00 8.40 1.68 6.00 10.3 1.72 TABLE B.8 Ratios of the Relative Rates of Leaching of Ferric Oxides and a, as a Function of a HCI Molarity (M/liter) Mean Activity (Molal) a-Fe 2 0 3 Michigan Relative' Rate Measured ( i f Table B.7) (Molal - 1 ) a-Fe 2 0 3 Single Crystal Bath a-Fe 2 0 3 Bath a-FeO-OH Surana Ferric Oxides Roach Calculated (Molal - 1 ) 0.2 0.15 0.087 0.318 0.5 0.35 0.372 0.615 0.6 0.42 0.631 0.663 1.0 0.72 0.896 0.906 0.902 1.2 0.91 1.10 1.10 1.10 1.09 1.5 1.21 1.17 1.24 1.8 1.54 1.33 1.47 2.0 1.78 2.18 1.55 1.56 2.4 2.36 1.78 1.82 1.84 3.0 3.39 3.32 1.92 2.00 2.08 3.6 4.83 2.26 2.22 2.46 4.0 5.89 3.70 2.75 2.59 2.59 4.8 9.10 2.69 2.13 3.00 5.0 10.0 3.67 2.96 3.08 5.4 12.6 3.56 3.26 6.0 16.9 3.35 3.73 3.27 3.43 7.0 27.0 3.41 3.68 7.2 29.8 4.00 3.76 TABLE B.9 Experimental and Calculated Rates of Leaching of Ferric Oxide (Michigan) in HCl-LiCl, HCl-NaOH and HCl-HC-10̂  Solutions (Table 7, Figures 10 and 11). HCl Molarity (M/liter) L'iCl Y;+(80°C-) -«- _̂ Mean Activity Coefficient mH+ . Molality (Molal) Cl -. -«- Relative Rate Measured Calculated Equation(5.6) 2.4 0 0.94 2.5 2.50 2.36 2.36 4.17 4.44 2.4 0.6- 1.06 2.5 3.20 2.66 3.40 5.97 6.18 2.4 1.2 1.28 2.5 3.89 3.16 4.83 8.63 8.80 2.4 2.4 1.70 2.5 5.35 4.25 9.10 17.0 15.50 HCl Molarity (M/liter) HC10. <- 4 <- NaOH -«- Y+(80°C) Mean . Activity ' Coefficient' . SH+ Molality (Molal) m c i - -«- +• aH+ •<r- Relative Measured Rate Calculated Equation(5. i 2.4 0 1.8 0.70 0.6 2.5 0.42 1.75 1.14 1.03 2.4 0 1.2 0.74 1.2 2.5 0.89 1.85 1.93 1.85 2.4 0 0 0.94 2.5 2.5 2.36 2.36 4.17 4.44 2.4 0.9 0 1.24 3.3 2.5 4.10 3.10 7.55 7.80 2.4 1.2 0 1.44 3.9 2.5 5.61 3.60 9.30 11.0 2.4 1.8 0 1.90 4/6 2.5 • ' 1 8.75 4.75 18.5 18.6 TABLE B.10 Calculated Mean Activities of H„SO. Solutions The mean activity of HoS0. was calculated 2 4 as a + = y ± ' m+ where Y + and m+ respectively the mean activity coefficient and molality of HnS0.. It is assumed that 2 4 a± = aHS0T aH+ 4 and m^ = m^- = = m + H2iS°4 H2S04 Y±(12°C) Y+(85°C) a+(12°C) a±(85°C) Molarity Molality Mean Mean Mean Mean (M/liter) (Molal) Activity Activity Activity Activity Coefficient Coefficient (Molal) (Molal) 0.036 0.036 0.735 0.452 0.0264 0.0163 0.09 0.09 0.515 0.307 0.0463 0.0276 0.18 0.18 0.403 0.248 0.0726 0.0446 0.27 0.27 0.366 0.099 0.36 0.36 0.326 0.191 0.117 0.0690 0.54 0.55 0.274 0.152 0.151 0.0836 0.72 0.74 0.259 0.151 0.192 0.112 0.90 0.93 0.245 0.146 0.228 0.136 1.00 1.03 0.142 0.146 1.08 1.12 0.242 0.271 1.8 1.93 0.132 0.255 2.0 2.17 0.131 0.284 3.0 3.42 0.148 0.506 3.8 4.45 0.164 0.730 4.0 4.80 0.169 0.810 5.0 6.25 0.197 1.23 7.4 10.7 0.296 3.17 9.0 15.0 0.435 6.55 TABLE B.ll Experimental and Calculated Rates of Leaching of Metal Oxides in H2S04 Solutions (Table 9, Figures 15 and 16). Rates of Leaching Oxide H2 S°A a ± Measured Calculated (origin, Molarity Mean (Equation 5.17, Table 9) temperature Activity Absolute Relative Absolute Relative of leach) (M/liter) (Molal) / 'mg Metal\ | 'mg Metal| 1 , min-gm / \ i min • gm ' Cu20 0.036 0.0264 21.6 0.288 21.6 0.288 Synthetic 0.09 0.0463 25.0 0.333 27.3 0.364 12°C 0.18 0.0726 33.2 0.442 33.2 0.442 0.54 0.151 49.0 0.653 49.0 0.653 1.08 0.271 75.0 1.000 72.3 0.965 CuO 0.036 0.0264 2.50 0.245 2.94 0.288 Synthetic 0.09 0.0463 3.80 0.372 3.71 0.364 12°C 0.36 0.117 5.46 0.535 5.76 0.565 0.54 0.151 6.70 0.653 6.70 0.653 0.72 0.192 7.60 0.745 7.88 0.772 1.08 0.271 10.2 1.000 9.85 0.965 MnO 0.036 0.0264 38.2 0.151 38.0 0.150 Synthetic 0.09 0.0463 56.0 0.221 63.5 0.250 12°C 0.18 0.0726 89.5 0.353 90.0 0.355 0.27 0.099 115.0 0.455 118.0 0.466 0.36 0.117 143.0 0.565 135.0 0.533 a-Fe203 0.09 0.0276 1.93x10"^ 0.305 1.64x10";* 0.259 Michigan 0.18 0.0446 2.58x10 ̂  0.407 2.58x10"̂  0.407 85°C 0.36 0.069 3.44x10 0.543 3.84x10 0.605 continued TABLE B.ll continued Oxide HoSO a Rates of Leaching Absolute Relative Absolute Relative a-Fe 0 0.90 0.0276 1.93xl0_J 0.305 1.64xl0_^ 0.259 7 1.80 0.255 V 1.08xH)J 1.70 1.09xl0_3 1.72 Michigan 3^g- 0 > 7 3 1.88xl0_^ 2.96 1.93xl0_3 3.05 7.40 3.17 2.84xl0_^ 4.48 3.10xl0_^ 4.89 9.00 6.55 3.64x10 5.74 3.64x10 5.74 a-Fe0-0H 1.0 0.146 , 1.08 1.08 Natural 80°C -=2.0 0.284 1.88 1.85 (After Surana) 3.0 0.506 2.40 2.00 4.0 0.810 3.44 3.25 5.0 ' 1.23 4.15 3.80 Ferric Oxides 1.0' 0.146 1.08 1.08 Natural 2.0 0.284 2.05 1.85 80°C 4.0 0.810 3.66 3.25 (After Roach) 5.0 1.23 4.60 3.80 Average Results 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 [c 2 o=i ° [HC204] tH2C204] Rate of Leaching (M/liter) (M/liter) (M/liter) Measured / ng'Fe ) i Calculated (Equation (5.19), Table 10) 1 mg Fe\ Imin•gm / \ , min-gm/ 0.35 1.75xlO~6 2.46xl0"2 2.76xl0_1 21.6xl0"3 20.25xl0~3 0.75 9.77xl0"6 5.50xl0"2 2.45xl0-1 11.05xl0"3 11.28xl0~3 1.05 3.28xl0~5 9.24xl0"2 2.07xl0_1 7.50xl0"3 8.14xl0"3 1.30 8.36xl0~5 1.33xl0_1 1.67xl0_1 6.80xl0"3 6.71xl0~3 2.80° 5.64xl0~3 • 2.82xl0_1 . 1.12xl0"2 4.40xl0~3 4.35xl0"3 4.00 7.20xl0~2 2.28xl0_1 5.73xl0~4 3.40xl0"3 3.32xl0~3 4.50 1.50xl0_1 1.50xl0_1 1.89xl0~4 2.16xl0~3 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 Ti Content (Figure 23) Sample Rate of Leaching in Rate of Leaching Relative (cf Table 5) 2.4 N HCl at 80°C in 0.2 M Oxalic - Rate ' Acid at 80°C Ti (mg Fe/min.gm.) (mg Fe/min.gm.) (wt %) A 0 1.64 X io" 1 6.55 X io" 1 4.00 D 0 1.96 X io" 1 6.76 X io" 1 3.46 G 0 1.83 X io" 1 7.12 X io" 1 3.90 H 0 1.71 X io" 1 6.46 X IO"1 3.77 P 0 1.35 X io" 1 4.94 X io" 1 3.66 B 0 3.31 X io" 1 1.08 3.26 I 0.1 1.37 X io" 1 1.51 11.0 J 0.2 1.36 X io" 1 2.38 17.5 K 0.4 3.-0 X io" 2 5.62 X IO"1 18.7 F 0.5 4.3 X io" 2 9.00 X IO"1 21.0 L 0.8 2.0 X IO"2' 4.44 X io" 1 •>-. 22.2 C 1.3 6.6 X io"2. 1.20 18.0 E 3.0 7.8 X io" 2 1.15 15.4 X 0.5 2.37 X io" 1 7.96 X IO"1 3.00 159 TABLE B.14 The Effect of Added Ferrous Oxalate on the Leaching of a-Fe203 in 0.2 M Oxalic Acid at 80°C and pH 2.8 (Figures 20 and 21) [ F e + + ]added log^tFe4"*"] Rate of Leaching / mg Fe \ / mg Fe | ( liter / i miri' Sample Q (Table 5) • gm ' Sample C (Table 5) Rate 1o2;LO R A T E Rate lo g 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 • . A -0.022 + 0.482 _ 0.460 _ „ ,, with: Sample Q: n = 1 - 0 7 9 _ 0 . 4 7 7 " 7Ĵ 98 " °' 6 6 n  n 0.278 + 0.142 0.420 _ Sample C: n =  1 > 0 7 9  _ . Q > 4 7 7  = Q^98 = °'6° 160 TABLE B.15 Effect of Sample Weight (Sample Q, Table 5). Leaching of a-Fe 0 in 0.2 M Oxalic Acid at 80°C and pH 2.8, ++ with 6 mg Fe /liter. Time Rate (min) / mg Fe \ /-mg -Fe \ i mg Fe\ 'l i t e r / 1 liter/ (min-gm/ 1 gm sample 2 gm sample 1 gm sample 2 gm sample 10 5.7 10.1 20 11.6 19.9 0.590 0.530 30 17.8 28.9 40 23.6 42.4 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  Sample D (Table 5) Sample E (Table 5) (0°C) ( ° K ) (°K) _1  Rate l o g i o  R a te l o g 1 0 Rate log 1 0 Rate (50-65 mesh) Rate Rate /6mg Fe"̂ 1 " ] /24mg Fe 4-1- ] / 6mg Fe 4-1- ) \ liter ' \ liter ' 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 Activation energies: Sample D (6mg Fe /liter) 12.2 kcal/mole I 0.5 Sample D (24mg Fe^/liter) 10.5 kcal/mole t 0.5 Sample E (6mg Fe^/liter) 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 stability- constants K , K„, K in Table B.3). PH Fe""" Fe(C204) Fe(C 20 4) 2~ Fe(C 20,) 4 (%) (%> (%; Z (%; 3 1.00 8. 47xl0 _ 1 1.53xl0_1 5.50xl0~4 6.23xlO~® 1.25 6. 82xl0 _ 1 3.19xl0_1 2.97xl0_3 8.73xl0"7 1.50 4. 70xl0 _ 1 5.22xl0_1 1.16xl0~2 8.10xl0~5 1.75 2. 81xl0 - 1 6.86xl0_1 3.35xl0_2 5.15xl0"5 2.00 1. 53xl0 _ 1 7.72xl0_1 7.80xl0"2 2.48xl0"4 2.25 7. 80xl0 _ 1 7.70xl0_1 1.51xl0_1 9.43xlO,"4 2.50x 3. 80xl0 - 1 7.03xl0_1 2.60xl0_1 3.03xl0~3 2.75 1. 76xl0~2 5.90xl0-1 3.95xl0-1 8.35xl0~3 3.00 7. 46xl0"3 4.47xl0_1 5,36xl0-1 2.02xl0"2 3.25 - 3.09xl0_1 6.50xl0.-1 4.30xl0~2 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 5.00 - 1.66xl0"2 5,03xl0_1 4.82xl0_1 5.50 — 1.35xl0~2 4.50xl0_1 5.25xl0_1 6.00 1.16xl0~2 4.40xl0_1 5.48xl0_1 Note: 100% = 1 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(C204)22-J r F e(C 20 4) 4-] Measured Calculation [Equation (5.30)] Absolute Normalized ..Normalized Normalized (M/liter) (M/liter) /.mg Fe ] 1 a =a =a1=a'=0 .6 a =a =a'=a'=l Imfnvgm / 1 a c a c k, .=1.31 x kj2j=5.00 x < 10 a c a c , k .=9.84 x 10^ kj2j=1.97 x 10 • 0.50 9.80xl0-1° -14 1.38x10 0.026 • 0.003 1.05 7,82xl0_8 1.08xl0 _ 1 1 0.113. 0.185 0.170 0.069 1.30 4.00xl0_7 1.40xl0 _ 1 0 0.182 0.298 0.330 0.196 1.60 1.85xl0~6 1.80xl0_9 0.256 0.420 0.563 0.465 1.90 . 5.70xl0-6 1.36xl0"8 0.370 0.606 0.760 0.740 2.50 2.60xl0_5 3.03xl0_7 0.610 1.000 0.942 1.000 2.80 4.18xl0~5 9.94xl0_7 0.603 0.988 0.920 0.960 3.65 7.37xl0_5 1.14xl0~5 0.400 0.655 0.676 0.664 4.10 6.91xl0_5 2.48xl0-5 0.220 0.360 0.445 0.440 4.70 5.50xl0_5 4.25xl0-5 0.146 0.239 0.236 0.182 5.00 5.03xl0~5 4.82xl0_5 _ — 0.163 0.100 164 TABLE B.19 Effect of Oxalic Acid Concentration on the Rate of Leaching of «-Fe203 (Sample Q, Table 5) at 80°C and at pH 2.8 (Figure 22). [H2C04] • : [Fe(C204)2-] [Fe(C 20 4) 4-] Rates of Leaching Molarity Molarity Molarity Measured Calculated [Equation (5.30)]* (M/liter) „, (M/liter) (M/liter) / mg Fej Vmih.gm/ / mg Fe\ 1 min.gm/ 0.05 1.46xl0~5 8.65xl0 _8 0.263 0.268 ;0.10 2.62xl0"5 3.31xl0~7 • 0.421 ;.. 0.415 0.15 3.50xl0~5 6.25xl0_7 0.560 0.519 0.20 4.18xl0~5 9.94xl0~7 0.605 0.605 0.30 5.18xl0-5 1.85xl0~6 0.723 0.740 0.40 5.85xl0~5 2.78xl0~6 0.780 0.845 0.60 6.61xl0~5 4.70xl0"6 0.870 1.000 * a = a = a ' = a ' = 0.6 a c a c 165 •TABLE B.2G Rate of Leaching of a-Ie^^ ( S a m P l e H> Table 5) in 0.5 M Malonic Acid at 80°C versus pH in 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"3 2.7 2.33 X i o " 2 3.2 3.75 X i o " 2 4.3 7.50 X i o " 2 5.0 8.35 X i o " 2 5.9 5.78 X i o " 2 6.4 3.67 X i o " 2 166 "TABLE B.21 Effect of Adding Ferrous Ion on the Leaching of a-Fe203 (Michigan) in HCl.Solutions at 80°C (Figure 29) [HCl] [Fe l a d d e d Rates of Leaching Molarity (M/liter). 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