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UBC Theses and Dissertations

The dissolution of goethite in aqueous solutions of sulphur dioxide Hay, Malcolm Geoffrey 1973

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THE DISSOLUTION OF GOETHITE IN AQUEOUS SOLUTIONS OF SULPHUR DIOXIDE BY MALCOLM GEOFFREY HAY B.A., Univ e r s i t y of Oxford, 1968 A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF SCIENCE i n the Department of Metallurgy We accept t h i s thesis as conforming to the required standard THE UNIVERSITY OF BRITISH COLUMBIA February, 1973 In p r e s e n t i n g t h i s t h e s i s i n p a r t i a l f u l f i l m e n t o f the requirements f o r an advanced degree a t the U n i v e r s i t y o f B r i t i s h Columbia, I agree t h a t the L i b r a r y s h a l l make i t f r e e l y a v a i l a b l e f o r r e f e r e n c e and study. I f u r t h e r agree t h a t p e r m i s s i o n f o r e x t e n s i v e c o p y i n g o f t h i s t h e s i s f o r s c h o l a r l y purposes may be grant e d by the Head o f my Department or by h i s r e p r e s e n t a t i v e s . I t i s understood t h a t c o p y i n g or p u b l i c a t i o n o f t h i s t h e s i s f o r f i n a n c i a l g a i n s h a l l not be allowed without my w r i t t e n p e r m i s s i o n . Department o f MFITALLTTROY  The U n i v e r s i t y o f B r i t i s h Columbia Vancouver 8, Canada Date TTt.h AFRTT. Tp7? -- i i -ABSTRACT The leaching of goethite in perchloric acid ("direct leaching") and i n acidified sulphur dioxide solutions ("reductive leaching or dissolution") has been investigated. The effects of temperature, hydrogen ion concentration, sulphur dioxide partial pressure, s t i r r i n g speed and sample weight on the reductive leach rate have been studied. The enhancement of the reductive leach rate as a consequence of adding cupric ions was also investigated. A leaching mechanism was proposed for the direct and uncatalyzed reductive dissolution. Protonation of the oxide surface sites, described by a Langmuir Adsorption Isotherm, resulted in the formation of ferric-hydroxy species. These either desorbed (the single solution step i n direct leaching) or underwent the adsorption of a sulphur dioxide molecule possibly followed by another hydrogen ion. The overall reductive leach rate was a function of the desorption of the intermediate and the f i n a l surface species. Cupric ions did not appear to directly catalyze a step in the postulated reductive dissolution mechanism. Rather, the essential feature of their catalytic role was thought to be an electron exchange reaction between Cu^ and Fe^* on the oxide surface to form cupric and ferrous ions i n solution. An anionic cuprous species capable of adsorbing on the oxide surface was formed by cupric ion reduction with aqueous sulphur dioxide and subsequent complexing with bisulphite and/pr sulphite ions. A comparison of the reductive leach rates of magnetite and hematite with goethite was made both with and without the presence,of cupric ions. - i i i -TABLE OF CONTENTS Page 1. INTRODUCTION 1 1.1 General , 1 1.2 Iron Oxides 2 1.3 Review of Related Work 2 1.3.1 Reductive Dissolution of Iron Oxides by Aqueous Sulphur Dioxide 2 1.3.2 Reductive Dissolution of Manganese Dioxide by Aqueous Sulphur Dioxide 10 1.3.3 Catalysis of the Dissolution of Oxides ..... 11 1.3.3.a Iron Oxides 11 1.3.3.b Manganese Dioxide 12 1.3.3.C Zinc Ferrite 13 1.3.3.d Uranium Dioxide 14 1.4 The Sulphur Dioxide-Water System 16 1.5 Scope of the Investigation . , 23 2. EXPERIMENTAL . 25 2.1 Materials 25 2.1.1 Preparation 25 2.1.2 Identification 26 2.1.3 Analysis 26 2.2 Equipment , 30 2.3 Reagents ., 37 2.4 Experimental Procedure 37 2.5 Analytical Techniques , 39 2.5.1 Iron 3 9 2.5.2 Copper 40 - i v -Page 3. RESULTS 41 3.1 Goethite 41 3.1.1 Direct Dissolution in Perchloric Acid ...... 41 3.1.2 Reductive Dissolution i n Aqueous Sulphur Dioxide Solutions 41 3.1.2.a Variation i n Stirring Speed ....... 41 3.1.2.b Variation in Sulphur Dioxide Partial Pressure . i 43 3.1.2.C Variation i n Hydrogen Ion Concentration 43 3.1.2.d Variation i n Sample Weight 43 3.1.2.e Variation in Temperature 47 3.1.2. f Effect of Ferrous Ions i n Solution 47 3.1.3 Catalysis of the Reductive Dissolution by the Addition of Cupric Ions 47 3.1.3. a Variation in Concentration of Cupric Ions 49 .3.1.3.b Variation i n Partial Pressure of Sulphur Dioxide 51 3.1.3.C Variation i n Hydrogen Ion Concentration , 51 3.1.3.d Variation in Sample Weight. 51 3.1.3.e Variation in Temperature 55 3.1.4 Effect on the Reductive Dissolution of the Addition of Other Cations 55 3.1.5 Effect on the Reductive Dissolution of the Addition of Oxalate Anions 58 3.2 Magnetite 58 3.2.1 Reductive Dissolution i n Aqueous Sulphur Dioxide Solutions ,. 58 3.2.2 Catalysis of the Reductive Dissolution by the Addition of Cupric Ions 58 - V -Page 3.3 Hematite 58 3.3.1 Reductive D i s s o l u t i o n i n Aqueous Sulphur Dioxide Solutions 58 3.3.2 Catal y s i s of the Reductive D i s s o l u t i o n by the Addition of Cupric Ions 58 4. DISCUSSION 1 62 4.1 Goethite 6 2 4.1.1 Dir e c t D i s s o l u t i o n i n P e r c h l o r i c Acid 62 4.1.2 Reductive D i s s o l u t i o n i n Aqueous Sulphur Dioxide Solutions 65 4.1.3 Ca t a l y s i s of the Reductive D i s s o l u t i o n by the addition of Cupric Ions .78 4.1.4 E f f e c t on the Reductive D i s s o l u t i o n of the Addition of Other Cations 89 4.1.5 E f f e c t on the Reductive D i s s o l u t i o n of the Addition of Oxalate Anions 90 4.2 Magnetite T 9 1 4.3 Hematite 9 1 5. CONCLUSIONS 9 3 5.1 Summary ' 9 3 5.2 Suggestions f o r Future Work 9 ^ 6. REFERENCES 96 - v i -LIST OF TABLES Table Page I Variation of the f i r s t dissociation constant of aqueous sulphur dioxide with temperature 20 II X-ray diffraction data for goethite ore 27 III Gravimetric analysis for principal ore components.. 28 IV Spectrographic analysis of sized ores 29 V Leach rates of 1 g goethite samples at 110°C , 2^ VI Variation of leach rate with s t i r r e r speed ........ 44 VII Concentrations of aqueous species at 110°C 45 VIII Variation of leach rate with sample weight 48 IX Variation of leach rate with temperature .......... 48 X Variation of leach rate with cupric ion concentration at 110°C 5 0 XI Variation of leach rate with cupric ion concentration and sulphur dioxide partial pressure at 90°C 50 XII Concentrations of aqueous species at 90°C 52 XIII Variation of leach rate with pH at 110°C with a cupric ion concentration of 2(10~3) M 53 XIV Variation of leach rate with pH at 90°C with a cupric ion concentrati on of 10~3 M 53 XV Variation of leach rate with sample weight at 110°C with a cupric ion concentration of 2(10~3) M 54 XVI Variation of leach rate with sample weight at 90°C with cupric ion concentrations of 10~3 and 2(10~3) M 56 XVII Variation of leach rate with temperature with a cupric ion concentration of 10" 3 M 57 XVIII Effect on leach rate of the addition of various cations , 59 XIX Effect on leach rate of the addition of oxalate anions 60 - v i i -Table Page XX Comparison of the leach rates of magnetite and goethite under identical conditions , 61 XXI Comparison of the leach rates of hematite and goethite under identical conditions 61 - v i i i -LIST OF FIGURES Variation of the f i r s t dissociation constant of aqueous sulphur dioxide with temperature Variation of the water solubility of sulphur dioxide with temperature at various partial pressures Cross-sectional view of the autoqlave head; north-south orientation , Cross-sectional view of the autoclave head; east-west orientation Schematic arrangement of the gas lines Amounts of iron leached from 1 g goethite samples with time at 110°C, 25 p.s.i.a. S0 2 > different acidities Variation of the direct perchloric acid leach rate with hydrogen ion concentration Arrhenius plot of log-^ Q (reductive leach rate) against the reciprocal of the absolute temperature. Reductive leach rates plotted against, aqueous sulphur dioxide concentration Reductive leach rates at 25 p.s.i.a. SO2 plotted against hydrogen ion concentration Comparison of the experimental and theoretical leach rates - variation with aqueous sulphur dioxide concentration Comparison of the experimental and theoretical leach rates - variation with hydrogen ion concentra-tion at 25 p.s.i.a. S0 2  Variation of leach rate with cupric ion concentration at 110°C Variation of leach rate with cupric ion concentration at 90°C, 14.7 p.s.i.a. S0 2  - ix -Figure 15 Variation of leach rate with pH at 90°C; 14.7 p.s.i.a. S0 2; 10"3 M [Cu 2 +] 16 Variation of leach rate with sample weight at 90°C; 10-3 M [Cu 2 +] 17 Arrhenius plot of log^g(catalyzed leach rate) against the reciprocal of the absolute temperature. 0 ACKNOWLEDGEMENTS The author wishes to express h i s gratitude to Dr. I.H. Warren for h i s help i n the t h e o r e t i c a l aspects of the study. The assistance of the members of the departmental workshop was much appreciated i n the design and construction of the co r r o s i o n - r e s i s t a n t autoclave heads, without which the work could not have been attempted. Thanks are also extended to those f a c u l t y and s t a f f members and graduate students who have helped i n the experimental and t h e o r e t i c a l aspects of the in v e s t i g a t i o n . The figures are cl e a r and l e g i b l e because Ms. Jacqueline Hay was kind enough to do them. , F i n a n c i a l support from the Government of Canada i n the form of a National Research Council research a s s i s t a n t s h i p was appreciated. - 1 -1. INTRODUCTION 1.1 General The leaching of iron oxides from s i l i c a t e minerals i s a widely practised industrial process (1). The presence of iron oxides in s i l i c a and alumino-silicates lowers their commercial value in such processes as high quality glass-making ( s i l i c a sands) and paper f i l l i n g and coating (kaolinite). Potential applications for the leaching of iron oxides also exist in the metallurgical industry for treating residues from various processes, and possibly for upgrading low grade iron ores T Solution mining of hydrated iron oxide ores with aqueous sulphur dioxide solutions has also been proposed (2). The increase in demand for iron powders may be met by a leach and gaseous reduction reaction, with iron oxide ores as the raw material (3). Goethite, a-FeOOH, i s the most common iron mineral contaminating industrially important s i l i c a t e minerals. Although many studies on the direct and reductive leaching of other iron oxides have been reported (4-14), there has been less work done on the kinetics and mechanism of goethite dissolution (1,10,11,15-18). , Industrially, goethite has been leached from acidified clays by the - 2 -addition o:f sodium dithionite, ^a^S^O^, at room temperature (1). Dithionite was either added to the clay suspension, or produced in i t by a variety of ways. This method of iron removal suffered from the poor st a b i l i t y of dithionites in acidic aqueous solutions, the relative chemical inertness of dehydrated iron oxides to dithionites, and the relative expensiveness of dithionites. This investigation was undertaken to determine the kinetics and. mechanisms involved in the reductive dissolution of goethite using acidified sulphur dioxide solutions, with arid wi thout catalysts. 1.2 Iron Oxides The preparation and properties of iron oxides have been reviewed (15,17). Goethite is equivalent to f e r r i c oxide monohydrate, ot-Fe202'H20. Dehydration to a-Ye^O^ is reported as occurring between 120-150°C (19-22). The transition is believed to be pH dependent, goethite being converted to hematite after a few weeks in 0.1 N HC1 at 100°C (23), As a consequence of the reported transition temperature, the majority of experiments were carried out at 110°C or less. 1.3 Review of Related Work 1.3.1 Reductive Dissolution of Iron Oxides by Aqueous Sulphur Dioxide Monhemius (15) has reviewed the reactions between f e r r i c ions (Fe*^) and aqueous sulphur dioxide in homogeneous systems and) in a heterogeneous system using freshly preci pitated f e r r i c hydroxide (4). - 3 -Monhemius (15,16) studied the reductive dissolution of goethite in acidified sulphur dioxide solutions. Perchloric acid was used to acidify solutions since the perchlorate anion has very weak complexing powers, and the strong acid merely behaves as a source of hydrogen ions. His experiments suffered from popr reproducibility and corrosion of his (steel) equipment. He found that the le aching rate at first increased with time, becoming linear after one hour. The leaching rate increased with aqueous sulphur dioxide concentration in. a non-linear fashion. At low (< 0.2 Molar) aqueous sulphur dioxide concentrations (ISO^^ ] ) f the rate of leaching was approximately independent of sample weight, perchloric acid concentration (over the range 0.07 to 0.28 M) and aqueous sulphur dioxide concentration. At higher aqueous sulphur dioxide concentrations the reaction became heterogeneously controlled, with the leach rat;§ becoming inyersely proportional to the hydrogen ion concentration. The reaction rate increased with increase in aqueous sulphur dioxide concentration. •> Attempting to develop a leaching mechanism common to both goethite and hematite, for both direct and reductive dissolution, Monhemius postulated the following mechanism: (slightly revi sed as a consequence of later information). 1. Hydration of the hematite surface to become similar to goethite: K /Feo0o + H.O •» 2/FeOOH Z 3 2 —— / denotes the oxide surface. - 4 -With synthetic hematite t h i s r e a c t i on was reported as having been very slow (21), whereas i t was rapid with natural hematite (24). Infrared absorption studies had i d e n t i f i e d hydroxyl (OH) groups on the a - f e r r i c oxide surface (25). Two hydroxyl groups were formed as a consequence of the reaction between a chemisorbed water molecule and a surface oxide ion. The chemisorbed water prevented the complete surface coverage of a - f e r r i c oxide by a monolayer of physisorbed water (26). I t may be removed by high temperature a c t i v a t i o n . Goethite exhibited complete surface coverage by water and per unit surface area was more rea c t i v e towards water than hematite. Hydroxyl groups i n goethite had been detected by i n f r a r e d spectros copy (27). Hydrated hematite surfaces had been found to have the properties of goethite (24). 2. Protonation of the hydrated surface : /FeOOH + H . /Fe(OH )„ ^inr 2 ® denotes a p o s i t i v e surface charge. The zero points of charge (z.p.c.) of natural hematite and n a t u r a l goethite were reported to be at approximately pH 6.7 (27,28). The actual value of the z.p.c. depended on such factors as p u r i t y and hydration state. Synthetic hematite had a z.p.c. at pH 8.5. At hydrogen ion concentrations greater than that corresponding to a pH of 6.7 the surfaces developed p o s i t i v e charges. Proton adsorption and p o s i t i v e charge development on hydrated i r o n - 5 -oxide surfaces had been investigated by several authors (28-31). The existence of a goethite-like layer on hematite immersed in water was inferred from the results of one investigation (30). 3. Double prbtonation of the surface: K /Fe(0H)2® + H + ^ 3 - /FeOH 9 8 + H20, This had been postulated as occurring at pH values less than about 2 on hematite (32). 4. Anion adsorption: K /FeOH® 8 + X" 4 ^ /Fe(0H)X® The a b i l i t y of anions to adsorb on goethite and hydrated f e r r i c oxide had been extensively reported (5,6,33-35). It had been shown that when some anions were adsorbed on goethite they might be coordinated with 3+ Fe (35). The rate of dissolution of hematite had been shown to decrease in the order HF > HC1 > HoS0. > HC10., in accord with the 2 4 4 3+ decrease in complexing power of the anion for Fe (5). A subsequent investigation had shown that other factors were also important i n determining the dissolution rate (6). The anions adsorbed as a consequence of the net positive charge on the hematite/goethite surface in low pH solution. Adsorption of anions was pH dependent. In a system with more than one anion present, the one which was most able to increase the negative charge of the surface was preferentially adsorbed (35). - 6 -5. Desorption of a hydroxy-ferric complex: /Fe(OH)X® -^=> Fe(OH)x|aq) Rate equation: ^ = k^K4K3[H+][X"][/Fe(OH)2W] By postulating that [/Fe(OH)2 ] was a constant, si nee the establish-ment of equilibria 1 and 2 (K^ and K2) were assumed to be fast, this expression accounted for the results of the experiments on the direct dissolution of goethite and hematite. Monhemius considered that aqueous sulphur dioxide molecules were the leaching species in the reductive dissolution experiments, and proposed a dual dissolution mechanism based on the direct dissolution model (16): Mechanism (a). After steps 1 and 2: 6. Adsorption of aqueous sulphur dioxide molecules: /Fe(OH) ® + 2S0. a n 6 " /Fe(SO.)„e + 2H+ 2. 1 aq i £ © denotes a negative surface charge. 7. Desorption of a ferric sulphite complex: /Fe(SO3)20 -^-> Fe(S0 3) 2- ( a q ) Ferric ions were known to complex with bisulphite and sulphite anions (36). - 7 -Rate equation: k 7 K 6 [ S ° 2 a g l 2 [ / F e ( ° < ]  d t " t H + ] 2 Mechanism (b). A f t e r steps 1 and 2: 8. Adsorption of aqueous sulphur dioxide molecules: /Fe(OH) ® + 2S0_ — ^ /Fe(HSO-) ® 2 2 • aq =^ j 2 9. Desorption of a f e r r i c b i s u l p h i t e complex: / F e ( H S 0 3 ) 2 e — F e ( H S 0 3 ) 2 + a q Rate equation: d[Fe aq] o ffi - H E — = k 9 K 8 [ S ° 2 aq ] [ / F e ( ° H 2 ] The rate of d i s s o l u t i o n was the sum of these two rate equations (mechanisms (a) and (b)), Mechanism (a) was considered to be more important at low [S0_ ], low [H +], since the anionic f e r r i c s u l p h i t e i aq complex formed could readsorb onto the p o s i t i v e l y charged surface. Homogeneous control then resulted, the rate determining step being the reduction of f e r r i c to ferrous i n s o l u t i o n (4). Surana (17,18) continued the study of the d i r e c t and reductive d i s s o l u t i o n of goethite. Using a c i d i f i e d aqueous solutions of sulphur dioxide, he observed a continually increasing rate of d i s s o l u t i o n f o r - 8 -experiments of 2 1/2 hours duration. He found that the rate was l i n e a r with sulphur dioxide p a r t i a l pressure over the range 2.3 to 23.3 p . s . i . The rate was always greater f o r sulphur dioxide solutions a c i d i f i e d with 0.5 M HCIO^ than for those with 0.15 M HCIO^. The rate c o n t r o l l i n g step was heterogeneous under a l l conditions studied. The c o n t i n u a l l y increasing rate of d i s s o l u t i o n was not due to autocatalysis by the i r o n leached into s o l u t i o n . Surana postulated that the increasing rate of leaching was due to surface morphology changes on the d i s s o l v i n g oxide. His d i s s o l u t i o n mechanism involved the following steps: 1. Protonation of the surface: /FeOOH + H + — ^ . /FeO® + ^ 0 . Anion adsorption: /FeO® + X~ — ^ /FeOX 3. Desorption of a complex f e r r i c ion: (at low acid concentrations) k 3 /FeOX — ^ FeOX (aq) Rate equation: d[Fe ] - j J S - = k 3K 1K 2[/FeOOH]a H fa x-a denotes a c t i v i t y . [/FeOOH] was assumed to be constant. - 9 -The complexing power of the anion f o r f e r r i c ion determined the value of and thus the o v e r a l l leaching rate. Strong complexing anions r e s u l t e d i n f a s t leach rates. At high a c i d concentrations: 4. Double protonation of the surface a f t e r anion adsorption: /FeOX + H + 4 ^ /Fe (OH)X® 5. Desorption of a f e r r i c complex: -> Fe(OH)X" /Fe(OH)X® " 5 ^ „ - , ~ , w + (aq) Rate equation: d[Fe 1 "DT3" = K 5K 4[/FeOX] a H + = k ^ K ^ l / F e O O H ] a H + 2 a x _ Surana postulated that f o r reductive d i s s o l u t i o n with a c i d i f i e d aqueous sulphur dioxide s o l u t i o n s , the active anion was b i s u l p h i t e , i . e . X = HSO^ . The rate of leaching was therefore a function of a^+ and a^g^ and, since these were rel a t e d to the concentration of aqueous sulphur dioxide, the sulphur dioxide p a r t i a l pressure. - 10 -1.3.2 Reductive D i s s o l u t i o n of Manganese Dioxide by Aqueous  Sulphur Dioxide The a b i l i t y of aqueous solutions of sulphur dioxide to leach manganese dioxide has been known for some time. The k i n e t i c s and mechanism have been investigated (4,37-39). The d i s s o l u t i o n was found to proceed by two reaction paths (37): 1. A slow re a c t i o n i n v o l v i n g b i s u l p h i t e ions; 2. A very f a s t reaction i n v o l v i n g n e u t r a l aqueous sulphur dioxide molecules. Under the experimental conditions u t i l i z e d , a slow chemical reaction was r a t e - l i m i t i n g i n 1; d i f f u s i o n of sulphur dioxide molecules through the oxide boundary layer was r a t e - l i m i t i n g i n 2. The rate equation at 25°C, with a s t i r r i n g speed of 500 r.p.m. was:-d[Mn ] 590[S0 o 1 + 60[HSO ~] L aq 2 aq 3 1 d t 1 + 15[HS0 3"] + 8[H +] The leach rate at 25°C, pH 1.24, [SO. ] = 0.199 M, [HSO ~] = 2 aq 3 -5 - 2 - 1 0.058 M, was 8.3 (10 ) moles Mn cm min . This i s comparable with the leach rates of metals. The adsorption of H + and HSO^- ions had an i n h i b i t i n g e f f e c t on the rate under the low s t i r r i n g speeds u t i l i z e d . Under conditions of low pH the calculated a c t i v a t i o n energy was 4.5 + 0.2 k c a l mole ^. This was i n good agreement with the t h e o r e t i c a l value of the a c t i v a t i o n energy fo r d i f f u s i o n of sulphur dioxide molecules i n water at i n f i n i t e d i l u t i o n , 4.3 k c a l mole \ Experiments (38) at much higher s t i r r i n g speeds (2,300 r.p.m.) showed that for d i s s o l u t i o n by mechanism 2, the rate became dependent - l i -on the chemical reaction between undissociated aqueous sulphur dioxide molecules and manganese dioxide. The activation energy was estimated to be 7.3 kcal mole ^. The rate equation at 25°C, with s t i r r i n g speeds in excess of 2,300 r.p.m. was:-d[Mn ] 1200[SO. ] + 60[HS0 "] aq _ 2 aq 1 3 J  d t 1 + 0.65[SO„ ] + 0.8[HSO ~] 2 aq 3 1.3.3 Catalysis of the Dissolution of Oxides 1.3.3.a Iron Oxides Monhemius (15,16) added copper ions to his goethite/acidified sulphur dioxide solution system in the form of cupric perchlorate. Cupric ions were known to catalyze the homogeneous reduction of f e r r i c ions to ferrous ions (40). He found that copper additions could catalyze the dissolution under certain conditions, and found an activation energy of 23.1 kcal mole ^. The rate control at low [SO ] changed from homogeneous to heterogeneous control on the addition of cupric ions. Monhemius postulated that the cupric ions catalyzed the reduction of aqueous complex fe r r i c ions. The catalyzed rate-determining step was then the desorption of species from the surface, and not as before, the reduction of f e r r i c complex ions in solution to prevent their readsorption. In addition, the increase in leach rate observed with large additions of copper was, thought to be due to some surface reaction involving cuprous ions. - 12 -1.3.3.b Manganese Dioxide The direct dissolution of manganese dioxide in sulphuric and perchloric acids has been investigated in the presence of ferrous ions (41). The overall chemical reaction has been shown to be: Mn02 + 2Fe 2 + + 4H+ — > Mn 2 + + 2Fe 3 + + 2H20 Experiments were performed under conditions where th e reaction rate was independent of the concentration of ferrous, f e r r i c , manganous and hydrogen ions and bulk d i f f u s i on in solution. The reaction rate in a perchlorate medium was much less than that in a sulphate medium under the same conditions. When the ferrous ion concentration was less than 0.01 M the rate was f i r s t order in respect of this component. The adsorption of added cupric ions on the manganese dioxide surface was found to retard the reaction. The reaction was heterogeneously controlled. The experimental evidence suggested the following mechanism: the rate-determining step occurred on the surface, and depended on the surface concentration of ferrous ions at active sites, or the ease of ferrous ion migration to them. The observed anion effect was explained IV 2+ by postulating that an electron exchange reaction between Mn and Fe resulted in the formation of Mn*** and Fe^ +. In perchloric acid media the M^O^ formed inhibited further reaction at the active sites and dissolution occurred uniformly over the surface as observed. Sulphate III anions, adsorbed in sulphuric acid media, complexed the Mn and permitted a further electron exchange reaction. The dissolution reaction then - 13 -proceeded unhindered at such active s i t e s as c r y s t a l boundaries, causing the surface p i t t i n g observed. 1.3.3.C Z i n c - F e r r i t e , ZnFe„0. z 2—4 — The d i s s o l u t i o n of z i n c - f e r r i t e i n sulphuric acid has been investigated with and without the presence of ferrous ions i n the leach s o l u t i o n (42). Ferrous ions were shown to increase the rate of d i s s o l u t i o n , as did cathodic p o l a r i z a t i o n of the sample. Ferrous ions were oxidized to f e r r i c during the experiments; the presence of f e r r i c ions lowered the rate of d i s s o l u t i o n . The rate of d i s s o l u t i o n was proportional to the square root of the concentration of ferrous ions under the conditions studied. The researchers explained t h e i r r e s u l t s i n the following way. Cathodic p o l a r i z a t i o n of the z i n c - f e r r i t e surface, or the adsorption of a reducing agent on the surface, resulted i n the formation of an anion vacancy: F e X I ( H 2 0 ) 6 2 + ( a d s ) ^ F e m O H ( H 2 0 ) 5 2 + + H 0 2~ , + H -M> OH . + e s o l i d s o l i d Fe + e — > Fe — > Fe, ' s o l i d s o l i d (aq) k was low and rate-determining. The presence of f e r r i c ion was said to decrease the adsorption of hydrogen free r a d i c a l s (H) or ferrous ions. The mechanism of d i s s o l u t i o n - 1 4 -rate enhancement by anion vacancy increase had been postulated to explain the results of their earlier leaching experiments on specially heat-treated ferrites without the addition of ferrous ions or cathodic polarization. No experiments were undertaken with a different anion present. 1.3.3.d Uranium Dioxide The oxidative acid dissolution of uranium oxides i s an important industrial process which has undergone much investigation. In particular, the use of oxidants other than oxygen has been investigated in an effort to improve on the leaching kinetics. Laxen (43) has made an extensive study of the effect pf adding fe r r i c ions to the leach solution, and has also reviewed earlier work in this f i e l d . Laxen measured the increase in acid dissolution rate of uranium dioxide upon addition of fer r i c ions with the following anions present: perchlprate - no effect, very low dissolution rate; nitrate - no effect; chloride - slight rate increase; sulphate - large rate increase. The activation energy with sulphate anions present was 15.8 kcal mole *. No effect was observed when the st i r r i n g speed was varied from 200 to 600 r.p.m. The dissolution rate reached a maximum at pH 2 for leach solutions with constant perchlorate ion concentration. A similar maximum was found using leach solutions with a fixed i n i t i a l sulphate ion concentration. In these solutions the pH was varied by the addition of perchloric acid. When the ferric ion concentration was raised the dissolution rate increased in a manner described by a Freundlich adsorption isotherm. The a d d i t i o n of ferrous ions reduced the d i s s o l u t i o n rate by a greater 2+ amount than did the a d d i t i o n of equivalent concentrations of Ni , 2+ 2+ 2+ Co , Cu , and Mn These r e s u l t s indicated that the rate-determining step was a chemical reaction occurring on the uranium dioxide surface. The form of the dependence of the d i s s o l u t i o n rate on the f e r r i c i on concentration showed that the adsorption of a f e r r i c - c o n t a i n i n g species was important. Laxen drew attention to the strong s i m i l a r i t i e s between h i s system 2+ 3+ and homogeneous el e c t r o n exchange reactions such as the Fe /Fe redox couple. In t h i s system the rate constant f o r electron exchange i s markedly affected by the presence of c e r t a i n anions i n s o l u t i o n . The rate constant for electron exchange between the aquo-complexed ions i s very small. On addition of anions i t increases i n the order: 1^0=010^ < - 2-C l <C S0^ . Sulphate anions are p a r t i c u l a r l y favorable f o r conducting 3+ 2+ electrons between Fe and Fe or f o r reducing s t e r i c hindrances to electron exchange. Since these various e f f e c t s p a r a l l e l e d the r e s u l t s of the experiments on the heterogeneous d i s s o l u t i o n of uranium dioxide, i t was considered that the b a r r i e r s hindering the electron exchange i n the Fe /Fe system were at l e a s t s i m i l a r to those a f f e c t i n g the f e r r i c ion promoted d i s s o l u t i o n of uranium dioxide. Uranium dioxide was leached by o x i d i z i n g the U*^ to the more soluble U^*. The pH of the s o l u t i o n and the anion concentration determined what f e r r i c complexes were formed and t h e i r concentration. The peak d i s s o l u t i o n rates coincided with the maximum concentration of the f e r r i c complex considered tc be most capable of adsorbing IV V and undergoing an electron exchange reaction with the U and U species - 16 -on the oxide surface. Adsorption of cations other than f e r r i c reduced the surface area a v a i l a b l e f o r reaction and consequently lowered the d i s s o l u t i o n rate. Ferrous ions were presumed to be more e f f i c i e n t at adsorbing on active surface s i t e s . Since no attempt was made to regenerate the f e r r i c i o n s , t h i s method of promoting the d i s s o l u t i o n rate could not be considered to be a t r u l y c a t a l y t i c r e a c t i o n . 1.4 The Sulphur-Dioxide-Water System Sulphur dioxide i s moderately soluble i n water and the following major species have been detected: 1. SC^'Xl^O, a cla t h r a t e of the gas hydrate type (44-46) with x — 7. Undissociated, aqueous sulphur dioxide w i l l be referred to as SO, 2 aq 2. H +; HSO^ . Ioniza t i o n leads to the formation of th e b i s u l p h i t e anion, HSO^ : K l + S0O + HLO ^ H + HS0O 2 aq I -v: 3 2- 2-3. SO^ . Further i o n i z a t i o n gives the su l p h i t e anion, SO^ - K2 + 2-HSO„ — ^ H + SO. 3 -"v J - 2-4. HS.O_ ; SO0_ . With a high concentration of SO both the i. j z J £. aq 2-pyrobisulphite anion, ^^O,. , and the pyrosulphite anion, S2<")5 ' a r e ^ o r m e ( ^ : - 17 -HS03 + 8 0 2 ^ ^ ^ HS 20 5 K4 2-2HS0- ^ S.0,. + H„0 3 ID I Evidence for the existence of unchanged aqueous sulphur dioxide molecules was obtained from ultraviolet absorption spectra (47-50) and Raman spectroscopy (51T53). Simon and Waldmann (54-56) undertook a Raman spectroscopic study of aqueous sulphur dioxide solutions and identified the bisulphite, pyrobisulphite and pyrosulphite anions, though in very [SO 2-] low concentrations. At 20°C, K = — — = 7(10 ) (57). * [HSO3-] 2 Studies of aqueous solutions of sulphur dioxide and of solid mixtures of sulphur dioxide and water by infrared spectroscopy (58,59) had, like a l l previous studies, failed to detect "sulphurous acid", I^SO^. It was therefore considered that ^SO^ either did not exist or was present in very small amounts. The equilibrium between S0„ and H+, HS0„ has been reported as n 2 aq ' 3 being established instantaneously - within 10 ^ second (60). Aqueous solutions of sulphur dioxide possess reducing properties (44)! SO,2" + 4H+ + 2e~ = S0 o + 2H.0 E° = 0.17 v 4 2 aq 2 S.O 2~ + 4H+ + 2e~ = 2S0o + 2Ho0 E° - 0.57 v. 2 6 2 aq 2 Ostwald (61) and Barth (62) showed that aqueous sulphur dioxide solutions behave as monobasic acids even at high dilutions. The f i r s t dissociation constant for aqueous sulphur dioxide is given by: - 18 -[H+][HS03"] K l = ^ 2 aq ] or, more correctly, in activities: • a, hr*- ' °HSOI" 1 *S°2 aq' V i s equivalent to the "apparent" dissociation constant, K&, of Campbell and Maass (63): [H +][HSO~] K ~ [S0 2 a q]+tH 2S0 3] since [H 2S0 3] -> 0. Campbell and Maass (63) calculated values of from conductivity measurements made over a range of aqueous sulphur dioxide concentrations and at temperatures up to 90°C. decreased with temperature but was virtually independent of [S0_ ] up to 8%. Johnstone and Leppla (64) recalculated from the conductivity data of Campbell and Maass (63) and Morgan and Maass (65) using activities rather than concentrations. Tartar and Garretson (66) calculated K and K-2 at 25°C from E.M.F. _2 measurements on suitable c e l l s . They found = 1.72 (10 ), K 2 = -8 6.24 (10 ). Their value of K 2 was in reasonably good agreement with the results of other workers which they enumerated. Subsequently Cuta —8 et a l . (67) obtained a value of 7.10 (10 ) at 25°C from potentiometric and spectrophotometrie measurements. - 19 -E l l i s and Anderson (68) calculated values of at 25°C at pressures -2 up to 2000 atm. by conductance measurements. They found = 1.39 (10 ) at 1 atm and no appreciable change at pressures up to at least 5 atm. Rabe and Harris (69) developed an expression for calculating at any temperature, T°A, from earlier conductivity measurements (64,68). Their equation: K x = exp ( 1 9^ 2 , 5 - 10.967). . Arklipova et a l . (70) gave values for K.^ at 10, 25 and 35°C of 1.8 (10 1.3 (10~2) and 1.0 (10~2) respectively. Table I gives values of at different temperatures according to the various authors. Values of are plotted against temperature in Figure 1. From this graph were taken the values of used to calculate bisulphite and hydrogen ion concentrations at 80, 90, 100, 110 and 120°C; namely 3.95 (10 _ 3), 3.20 (10 - 3), 2.70 (10~ 3), 2.35 (IO - 3) and 2.05 (10~3) respectively. Although various authors gave sulphur dioxide solubility data, only several (63,69,71,72) had conducted studies over a large temperature range. Of these, only Campbell and Maass (63) had also gathered data at high concentrations of dissolved sulphur dioxide. The available data was used to construct the isobars i n Figure 2, relating solubility, C , to temperature. The density of sulphur dioxide at N.T.P. was taken to be 2.9269 g.l 1 (73). From a l l this data i t was possible to calculate the equilibrium molar concentrations of S0o , H*" and HS0„ at the 2 aq 3 temperatures, pH values and the partial pressures of sulphur dioxide of interest. - 20 -Temp. °C Campbell and Maass Johnstone and Leppla Tartar and Garretson E l l i s and Anderson Rabe and ^ Harris Arklipova et a l . 0 31.3 23.2 - - - -5 27.9 - - - - -10 - 18.4 - - 18.2 18.0 15 22.1 - - - - -18 — 15.4 - - - -25 17.3 13.0 17.2 13.9 - 13.0 30 15.1 - - - 11.5 -35 10.5 - - - 10.0 50 8.6 7.6 - - 7.7 -70 4.6 - - 5.4 -80 - - - - 4.6 -90 2.5 - - - 3.9 -100 2.0 - - - 3.4 -110 - - - - 3.0 -120 - - - 2.6 -Calculated. 3 TABLE I. Variation of the f i r s t dissociation constant (x 10 ) of aqueous sulphur dioxide with temperature. - 21 -40 60 80 TEMPERATURE °C 100 120 FIGURE I. Variation of the first dissociation constant of aqueous sulphur dioxide with temperature. - 22 -FIGURE 2. Variation of the water solubility of sulphur dioxide with temperature at various partial pressures. - 23 -In these calculations i t was assumed that at a l l temperatures and pH values the concentrati pns of a l l other sulphur-containing ions except HSO ~ were zero, i.e. C = [SO, ] + [HSO " ] . It was also assumed that the concentration of a l l aqueous sulphur dioxide species, C „ n , was constant for a given par t i a l pressure of sulphur dioxide at a fixed temperature and was not influenced by pH changes nor by the addition of sulphite salts. In a l l calculations concentrations were used rather than activities since activity coefficients were only known at 25°C (74). Mean activity coefficients of aqueous sulphur dioxide have been estimated at temperatures up to 130°C, over a wide concentration range (69). 1,5 Scope of the Investigation The main aim of the investigation was to gain an understanding of the mechanism and kinetics of the reductive dissolution of goethite in aqueous sulphur dioxide solutions. To accomplish this i t was f i r s t necessary to extend our understanding of the mechanism and kinetics of the direct dissolution of goethite in- perchloric acid, which was subsequently used to vary the pH of the leaching solutions in the main study. Since the reductive dissolution was kinetically slow, various cationic additions were made in order to catalyze the dissolution. Of these, only the addition of copper was found to be of value. A further investigation was undertaken to try to understand the mechanism by which the copper could catalyze the reductive dissolution. r 24 -A comparison of the uncatalyzed and copper-catalyzed rates were also made for two other important naturally occurring iron oxides, magnetite and hematite. - 25 -2. EXPERIMENTAL 2.1 Materials Natural oxide ores were used in a l l experiments. Goethite. The goethite originated from Biwabik, Minnesota, and was supplied by Ward's Natural Science Establishment, Inc., New York (10). Magnetite. The magnetite originated from the Queen Charlotte Islands, British Columbia, and was obtained from the Mineral Engineering Department of the University of British Columbia, Vancouver (14). Hematite. The hematite originated from Ishpeming, Michigan, and was also supplied by Ward's (14). 2.1.1 Preparation The as-received ores were f i r s t mechanically-or hammer-crushed and then hand-mulled to the required particle size range. For gravimetric and spectrographic analysis, and for a l l leaching experiments, the ores were wet-screened to -70+140 mesh number, U.S. standard seives. The magnetite particulate feed was mechanically purified by discarding material not adhering to a permanent magnet, and then treated with 20% n i t r i c acid for about -16 hours at room temperature to remove the exposed sulphides. After subsequent re-screening only a trace of sulphides remained. - 26 -2.1.2 Identification Goethite. X-ray diffraction analysis using Co Kc*^  radiation in conjunction with an Fe f i l t e r confirmed that the ore contained iron in the form of goethite together with a sizeable amount of<K-quartz. Table II. Magnetite and Hematite. These ores had been the subject of a direct acid leaching investigation by Devuyst (14). His X-ray diffraction analysis had shown them to be very pure. 2.1.3 Analysis Sized samples of the ores were analyzed gravimetrically for their J.ron, s i l i c a and moisture contents and sp ectrographically for other impurities by Can Test Ltd., Vancouver. Gravimetric results are shown in Table III, with the iron values being converted to the pure relevant oxide. The spectrographic results in Table IV give weight percentage concentrations for those elements detected except iron and s i l i c o n . A l l other elements were either not detected (N.D.), were present only in trace amounts, or were non-detectable. The gravimetric results showed that for the goethite ore, the iron was not a l l present in the form of goethite. Analysis showed only 7.09 g contained water in 100 g sample, corresponding to 69.93 g goethite and 43.95 g iron, i f i t was a l l contained as goethite. If the remaining 8.26 g iron were a l l i n the form of hematite, then i t would weigh 11.81 g. The contribution of goethite, s i l i c a , hematite and uncontained water to the ore's composition then would be 99.37%. - 27 -Experimental Data ASTM powder diffraction f i l e i-270 mesh ore 3.0249; c<-Fe00H 5.0490; <X-Si02 e d; A order of intensity d; A relative intensity d; A relative intensity 4.98 6 4.97 60 4.18 1 4.18 100 4.26 35 3.34 2 3.36 60 3.34 100 -2.69 3 2.69 70 2.58 7 2.58 55 2.49 8 2.48 40 2.45 3 2.44 80 2.25 5 2.25 60 2.19 4 2.18 60 1.92 9 1.92 40 1.82 7 1.82 17 1,80 7 1.80 50 1.72 4 1.72 70 1.69 6 1.69 50 1.61 9 1.60 5Q 1.56 5 1.56 65 1.54 ' 7 1.54 15 1.51 5 1.51 60 1.45 5 1.45 60 TABLE II. X-Ray diffraction data for goethite ore. - 28 --70 + 140 Mesh Ore Goethite Magnetite Hematite Analysis For:- Weight Percent Iron 52.21 6? .40 47.15 Iron Oxide 83.07 95.91 67.42 Silicon - * 5 Silicon Dioxide 17.52 * 10.70 Moisture Loss; 105°C 0.11 Total Moisture Loss 7.20 Total Compounds 100.70 106.61 * Spectrographic TABLE III. Gravimetric analysis for principal pre components. - 29 -Ore Goethite Magnetite Hematite Element Weight Percent Al 0.005 0.2 Ba N.D. 0.002 Ca N.D. 2.0 Co N.D. 0.003 Cu 0.003 0.1 Pb N.D. 0.001 Mg 0.01 0.2 Mn 0.3 0.5 Mo 0.001 Trace NI N.D. 0.002 Sr N.D. 0.001 Ti N.D. 0.03 V N.D. 0.001 TABLE IV. Spectrographic analysis of sized ores. - 30 -Since 30% of the goethite ore was contaminants i t was necessary to prepare a very large, well-mixed quantity of sized ore to reduce the likelihood of sample and leach rate irreproducibility. Depending on the leaching charapteristics of the hematite ore, i t might be necessary to apply some correction to the leach rates of the goethite ore to take into account the 12% hematite present. The magnetite and hematite analyses were incomplete. The weight percentage of s i l i c o n dioxide given for the magnetite ore was not very accurate since i t was obtained spectrographically. The total goethite surface area in a 1 g goethite ore sample was calculated assuming the goethite particles were pure and perfectly -3 spherical, and had a density of 4.3 g cm (73). If a l l the goethite spheres had a diameter equivalent to a mesh number of 70, then the total 2 surface area would have been 46.5 cm . Had the diameter been equivalent to a mesh number of 140, then the total surface area would have been 92.9 cm2. In practice the goethite was not a l l present as spheres, and there was a spectrum of particle sizes between mesh numbers 70 and 140. If the 2 total goethite surface area was taken as 75 cm , then to convert the leach rates expressed i n the text as mg Fe hr (g goethite ore ^) to -2 -1 absolute rates moles Fe cm min , the multiplication factor would be 3.98 (10~ 9). 2.2 Equipment A l l experiments were performed in either one of two essentially, identical totally corrosion-resistant autoclaves. The autoclave bodies were of 2 1. capacity and were fabricated by the Pfaudler Company, New York. - 31. -They consisted of round-bottomed cylindrical mild steel pressure vessels internally lined with Glasteel 3315. An autoclave head was made of 316 stainless steel, machined to allow a Teflon plate to be firmly secured by four tantalum bolts and Teflon washers to the internal surface, Figures 3 and 4. The bolts were equally spaced around the circumference of a circ l e of radius 1 1A6"> measured from the centre of the autoclave head. A l l the tantalum used in the construction of the autoclave was supplied by Teledyne Wah Chang Albany, Oregon. Low pressure stainless steel fitt i n g s were used to connect two 1/16" stainless steel gas lines to the outward facing ends of two of the tantalum bolts. A 1/8" diameter zirconium sampling tube and a thermistor probe were attached with similar, fittings %o the other bolts. The model 410 thermistor probe was supplied by the Yellow Springs Instrument Co., Inc., Ohio. The gas inlet and gas outlet tantalum bolts were diametrically opposite each other and had small holes d r i l l e d through their centres to permit the flow of gas. The other tantalum bolts had holes d r i l l e d through their centres wide enough for the.sampling tube or thermistor probe to pass through. In addition, they were internally threaded to allow a Teflon sleeve to f i t over the tube and a Teflon well to f i t over the probe, each forming a gas-tight seal. The Teflon sleeve, well and the thermistor probe were a l l long enough for their ends to be well below the leach solution level. In order to prevent the passage of fine ore into the sampling tube, a f r i t t e d glass f i l t e r was affixed to the lower end of the Teflon sleeve. - 32 -ZIRCONIUM SAMPLING TUBE-STIRRING SHAFT-THERMISTOR PROBE SLEEVES FRITTED GLASS-FILTER STIRRING SHAFT PLUS IMPELLER AND PROPELLER WELL 316 S.S. TEFLON. TANTALUM. * DRAWN TO SCALE. FIGURE 3. Cross-sectional view of the autoclave head; north-south orientation. - 33 -DRAWN T O S C A L E . FIGURE 4. Cross-sectiona! view of the autoclave head; east-west orientation. - 34 -A simple tantalum sampling valve screwed onto the other end of the zirconium tube enabled samples to be taken using the internal pressure to expel the solution. A cylindrical copper jacket was a f f i xed to the zirconium tube between the autoclave head and the sampling valve. Cold water was passed through this jacket to cool the leach solution sampled. A Yellow Springs Instrument Co., Inc. model 71 A Thermistemp temperature controller, together with a Parr Instrument Co., I l l i n o i s , s t i r r e r type pressure reaction apparatus - model 4501 - was used to maintain the leach solution within + 1.5°C of the desired temperature. The model 4501 consisted of a 1500 watt output electric heater, Variac type W 10 M autotransformer (General Radio Co., Massachusetts), and a Bodine Electric Company, Chicago, a.c. motor to provide s t i r r i n g . Driving pulleys of different di ameters were used to vary the sti r r i n g speed. Solution sti r r i n g for gas, solid and liquid dispersion was provided by a Teflon-covered tantalum shaft f i t t e d with a Teflon impeller and Teflon propeller. The tantalum shaft passed through a Teflon cord-packed gland at the centre of the outer side of the autoclave head. The zip joint Teflon cord was supplied by the Crane Packing Company, I l l i n o i s . A pressure seal was kept by tightening this gland by means of a bolt whenever necessary. The st i r r i n g shaft was kept rotating about the central axis of the autoclave by two bearings at either end of . a housing above the packing gland. Cooling water was circulated in a copper pipe wrapped around the outer surface of the packing gland. The head was clamped to the autoclave body by means of a hardened steel s p l i t ring with six compression bolts. An excellent pressure seal - 35 -was obtained by coating the area of the steel pressure vessel in contact with both the Teflon plate of the head and the glass lining with Silicon^ Sealant, manufactured by Pow Corning Silicones, Vancouver. This procedure had the added advantage of ensuring no corrosion at the glass/metal boundary. Teflon tape was used on a l l threaded connections. A pressure gauge was situated on the dry, inlet gas line side of an autoclave to minimize possible corrosion. The Duraguage (Ashcroft, U.S.A.) was graduated i n 1 p.s.i. divisions from 0 to 100 p.s.i.g. During an experiment a low, continuous flow of sulphur dioxide gas was allowed to pass through the autoclave to prevent any possible gas line corrosion products from entering the autoclave. Dry, cold sulphur dioxide was non-corrosive, so any corrosion would only have occurred in the exit gas line where the sulphur dioxide was wet and hot. These corrosion products would be swept away into the bubblers by the gas stream. Needle valves situated at convenient points allowed flushing of the autoclave with nitrogen or sulphur dioxide and were used for fine pressure control, Figure 5. "Blank runs", with a l l reagents present except ore, were often carried out on both autoclaves to ensure that no iron was entering the leach solution as a consequence of corrosion of the apparatus. These runs showed that there was no detectable corrosion. An experiment performed in one autoclave was normally repeated in the other. No systematic difference in leach rates could be detected for those experiments performed in both autoclaves. - 37 -2.3 Reagents D i s t i l l e d water was used for a l l solutions. A l l chemicals used were reagent grade. The nitrogen gas was supplied by Canadian Liquid Air. Matheson of Canada supplied the anhydrous grade sulphur dioxide. Gases were used direct from their cylinders. 2.4 Experimental Procedure A weighed (1, 2, 3, or 5 g) sample of ore was placed i n a thoroughly cleaned autoclave body together with reagents and sufficient d i s t i l l e d water to make a solution volume of 1195 ml. The autoclave head and body were sealed together arid placed in the furnace. During rapid heating to the desired temperature the autoclave was purged with nitrogen at least six times. A purging cycle consisted of f i l l i n g the autoclave with nitrogen to 50 p.s.i.g., with solution s t i r r i n g to remove dissolved oxygen; then venting to a small positive pressure. When the desired solution temperature was reached the Variac was used as a fine temperature controller in conjunction with the Thermistemp. The stir r i n g was stopped and the autoclave was twice flushed out with sulphur dioxide to the maximum attainable pressure (about 35 p.s.i.g.). Stirring was then restarted and sulphur dioxide was reintroduced to the desired partial pressure, with a very low bleed of gas leaving the autoclave. A waiting period of at least thirty minutes was allowed in order to ensure equilibrium between the gaseous and liquid phases and between the furnace and the autoclave and i t s contents. - 38 -A 10 ml sample of the leach soluti on was then removed from the autoclave and discarded. This procedure flushed out the sampling tube and the 10 ml graduated cylinder used to measure sample volumes. A further 5 ml sample was taken and saved. When analyzed for iron, this sample was taken as a blank (zero time, zero dissolved iron) for subsequent leaching rate calculations. Further samples were taken in an identical manner at th i r t y minute intervals. A leaching run was at least four hours in duration, with no change in conditions except the reduction i n leach solution volume as a consequence of sampling. Water vapour loss as a consequence of the slow bleed of gas through the autoclave was negligible. Corrections were applied when calculating the leach rates to take into account the iron in solution lost by sampling. Apart from the st i r r i n g speed variation experiments, the st i r r e r was always rotated at 1100 r.p.m. At this speed a l l the ore was in suspension and sulphur dioxide was. rapidly absorbed by the solution. A disadvantage of this experimental arrangement was the limited range of sulphur dioxide partial pressure available. The maximum pressure available was the vapour pressure of the liquid sulphur dioxide in the cylinder. At the ambient laboratory temperature, this was about 36 p.s.i.g. maximum. Higher partial pressures could have been obtained by heating a suitable reservoir vessel containing liquid sulphur dioxide transferred from the cylinder, but this was not attempted. Alternatively, the aqueous sulphur dioxide concentrati on could have been increased by adding certain chemicals, soluble i n the leach solution, which could - 39 -increase the solubility. Chemicals which are capable of doing this and which can remain inert to the effects of sulphur dioxide at high temperatures include dimethyl formamide and dimethyl sulfoxide. This approach was also not attempted. Perchloric acid and sodium sulphite were used to adjust the pH of the aqueous sulphur dioxide solutions. The maximum sampling error was estimated to be + 3%, whilst the maximum error i n calculating the leach rate was estimated to be + 5%. 2.5 Analytical Techniques 2.5.1 Iron The concentration of iron i n the sample solutions was measured colpri-metrically by means of the orange-red ferrous complex of 1,10-ortho-phenanthroline. Hydroxylamine hydrochloride was used to reduce a l l iron in solution to the ferrous state. A sodium acetate-acetic acid buffer solution was used to keep the hydrogen ion concentration in the range 4-5 pH units. A composite reagent solution was made up containing 163.2 g l / anhydrous sodium acetate, 162 ml 1 * glacial acetic acid, 2 g 1 * , \ -1 hydroxylamine hydrochloride, and 0.3 g 1 1,10-orthophenanthroline. A few drops of concentrated hydrochloric acid was added to the d i s t i l l e d water used to dissolve the 1,10-orthophenanthroline. 90 ml of the composite reagent was then added to a suitable amount of the 5 ml sample of the leach solution. This was made up to 100 ml and after thorough mixing was l e f t for at le ast thirty minutes to allow the colour to completely develop. The optical density of each solution - 40 -was measured on a Beckman Model B Spectrophotometer using light of wavelength 512 my. The concentration of iron was read directly from a calibration curve prepared by using standard iron solutions. Since the presence of certain anions and cations interferes with the determination of iron by this method, i t was necessary to find the maximum permissible concentrations of perchloric acid, cupric perchlorate and sodium sulphite that could be present in a sample without i l l effect. This was accomplished by doping standard iron solutions with the above chemicals. The maximum permissible concentrations of these compounds were well above those used in subsequent experiments. 2.5.2 Copper The concentration of copper.in the sample solutions was measured with a Unicam Model SP90 Atomic Absorption Spectrophotometer. A copper lamp was used to provide light of wavelength 324.8 my. An air/acetylene flame was used to dissociate the copper in the sample into the atomic and excited states. The absorption of light by the ground state atoms was measured and knowing the absorption resulting from standard copper solutions i t was possible to measure the sample copper concentration. - 41 -3. RESULTS Graphs of the amount of iron leached into solution against time were linear for most runs with a l l ores. Under conditions of high aqueous sulphur dioxide concentration and high aci dity, some runs experienced an i n i t i a l increase in leach rate which lasted up to one hour unless the perchloric acid concentration was greater than 0.75 Molar. Under those conditions the rate of leaching only became constant after three hours. . The rates of leaching under the various conditions investigated have been tabulated and graphs have been drawn to i l l u s t r a t e important trends. 3.1 Goethite 3.1.1 Direct Dissolution i n Perchloric Acid Experiments were undertaken on 1 g samples at 110°C, under a nitrogen atmosphere. The measured rates at different concentrations of perchloric acid up to 6 Molar are given in Table V. 3.1.2 Reductive Dissolution in Aqueous Sulphur Dioxide Solutions 3.1.2.a Variation in Stirring Speed One gram samples were leached at 110°C, 25 p.s.i.a. S0„, 0.21 M HC10 p.s.i.a. 'so, Molar HC10, Molar 4 0 0.105 0.21 0.75 1.5 3.0 4.5 6.0 0 0 - - -1.48 . 1.64 2.35 2.39 2.75 2.96 3.10 3.43 2.93 3.52 3.68 3.25 3.48 4.35 5 0.067 - 4.0 4.5 4.9 5.1 5.5 - - - - -10 0.134 4.6 7.2 7.5 7.3 7.3 9.4 9.8 - - - -15 0.200 5.23 9.65 11.4 11.6 12.2 14.0 - - - -20 0.266 7.6 12.6 12.9 12.9 17.2 - - - -25 0.330 8.75 9.35 13.5 13.7 14.3 15.2 15.8 16.3 19.6 19.8 20.6 21.5 21.8 22.5 27.6 27.3 33.4 33.4 -30 0.395 10.40 16.15 17.9 18.0 22.0 - - -TABLE V. Leach rates of 1 g goethite samples at 110°C; mg Fe hr . - 43 -under three different s t i r r i n g conditions, 620, 1100, and 2000 r.p.m., Table VI. 3.1.2.b Variation in Sulphur Dioxide Partial Pressure Four series of experiments were undertaken at 110°C on 1 g samples where the sulphur dioxide par t i a l pressure was increased by 5 p.s.i.a. intervals up to a maximum of 30 p.s.i.a. Each series had a different perchloric acid concentration: series 1 - no added acid; 2 - 0.105 M HC104; 3 - 0.21 M HCIO^; 4 - 0.75 M HCIO^. The corresponding concentra-tions of molecular and ionic species present under these conditions are given i n Table VII. Leaching rates are in Table V. 3.1.2.C Variation in Hydrogen Ion Concentration In addition to the four different acid concentration series (3.1.2.b), higher perchloric acid concentration experiments were under-taken on 1 g samples at 110°C, 25 p.s.i.a. SO2, Table V. In these experiments i t was found that the leach rate increased with time, only becoming constant after about 3 hours. The leach rates quoted i n Table V were the f i n a l observed rates, as illustrated by the graph, Figure 6. 3.1.2.d Variation in Sample Weight In order to determine whether a homogeneous or heterogeneous reaction was rate-controlling, two series of experiments were conducted with different sample weights at 110°C, 10 p.s.i.a. SO2. In one series 1 g and 2 g samples were leached in 0.105 M HC10,. In the other series 1, - 44 -R.P.M. 620 1,100 2,000 15.2 Leach Rate, 14.6 -1 15.8 13.4 mg Fe hr 14.7 16.3 TABLE VI. Variation of leach rate with s t i r r e r speed. 1 g goethite samples; 110°C; 25 p.s.i.a. S0 2; 0.21 M HC10.. cso 2 Molar [H +], t S 02 aq^ Molar HCIO. Molar 4 0 0.105 0.21 0.75 1.5 3.0 4.5 6.0 0 [H+] 0 0.1050 0.2100 0.7500 1.5000 3.0000 4.5000 6.0000 0.067 [H+] 0.0114 0.1064 0.2107 - - - -tS°2 aq ] 0.0556 0.0656 0.0663 - - - - -0.134 tH+] 0.0166 0.1079 0.2115 0.7504 - - - -2 aq 0.1174 0.1311 0.1325 0.1336 - - - -0.200 ' [H+] 0.0205 0.1092 0.2122 0.7506 - - - -IS°2 aq ] 0.1795 0.1958 0.1978 0.1994 - - - -0.266 [H+] 0.0239 0.1105 0.2129 0.7508 - - - -[ S°2 aql 0.2421 0.2605 0.2631 0.2652 - - - -0.330 [H+] 0.0267 0.1118 0.2136 0.7510 1.5005 3.0003 4.5002 -C S°2 aq^ 0.3033 0.3232 0.3264 0.3291 0.3295 0.3297 0.3298 -0.395 [H+] 0.0293 0.1130 0.2143 0.7512 - - - -[ S°2 aq^ 0.3657 0.3870 0.3907 0.3938 - - - -TABLE VII. Concentrations of aqueous species at 110°C. - 46 -2 3 4 T IME , hours. FIGURE 6 . Amounts of iron leached from I g. goethite samples with time at I I O ° C , 25 RS. I .A.S0 2 different acidities. - 47 -2 and 3 g samples were leached in 0.21 M HCIO^, Table VIII. 3.1.2.e Variation in Temperature Experiments were carried out to enable an apparent activation energy to be calculated. One gram samples were leached i n 0.21 M HCIO^ containing a constant C c n (= 0.266 M). Leach rates were measured at 80, 90, 100, 110 and 120°C, Table IX. 3.1.2.f Effect of Ferrous Ions in Solution In"order to determine whether the dissolution reaction was auto-catalyzed by the ferrous ions formed by the reductive leaching of goethite, experiments were conducted at low and high acidities with a large i n i t i a l addition of ferrous ions. One gram samples were used at 110°C, 25 p.s.i.a. SO2, in both experiments. With no perchloric acid -3 present 335 mg of iron (5(10 ) M) was added as ferrous sulphate. With -3 3.0 M perchloric acid 134 mg of iron (2(10 ) M) was added as ferrous perchlorate. In the high acidity experiment the rate of leaching increased during the f i r s t three hours, becoming constant after this time. The leach rate quoted in Table XVIII was the f i n a l observed rate. This behaviour was also noticed for the identical experi ment with no i n i t i a l added ferrous ions. 3.1.3 Catalysis of the Reductive Dissolution by the Addition of Cupric  Ions Cupric ions were added as cupric perchlorate, Cu^lO^^^I^O. The samples of leach solution were analyzed for both iron and copper - 48 -HC10. Molar 4 0.105 0.21 Sample Weight 1 g 2 g 1 g 2 g 3 g Leach Rate, 7.2 7.3 15.9 24.2 14.8 mg Fe hr 7.5 7.3 15.4 24.2 TABLE VIII. V a r i a t i o n of leach rate with sample weight. Goethite samples; 110°C; 10 p . s . i . a . SO . Temperature °C 80 90 100 110 120 p s o 2 P - 8 - 1 - * - 12.2 14.7 17.7 20.0 22.4 [SO- ] Molar L 2 aq 0.2612 0.2621 0.2627 0.2631 0.2635 [H+] Molar 0.2148 0.2139 0.2133 0.2129 0.2125 [HS03~] Molar 0.0048 0.0039 0.0033 0.0029 0.0025 Leach Rate 2.25 4.13 8.05 12.9 25.7 mg Fe hr * 2.38 4.28 8.20 12.9 25.7 TABLE IX. V a r i a t i o n of leach rate with temperature. 1 g goethite samples; C = 0.266 M; 0.21 M HC10 - 49 -concentrations. In a l l experiments some copper was lost. Unless otherwise stated, the loss never exceeded 10%, and this loss was not considered when calculating the leach rate. Unlike the uncatalyzed experiments, the copper catalyzed experiments resulted in the precipitation of a very fine black powder on the Teflon and glass surfaces below the leach solution level. The maximum concentration of copper added was kept as low as possible for two reasons. F i r s t l y , too high a concentratiqn of copper interfered with the colorimetric method used for the estimation of iron. Secondly, i t was necessary to show that copper did i n fact behave as a catalyst and not as a reagent being consumed by the dissolution reaction. ^ 3.1.3.a Variation in Concentration of Cupric Ions Experiments were undertaken where the concentration of copper was varied, keeping a l l other conditions constant. At 110°C, 25 p.s.i.a. S0 2, 0.21 M HCIO^, the cupric ion concentration -3 -2 was varied over an order of magnitude, from 10 M to 10 M. At the same temperature and sulphur dioxide partial pressure, but with no added acid, a smaller cupric ion concentration range was investigated. Table X. -3 -3 At 90°C, no added acid, cupric ion concentrations of 10 , 2(10 ), -3 and 3(10 ) M were investigated at three sulphur dioxide partial pressures: 14.7, 24.7, and 34.7 p.s.i.a. Table XI. - 50 -HC10. Molar 4 0.21 0 [Cu 2 +] Molar Leach Rate mg Fe h r ~ l 0 15.2 15.8 16.3 8.75 9.35 10"3 - 129.0 2(10 - 3) 43.9 235.0 4(10 _ 3) 71.9 -6(10 - 3) 96.1 — 8(10 _ 3) 121.0 -10" 2 147.0 -TABLE X. Variation of leach rate with cupric ion concentration at 110°C. 1 g goethite samples; 25 p.s.i.a. SO„. p-2 p.s.i.a. [Cu 2 +] Molar 0 io" 3 2(10 3) 3(10 3) 14.7 2.6 48.5 52.0 54.7 73.2 76.5 96.0 108.0 24.7 - 39.0 52.3 80.0 84.6 -34.7 - 48.3 49.0 84.0 91.4 -TABLE XI. Variation of leach rate with cupric ion concentration at and sulphur dioxide partial pressure at 9 0PC. 1 g goethite samples; mg Fe hr ^ . - 51 -3.1.3.b Variation In Partial Pressure of Sulphur Dioxide The effect of increasing the sulphur dioxide partial pressure was investigated at 90°C, no added acid, for cupric ion concentrations of - 3 - 3 1 0 and 2 ( 1 0 ) M. Experiments were performed at 14.7, 24.7, and 34.7 p.s.i.a. SO2. Table XI. The concentrations of species in solution at 90°C are given in Table XII. 3.1.3.C Variation in Hydrogen Ion Concentration At 110°C, 25 p.s.i.a. SO2, the added cupric ion concentration was -3 2 ( 1 0 ) M, and leaching rates were measured at no added acid, 0.105 M HC10, and 0.21 M HC10.. Table XIII. 4 4 At 90°C, 14.7 p.s.i.a. SO2, the added cupric ion concentration was _3 1 0 M and a much wider pH range was studied. At high pH (> 2 ) sample analysis showed that copper was continually lost from solution during a run. Metallic copper was precipitated from soluton and covered the Teflon and glass surfaces under the solution level with a lustrous bronze-coloured film. Graphs of amount of iron leached into solution against time showed a dissolution rate which decreased with time. The rates quoted in Table XIV for these high pH experiments (marked with an asterisk) were i n i t i a l leach rates, taken when the loss in copper from solution was s t i l l less than 1 0 % . They were only considered to be minimum possible rates. 3.1.3.d Variation in Sample Weight Experiments with 1, 2, and 3 g samples were conducted at 110°C, 25 p.s.i.a. S 0 „ , 0 . 2 1 M HC10., Table XV. At this pH and temperature, the 'so p.s.i.a. s Molar [H +], 2 aq Molar .-• Na S 0 M o i a r HC10. Molar 2 3 4 0.10 0.05 0 0.105 0.21 14.7 0.266 TH+] 0.0010 0.0049 0.0276 0.1124 0.2139 CS02 aq ] 0.0650 0.1611 0.2384 0.2586 0.2621 24.7 0.441 [H+] - - 0.0360 - -tS°2 aq* - - 0.4050 • - -34.7 0.609 [H+] - - 0.0426 - -[ S°2 aql - 0.5664 - -TABLE XII. Concentrations of aqueous species at 90°C. - 53 -HC10. Molar 4 0.21 0.105 0 Leach Rate mg Fe h r - l 43.9 76.2 235.0 TABLE XIII. Variation of leach rate with pH at 110°C with a cupric ion -3 concentration of 2(10 ) M. 1 g goethite samples; 25 p.s.i.a. S0 0. HC10. 4 Molar Na 2S0 3 Molar PH -log 1 0[H +] Leach Rate, mg Fe hr * 0.21 0 0.670 7.2 7.7 0.105 0 0.949 12.1 13.1 ; 0 0 1.559 48.5 52.0 54.7 0 0.05 2.310 * 74.1 87.0* 0 0.10 3.000 * 76.1 86.7* I n i t i a l rates; Cu° precipitated during run. TABLE XIV. Variation of leach rate with pH at 90°C with a cupric ion -3 concentration of 10 M. 1 g goethite samples; 14.7 p.s.i.a, SO - 54 -[Cu 2 +] Molar Sample Weight 1 g 2 g 3 g 2(10" 3) Leach Rate, mg Fe hr* 43.9 60.8 78.5 0 *1 15.8 *2 31.6 *2 47.4 D i f f e r -ence 28.1 29.2 31.1 *1 Average of three rates. *2 Rates calculated from known heterogeneous co n t r o l of uncatalyzed reaction. TABLE XV. V a r i a t i o n of leach rate with sample weight at 110°C with -3 a cupric ion concentration of 2(10 ) M. Goethite samples; 25 p . s . i . a . S0 0; 0.21 M HC10.. uncatalyzed rate was appreciable. Als o included in Table XV were the calculated uncatalyzed rates for 1, 2, and 3 g samples based on the averaged measured 1 g rates, assuming heterogeneous control, as indicated by the results i n Table VIII. The differences between the measured catalyzed rates and the calculated uncatalyzed rates were also tabulated. At 90°C, 14.7 p.s.i.a. SC^, experiments were undertaken with 1, 2, an4 5 g samples with solutions at three different pH values. One series -3 -3 had an added cupric ion concentration of 10 M, the other 2(10 ) M. Table XVI. 3.1.3. e Variation i n Temperature Experiments were conducted at 80, 90, 100 and 110°C at constant C n (= 0.266 M) and almost constant hydrogen ion concentration. Table 2 XVII. 3.1.4. Effect on the Reductive Dissolution of the Addition of  Other Cations The cations that were individually added to a typical leach solution f u l f i l l e d one or both of these conditions: 1. They had, like copper, two oxidation states whose redox potential was close to that of the 2- 2+ 3+ S02/S0^ or Fe /Fe couple. 2. They had, like copper cations, the a b i l i t y to form complex anions with bisulphite or sulphite ions, in one or more oxidation states. whenever possible, the cation was added as the perchlorate. Experiments were conducted at 110°C, 25 p.s.i.a. S0_, no added acid HC10, Molar 4 0.21 0 0 Na 2S0 3 0 0 0.05 [Cu 2 +] Molar Sample Weight 1 g 2 g 1 g 2 g 5 g 1 g 2 g IO" 3 Leach Rate, mg Fe hr"* 7.2 7.7 11.1 48.5 52.0 54.7 73.7 78.7 118.0 * 74.1 87.0* * 143.6 2(10~3) - - 73.2 76.5 95.5 - - -I n i t i a l rates - Cu° precipitated during run. TABLE XVI. Variation of leach rate with sample weight at 90°C with cupric ion concentrations -3 -3 of 10 and 2(10 ) M. Goethite samples, 14.7 p.s.i.a. S0_. - 57 -Temperature °C 80 90 100 110 pso 2 P- 3- 1- 3' 12.2 14.7 17.7 20.0 HC10. Molar 4 0 0 0.0046 0.0078 Na 2S0 3 Molar 0.0034 0 0 0 [H +] Molar 0.0271 0.0276 0.0280 0.0282 [S0_ ] Molar 2 aq 0.2321 0.2384 0.2426 0.2456 Leach Rate, mg Fe hr 28.1 29.8 48.5 52.0 54.7 79.9 142.0 TABLE XVII. V a r i a t i o n of leach rate with temperature with a -3 cupric ion concentration of 10 M. 1 g goethite samples; C . - 0.266 M. - 58 -(unless stated). Table XVIII. 3.1.5 Effect on the Reductive Dissolution of the Addition of  Oxalate Anions Oxalate anions were added either as uranyl oxalate, U0o.C„0,.3H_0, 2 2 4 2 or as oxalic acid, H ^ O ^ ^ O . Experiments at 110°C, 25 p.s.i.a. S0 2 > were continued for 5 1/2 hours. Table XIX. 3.2 Magnetite 3.2.1 Reductive Dissolution in Aqueous Sulphur Dioxide Solutions A 1 g sample of magnetite was leached at 110°C, 25 p.s.i.a. S0 2, 0.21 M HC104? Table XX. 3.2.2 Catalysis of the Reductive Dissolution by the Addition of  Cupric Ions A 1 g sample of magnetite was leached at 90°C, 14.7 p.s.i.a. S0 2 > -3 with an i n i t i a l cupric ion concentration of 10 M. Table XX. 3.3 Hematite 3.3.1 Reductive Dissolution in Aqueous Sulphur Dioxide Solutions A 1 g sample of hematite was leached at 110°C, 25 p.s.i.a. S02-Table XXI. 3.3.2 Catalysis of the Reductive Pis solution by the Addition of  Cupric Ions A 1 g sample of hematite was leached at 110°C, 25 p.s.i.a. S0 2 > with -3 an i n i t i a l cupric ion concentration of 10 M. Table XXI. Cation Compound Type Concn. Molar Rate . -1 mg hr Comments - - - - 8.75 9.35 No Additions Mn 1 1 Perchlorate 1. 10"3 Falls from 7.0 to zero pH rose to 3.5 units Sn 1 1 Chloride 1. 5(10"3) 11.7 +0.21 M HC10A to acidify V (vv) Sodium Vanadate 1. 5(10"3) 9.7 Co 1 1 Perchlorate 1,2. 10"3 7.05 Co" ^unstable C r 1 1 1 Perchlorate 1,2. 5(10"3) 4.4 C r " unstable Hg" Perchlorate Solution 1,2. 3(10"2) 15.0 Hg° formed Ag 1 Sulphate 2. 10"3 7.1 Ag* ^"unstable Cd 1 1 Sulphate 2. 5(lO" 3) 7.55 Cd 1 unstable T I 1 Nitrate 2. io" 3 7.3 (uV I) Uranyl Nitrate 1?2. 5(10- 3) 9.2 F e 1 1 Ferrous Sulphate 5(10- 3) 10.4 Autocatalysis Experiment - <- - - 27.6 27.3 3.0 M HCIO^; F e 1 1 Ferrous Perchlorate - 2(IO- 3) 29.6 Autocatalysis Experiment TABLE XVIII. Effect on leach rate of the addition of various cations. 1 g goethite samples; 110°C; 25 p.s.i.a. 60.. - 60 -Compound Concn. Molar Rate . -1 mg hr Comments - - 8.75 9.35 'No Additions H2C204.2H 20 IO" 3 28.8 H2C204.2H 20 5(10~3) 68.0 U0 2C 20 4.3H 20 5(10"3) 26.8 U0 2C 20 4.3H 20 5(10"3) 41.1 0.14 N H.SO. 2 4 TABLE XIX. Effect on leach rate of the addition of oxalate anions. 1 g goethite samples; 110°C; 25 p.s.i.a. SO„. - 61 -Temp. °C 'so p . s . i . a . HC10. 4 Molar [Cu 2 +] Molar Leach Rate mg Fe hr-''-Magnetite Goethite 110 25.0 0.21 0 19.3 15.2 , 15.8 16.3 90 14.7 0 0.001 111.7 48.5 52.0 54.7 TABLE XX. Comparison of the leach rates of magnetite and goethite under i d e n t i c a l conditions. 1 g samples. Temp. °C % 2 p.s.i.a. HC10. 4 Molar [ C u 2 + ] Molar Leach Rate mg Fe hr Hematite Geothite 110 25 0 0 0.90 8.75 9.35 110 25 0 0.00! 32.7 129.0 TABLE XXI. Comparison of the leach rates of hematite and goethite under i d e n t i c a l conditions. 1 g samples. - 62 -4. DISCUSSION The constant leach rates observed in a l l low acid (^0.75 M [H +]) experiments showed that anisotropic leach ing was not observable. The constant leach rates also showed that autocatalysis by ferrous ions did not occur to any appreciable extent. Since the dissolution rate of the hematite was so low, no correction was applied to the leach rates of goethite samples to take into account the hematite present in that ore. Because of the absence of high temperature data on the activity coefficients of perchloric acid solutions, and aqueous sulphur dioxide solutions, concentrations were used throughout the discussion. 4.1 Goethite 4.1.1 Direct Dissolution in Perchloric Acid The leach rates have been plotted against hydrogen ion concentration, assuming 100% ionization, in Figure 7. The form of a curve that best f i t s the experimental results is an adsorption isotherm i n hydrogen ion concentration: constant^ [H+] rate = — (Langmuir) 1 + constant,, [H ] 2 3 HYDROGEN ION 4 5 6 CONCENTRATION, Molar. FIGURE 7. Variation of the direct perchloric acid leach rate with hydrogen ion concentration. Ig. goethite samples; 110°C. - 6 4 -Postulated mechanism: 1. Protonation of a surface site: /FeOOH + H + — ^ /Fe(OH) ® ^ I 61 °2 la . Desorption of an hydroxy-ferric species: k s /Fe(OH)® F e ( 0 H ) 2 + a q (28) 92 / denotes the surface of the oxide. K and k are equilibrium and rate constants respectively. ^ denotes a surface charge 6^ and 6 2 are the fractions of the total active surface area occupied by FeOOH and Fe(OH)2 species respectively, i.e., 6 + 6 2 = 1 6 2 = K 1tH +]6 1 e 1 1 + K 1tH +] d[Fe ] Rate of Dissolution = — -ac* = k.,6,, dt 1 I d t 1+ K 1[H +] The constants k^ and were evaluated by substituting the averaged leach rates measured at two different hydrogen ion concentrations into the above equation and solving simultaneously. Using the averaged rates - 65 -at 0.75 M and 6.0 M [H ],the constants had the following values: k = 4.584; ^ = 0.688 i - 3.153[H+] Theoretical rate = 1 — ' — — 1 + 0.688[H ] This curve was also plotted on Figure 7, and fi t t e d the experimental data at other hydrogen ion concentrations. The good correlation between the observed and theoretical rates j u s t i f i e d the assumption in the postulated mechanism that perchlorate anions did not affect the leach rate, even i f they adsorbed on the surface. This was also j u s t i f i e d by their known low complexing power for f e r r i c ions (37). 4.1.2 Reductive Dissolution i n Aqueous Sulphur Dioxide . Solutions The variable s t i r r i n g speed experiments (Table VI) indicated that, under those conditions, the rate of the dissolution reaction was not controlled by the diffusion of species to or away from the surface of the oxide. If the reaction were diffusion controlled, then a large variation in leaching rate with s t i r r i n g speed should have been observed. The variable sample weight experiments (Table VIII) showed that the dissolution reaction was heterogeneously controlled, i.e. occurred at the oxide surface. If the reaction were homogeneously controlled, i.e. the slow step was one which occurred in solution, then the rate with 2 or 3 g samples should have been equal to the rate with a 1 g sample. Since - 66 -the rates with 2 and 3 g samples were twice and t h r i c e the rate with a 1 g sample r e s p e c t i v e l y , the slow step was one which occurred e x c l u s i v e l y at the oxide surface. The apparent a c t i v a t i o n energy was c a l c u l a t e d from the experimental data at 80, 90, 100, 110 and 120°C (Table IX) using the Arrhenius equation: EA Rate = K exp(- — ) where K = a pre-exponential faator. R =• the gas constant, 1.987 cal.°A o T = temperature, A c - 1 E = apparent a c t i v a t i o n energy. A ^ EA 1 Rearranging, l o g 1 Q ( r a t e ) &< - 2 3 0 2 5 R ' (j) • A 8 r aPh, o f log 1 Q!?rate) against ^ was a good s t r a i g h t l i n e (obtained by l i n e a r regression EA a n a l y s i s ) , the gradient being equal to - ^ 3026R ( F i 8 u r e 8 ) • Hence E. = 16.38 + 0.75 k c a l mole - 1. A — This value was w e l l above the normal maximum-observed apparent a c t i v a t i o n energy for a d i f f u s i o n - c o n t r o l l e d reaction, 6.5 k c a l mole ^. The experiments performed with additions of ferrous ion had s i m i l a r leach rates to i d e n t i c a l experiments with no i n i t i a l ferrous ions (Table XVIII). This was a d d i t i o n a l proof that no autocatalysis occurred. A graph of leach rate against aqueous sulphur dioxide concentration, Figure 9, showed that the rate was f i r s t order i n [SO ] under the JL 3 C J - 67 -1.6 LU < CC X o < O O 0.8h 0.6 0.4 0.2 2.5 SLOPE = - 3 , 5 8 0 E . = 16.38 kcal .mole"" 1 A 2.6 2.7 2.8 •^(XIO 3), °A' FIGURE 8. Arrhenius plot of log, 0 (reductive leach rate,) against the reciprocal of the absolute temperature. ON CO 0.1 0.2 0.3 0.4 AQUEOUS SULPHUR DIOXIDE CONCENTRATION, Molar. FIGURE 9. Reductive leach rates plotted against aqueous sulphur dioxide concentration. Ig. goethite samples; IIO°C. - 69 -conditions studied. In addition, the leach rate increased for a given [S0„ ] as the perchloric acid concentration was increased. The I aq incremental increases i n leach rate as [H+] increased at constant [SO. ] were greater at higher [SO- ]. The increase in leach rate with [H +], at con stant sulphur dioxide partial pressure, was typified by the graph of the results at C c- = 2 0.33 M. Figure 10. The increase in leach rate with increasing [H+] f e l l until the rate became f i r s t order in [H ] at ^> 0.75 M [H ]. However, at hydrogen ion concentrations greater than this value the leaching behaviour was anomalous (Figure 6), These curved plots were not a consequence of autocatalysis since _3 the experiment at 110°C, 25 p.s.i.a. S0 2, with 3,0 M HCIQ^ and 2(10 ) M added ferrous ion was similarly curved, and had a f i n a l leach rate comparable with that of an identical experiment wit h no i n i t i a l ferrous ion addition. Table XVIII. The curved plots might'- be the results of dissolution anisotropy at these high hydrogen ion concentration experiments. It was considered unlikely that the curved plots could be a consequence of a slow rate of attainment of the equilibrium concentration of SO- 'in solution 2 aq at the high [H+] which prevailed. An experiment was performed on a 1 g goethite sample at 110°C, 25 p.s.i.a. S0 2, 3.0 M HCIO^ to investigate this possibility. After four hours leaching, the leach rate had become constant at 26.4 mg hr *. The autoclave and contents were cooled to room temperature and the sulphur dioxide was expelled from the autoclave under continuous sti r r i n g and flushing with nitrogen. When i t was considered that sufficient time had elapsed for the 2 3 ION CONCENTRATION, MOLAR. FIGURE 10. concentration. HYDROGEN Reductive leach rates at 25PS.I.A.S02 plotted against hydrogen ion g. goethite samples; IIO°C. - 71 -expulsion of almost ( i f not) a l l the sulphur dioxide, the autoclave and contents were again heated to 110°C. A f t e r readmission of sulphur dioxide to a p a r t i a l pressure of 25 p . s . i . a . , the leaching experiment was continued. No curving was observed, and f o r the three hours of the continued experiment the leach rate was constant at 28.0 mg hr *. The s o l u b i l i t y of sulphur dioxide i n the high hydrogen ion concentration experiments was assumed to be unaffected by the hydrogen ions due to a lack of data to the contrary. Because of a l l these u n c e r t a i n t i e s , the high p e r c h l o r i c aci d concentration experimental data was not taken i n t o consideration i n the development of a leaching mechanism. A l l the s t r a i g h t l i n e s i n Figures 9 and 10 were obtained by l i n e a r regression analysis of the relevant experimental r e s u l t s . In p o s t u l a t i ng the following mechanism, a l l the above f a c t s , the theory reviewed e a r l i e r , and the observed hydrogen ion adsorption isotherm f o r d i r e c t p e r c h l o r i c acid leaching, have been taken i n t o consideration. Postulated Mechanism: 1. Protonation of a surface s i t e : /FeOOH + H+ K. 1 e e l 2. Adsorption of an aqueous sulphur dioxide molecule on a protonated surface s i t e : K, "2 e 2 aq ^ - 72 -3. Second protonation of a protonated site already containing an adsorbed sulphur dioxide molecule: K /Fe(0H).S0® + H + — / F e ( O H ) ^HSO.8® 2 2 z 2 63 64 The following leaching reactions are possible i f solution of iron from a l l surface sites is considered: l a . / F e ( O H 2 ) ® - % F e< 0 H>2(aq) 2a. /Fe(OH)2S02® —^> FeS0 3* a q ) + H20 3a. /Fe(OH) 2HS0 2 9® —h> FeHSC^2/^ + H20 94 4. Reduction of fe r r i c ions in solution: Ferrous ions were formed in solution by the reductive action of aqueous sulphur dioxide or anions derived from i t s dissociation (4). No fer r i c ions were detected in leach solutions analyzed for both f e r r i c and ferrous ions independently. Consequently the reduction in solution was very fast. e. + e. + e. + e. = 1 1 2 3 4 e 2 = K1[H+]61 e 3 = K 1 K 2 [ H + ] [ S O 2 A Q ] B L 6 4 = K 1 K 2 K 3 [ H + ] 2 [ S 0 2 A Q ] 6 1 /Fe(OH)2HS02 sites were expected to be highly transitory and thus 6, ->- 0. 4 Hence 6 1 1 + K^[H+] + K 1K 2[H +][S0 2 • ] d[Fe ] Rate of dissolution = r 3 ^ — = k .e„ + k_e o + k_6, dt 1 2 2 3 3 4 d t F e a q ] _ K 1 [ H + ] ( k 1 - r k 2 K 2[S0 2 a q ] + k 3 K 2 K 3 [ H + ] [ S 0 2 a q 3 ) d t 1 + K^H"1"] + K 1 K 2 [ H +][S0 2 a q ] The groups of constants were evaluated by substituting the leaching rates (as predicted by the linear regression analysis) at the low [H+] experiments with constant [S0"2 ] = 0.30 M, into the above equation and solving simultaneously. When the values of and k^ were taken to be those of the direct perchloric acid leaching mechanism, the theoretical rate became: [H+](3.153 + 2705.67[SO ] + 984.53[H+][SO 1) = — 23 _ SS t h * 1 + 0.688[H+] + 183.99[H+][S0„ ] Z aq Although this theoretical expression accurately predicted the experimentally observed leach rates at [S00 ] =0.30 M, i t failed - 74 -when applied to other hydrogen ion and aqueous sulphur dioxide concentrations. The equilibrium constant i n the direct perchloric acid leaching, K^ , should have been expressed in ac t i v i t i e s : a/Fe(OH)2® K. 1 a ^ . a/FeOOH Y/Fe(QH)2®[/Fe(OH)2e] ^ H + ] ' V F e O O H t / F e O O H ] The presence of aqueous sulphur dioxide, bisulphite anions and other sulphur-containing species in solution might result i n their adsorption without significant reaction on unprotonated /FeOOH surface sites. This would reduce the activity of such surface sites i n comp arison with the direct perchloric acid leaching. The activity of the protonated/Fe (OH ) 2 sites might change for a similar reason. Consequently the value of in the reductive dissolution experiments might be expected to have an apparently different value from that in the direct leaching investigation. Since was related to a fixed concentration of the protonated sites (1 g sample of ore) i t also would be expected to have a different apparent value. The theoretical expression was re-solved for a l l groups of constants, using results from the linear regression analysis at [S0 2 ] = 0.30 M, low [H +], and [SO ] = 0.1994 M, [H+] = 0.7506 M. Analysis showed that the denominator term K1K_[H+][S0„ 1 »• zero, i.e. 6. -+ zero. 1 2. l aq j This was expected since no adsorption isotherm in [S0„ ] was observed. - 75 -Rate [H +](195.19 + 2063.27[SO_ ] + 984.50[H +][SO. ] ) th. 1 + 55.185[H +] with K 55.185 3.537 37.388 17.840 This expression was pl o t t e d against [SO 2 aq ] on the graph, Figure 11. I t described the l i n e a r regression analyses of the experimental r e s u l t s at the four d i f f e r e n t low p e r c h l o r i c a c i d concentrations f a i r l y accurately. With constant C = 0.33 M, the t h e o r e t i c a l rate expression was pl o t t e d against [H +] on the graph, Figure 12. At high p e r c h l o r i c acid concentrations (> 0.75 M), the t h e o r e t i c a l expression predicted greater leach rates than those experimentally measured under somewhat obscure conditions. Although many other d i f f e r e n t reaction mechanisms were t r i e d , none was even remotely successful i n accounting f o r the experimental r e s u l t s . Included i n the mechanisms t r i e d was one based on the adsorption of an aqueous sulphur dioxide molecule followed by protonation. P h y s i c a l l y , the postulated mechanism seemed d i f f i c u l t to comprehend since i t required that a ne u t r a l , though solvated and possibly p o l a r i z e d , sulphur dioxide molecule should adsorb on or adjacent to a p o s i t i v e l y charged s i t e . However, neutral aqueous sulphur dioxide molecules have been found to leach manganese dioxide (37,38), although a mechanism was not proposed. 20 Li_ d» E LU < 10 or x u < LU EXPERIMENTAL FI6.9. THEORETICAL ON 0 0.1 0.2 0.3 0.4 AQUEOUS SULPHUR DIOXIDE CONCENTRATION, Molar. FIGURE II. Comparison of the experimental and theoretical leach rates variation with aqueous sulphur dioxide concentration. 0 •^ 1 EXPERIMENTAL - THEORETICAL _L I 2 3 4 5 HYDROGEN ION CONCENTRATION, MOLAR. FIGURE 12. Comparison of the experimental and theoretical leach rates—variation with hydrogen ion concentration at 25 RS.I.A. S0 2 . - 78 -The b i s u l p h i t e anion was ruled out as a leaching species on account of the high rates of d i s s o l u t i o n measured when i t s concentration was extremely small. However, i n the postulated t h e o r e t i c a l rate [H+] [HSOJ] expression [SC^ ] could be replaced without e f f e c t by , where K = f i r s t d i s s o c i a t i o n constant of aqueous sulphur dioxide. Such an expression required that an extra protonation step occurred i n the reaction mechanism, followed or preceded by the adsorption of a b i s u l p h i t e anion instead of a sulphur dioxide molecule. This was considered highly u n l i k e l y i n view of the apparent lack of such an occurrence even at high [H +] i n the d i r e c t p e r c h l o r i c acid leaching. 4.1.3 C a t a l y s i s of the Reductive D i s s o l u t i o n by the Addition of Cupric Ions Cupric ions are reduced i n aqueous sulphur dioxide: 2 C u 2 + + S0„ + 2H„0 — ^ 2Cu + + SO. 2" + 4H + 2 aq 2 ^ — 4 The formation of cuprous ions i s favoured by a high concentration 2- + of SO^ , and low concentrati ons of SO^ and H (high pH). Cuprous ions are unstable i n aqueous s o l u t i o n unless complexed: 2Cu + ^ C u 2 + + Cu°+ At high pH, high [S0» ], cuprous ions may be reduced to copper metal: - 79 -2Cu + + S0„ + 2Ho0 — ^ 2Cu°+ + SO 2 _ + 4H + 2 aq 2 *;— 4 Copper metal dissolves slowly in aqueous sulphur dioxide with the formation of insoluble, black cupric sulphide (75): 2Cu° + 2S0„ — > CuS + Cu + SO. 2 aq 4 It was concluded that the fine black powder found adhering to the Teflon and glass surfaces below solution level in copper catalyzed experiments was cupric sulphide, formed in this manner. - 2-Cuprous ions are strongly complexed by HSO^ and SO^ anions (36), forming species of the form: Cu(HS0 3)^ n _ 1 ) - and Cu^O^ 2 1" - 3^ -. It was thought that the formation of these complexes stabilized the cuprous ions, preventing excessive disproportionation and reduction in low pH leach solutions. At high pH and possibly at high [S0_ ] (15), the cuprous complexes were unable to prevent reduction to metal. Since the system was so complex, i t was d i f f i c u l t to obtain a meaningful activation energy for the reduction of cupric ions to cuprous ions by aqueous sulphur dioxide. However, a method of recovering copper as cuprous chloride from leach solutions containing cupric ions is of relevance. The "Hunt and Douglas" process (circa 1881) involved reducing a solution of cupric chloride with aqueous sulphur dioxide and collecting the cuprous chloride precipitated from solution: 2CuCl„ + S0„ + 2H.0 — > 2CuCl4- + 2HC1 + H„SO, 2 2 aq 2 2 4 - 80 -It was noted that the reaction was slow at room temperature but very rapid between 80 and 90°C (76). This implie d that the activation energy for the reduction of cupric ions to cuprous ions by aqueous sulphur dioxide was large. At 110°C the leach rate was found to have a f i r s t order dependence on the i n i t i a l concentration of cupric ions. Table X; Figure 13, At 90°C, 14.7 p.s.i.a. SO2, a similar f i r s t order dependence was observed 2+ -3 for [Cu ] > 10 M. Table XI; Figure 14. The straight lines were obtained by linear regression analysis of the experi mental results. The pH of the leach solution was found to have a great effect on the rate of dissolution. Tables XIII, XIV. At 110°C the leach rate became so great at high pH that i t was found necessary to investigate the cupric catalyzed system at 90°C. Also, the runs at 90°C did not suffer from having a high rate of uncatalyzed dissolution. The observation that at 110°C, C_n =0.33 M, [Cu 2 +] = 2(10"3) M, + Z [H ] = 0.0267 M, a l l the iron was in solution within three hours proved the existence of a fast catalytic mechanism. A graph of pH against leach rate for the experiments at 90°C, 2+ - 3 [Cu ] = 10 M, C = 0.266 M showed how the leach rate rose dramatically as the pH increased from 1 to 2 units. Figure 15. The constant leach rate observed between pH 2 and 3 was probably genuine, but the rates measured might have been low due to the precipitation of copper metal from solution. The large decrease in [S0„ ] probably had some effect. The results obtained from the experiments varying the partial pressure of sulphur dioxide were complicated by the fact th at there - 81 -200-£ < or x o < UJ o 100 0 2 4 6 8 10 CUPRIC ION CONCENTRATION (xiO3), MOLAR. FIGURE 13. Variation of leach rate with cupric ion concentration at 110°C, 25 PS.I.A. S0 2 . Ig.goethite samples. - 82 -CUPRIC ION CONCENTRATION (xiO3), MOLAR. FIGURE 14. Variation of leach rate with cupric ion concentration at 90°C, 14.7 PS.I.A. S0 2 . Ig. goethite samples. - 83 -- 84 -was a slight pH variation (from 1.371 at 34.7 p.s.i.a. to 1.559 at 14.7 p.s.i.a.). This was i n the range where there was a maximum change in leach rate with hydrogen ion concentration. From Figure 15, estimates 2+ -3 were made of the rate of leaching with [Cu ] = 10 M, 14.7 p.s.i.a. S0 2, but with [H+] = 0.0360 M (40.5 mg hr" 1) and [H+] = 0.0426 M (34.0 mg hr * ) , corresponding to the pH of the solutions with sulphur dioxide partial pressures of 24.7 and 34.7 p.s.i.a. respectively. Since these values were lower than the averaged experimentally obtained rates at the same pH but with higher sulphur dioxide par t i a l pressure, i t was concluded that the higher the sulphur di oxide partial press ure the greater the leach rate. The increase in the rate of dissolution may - 2-have been due to the higher {SO ] and/or the higher [HSO ] and [SO ]. The variable sample weight experiments at 110 °C, Table XV, indicated that the catalyzed and uncatalyzed dis solution reactions were able to occur simultaneously. At the high [H +l and [SO„ 1 of these J 2 aq experiments the catalyzed reaction was homogeneously controlled. At 90°C, Table XVI, the catalyzed rate was also homogeneously controlled at high [H +]. As the pH was increased the control became mixed, and the rate-controlling step tended to become more heterogeneous for low sample weights. However, the experiments on 1, 2, and 5 g + 2+ -3 samples at [H ] = 0.0276 M, [Cu ] = 10 M showed that the control tended to become homogeneous for sample weights greater than or equal to 5 g. Figure 16. The apparent activation energy was calculated from the results in Table XVII using the Arrhenius equation. A value of 13.94 + 1.00 kcal mole * was obtained from the slope of a line obtained by linear regression - 85 -2 SAMPLE 3 4 WEIGHT, g. FIGURE 16. Variation of leach rate with goethite sample weight at 90°C; I4.7PS.I.A.S02; I0"3 M [< - 86 -analysis of the experimental data. Figure 17. This . v a l u e of the apparent activation energy indicated that the diffusion of species i n solution was not the rate-controlling step. Because of the lack of data on the many equilibria present in the copper catalyzed system and the apparent complexity of the experimental results, a quantitative analysis was impossible. However, the following qualitative mechanism was postulated, bearing in mind the experimental trends and their similarity to the oxidative, f e r r i c ion promoted leaching of uranium dioxide (43). The role of aqueous sulphur dioxide was four-fold: 1. It leached the goethite by the same mechanism as in the uncatalyzed system. 2. It reduced cupric ions to cuprous ions, which were considered to be the "catalyzed" leaching species. 3. By dissociating, i t provided bisulphite and/or sulphite anions which formed stable anionic complexes with the cuprous ions. In the form of such an anionic complex, a cuprous ion adsorbed on the surface of the oxide and underwent an electron-exchange reaction with a surface Fe*** species. The Fe** and Cu** species produced desorbed, forming ferrous and cupric ions in solution. The cupric ion was again reduced by the aqueous sulphur dioxide and this behaviour j u s t i f i e d the use of the term "catalysis". 4. It redissolved part of the copper metal precipitated from solution: 2+ + SO 2-2 aq 4 - 87 -2.6 2.4 1x1 < 2.2L or x o 2 2.0 Q Id </> 1.8 >-< o 1.6 o ef 9 1.4, 1.2 1.0 2.5 SLOPE = - 3 , 0 5 7 E A « I 3 . 9 4 kcal . mole" 2.8 2.6 2.7 - f (xiO3), 0 A"' FIGURE 17. Arrhenius plot of log | 0 (catalysed leach rate) against the reciprocal of the absolute temperature. - 88 -The leach rate was therefore directly proportional to the concentration of cuprous ion, which in turn was proportional to the i n i t i a l concentration of cupric ion. A f i r s t order dependence of the leach rate on the i n i t i a l cupric ion concentration was observed. At a fixed i n i t i a l cupric ion concentration, the leach rate depended directly upon the concentratio n or rate of formation of cuprous ions. Factors which theoretically increased the concentration of cuprous ions were high temperatures and, according to the Law of Mass Action applied to the equilibrium: 2Cu 2 + + S0„ + 2Ho0 — ^ 2Cu+ + SO.2" + 4H +, 2 aq 2 K — 4 + 2-high [S0„ ], low [H ] and low [SO. ]. Leach rates were increased 2. aq 4 by raising the temperature and the aqueous sulphur dioxide concentration, and by reducing the hydrogen ion concentration. The leach rate also depended upon the formation of an anionic cuprous complex which: •, 1. was strong enough to stabilize the Cu 1 state against further reduction or disproportionation; 2. had a favourable charge and stqich:Lometry for ads orption on the oxide surface and subsequent electron exchange between Cu* and Fe* 1* (possibly via a suitable bridging ligand). High pH and [ S 0 0 ] c, aq favoured the formation of complex anionic cuprous species because of the greater concentrations of bisulphite and/or sulphite ions. However, at the higher pH and [SO2 ] the reduction of cuprous ions to copper - 89 -metal was favoured. The effects that these considerations had on the leach rate were not apparent because the leach rate was increased by the greater concentration of cuprous ions formed under identical circumstances. The reduction of cupric ions to cuprous ions was shown to be taking place exclusively i n solution by the variable sample weight experiments. At low pH the rate of formation of cuprous ions was slow and this was the homogeneous catalyzed-rate-determining step. As the pH was raised the rate of formation of cuprous ions became fast.enough at 90°C for a surface reaction to become the slow, purely heterogeneous, rate-controlling step. 4.1.4 Effect on the Reductive Dissolution of the Addition of Other  Cations « Of the other cations added to a typical leach solution only mercury was found to promote the dissolution of goethite. Table XVIII. However, the conditions of the experiment were too reducing and/or the mercurous-bisulphite or -sulphite complexes were not strong enough to prevent the precipitation of mercury metal. Cations (type 2) capable of forming bisulphite or sulphite species which could adsorb on the goethite surface were added with no resultant increase in leach rate. This was consi dered as additional evidence for the assumption that increasing the surface concentration of bisulphite or sulphite ions does not increase the rate of dissolution. The reasons for the failure of the cations to promote or catalyze the leach rate were not investigated. - 90 -However i t was possible to speculate that they may have f a i l e d because of one or more of these reasons: 1. No stable higher pr lower oxidation state i n the leach s o l u t i o n . 2. Thermodynamics and/or k i n e t i c s were not favourable for the reduction of the cation i n the higher oxidation state by aqueous sulphur dioxide. 3. Thermodynamics and/or k i n e t i c s were not favourable f o r the I I I oxidation of the cation i n the lower oxidation state by Fe on the goethite surface. 4. I n a b i l i t y of the cation i n the lower oxidation state to form an anionic complex capable of adsorbing on the goethite surface. 5. Reduction of the cation to metal. It may be possible that any of these cations might catalyze the d i s s o l u t i o n of goethite i f c e r t a i n s p e c i f i c anions were present i n the s o l u t i o n capable of complexing the cation i n such a manner as to overcome the thermodynamic, k i n e t i c , or s t a b i l i t y b a r r i e r s . The d i r e c t leaching of hematite by o x a l i c acid i n the absence of oxygen was catalyzed by the addition of a small amount of ferrous oxalate, but was unaffected by the addition of ortho-phenanthroline, a stronger ferrous ion complexer (13,14). 4.1.5 E f f e c t on the Reductive D i s s o l u t i o n of the Addition of Oxalate  Anions The c a t a l y t i c e f f e c t of oxalate anions was observed when uranyl VI oxalate was used to investigate the p o s s i b i l i t y that U might catalyze the reductive d i s s o l u t i o n of goethite. Table XIX. Although the e f f e c t - 91 -was not investigated thoroughly, the observed pH and oxalate concentration dependency of the leach rate was reminiscent of Devuyst's ferrous oxalate catalyzed, direct oxalic acid leaching of hematite (13,14). In the oxalate catalyzed, reductive dissolution of goethite, the necessary ferrous ions were provided by the normal uncatalyzed leaching mechanism. Apparently, as found by Devuyst, the presence of ferrous ions over and above a certain low concentrati on had no extra catalytic effect. In his work the limit was the maximum solubility of ferrous oxalate. In this work the limit might have been the ferrous ion concentration at equilibrium with the much lower oxalate ion concentration. 4.2 Magnetite A detailed investigation of the reductive dissoluti on of magnetite by aqueous sulphur dioxide solutions was not carried out. The two experiments undertaken with and without the addition of cupric ions (Table XX) showed that in both systems the magnetite had comparable leach rates to goethite after normalisation of the different ores' iron contents. 4.3 Hematite Although a f u l l investigation of the reductive diss olution of hematite by aqueous sulphur dioxide soluti ons was not carri ed out, the two experiments performed showed anomalous behaviour compared with goethite (Table XXI). - 92 -The normalized uncatalyzed leach rate was an order of magnitude lower than the comparable rates for goethite, even though the surfaces were reported to be similar i n aqueous solution, Section 1.3.1. The slow rate of dissolution of hematite could have been due to a slow rate of hydration of hematite to produce a surface similar to goethite. However, although this could have been true for synthetic hematite (21), i t i s supposedly not true for natural hematite (24). The copper catalyzed leach rate of hematite was more comparable with that of goethite. Since a quantitative analysis was not made of the catalyzed dissolution of goethite, no attempt was made to take into account the catalyzed dissolution of the hematite present i n the goethite ore. - 93 -5. CONCLUSIONS 5.1 Summary 1. A mechanism has been proposed to explain how goethite was leached by perchloric acid. The rate of dissoluti on at 110°C was found to be described by the equation: d [ F e a q ] 3.153[H+] dt 1 + 0.688[H4T 2. A mechanism was also proposed to account for the reductive dissolution of goethite i n acidified aqueous sulphur dioxide solutions. The rate of dissolution at 110°C was found to be described by the equation: d[Fe ] [H+](195.19 + 2063.27[S0o ] + 984.50[H+][S0o ]) 1 aq J _ 2 aq J 1 2 aq d t 1 + 55.185[H+] The rate controlling step was heterogeneous. 3. A qualitative mechanism was proposed to describe the role of cupric ions in increasing the reductive leach rate of goethite. Cuprous ions, formed by the reduction of cupric ions by sulphur dioxide molecules in solution, were thought to have been adsorbed on the - 94 -goethite surface as a suitable anionic complex. Electron exchange between Fe*** and Cu* was followed by the desorption of ferrous and cupric ions into solution. Cupric ions were therefore catalytic agents. The reduction of cupric ions to cuprous ions was the rate controlling step at low pH. At higher pH the control was mixed for low sample weights of ore. 4. Both magnetite and hematite leached at comparable rates with goethite i n the cupric ion catalyzed system. Although the reductive leach rate of magnetite was comparable with goethite's in the uncatalyzed system, that of hematite was an order of magnitude lower. 5.2 Suggestions for Future Work 1. The uncatalyzed reductive leaching of goethite could be examined at high acidity and at much higher aqueous sulphur dioxide concentrations. 2. A detailed investigation of the reductive leaching of both magnetite and hematite i n aqueous sulphur dioxide s olutions could be made. This might explain the anomalous dissolution behaviour of hematite in the uncatalyzed system. 3. The cupric ion catalyzed reductive dissolution of goethite could be investigated in the presence of an anion having a strong complexing power for cuprous ions. This might have the effect of stabilizing the cuprous ion so that further reducti on or disproportionation was impossible, even at high pH. The anionic cuprous complex would have to be able to adsorb on the oxide surface and permit electron exchange between Cu* and Fe***. - 95 -4. 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