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Reductive dissolution of goethite and pyrolusite in alkaline solution Devuyst, Eric 1970

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REDUCTIVE DISSOLUTION OF GOETHITE AND PYROLUSITE IN ALKALINE SOLUTION by ERIC DEVUYST Ing. Civ. Mines, U.L.B. Brussels, 1968 A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF APPLIED 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 October, 1970. In p r e s e n t i n g t h i s t h e s i s in p a r t i a l f u l f i l m e n t o f t h e r e q u i r e m e n t s f o r an advanced d e g r e e a t t h e U n i v e r s i t y o f B r i t i s h C o l u m b i a , I a g r e e t h a t t h e 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 s t u d y . I f u r t h e r a g r e e t h a 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 p u r p o s e s may be g r a n t e d by t h e Head o f my Depar tment o r by h i s r e p r e s e n t a t i v e s . I t i s u n d e r s t o o d t h a t c o p y i n g o r 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 a l l o w e d w i t h o u t my w r i t t e n p e r m i s s i o n . Depar tment o f The U n i v e r s i t y o f B r i t i s h C o l u m b i a V a n c o u v e r 8, Canada i ABSTRACT The s e l e c t i v e d i s s o l u t i o n of p y r o l u s i t e (MnC^) i n presence of goethite (FeO.OH) i n aqueous ammoniacal ammonium carbamate s o l u t i o n has been investigated; various reducing agents were studied, i n p a r t i c u l a r sodium s u l f i t e (Na^SO^) and hydrazine hydrate (N^H^.H^O). Sodium s u l f i t e was found to be an unsatisfactory reducing agent for both goethite and p y r o l u s i t e , f or the leaching rate decreased s t e a d i l y with time to an impractical l e v e l because of the l i m i t e d s o l u b i l i t y of the reactions products. Hydrazine hydrate was an e f f e c t i v e reducing agent for p y r o l u s i t e even at low temperature, whereas goethite required much higher temperatures f o r any appreciable d i s s o l u t i o n . A one hour leach at 35°C dissolved 90 per cent of the p y r o l u s i t e but only 0.1 per cent of the goethite. Additions of ammonium phosphate had no e f f e c t on the rates, there being no phosphate detected i n the leach l i q u o r . An a c t i v a t i o n energy of 17.0 K cal/mole was found for the d i s s o l u t i o n of p y r o l u s i t e with hydrazine. A r e a c t i o n mechanism f o r the leaching of p y r o l u s i t e with hydrazine was proposed i n which a reduction reaction on the surface of the mineral was the rate-determining step. I t was possible to recover manganese from s o l u t i o n with oxygen as an amorphous Mn0„ p r e c i p i t a t e . ACKNOWLEDGEMENT The author wishes to express h i s sincere thanks to Dr. I.H. Warren for h i s continued i n t e r e s t , i n s p i r a t i o n , and encourage-ment necessary to bring t h i s work to i t s f i n a l form. Thanks are also extended to members of f a c u l t y , fellow graduate students and research a s s i s t a n t s , for h e l p f u l discussions, and to the te c h n i c a l s t a f f f o r t h e i r kind assistance with p r a c t i c a l aspects of the work. F i n a n c i a l support from the National Research Council of Canada i n the form of a Research Assistantship i s g r a t e f u l l y acknow-ledged. i i i TABLE OF CONTENTS Page GENERAL A. GENERAL 1 B. REVIEW OF LITERATURE 4 1. Physico-Chemistry of Goethite 4 2. Physico-Chemistry of Manganese Dioxides 6 2.1 C r y s t a l Structures 6 2.1.1 Well C r y s t a l i z e d V a r i e t i e s 6 2.1.2 A r t i f i c i a l Products and Non-Stoichiometric 8 Natural Forms 2.1.3 Chemical Composition 8 2.3 Chemical Properties 8 2.3.1 Thermal D i s s o c i a t i o n of Manganese Dioxide 8 2.3.2 Equilibrium Diagram of the System Mn-H20-CO2 10 at 25°C and 1 Atmosphere 2.3.3 Colloid-Chemical Properties of Manganese Dioxide 12 3 Hydrometallurgy of Manganese 12 3.1 Sulfur Processes 13 3.2 Ammonia Processes 13 3.2.1 Bradley-Fitch Process 13 3.2.2 Dean Process 13 3.3, N i t r i c Acid Process 14 3.4 E l e c t r o l y t i c Production of y-MnO^ 14 4. Hydrazine as a Reducing Agent 15 4.1 General 15 Page i v 4.2 Oxidation Mechanisms of Hydrazine 16 4.3 Reduction of Manganese Dioxide by Hydrazine 19 5. The NH3-CO2-H20 System 20 5.1 Equilibrium i n the System NH3~CO2-H20 20 5.2 K i n e t i c s and Equilibrium of the Hydrolysis of 22 Ammonium Carbamate i n the NH^-CO^^O System 5.2.1 Determination of Ammonium Carbamate 22 5.2.2 K i n e t i c s and Mechanism of Hydrolysis of Ammonium 23 Carbamate 3.2.3 K i n e t i c s of Carbamate Formation and Breakdown 24 5.3 Applications of the NH3~CO2-H20 System 24 5.3.1 Manganese from Low-Grade Ores 24 5.3.2 E l e c t r o l y t i c D i s s o l u t i o n of Iron 26 C. SCOPE OF THE PRESENT INVESTIGATION 29 EXPERIMENTAL A. MATERIALS 30 1. Reagents 30 2. Minerals 30 2.1 Goethite 30 2.2 P y r o l u s i t e 30 B. AUTOCLAVE DESIGN 32 C. EXPERIMENTAL PROCEDURE 33 D. ANALYTICAL METHODS 34 1. Iron 34 2. Manganese 34 3. Phosphate 35 4. Ammonium Carbamate 36 E. SYNTHETIC MINERALS 36 1. Ferrous Oxide 36 2. Manganous Oxide 36 Page v RESULTS AND THEIR ANALYSIS A. THE NH3-CO2-H20 SYSTEM 1. E f f e c t of Temperature on the Equilibrium H„0 + NH.CCi NH. = (NH.)_ CO. I 4 I I 4 2 3 2. Equilibrium Pressure at High Temperature i n the System NH.J-CO -H 0 B. REDUCTIVE LEACHING OF GOETHITE 1. Leaching of Iron (130°C) S o l u b i l i t y of FeO Ef f e c t of Temperature E f f e c t of Adding Sodium Sulfate at 130°C E f f e c t of Adding Sodium S u l f i t e at 130°C Reductive Leaching of Goethite at 130°C 2. 2.1 2.2. 2.3 3. 3.1 Reductive Leaching of Goethite Using Sodium S u l f i t e i n Presence of Anthraquinone Catalyst Compounds 3.1.1 E f f e c t of Sodium S u l f i t e Concentration 3.1.2 E f f e c t of Anthraquinone Concentration 3.1.3 E f f e c t of the Nature of the Catalyst 3.1.4 E f f e c t of Leaching the Residue 3.1.5 E f f e c t of Doubling the Amount of Goethite 3.1.6 E f f e c t of Adding Sodium Sulfate at the Start 3.1.7 E f f e c t of Adding Sodium S u l f i t e dn the Middle of the Run 3.1.8 E f f e c t of Adding Sodium g Sulfonate Anthraquinone i n the Middle of the Run 3.2 Reductive Leaching of Goethite Using Hydrazine Sulfate at 130°C 3.2.1 E f f e c t of Hydrazine Sulfate Concentration 3.3 Reductive Leaching of Goethite Using Hydrazine Hydrate at 84°C C. REDUCTIVE LEACHING OF PYROLUSITE 1. S o l u b i l i t y of MnO 38 38 39 39 39 44 44 47 47 50 50 50 52 52 52 57 57 57 59 59 59 63 63 1.1 E f f e c t of Time at Constant Temperature 1.2 E f f e c t of Temperature 1.3 E f f e c t of Adding Sulfate 1.4 E f f e c t of Adding S u l f i t e 2. Reductive Leaching of P y r o l u s i t e 2.1 Reductive Leaching of P y r o l u s i t e Using Sodium S u l f i t e 2.2 Reductive Leaching of P y r o l u s i t e Using Hydrazine Hydrate as a Reducing Agent 2.2.1 Leaching of P y r o l u s i t e with a Large Excess of Hydrazine Hydrate at 85°C 2.2.2 E f f e c t of Temperature 2.2.3 E f f e c t of Hydrazine Hydrate Concentration on the Reductive Leaching of P y r o l u s i t e 2.2.4 E f f e c t of Doubling the I n i t i a l Amount of Py r o l u s i t e D. SOLUBILITY OF AMMONIUM PHOSPHATE IN THE SYSTEM NH 3-C0 2-H 20 1. E f f e c t of Temperature 2. E f f e c t of Adding FeO 3. E f f e c t of Adding MnO E. RECOVERY OF MANGANESE FROM SOLUTION WITH OXYGEN 1. E f f e c t of the Oxygen Pressure on the P r e c i p i t a t i o n Rate of Complexed Manganese out of Solution 2. E f f e c t of the S t i r r i n g Rate on the P r e c i p i t a t i o n . Rate DISCUSSION AND CONCLUSIONS A. DISSOLUTION OF FeO AND MnO B. REDUCTIVE DISSOLUTION OF GOETHITE I Sodium S u l f i t e II Hydrazine Sulfate III Hydrazine Hydrate v i i Page C. REDUCTIVE DISSOLUTION OF PYROLUSITE 96 I Sodium S u l f i t e 97 II Hydrazine Hydrate 97 D. DISSOLUTION OF PHOSPHATE 106 E. RECOVERY OF MANGANESE 107 v i i i LIST OF TABLES No. Page 1. L a t t i c e Parameters of Manganese Dioxide 6 2. Thermal Transformation of Manganese Dioxide 9 3. E l e c t r o l y t i c Mn0 2 P u r i t y 15 4. Chemical Composition of Goethite 30 5. Chemical Composition of P y r o l u s i t e 31 6. X-ray D i f f r a c t i o n Pattern of P y r o l u s i t e 31 I E f f e c t of [Na„S0-] Concentration on the Leaching -Q3 of Goethite at 130°C II E f f e c t of [N H..B SO,] Concentration on the Leaching of Goethite at 130°C II I E f f e c t of Temperature on the S o l u b i l i t y of FeO 1 1 5 and MnO IV Log-log P l o t between [SO^ ^±n±^±ai~^n complex] and -Qg time V Pressure V a r i a t i o n i n the Leaching of P y r o l u s i t e with -j.17 Hydrazine Hydrate VI E f f e c t of [Na ?S0„] Concentration on the Leaching of Py r o l u s i t e at 130°C H 8 VII E f f e c t of Temperature on the D i s s o l u t i o n Rate of ^_19 P y r o l u s i t e (50-70 mesh) i n 0.0874 M N ^ . ^ O VIII E f f e c t of Temperature on the D i s s o l u t i o n Rate of Py r o l u s i t e (65-150 mesh) i n 0.0874 M N ^ . ^ O 120 IX E f f e c t of [N_H,.H_0] Concentration on the D i s s o l u t i o n Rate of P y r o l u s i t e (65-150 mesh) at 30°C 1 2 1 LIST OF FIGURES The Structures of P y r o l u s i t e , Ramsdellite, Cryptomelane and Psilomelane S t a b i l i t y Relations Among Manganese Compounds i n Water at 25°C and 1 Atmosphere T o t a l Pressure. T o t a l Dissolved Carbonate Species = 10 -1»4 E f f e c t of Temperature on the Equilibrium NH 4C0 2NH 2 + H 20 = (NH 4) 2 C0 3 E f f e c t of NH3 Concentration on the Equilibrium NH 4C0 2NH 2 + H 20 = (NH 4) 2 C0 3 Equilibrium Pressure Versus Temperature Mole percent of Ammonium Carbamate i n Solution versus Time D i s s o l u t i o n of M e t a l l i c Iron i n the System NH 3-C0 2-H 20 S o l u b i l i t y of FeO i n 18 moles/1 NHLy- 3 moles/1 C0 2 Solution at Room Temperature E f f e c t of Temperature on the S o l u b i l i t y of FeO i n the System NH3-CO2-H20 E f f e c t of Adding Sulfate on the S o l u b i l i t y of FeO at 130°C E f f e c t of S u l f i t e Concentration on the D i s s o l u t i o n of Goethite at Constant Anthraquinone Concentration E f f e c t of Anthraquinone Concentration on the D i s s o l u t i o n of Goethite at Constant S u l f i t e Concentration Log-log P l o t Between Anthraquinone Concentration and the Amount of Iron Dissolved at 130°C E f f e c t of the Nature of the Catalyst Leaching of Goethite Using S u l f i t e at 130°C. E f f e c t of Adding Su l f a t e , of Doubling the I n i t i a l Amount of Goethite and of Redissolving the Residue E f f e c t of Adding S u l f i t e or Sodium $ Sulfonate Anthraqui-none i n the Middle of the Runs i n the D i s s o l u t i o n of Goethite with S u l f i t e at 130°C E f f e c t of Hydrazine Sulfate Concentration on the Leaching of Goethite at 130°C Log-log P l o t Between Hydrazine Concentration and the Amount of i r o n Dissolved at 130°C. Constant Time Plots Reductive Leaching of Goethite Using Hydrazine Hydrate at 84°C E f f e c t of Time at Constant Temperature on the S o l u b i l i t y of MnO E f f e c t of Temperature on the S o l u b i l i t y of MnO E f f e c t of S u l f i t e Concentration on the Leaching of P y r o l u s i t e Log-log Plot Between (K^-Mn Complex) and Time Presssure V a r i a t i o n i n the D i s s o l u t i o n of P y r o l u s i t e Using Hydrazine Hydrate as Reducing Agent Di s s o l u t i o n of P y r o l u s i t e Using Hydrazine Hydrate E f f e c t of the S t i r r i n g Rate on the D i s s o l u t i o n of Pyr o l u s i t e Using Hydrazine at 30°C E f f e c t of Temperature on the Leaching of P y r o l u s i t e (50-70 mesh) with Hydrazine E f f e c t of Temperature on the Leaching of P y r o l u s i t e (65-150 mesh) with Hydrazine Arrhenius P l o t f o r Leaching of Py r o l u s i t e with Hydrazine E f f e c t of Hydrazine Concentration on the D i s s o l u t i o n of P y r o l u s i t e at 30°C Log-log P l o t Between Hydrazine Concentration and the Di s s o l u t i o n Rate of P y r o l u s i t e at 30°C E f f e c t of Doubling the I n i t i a l Amount of P y r o l u s i t e on the D i s s o l u t i o n of P y r o l u s i t e at 30°C S o l u b i l i t y of Phosphate; E f f e c t of Temperature and of I n i t i a l FeO or MnO E f f e c t of Phosphate on the D i s s o l u t i o n of FeO and MnO E f f e c t of Oxygen Pressure on the P r e c i p i t a t i o n of Manganese from Solution at 35°C No. 36. Log of Manganese Concentration i n Solution versus Time Plots for D i f f e r e n t Oxygen Pressures 37. Log-log Plot Between k 3 ( f 0 2 ) X and |>02 38. E f f e c t of the S t i r r i n g Eate on the P r e c i p i t a t i o n of Manganese from Solution at Constant Oxygen Pressure GENERAL A. GENERAL This work was undertaken as part of a program of studying the mechanism of d i s s o l u t i o n of oxides i n acid and a l k a l i n e s o l u t i o n s . The present study i s concerned with d i s s o l u t i o n of i r o n and manganese oxides i n a l k a l i n e s o l u t i o n . To be used f o r m e t a l l u r g i c a l purposes, manganese ore should be low i n s i l i c a plus alumina, low i n phosphorus (0.10 - 0.20 per cent) and have a manganese to i r o n content r a t i o of at le a s t 7:1 A t y p i c a l percentage analysis of the higher grade ferro-manganese obtained a f t e r smelting i s : Mn:79.1, C;7.07,iv Si:0..76, S: 0.027, P:0.20, Fe;12.84 Phosphorus and sulphur are undesirable impurities i n manganese ore. Phosphorus i n s t e e l i s usually l i m i t e d by s p e c i f i c a t i o n to a quantity less than 0.05 per cent. Low grade ores''" are too high i n objectionable impurities to serve as economic sources of manganese for the production of f e r r o -manganese i n the b l a s t furnace. They t y p i c a l l y contain l e s s than 10 per cent manganese, plus two to four times as much i r o n or s i l i c a or both; some are also high i n phosphorus; and the elements are intimately mixed so that the manganese cannot be concentrated s u c c e s s f u l l y by mineral dressing methods. Therefore pyro-or hydrometallurgical tech-niques are required to produce an acceptable synthetic ore. Manganese dioxide may be used as a depolarizer on the p o s i t i v e '2 pole of dry c e l l s . Natural manganese dioxide ore must have an oxide content which i s as high as possible and must be free from impurities. Soluble impurities which are electronegative to zin c (negative pole 2 of the c e l l ) such as copper, n i c k e l , cobalt and arsenic, are p a r t i c u l a r -l y harmful since i n s o l u t i o n they are deposited on the zinc and cause corrosion and d e t e r i o r a t i o n of the c e l l . I f the impurities are i n an insolub l e form they increase the i n t e r n a l resistance of the c e l l which i s also undesirable. Iron oxide i s i n e r t and can be tolerated to an ex-3 4 tent of 3-4 per cent. Synthetic manganese dioxide ' f u l f i l s the above mentioned conditions and can be obtained from leach solutions by pre-c i p i t a t i o n or by e l e c t r o l y s i s . In the future, leaching of manganese sea nodules r i c h i n copper, n i c k e l and cobalt w i l l be of great i n t e r e s t . Sea-floor manganese nodules'* are intimate mixtures of manganese and i r o n oxides, coprecipitated or absorbed metals, various organic constituents and s i l i c a t e and oxide minerals. The i r o n oxides are very poorly c r y s t a l l i n e while the manganese oxides are present i n a number of recognizable cry-s t a l forms, probably layered, the layers allowing the incorporation of a group of cations - such as n i c k e l , copper, cobalt, lead, z i n c , molybdenum - which are found i n rather high concentrations i n the nodules. It may be expected that i f the nodules are leached under reducing conditions, n i c k e l , copper and cobalt w i l l remain behind as a residue, whilst manganese and p o s s i b l y i r o n are leached. This might be a commercially interesting process. Generally, i n hydrometallurgical p r a c t i c e , a l k a l i n e leaching i s very often preferred to acid leaching. Basic solutions are usually more s e l e c t i v e i n d i s s o l v i n g minerals and are generally l e s s corrosive. Ammoniacal leaching usually enables the ammonia to be recovered and r e -cycled or used to make f e r t i l i z e r s . The ammines of n i c k e l , cobalt (111), copper, zinc and cadmium are soluble i n water either alone or on addition of ammonia, a con-siderable excess of, ammonia i s necessary for the s o l u t i o n of the c o b a l t ( l l l ) ammines, When the ammines of the s a l t s of magnesium, manganese, divalent cobalt and i r o n are dissolved i n water, p r e c i p i t a t e s of the hydrox-6 8 9 ides are formed ' ' . Excess of ammonia Increases the s t a b i l i t y of the complexes but also diminishes the s o l u b i l i t y of the ammines. The s t a b i l i t y of complexes with ammonia diminishes i n the order Pd, Cu, N i , Pb, Co, Zn, 7 Cd, Fe, Mn, Mg, independently of the l i q u i d . Werner pointed out that ammonia molecules can be displaced, one by one, from the coordination sphere, either by other neutral groups such as water, or by negative groups. Dean"^ showed that soluble manganous s a l t s such as chloride and su l f a t e could be dissolved to clear solutions In concentrated ammonium hydroxide provided a i r was eliminated. He suggested that at high ammonia concentrations, a new type of manganese ammonia complex was formed having manganese i n the anion. Such a compound might, according to Dean, have the formula (X - Mn (NH 3) x - 0)~ (NH 4 +) where X i s a monovalent anion. The i n t e r e s t i n g f a c t i s the discovery by the author that such complexes can be b u i l t up to high concentrations i n sol u t i o n whereas i n acid s o l u t i o n and i n absence of s p e c i f i c X anions manganous hydroxide i s p r e c i p i t a t e d out of s o l u t i o n on addition of ammonia. The postulated anionic character of the new manganese-ammonia complex was regarded as confirmed by Dean by e l e c t r o l y t i c transport experiments. Although many ammonium s a l t s can be used to obtain manganese i n so l u t i o n , there are operational and economic advantages attached to the use of ammonium carbamate, which i s obtained when CC^ i s passed into con-centrated aqueous solutions of ammonia. It appears highly probable that 4 a manganese-ammonia complex i s farmed with the carbamate ion and that carbonate ion i s present i n r e l a t i v e l y small amounts. I t should be noted that also ferrous oxide i s soluble to a large extent i n the above mentioned carbamate s o l u t i o n . Since f e r r i c and manganic s a l t s are i n s o l u b l e i n these s o l u t i o n s , goethite, hematite and manganese dioxides must f i r s t be reduced to y i e l d divalent metal ions before they can form soluble carbamate complexes. This reduction may be performed i n a preliminary pyrometallurgical step or simultaneously i n a leaching operation. S u l f i t e and hydrazine are power-f u l reducing agents i n a l k a l i n e s o l u t i o n . S u l f i t e , however, introduces s u l f a t e into the processing solutions and a d d i t i o n a l l y t h i s anion might complicate the NH^ - CO^ - carbamate equilibrium at high temperature. Hydrazine changes the oxidation-reduction p o t e n t i a l of the s o l u t i o n and leads to the possible by-products ammonia and nitrogen. It i s therefore intended to give a general review of the l i t e r -ature on goethite, manganese dioxides, hyrazine and the properties of the leaching s o l u t i o n , which was used, including some applications of ammoniacal carbamate solutions which have been studied previously. B. REVIEW OF LITERATURE 1. Physico-Chemistry of Goethite For the physico-chemistry of goethite reference can be made to 11 12 13 the work of Monhemius , Bath and Surana who investigated the d i s -s o l u t i o n of i r o n oxides i n acid solutions. A mechanism of d i s s o l u t i o n i n -volving hydration and production of a p o s i t i v e l y charged oxide surface was invoked to explain the observed k i n e t i c s of d i s s o l u t i o n of goethite and hematite by the above workers. E a r l i e r Gayer"^ and P a r k s ^ studied the surface behaviour of i r o n oxide i n acid-base solutions. They suggest that the mechanism by which the surface charge i s established may be viewed q u a l i t a t i v e l y as a two-step process; surface hydration followed by d i s s o l u t i o n of the surface ''hydroxide". In any event experimental adsorption measurements cannot d i s t i n g u i s h between the adsorption of a hydroxy! ion or the desorption of a hydrogen ion. The zero point of charge may be rela t e d to the pH of the s o l u t i o n when the adsorption d e n s i t i e s of H + and OH ions are equal or: F e 2 0 3 + 3H20 = 2Fe (0H>3 (1) and was measured at pH 8.5. The surface reactions involved i n the establishment of a sur-face charge are represented formally as: |Fe(OH) 3 + H* . = |Fe(OH) 2 + H 20 pH<8.5 (2) s q s |Fe(OH) 3 = JFe 0~ + ^0+ pfl >8.5 (3) s s By adding (2) and (3) i t i s found that: h |Fe(0H) 2 = |Fe 0 2 + 2H + s s with - [|Fe 0~ ] K, = — ~Z x ( a „ + ) 2 - l(f17 [|Fe(0H) 2l s at 25°C, and assuming that a," - of ~ 1 ( a c t i v i t y c o e f f i c i e n t s of the two surface s i t e s ) . The presence of Fe 0 2 has only been known i n highly a l k a l i n e s o l u t i o n s . Below pH 1 a second protonation of the oxide surface 6 may occur; | FeCOH)* + . H + H . | Fe(QH)"^" + ^ 0 s s 2. Physico-Chemistry of Manganese Dioxides 2.1 C r y s t a l structures 2.1.1. Well C r y s t a l i z e d V a r i e t i e s Mineralogists found a natural tetragonal r u t i l e type v a r i e t y c a l l e d p y r o l u s i t e . An orthorhombic v a r i e t y was mentioned by Strunz^^, c a l l e d ramsdellite. Wadsley"*"^ described a monoclinic. form of manganese 17 18 dioxide, psilomelane. Bystrom ' gave a complete d e s c r i p t i o n of crypto-melane which he c l a s s i f i e d as pseudo-tetragonal with a monoclinic character. The structures are shown i n Figure 1. The parameters of the above mentioned v a r i e t i e s are given i n Table I. TABLE I L a t t i c e Parameters of Manganese Dioxides Ramsdellite P y r o l u s i t e Psilomelane Cryptomelane a 4.53 4.39 9.56 9.82 b 9.27 4.39 13.85 9.27 c 2.87 2.87 2.88 2.87 o Parameter c has a constant value of 2.87 A f o r the four forms of n atural manganese dioxide. 2.1.2. A r t i f i c i a l Products and Non-Stoichip-metric Natural Forms 19 P. Dubois has studied a -r MnC^ (similar to cryptomelane) and g - Mn02 ( s i m i l a r to pyrolus i t e ) which, can be obtained with a perfect structure. The y " ^ nQ^ v a r i e t y (approaching the ramsdellite structure) 20 was studied by Fleitknecht . 6 - Mn0 2 was obtained by e l e c t r o l y s i s . 21 Matsuno described i t s structure as a defective structure of cryptomelane. Other more complex v a r i e t i e s e x i s t . 2.1.3. Chemical Composition 1) P y r o l u s i t e : i t i s the most important manganese ore. P y r o l u s i t e can contain, besides manganese oxides, elements such as Fe, Al , Ba, S^, Ca, Pb, Cu, N i , Co and Zn. 2) Psilomelane: Wadsley^ proposed the formula. (Ba, H 0) 2' Mn^O^^ with Ba/H^O approximately equal to 0.5. 3) Cryptomelane: Bystrom"^ gave a general formula, A. B„ X. r , A 2-y 8-z 16 being large ions - (Ba , Pb , K , Na ) B small ions as Mn , Mn and X, 2- -0 or OH ions. 4) Ramsdellite: i t s c o n s t i t u t i o n i s s i m i l a r to that of p y r o l u s i t e with res-pect to oxygen configuration i n the l a t t i c e . However, i n the a b d i r e c t i o n " 2 + ramsdellite shows Mn occupied octahedra a l t e r n a t i n g with unoccupied octahedra, while f o r p y r o l u s i t e one out of two octahedra i s occupied by Mn ions. 2.3 Chemical Properties 2.3,1. Thermal D i s s o c i a t i o n of Manganese Dioxide 22 K. Traore studied the thermal decomposition of 3 - Mn02 a n c^ Y - Mn©2 by thermogravimetric analysis. On heating, i n presence of oxygen, between 20°C and 300 oC,y- MnO^ transforms into g - Mn02 due to the decrease • 4+ of the amount of.Mn holes i n the l a t t i c e by loss of OH ions i n the form 9 of water and oxygen. Between 300°C and 500°C, 8 - MnO,, goes to MnO^ which i s an intermediate compound containing 0^ i n s o l i d s o l u t i o n i n the M n ^ l a t t i c e . Between 500°C and 600°C Mn0 1 5 5 transforms i n t o M^O-j by loss of the dissolved 0,,. If a l l the transformations were carried out under vacuum the Mn01 R (. intermediate did not appear. 23 Brenet and coworkers studied the same transformations by d i f f e r e n t i a l thermal/analysis. Their r e s u l t s are described i n Table 2. TABLE 2 Thermal Transformations of Manganese Dioxide P y r o l u s i t e : 3 - MnO^ T° Transformation A c t i v a t i o n Energy (°C) . (K cal/mole) 20-120 Loss of adsorbed H^ O 750 3 Mn02 * a M n ^ 105 860-1000 a Mn„0„ Mno0. 270 2 3 3 4 E l e c t r o l y t i c : y - MnO^ T° Transformation A c t i v a t i o n Energy (K cal/mole) 24 160 50 270 20-120 Loss of adsorbed H 20 120-400 Loss of bonded H 20 Y Mn02 •+ 3 Mn0 2 400-600 3 Mn02 -> Mn0 1 5 5 600-750 rniQ± 5 5 -> a Mn 20 3 860-1100 a Mn o0 o -> Mn„G. 2 3 3 4 10 2.3.2. Equilibrium Diagram; of the System Mn - H^ O - C0 2 at 25°C  and 1 Atmosphere 24 The diagram for p y r o l u s i t e as Mn0 2 i s given i n Figure 2. The s t a b i l i t y region of Mn° i s f a r below that of water, of which the lower l i m i t i s given by l i n e (a). Manganese i s a strong reducing agent and reacts with water giving H 2 > However, the l a s t r e action i s slow due to the high hydrogen overvoltage. Manganese dissolves e a s i l y i n acid and neutral s o l u t i o n with a low 0 2 content; 2+ Mn ions being formed. These manganous solutions can be oxidized i n acid or a l k a l i n e medium p r e c i p i t a t i n g out Mn^O^, Kn^O^ or Mn0 2 de-pending on the oxidation p o t e n t i a l . Strong e l e c t r o l y t i c oxidation i s 2- -able to produce Mn.O^  ions at high pS and, MnO^ ions at any pH. The s t a b i l i t y region of MnO^ i s situated above the oxygen-water l i n e b at atmospheric pressure. In presence of water MnO^-decomposes slowly to Mn02 and 0^. The Kn^O^ region only e x i s t s i n a l k a l i n e s o l u t i o n s . 2+ If such solutions become acid Mn20.j diss o c i a t e s i n Mn ions and Mn0 2 > 2+ Also, by a l k a l i n i z a t i o n of Mn solutions under reductive conditions a MnCO^ (rhodochrosite) p r e c i p i t a t e i s formed or even a white Mn(0H) 2 (pychroite) p r e c i p i t a t e may appear. (Mn(0H) 2 i s not l i k e l y to be stable when excess C0 2 i s present). In very strong a l k a l i n e s o l u t i o n Mn(0H) 2 dissolves by giving HMn02 ions. The Mn(0H) 2 and MnCO^ region contains the whole MnO s t a b i l i t y region so that MnO i s not stable i n presence of C0 2 - water and forms MnCO^ or Mn(0H) 2 < The large f i e l d of the carbonate (rhodochrosite) i s i n marked contrast to the much smaller f i e l d of s i d e r i t e (FeCO^) under the same conditions. The Mn(0H) 2 s t a b i l i t y region i s s h i f t e d to the a l k a l i n e side by the presence of CX>2 and i s very much smaller. 11 12 2.3.3. Colloid-Chemical Properties of Manganese Dioxide As has been shown to be the case f o r several polyvalent metal oxide hydrates the surface s i t e s of c o l l o i d a l hydrous manganese dioxide have amphoteric properties. The i s o e l e c t r i c point, or more p r e c i s e l y the zero point of charge with respect to net OH and H + bound of 25 c o l l o i d a l hydrous manganese dioxide was measured by Gordon at a pH of 2.8 - 0.3. Within the pH range of p r a c t i c a l i n t e r e s t c o l l o i d a l MnO^ i s negatively charged. The potential-determining r o l e of OH ions at pH values above the zero point of charge may be v i s u a l i z e d either as a binding of OH ions or as a d i s s o c i a t i o n of H + ions from surface OH groups as following: ( M n 0 2 ) J H 20 + 0 H " a q ) - (MnO^J ^ 0 . 0 ^ s s (Mn0 2) x| H 20 - ( M n 0 2 ) J OH" + H+ s s ^ Morgan showed that the cation exchange capacity of the hydrous oxides i s strongly pH dependent and increases with increasing pH (increasing negative charge of the oxide surface) and the a f f i n i t y of the Mn0 2 for H + and multivalent cations i s larger than that for a l k a l i ions. At pH 7.5 the capacity of the manganese dioxide f o r 2+ 2+ Mn sorption i s 0.5 mole of Mn per mole of Mn02; capacities of 2 moles per mole are r e a l i z e d at pH values near 9. 3. Hydrometallurgy of Manganese As mentioned above hydrometallurgical processes apply e s s e n t i a l l y to low grade manganese o r e ^ . 13 3.1.Sulphur Processes Manganese dioxide may be r e a d i l y converted to manganese s u l -fate by treatment by gaseous s u l f u r dioxide or by sulphuric or s u l -furous acid solutions at concentrations which do not convert i r o n oxide to s u l f a t e . The chemical separation of manganese from i r o n , s i l i c o n and other elements present i n the ores i s r e l a t i v e l y simply accomplished but the development of a successful commercial process has not been achieved. 3.2. Ammonia Processes 3.2.1. Bradley-Fitch Process A preliminary reducing roast i s c a r r i e d out on the ore to convert the manganese dioxide, which i s insolu b l e i n ammonia, to man-ganous oxide, which i s soluble i n the presence of ammonium s u l f a t e . The f i l t r a t e could be used as an e l e c t r o l y t e f o r the production of e l e c t r o l y t i c manganese. This process remained on the scale of a p i l o t -plant. 3.2.2. Dean Process The process that has seemed most promising i s the Dean-Leute ammonium carbonate process, which involves reduction of manganese dioxide to manganous oxide and then leaching with a solvent composed of carbon dioxide and ammonia. More d e t a i l s on t h i s process are given i n the l a s t section of the l i t e r a t u r e review. I t should be noted that a f t e r ex-tensive research, Manganese Chemical Corporation erected a semi-commer-c i a l plant at Riverton, Minnesota, for treatment of Cuyuna range man-ganiferous ore. This plant closed early i n 1962, apparently as an un-economic venture. 14 3.3, N i t r i c Acid Process In the Nossen n i t r i c acid c y c l e , manganese n i t r a t e i s pro-duced and decomposed to manganese dioxide at 200°C. The manganese dioxide product can be used for b a t t e r i e s as well as i n production of ferro a l l o y s . 3.4. E l e c t r o l y t i c Production of y. - MnO^ Production of e l e c t r o l y t i c manganese dioxide has grown r a p i d l y i n recent years, p a r t i c u l a r y i n Japan since 1963. A good 27 summary of Japanese y - MnO^ production was given by Takuhashi . The two major uses of e l e c t r o l y t i c - MnO^ are as a depolarizer (cathode active material) i n dry ba t t e r i e s and as a component of manganese f e r r i t e s . The premium c e l l s (or high q u a l i t y Leclanche c e l l s ) use 100% e l e c t r o -l y t i c - Mn02 instead of natural ore i n the cathode. In the case of Mn02 ores, since the ore does not dis s o l v e i n l^SO^, the ore must be roasted with a reducing agent, coal or carbon, i n order to convert i t to MnO which dissolves e a s i l y i n ^SO^ solutions. The impurities are usually leached out into the s u l f u r i c acid s o l u t i o n , so that they must be removed by s u l f i d e addition or by 28 other ways The c e l l s are operated with d i r e c t current of current density 2 0.7-1.2 amp/dm at 88-98°C and with an e l e c t r o l y t e containing 0.5 - 1.2 M/l of MnSO and 0.5 - 1.0 M/l of H SO,. During e l e c t r o l y s i s , 4 * ' the MnSO, concentration decreases and the H„S0. concentration increases 4 2 4 i n the e l e c t r o l y t e as a r e s u l t of the following reactions: Anode: Mn + 2H20 Mn02 + 4H + 2e Cathode: 2H + + 2e~ -> Hg Ov e r a l l : Mn"^ + 2H20 H 2 + Mn02 + 2H + The Mn02 content and impurity l e v e l s of t y p i c a l e l e c t r o l y t i c -MnO„ are given i n Table 3. Element TABLE 3 E l e c t r o l y t i c MnO„ Weight % Pur i t y Element Wt% Mn02 91.65 Cr 0.0003 H 20 1.77 Cu 0.0002 Fe 0.008 Pb 0.0005 Ni 0.005 SO. 4 0.9 Co 0.002 80% by weight of the Mn0 2 has a p a r t i c l e s i z e of less than 74 y, 20% between 74 - 149u. The y-Mn02 has a non-stoichiometric formula M n 0 1 > 9 5 _ 1 > 9 ? . 4. Hydrazine as a Reducing Agent 29 4.1 General Hydrazine (N^^) i s a c o l o r l e s s , hygroscopic l i q u i d having an odor resembling that of ammonia, but with several of i t s p h y s i c a l prop-e r t i e s s i m i l a r to those of water. Hydrazine i s a powerful reducing agent and leaves no un-desirable residues, hence, i t s use as an oxygen scavenger i n the t r e a t -16 ment of b o i l e r water. With water, i t forms a monohydrate N^R^.E^O, Hydrazine i s abase s l i g h t l y weaker than NH^OH. I t forms s a l t s with both organic acids and inorganic acids. In the anhydrous form, hydrazine has one s l i g h t l y a c i d i c reacting H atom. 4.2. Oxidation Mechanisms of Hydrazine Oxidation of hydrazinium ions, N H + f has been extensively studied, The products of oxidation depend on the reagent and on the conditions of reaction. Complete oxidation gives nitrogen and water but incomplete oxidation may y i e l d ammonia or varying proportions of ammonia and hydrazoic acid. The stoichiometry i n the oxidation of aqueous hydrazine was 30 31 worked out by Browne and Bray and the i r co-workers. They postulated three d i s t i n c t reactions. N 2 H 5 + 4 e ~ + N 2 + 5 H + ( 1 ) N 2 H 5 + e" + NH 4 + + h N 2 + H + (2) N 2 H 5 + - ^ 2e~ +J5HN3 + \ NH 4 + + | H + (3) Reaction (1) occurs when the ox i d i z i n g agent i s acid iodate, neutral iodine or a l k a l i n e f e r r i c y a n i d e , permanganate and manganese dioxide. Reaction (2) i s approached by using f e r r i c , eerie or manganic ions i n neutral or acid s o l u t i o n . Reaction (3) i s the main contributor when hydrazine i s oxidized i n hot, concentrated s u l f u r i c acid with hydrogen peroxide. Three d i f f e r e n t mechanisms of reactions were proposed by Ki r k 32 and Browne which summarize very well the work achieved before 1928. 17 -(e~+ H+) (I) N 2H 4 NH2-NH e- NH2-NH-NH-NH2 «- 2NH3+N2 Reaction Dimerization Dismutation -2(e"+ H+) (II) N 2H 4-^ *~ NH2-N *~ NH2-N=N-NH2 *- N 2+N 2H 4 Reaction Dimerization Dismutation -2 (e"+ H +) . ( I l l ) N„H. -~ NH-NH — : r — NH=N-NH-NH NH„+HN„ 2 4 „ . . Dimerization 2 . 3 3 Reaction Dismutation These mechanisms i n v o l v e ^ a one electron and two d i f f e r e n t kinds of two-electron primary oxidations, and involve s i x d i f f e r e n t hydronitrogen intermediates. 33 Cahn and Powell investigated the mechanism of f e r r i c ions r e -15 duction by hydrazine s u l f a t e , using i s o t o p i c N . They proposed a d i f f e r e n t mechanism to mechanism ( I ) . The f i r s t product of oxidation i s a hydrazyl r a d i c a l , which then su f f e r s (a) further oxidation, (b) d i s -mutation or (c) dimerization. Dismutation Dismutation NH2-NH-NH-NH2 •> NH^+HN^N-NI^ -* 2NH3+N2 (c) Dimerization / m ^ N + "y (Tetrazone) (IV) N H " ^ ^ ^ N 2H 3 ^ ( 6 ) _ . Oxidation Reaction N ( a ) (Diazene) The production of the hydrazyl r a d i c a l i s the rate-determining step. This r e a c t i o n sequence i s i d e n t i c a l with that given by Higginson 34 and Sutton An a d d i t i o n a l reaction i s the dismutation of two hydrazyl r a d i c a l s , 2ljp3 *> HN=NH+N2H4 i N2 + | N2 H4 ( b ) 18 which regenerates hydrazine and contributes to a stoichiometric r e a c t i o n involving four electrons according to rea c t i o n 3+ + + ?+ 4F + N„H C N„+5H + 4Fe e 2 5 2 Although cupric ion does not react at an appreciable rate with i n acid s o l u t i o n , nor changes the rate of oxidation of hydrazine by f e r r i c ion, i t s presence greatly increases the r e l a t i v e amount of four-electron oxidation. This observation suggests that the re a c t i o n N 2H 3 + Cu 4 4" — » - Cu + + RN = NH + H + i s r e l a t i v e l y f a s t . Of course r e a c t i o n : 3+ + 2+ 2+ Fe + Cu Fe + Cu i s f a s t . 35 Huang investigated the oxidation of hydrazine by Mo (VI) i n the PH range 1.2 and.3.2. Since i n the oxidation by Mo (VI) only N 2 i s formed and has been q u a l i t a t i v e l y detected by mass spectrometry i t appears Mo (VI) i s acting as a two-electron oxidant. However, mechanism (II) proposed by Kirk and Browne and other authors i s not exactly followed as f a r as intermediate are concerned and becomes -(2e~ + 2H + ) (V) N 2H 4 NH = NH N 2 + N ^ Reaction Dismutation 3+ The stoichiometry and k i n e t i c s of Mn reactions with hydrazine and the methylhydrazines i n acid perchlorate s o l u t i o n were studied by 36 Davies and Kustin . Mechanism IV was meanly responsible for the ob-served reaction. Davies compared reactions iy 2Mn 3 + + 2N K r + : — N 4 H 8 ' + + + 2 H + 2 5 I N xx - , X, XX 3 2 4 w"3'2 4 4 and 4Mn 3 + + 2CH 0 „H. + (CH 0) 0 N.H. 2 + + 4Mn 2 + + 4H + 2+ His observation that MnOH i s more re a c t i v e towards protonated hydra-3+ zine molecules than i s Mn suggests that H-atom transfer i s the dominant oxidation mechanism. Methylatipn of hydrazine causes a decrease i n i t s rate of oxidation what suggests that i n t h i s l a t t e r case electron trans-f e r might be expected to be the preferred mechanism. 4.3 Reduction of Manganese Dioxide by Hydrazine Not very much i s known on the reduction mechanism of s o l i d p a r t i c l e s with hydrazine i n a l k a l i n e s o l u t i o n . 37 38 Feitknecht and Brenet showed that y - MnC^ can be r e -duced with i n buffered a l k a l i n e s o l u t i o n (pH = 10). Brenet ob-served very important f a c t s by analysing residues obtained a f t e r d i f f e r e n t times of reduction under the same conditions. It was found that hydrazine i s decomposed according to reaction: N.H. — N„ + 4H + + 4e 2 4 2 The volume of the unit c e l l increased approximately l i n e a r l y with the advance of the reduction. This would be due to the reduction of some 4+ 3+ 2+ Mn to Mn or Mn i n the l a t t i c e and also to formation of OH groups. This reduction i n s o l i d phase i s well known i n the cathodic reduction of T - Mn02» There i s probably no fundamental diffe r e n c e i n the e l e c t r o l y t i c reduction process and the chemical reduction mechanism i n the above chemical reacti o n . In order to explain such a mechanism Brenet suggested that 20 N^H^ molecules were adsorbed, followed by t r a n s f e r of electrons i n the s o l i d phase (no manganese was soluble i n the NaOH buffered a l k a l i n e s o l u t i o n ) . Feitknecht assumed that the reduction occurred i n a s i n g l e phase topochemical r e a c t i o n , electrons and protons migrating through the l a t t i c e according to the formula: MnO + xe" + xH + — - MnO (OH) ^(S) (S) (s) z ~* x ( s ) For x = 2, Mn(0H) 2 i s formed. A comparative study of a, 6 .and y - MnO^ was c a r r i e d out by 39 Brenet . The d i l a t i o n of the unit c e l l and reaction i n s o l i d phase hold for a and y ~ Mn0 2 which are a c t i v e and non stoichiometric forms of Mn02. However, B-Mn02 or p y r o l u s i t e are stoichiometric and do not act i n the same way. It was suggested that a l l reduction reactions occurred at the surface, a progressive transforsuxtion of 3Mn02 to MnO.OH being observed only on the surface. 5. The NH 3 - C0 2 - H 20 System 5.1 Equilibrium i n the Systems NH^-CO^^O For studying the liquid-and-vapor-phase equilibrium of the 40 CO2-NH3-H20 system under high temperature and pressure, Kawasumi gave a d e s c r i p t i o n and d e t a i l s of methods to determine mole r a t i o s of NH^, C0 2 and H 20 i n the vapor phase, the mole r a t i o of NH^ to C0 2 i n the l i q u i d phase, and volume-ratio of the vapor-to liquid-phase i n an . autoclave. Kawasumi expressed the loading density as the r a t i o of the sum of the weights of ammonia and carbon dioxide over the volume of water. Increasing the loading density from 0.2 gm/cc to 1.00 gm/cc at constant 21 temperature and loading mole r a t i o , i . e . 2NH3 f o r 1 CO^, r e s u l t s i n : 1) an increase of the t o t a l pressure 2) a decrease of the mole r a t i o NH^/CO^ i n the vapor phase. The r a t i o remains below two. 3) a mole r a t i o H^O/CO^ i n the vapor phase i n the range of 0.03 to 0.16. 4) an increase i n the mole r a t i o of NH^ to CO^ i n the l i q u i d phase from 2.18 to 2.72 5) a decrease i n the urea y i e l d (thus an increase i n ammonium carbamate). Increasing the temperature at constant loading density r e s u l t s i n : 1) an increase of the t o t a l pressure 2) an enrichment of the gas phase i n NH^ 3) a mole r a t i o H^O/CX^ ranging around 0.02 to 0.12 4) an enrichment of the l i q u i d phase i n NH^ 4 Adding water to the system at constant temperature and loading mole r a t i o of 2NH^ to lCC^ r e s u l t s i n : 1) a constant equilibrium pressure 2) an increase i n the CO^ composition of the vapor phase and the l i q u i d phase 3) a decrease i n the y i e l d of urea i n the l i q u i d phase 4) a constant mole r a t i o H^O/CO^ i n the vapor phase. 42 Kawasumi studied the influence of varying the loading mole r a t i o NH^/CX^ on the equilibrium pressure at a temperature of 160°C with a constant loading density of 0.6 gm/cc. and concluded that: 1) the equilibrium pressure r i s e s with increasing the excess amount of CO2 and reaches 150 atmospheres at a loading mole r a t i o of 1.60. 2) a minimum equilibrium pressure of 65 atmospheres i s reached at a 22 loading mole r a t i o of 2.425. 3) the equilibrium pressure increases with increasing the excess amount of NH^ and reaches 90 atmospheres at a loading mole r a t i o of 3.4. 4) both the r a t i o s NH^ /CO,, and H 20/C0 2 increase i n the vapor phase with increase of the loading r a t i o from 1.4 to 3.4. 43 Later on Blasiak showed that the e f f e c t of water on the pressure depended on the packing degree, defined by the r a t i o of the weight of ammonium carbonate to the volume of the system. Considering us as the r a t i o of water to ammonium carbamate, Blasiak and Baranski ob-served that when OJ v a r i e s from 0 to 1, at 181°C, the pressure v a r i e s from 127 to 116 atmospheres f o r a packing of 0.30 gm/cc , from 165 to 222 atmospheres for a packing of 0.92 gm/cc and remains constant at 172 atmospheres for a packing degree of 0.74 gm/cc. 5.2 K i n e t i c s and equilibrium of the Hydrolysis of Ammonium Carbamate  i n the NH., -CO ~ H 20 System . 5.2.1. Determination of Ammonium Carbamate 44 In 1885 Fenton made measurements on the equilibrium NH 4 C0 2 NH 2 + H 20 = (NH 4) 2 C0 3 (1) His method 'is based on the d i f f e r e n c e i n action of a l k a l i n e hypochlorite and hypobromite s o l u t i o n toward ammonium and amino groups, previous ex-45 periments having shown that the hypochlorite s o l u t i o n l i b e r a t e s nitrogen from ammonium groups only, while sodium hypobromite acts on both ammonium and amino groups. Reaction (1) goes to the r i g h t with increasing temperature, water content and time. 46 Burrow i n a seri e s of experiments involving the p r e c i p i t a t i o n and determination together of the carbamate and carbonate by addition of cold barium hydroxide s o l u t i o n drew attention to the slowness with 23 which the equilibrium between these s a l t s adjusts i t s e l f at low temper-atures, Only the ammonium carbonate reacts on B a C l 2 and the carbamate 47 does not hydrolyse to carbonate at low temperature. Faurholt used this important property i n his study of equilibrium reaction (1) and C0 2 + H 20 = H 2C0 3 (2) 5.2.2. K i n e t i c s and Mechanism of Hydrolysis of Ammonium Carbamate 48 Gorin suggested that the mechanism of the ammonium carbamate hydrolysis reaction (1) may involve, as the rate-determining step, the decomposition of carbamic ac i d , represented by equation (3) NH 2HC0 2 = NH 3 + C0 2 (3) The concentration of carbamic acid i s controlled by the i o n i z a t i o n constant, equation (4) [NH CO."] [H +] K = l- ^ (4) [NH 2 HC0 2] measured by Roughton^ as equal to 2.10 ^ . This value would give a carbamic acid concentration 10 ^ times smaller than the concentration of carbamate i n 0.1 M NaOH so l u t i o n . While the concentration of carbamic acid i s very small, i t s rate of decomposition i s large. In strong a l k a l i n e s o l u t i o n s , the C0 2 that i s l i b e r a t e d i n reaction (3) i s r a p i d l y neutra-l i z e d , so that the reaction i s rate determining. On the other hand, as the hydroxide concentration i s lowered, the n e u t r a l i z a t i o n of C0 2 be-comes slow and rate-determining, and, as C0 2 accumulates, the r e v e r s a l of equation (3) becomes important. 24 47 Faurholt hydrolyzed ammonium carbamate i n an NH^ - NH^Cl buffer s o l u t i o n . In these experiments the pH was probably around 9 and i t was found that the rate of hydrolysis was nearly independent of the NHg - NH^ "*" r a t i o , but inversely proportional to the NH^ con-centration. This, of course was due to the re v e r s a l of reaction (3) . In strongly acid s o l u t i o n , the re v e r s a l of reaction (3) i s prevented by the conversion of NH^ to NH^ "1", and the hydrolysis of carbamate again becomes very f a s t . 5 . 2 . 3 . K i n e t i c s of Carbamate Formation and Breakdown A mechanism for carbamate formation and breakdown was outlined by Caplow"^ which may be expressed by equation (5) k i + K k 2 H 2N + C 0 2 = H2NCO~ = H 2 NC0 2 = H2NCC<2" 4- H 3 0 + (5) i k_ i k_ H 1 H : 2 H : 1 + H-O-H H-O-H In the d i r e c t i o n of synthesis, following carbon-nitrogen bond formation the amine proton i s transferred to a water molecule to give charged products which d i s s o c i a t e at a d i f f u s i o n - l i m i t e d rate. Proton trans f e r occurs within a hydrogen bonded complex and i s exceeding-l y rapid i n the d i r e c t i o n i n which the equilibrium i s favorable. 5 . 3 . Applications of the NH^ - CO^ - System 5 . 3 . 1 . Manganese from low-grade ores Welsh and Ifeterson"^ worked out a new process for the hydro-m e t a l l u r g i c a l extraction of manganese. The process i s known as the "Dean-Leute Ammonium Carbamate Process". It i s based on the fact that 25 manganous oxide, MnO, forms a soluble complex i n aqueous solutions containing about 18 moles per l i t e r NH^ and 3 moles per l i t e r CO^, which was postulated to be 0 - Mn - 0 0 = C NH.+ 4 The complex, when broken down by heat and/or loss of NH^, r e s u l t s i n the p r e c i p i t a t i o n of manganese as a carbonate. The process consists of a primary roasting of the ore i n order to obtain MnO from Mn02, followed by d i s s o l u t i o n i n the ammonium carbamate s o l u t i o n and p r e c i p i t a t i o n of Mn CO^ when the s o l u t i o n becomes saturated. The ore reduction step has to be c o n t r o l l e d i n such a way that i r o n i s maintained as i n s o l u b l e Fe^O^ since any i r o n present at a reduction l e v e l equivalent to FeO i s also s o l u b i l i z e d by the leach l i q u o r . The formation of the carbamate complex gives of heat. This i s important since the complex i s unstable with respect to heat and c o n t r o l must be exercised to hold the temperature within proper l i m i t s . One l i t e r of f u l l strength leach l i q u o r can d i s s o l v e and I | r e p r e c i p i t a t e as much as 110 gm. Mn as Mn , while maintaining a f i n a l I | soluble manganese concentration of about 55 gm. per l i t e r as Mn . The f i r s t MnO taken i n s o l u t i o n raises the temperature, r e s u l t i n g i n a f a s t e r reaction evolving more heat u n t i l a p r e c i p i t a t e of MnCO^ r e s u l t s . This permits more MnO to be dissolved, with the carbamate functioning as a. t r a n s i t o r y intermediate. Any oxygen pick-up by MnO leads to a reduction i n leaching e f f i c i e n c y . Fortunately any FeO present as the complex acts as a mild reducing agent, the i r o n being removed as in s o l u b l e Fe(0H) 3 > Ammonium 26 s u l f i d e i n which the s u l f i d e ion acts as a reducing agent, i s added i n c o n t r o l l e d amounts (0.1 pet or l e s s ) . The manganese i s recovered from the soluble complex by heating the s o l u t i o n to 55-65°C. The ammonia content i s reduced to about 10-11 moles per l i t e r . Heat and loss of NH^ permits the p r e c i p i t a t i o n of v i r t u a l l y a l l the manganese. The MnCO^ (rhodochrosite) has a basic p a r t i c l e s i z e of 1 to 4 y. The extraction of manganese from poor manganese ores by r e -ductive roasting and leaching i n ammoniacal s o l u t i o n was also investigated 52 by Vucurovic . The investigations were aimed at f i n d i n g out optimal con-d i t i o n s of roasting and leaching. Therefore the e f f e c t s of roasting temperature, roasting time, quantity of reducing agent, and leaching conditions were investigated. The maximum % extraction (about 85% of the manganese) was obtained f o r a 90 minutes roasting at 700°C, followed by leaching with 120 gm/l (NH^CX^ and 20 gm/l. of free NH^ f o r one hour and using a s o l i d - l i q u i d r a t i o of 1 to 8. Their f i n a l product, however, contained 4 to 5% of i r o n . 5.3.2. E l e c t r o l y t i c D i s s o l u t i o n of Iron 53 Welsh worked out the p o s s i b i l i t y of d i s s o l v i n g i r o n i n aqueous ammoniacal ammonium carbamate s o l u t i o n with the production of a s o l u t i o n containing i n excess of 45 gm/l of i r o n i n the form of an ammonia complex, which l a t t e r he represented by the formula 0 0 - C - NH 2 Fe (NH ) j n NH 4 + (6) 27 where "n1' i s a whole number greater than one. This complex was found to be stable at higher temperatures, and/or at lower ammonia con-centration, than i s the corresponding manganese complex. It was believed that e i t h e r the number "n" of NH^ groups i s l e s s or else that they are more c l o s e l y bound to the i r o n atom. I t was found that no more than 45 gm/1 of i r o n could be leached from reduced i r o n ore i n a three-day leaching period, whereas solutions containing 90-100 gm/1 of manganese could be produced i n a leaching period of about 15 minutes with a strongly ammoniac al aqueous s o l u t i o n of ammonium carbamate, e.g. with an aqueous so l u t i o n made from 50-300 grams per l i t e r of NH^ and 38.5 - 90.0 grams per l i t e r of CO^. M e t a l l i c manganese r e a d i l y dissolved as MnO, i n the ammoniacal carbamate l i q u o r with rapid evolution of hydrogen, m e t a l l i c i r o n was not observed to go into s o l u t i o n i n a s i m i l a r manner. Welsh discovered that carbamate solutions containing a high concentration of i r o n i n complex can be produced by the a i d of anodic d i s s o l u t i o n of i r o n metal i n the carbamate l i q u o r . The d i s s o l u t i o n of i r o n proceeds with evolution of hydrogen from the cathode of the e l e c t r o -l y t i c c e l l . The process can be summarized as follows: 1) The anodic d i s s o l u t i o n of i r o n i n ammonium carbamate s o l u t i o n progresses i n a normal way i f the current density i s maintained below a c r i t i c a l maximum, which l a t t e r i s r e l a t e d to the formation of a highly passive oxide f i l m on the i r o n anode surface r e s u l t i n g i n the l i b e r a t i o n of oxygen, rather than i n the d i s s o l u t i o n of i r o n . Passivation of the anode surface -when the same occurs - can be removed by reversing the d i r e c t i o n of the current for 30 seconds or l e s s . In order to maintain r e l a t i v e l y high current de n s i t i e s (10-30 amperes/sq.ft.) the i r o n i n s o l u t i o n was l i m i t e d 28 to 40-70 g m / l . 2) Iron carbonate, FeCO^, can be p r e c i p i t a t e d from the s o l u t i o n , by lowering the ammonia concentration of the l a t t e r to 8 mols/1 or l e s s . Preferably, the complex - containing s o l u t i o n i s heated to 85°C to e f f e c t the disengagement of ammonia. When p r e c i p i t a t e d out of access to a i r the carbonate i s whitish, but upon drying by heating i n a i r the carbonate decomposes and a very f i n e l y divided, active form of very pure hydrated f e r r i c oxide i s produced. The p a r t i c l e s are s i n g u l a r l y uniform as to s i z e , about 3 to 4 microns i n diameter. 3) By bubbling a i r through the s o l u t i o n instead of heating, the pre-c i p i t a t e consists e s s e n t i a l l y of red Fe-COH)^. In t h i s way i t was made possible to produce pure i r o n oxide from a scrap i r o n anode. 29 C. SCOPE OF THE PRESENT INVESTIGATION As already o u t l i n e d , previous workers have studied the leaching of ferrous and manganous oxides i n aqueous ammoniacal ammonium carbamate solutions at room temperature. In a l l the cases a previous reductive roasting was c a r r i e d out to reduce both goethite and manganese dioxide to FeO and MnO r e s p e c t i v e l y . In t h i s work i t was proposed to study the p o s s i b i l i t y of leach-ing iron-manganese ore at elevated temperature i n NH^ - CO^ - H^ O solutions without employing a preliminary reductive roast. The favorable s h i f t of the carbonate-carbamate equilibrium to the carbamate side observed at room temperature had to be confirmed f o r elevated temperatures. It was then necessary to make a preliminary study of the d i s -s o l u t i o n of MnO and FeO i n carbamate s o l u t i o n at elevated temperature. This being s a t i s f a c t o r y , a s u i t a b l e reducing agent had to be found. I t was intended to study the use of organic reducing agents such as methanol, formaldehyde, and formic acid and inorganic reducing agents such as s u l f i t e and hydrazine. The mechanism of the reduction of p y r o l u s i t e with hydrazine was then studied. F i n a l l y , the recovery of manganese from s o l u t i o n with oxygen was investigated. EXPERIMENTAL A. MATERIALS 1. Reagents Reagent-grade chemicals were used e x c l u s i v e l y . Helium and oxygen were supplied by Canada L i q u i d A i r Ltd. 2. Minerals 2.1 Goethite The goethite mineral was purchased from Minesota . The chemical analysis of the goethite mineral i s given i n Table 4. TABLE 4 Chemical Analysis of Goethite Element Weight % Element Weight % S10 2 27.41 Mg 0.2 FeO.OH 70.61 Mn 0.3 A l 0.3 Mo 0.005 Ca 0.4 V 0.01 Cr 0.001 Cu 0.02 Pb 0.001 2.2 P y r o l u s i t e ro P y r o l u s i t e ore was obtained f^jpn New Mexico, Lake Valley. The chemical analysis of the ore i s given i n Table 5. 31 TABLE 5 Chemical Composition of P y r o l u s l t e Element Weight % Mn0 2 75.20 sio 2 5.89 Fe 0.02 CaO 3.84 P 0.045 (Al, Cr, Co, Cu, Pb, Mg, 15.0 Mo, N i , Zn, HO) The elements which are of importance f o r th i s study are manganese, s i l i c o n , i r o n and phosphorus. In Table 6 i s given a X-ray pattern of the p y r o l u s i t e ore which was used i n the experiments, compared with reported X-ray data. TABLE 6 X-ray D i f f r a c t i o n Pattern  of P y r o l u s i t e Reported This study o o dA 20 ^ i m dA 20 ^ i m 3.14 28.40 100 3.14 28.40 100 2.41 37.28 50 2.41 37.28 90 2.13 42.40 85 2.12 42.60 50 1.63 56.40 50 1.64 56.78 50 1.56 59.18 25 1.55 59.58 20 1.31 72.00 20 1.31 71.90 100 32 B. AUTOCLAVE DESIGN A 1 l i t e r s t a i n l e s s s t e e l (316 ss ) autoclave with a glass l i n e r was used. S t i r r i n g was provided by a Magnedrive u n i t , supplied by Auto-clave Engineers. This unit had a completely enclosed s t i r r i n g shaft, which was activated by means of an external r o t a t i n g magnet. The s t a i n -less s t e e l s t i r r i n g shaft was sheathed by Teflon rod, d r i l l e d to give a good f i t over the shaft. A g i t a t i o n was provided by a Te f l o n block which was threaded on to the bottom of the s t i r r e r shaft sheath. The s t i r r i n g speed used for most of the runs was 900 R.P.M. unless otherwise in d i c a t e d . V i s u a l tests showed that t h i s s t i r r i n g speed, plus the e f f e c t of the b a f f l e s (one on the sampling tube, another being the thermo couple s t a i n -less s t e e l i n l e t i n the s o l u t i o n i t s e l f ) provided extremely turbulent a g i t a t i o n of the s o l u t i o n . The thermocouple was an iron-constan.tan type, s u i t a b l e f o r use over a temperature range from 0 to 300°C. A Wheelco range 0 - 300°C ; temperature c o n t r o l l e r activated a solenoid valve connected to a c i r c u l a r gas burner. Using t h i s system, the temperature was co n t r o l l e d to wi t h i n - 2°C (due to the many contacts: s o l u t i o n - t e f l o n - thermocouple tubing, thermocawple wires). For some experiments at low temperature a constant manual control was preferable to an automatic control of the temperature. The pressure gauge, 0 - 1000 p s i from Autoclave Engineers Inc. was connected to the gas i n l e t tube, on the pressure side of the valve. An exhaust tube and valve were mounted separately. Sampling was ca r r i e d out through a sintered glass f i l t e r f i t t e d on the i n l e t end of the tube; the external portion of the sample tube was water cooled. 33 A l l tubing was 1/8" 316 ss. C. EXPERIMENTAL PROCEDURE The experimental procedure consisted of the following steps: 1) A two or f i v e gram sample of mineral, 480 cc of 15 Normal NH^OH and 140 gr. of ammonium carbamate, plus eventually a v a r i a b l e amount of re-ducing agent, to give . a s o l u t i o n volume of 600 cc were put into the autoclave. 2) The ves s e l was sealdd, the s t i r r e r started. 3) The gas i n l e t valve and the exhaust valve were opened and a f a s t flow of helium gas was flashed through the gas phase i n order to expel a l l oxygen of the gas phase. The system was closed a f t e r two minutes. 4) The beating was started and the s o l u t i o n temperature was rai s e d at a rate of 5°C per minute up to the desired temperature. 5) A f i r s t sample was taken as a blank for the c a l c u l a t i o n s of the rate curve f o r the run. 6) Samples f o r i r o n , or manganese and/or phosphate determination were taken at regular i n t e r v a l s throughout a run, usually f i v e samples per run. Sampling procedure involved f i r s t taking and discarding 15 ml of so l u t i o n i n order to clean the sampling tube, then further 5 ml sample was taken and used f o r the Mn, Fe or P determination. Corrections were applied to the r e s u l t s for the change i n the t o t a l volume by expressing the amount of Fe, Mn or P dissolved rather than the amount of each of them i n s o l u t i o n . 7) For i r o n analysis the sample was neutralized with hydrochloric a c i d . 34 8) For manganese and/or phosphate and i r o n the sample was neutralized with n i t r i c a c i d . The i r o n and manganese complexes are unstable with respect to oxygen and water so that p r e c i p i t a t i o n of Fef(OH) 3 or MnC^ occurs r e -l a t i v e l y f a s t outside the autoclave. The p r e c i p i t a t i o n i s accelerated by rapid loss of NH^ into the a i r . By a c i d i f y i n g the 5 ml samples both Fe and/or Mn return to s o l u t i o n . The a c i d i f i e d solutions were made up to 50 ml with water. D. ANALYTICAL METHODS 1. Iron 54 Iron was estimated by a co l o r i m e t r i c method using the orange-red ferrous complex of 1 - 10 orthophenanthroline. Hydroxylamine hydrochloride was added to ensure that a l l the i r o n i n the s o l u t i o n was reduced to the ferrous state. Buffering of the solutions at pH = 4.5 was obtained with a sodium acetate - ac e t i c acid buffer s o l u t i o n . A composite reagent was prepared by mixing 200 cc of a 0.15% orthophenanthroline s o l u t i o n , with 200 cc of a 1% hydroxylamine hydrochloride s o l u t i o n and with 600 cc of acetate b u f f e r s o l u t i o n . 5 ml aliquot's of the neutralized sample s o l u t i o n were pipetted into 90 cc of reagent s o l u t i o n and made up to 100 ml with water. The o p t i c a l density of each s o l u t i o n was measured on a Beckman Model B Spectrophotometer, using l i g h t of wavelength 510 mu. The concentration of i r o n was read d i r e c t l y from a c a l i b r a t i o n curve prepared by using standard i r o n s o l u t i o n s . 2. Manganese 2+ - 55 Mn was oxidized to MnO^ by potassium periodate Phosphoric acid a d d i t i o n prevented p r e c i p i t a t i o n of Mn0„ or manganese periodate 35 and decolorized the f e r r i c s o l u t i o n while s t a b i l i z i n g MnQ^ . The o p t i c a l density was measured at a l i g h t wavelength of 524 my. Some ions do i n t e r -f e r e , e s p e c i a l l y Mo (VI), V (V), T i (IV), and Cr (VI). A composite reagent was prepared by mixing 100 cc of con-centrated s u l f u r i c acid with 100 cc of concentrated phosphoric acid and with 800 cc of a 0.5% potassium periodate s o l u t i o n . 5 ml aliquots of the neutr a l i z e d samples were added to 90 ml of the reagent s o l u t i o n ; to enable 4 + -Mn to oxidize to MnO^ i t was necessary to heat up the s o l u t i o n to 90°C for f i v e minutes, the presence of hydrazine slowed down the oxidation of ++ Mn and more than 5 minutes were needed. A f t e r complete oxidation of I j _ Mn to MnO^ the s o l u t i o n was made up to 100 ml with water. The con-centration of manganese was read d i r e c t l y from a c a l i b r a t i o n curve pre-pared by using standard manganese sol u t i o n s . 3. Phosphate 56 Phosphate was measured by a colo r i m e t r i c method using a yellow phosphate complex formed i n presence of V (V) and Mo (VI) i n a c i d i c 3+ so l u t i o n . Some ions do i n t e r f e r e such as Fe and Cr (VI). Iron can be complexed with F . A composite reagent s o l u t i o n was prepared by mixing 100 cc of 6N n i t r i c acid with 100 cc of a 0.25% ammonium vanadate solution(2.5 g of ammonium vanadate i n 800 cc of water plus 200 cc of concentrated n i t r i c acid) and with 100 cc of a 5% ammonium molybdate s o l u t i o n . The t o t a l 300 cc of s o l u t i o n i s made up to 1000 cc with water. 5 ml aliquots of the neutralized solutions were added to 90 ml of reagent s o l u t i o n and made up to 100 ml with water. Af t e r two hours of s t a b i l i z a t i o n the o p t i c a l density was read on a Beckman Model B. Spectrophotometer, using l i g h t of wavelength 460 mu. The concentration of phosphorus was read d i r e c t l y from a c a l i b r a t i o n curve prepared by using standard phosphate s o l u t i o n s . 4. Ammonium Carbamate 47 45 Both Faurholt and Fenton re l a t e d methods f o r ammonium carbamate determination. However, Faurholt gave the most convenient method. In t h i s study the NH^ - CO^ - H^ O solutions were prepared using ammonium hydroxide, water and ammonium carbamate. Considering equilibrium (NH 4) 2 C0 3 = NH 4 C0 2 NH 2 + H 20 Faurholt found that NH, C0„ NH„ was not hydrolysed at 0°C even when a l l 4 2 2 (NH 4) 2 CO^ was removed from s o l u t i o n . He showed that only (NH 4) 2 CO^ p r e c i p i t a t e d out when Ba C l 2 was added to the system (at 0°C) giving a white Ba CO^ compound. To test the ammonium carbamate y i e l d obtained at higher temper-ature a known sample volume was added to an i c e cold Ba C l 2 - water s o l u t i o n . A l l (NH 4) 2 CO^ present reacted immediately with Ba C l 2 , leaving NH 4 C0 2 NH 2 i n s o l u t i o n . A f t e r f i l t e r i n g and washing Ba CO^ was dissolved i n an excess of HC1 and the excess back t i t r a t e d by Na OH. This method was found to give reproductible r e s u l t s with l e s s than 2% error . E. SYNTHETIC MINERALS 1. Ferrous Oxide FeO was purchased from Rocky Mountain Research Incorporation, o Denver, Colorado. 2. Manganous Oxide MnO was prepared from Mn0 2 powder by heat treatment at 650°C under cracked, ammonia atmosphere f o r 24 hours. The percentage of MnO ob-tained was estimated by weight l o s s . I t was concluded that v i r t u a l l y a l l MnO was reduced to MnO under these conditions. RESULTS AND THEIR ANALYSIS A. THE NH„ - C0„ -H-0 SYSTEM j L ~L 1. E f f e c t of Temperature on the Equilibrium H 20 + NH^ C0 2 NH^ = In ammonia-carbon dioxide-water solutions ammonium carbamate i s know to be i n equilibrium with both ammonium carbonate and urea. The r e l a t i v e quantity of each of them i n the system depends mainly on the composition of the s o l u t i o n , i n other words on the proportion of 40 41 42 NH 3 to C0 2 and to H 20 H V : (NH 4) 2 C0 3 = NH 4C0 2NH 2 + HO = NH 2 C0NH2 + 2H20 (7) For low ammonia concentration and high water content the equilibrium (7) i s s h i f t e d to the l e f t whereas f o r high ammonia con-centration and high carbon dioxide content the equilibrium i s s h i f t e d to the r i g h t . To avoid urea formation the water content of the system must be held high enough. A l l r e s u l t s i n th i s work were obtained using solutions con-taining 18 moles per l i t e r of NH 3 and 3 moles per l i t e r of C0 2 since, 43 as mentioned by Davies and Welsh , a high NH 4 C0 2 NH 2 y i e l d i s then obtained i n s o l u t i o n . Since the experiments were c a r r i e d out at high temperature and pressure i t had to be v e r i f i e d that the equilibrium (7) remained to the side of ammonium carbamate i n order to eventually maintain the s o l u b i l i t y of FeO and MnO. The percentage of ammonium carbamate (for a maximum of 3 mol« 39 per l i t e r ) versus . temperature i s p l o t t e d i n Figure 3. At constant NH^ and H^ O content of the s o l u t i o n the ammonium carbamate y i e l d decreased only s l i g h t l y when temperature increased from 20°C to 130°C. Figure 4 shows a p l o t of the percentage of ammonium carbamate versus the ammonia content of the s o l u t i o n , at 130°C and with a constant 3 moles per l i t e r C0 2 content. In order to obtain a high NH^ C0 2 NH 2 y i e l d i t was necessary to introduce an as high as possible ammonia amount i n s o l u t i o n . Only 60 per cent NH^ C0 2 NH^. (1.8 moles/1) was obtained f o r 6 moles/1 of NH 3 whereas 85% N'H^  C0 2 NH 2 (2.55 moles/1) was reached f o r 18 moles per l i t e r of NH^. 2. Equilibrium Pressure at High Temperature i n the System NHo-C0?-H 0 9 - 1 8 moles per l i t e r NH^ are plotted i n Figure 5. A pressure of 300 p s i was reached at 130°C. The gas phase was mainly composed of NH Only a l i t t l e CO and HO could be detected at 130°C. (7). At room temperature, equilibrium was obtained a f t e r 5 hours when ammonium carbamate was used at the s t a r t and only a f t e r 30 hours when ammonium carbonate was used at the s t a r t ? (Figure 6). Therefore, f o r a l l leaching experiments ammonium carbamate was taken rather than ammonium carbonate. The d i s s o l u t i o n of ammonium carbamate was endothermic so that the temperature dropped from 22°C to 10°C i f no heat was supplied. B, REDUCTIVE LEACHING OF GOETHITE 1. Leaching of i r o n (130°C) The leaching of iron was investigated to check i f there would be a corrosion problem since a s t e e l autoclave was to be used. Pressure curves corresponding to 3 moles per l i t e r CO and Equilibrium pressures are re l a t e d to equilibrium r e a c t i o n 100 80 ( NH4CO2NH2) mole % go 4 0 20 -® ——@— © 9 M/l NH. O 18 M/l NH, 100 % = 3 M/l N^COgNHg JL 50 100 Temperature (°C) 150 Figure 3 : E f f e c t o f temperature o n the equilibrium N H 4 C 0 2 I N T H ? + H ? 0 = = = ( N H 4 ) 2 C O 100 ( N H 3 ) ( M/l ) Figure h : E f f e c t of NH3 concentration on the equilibrium NH 4 C 0 2NH 2 + H 2 0 === (WH 4 ) 2 C 0 3 0 5 0 100 Temperature (°C) ^ to Figure 5 : Equilibrium pressure versus temperature. mole 7c IH4C02NH2 s O o starting with (NH ) CO 4 2 3 ° " " NHjOO^Hj 25 50 100 Time ( hours) Figure 6 : Mole percent of ammonium carbamate i n so l u t i o n versus time, 44 When i r o n f i l i n g s were sealed i n the autoclave with the carbamate s o l u t i o n a large over-pressure developed. Analysis of the gas phase showed that the overpressure was due to hydrogen. The amount of i r o n dissolved from 5 gm of f i l i n g s and n ^ / n ^ ( r a t i o of the number of moles of i r o n dissolved to the number of moles of B.^ present i n the vapor phase) are p l o t t e d i n Figure 7. The r a t i o n ^ / n ^ became approx-imately equal to one a f t e r a long leaching period. F i r s t , the reaction of i r o n with ammonium carbamate and water was slow but i n i t s l a t e s t stages i t appeared to be catalysed by the presence of the ferrous i r o n complex. The construction of the autoclave was modified so that no s t e e l portions were exposed to the carbamate s o l u t i o n . 2. S o l u b i l i t y of FeO 2.1. E f f e c t of temperature Experiments c a r r i e d out at room temperature using Rocky Mountain FeO and under He atmosphere flow enabled d i s s o l u t i o n of 61 g/1 of FeO a f t e r about 30 hours. (48 g/1 of i r o n ) . In the same time an NH^ loss of .5% could not be avoided. At l e a s t 70% and l e s s than 90% of the carbon dioxide content of the s o l u t i o n would have been present as carbamate i n the system (70% i s obtained for 9 M/l NH^ as seen i n A). A f t e r 30 hours the s o l u b i l i t y of FeO decreased because of fu r t h e r loss of ammonia (Figure 8). Welsh"^ mentioned that even under the most favorable conditions i t was not possible to obtain more than 45 g/1 of i r o n from ferrous oxide. Unfortunately, h i s favorable conditions were not delineated. By e l e c t r o l y s i s of a scrap i r o n anode Welsh was able to dissolve as much as 80 g/1 of i r o n i n complex form i n aqueous ammoniacal ammonium carbamate s o l u t i o n at room temperature. 5 F i g u r e 7 : D i s s o l u t i o n o f m e t a l l i c i r o n i n t h e sys t e m M 3 - C O P - H 2 0 . g/ l Fe dissolved Ti me ( h o u r s ) Figure 8 : S o l u b i l i t y of FeO i n l8 M./liter NH3 - 3 M / l i t e r C 0 2 s o l u t i o n at room temperature. ( open system , He atmosphere ). ON 47 This suggested that f a i l u r e to achieve:saturation of the carbamate sol--ution of ferrous i r o n when using FeO as s t a r t i n g material was most probably due to a decreasing rate of d i s s o l u t i o n as the i r o n content of the s o l u t i o n increased. This argument was supported by experiments car r i e d out i n t h i s work at elevated temperature i n an autoclave. Already at 40°C a s o l u b i l i t y of 70 g/1 of i r o n was detected from the d i s s o l u t i o n of excess FeO. However, i t was found that further increase of temperature up to 130°C did not increase the s o l u b i l i t y of FeO which remained constant at about 70 - 75 g/1 of i r o n . (Figure 9). 2,2. E f f e c t of Adding Sodium Sulfate at 130°C As shown i n Figure 10 the s o l u b i l i t y of FeO was not affected by the presence of s u l f a t e at 130°C, even when 0.58 moles/1 of s u l f a t e were added f o r .only 0.34 moles/1 of FeO. 2.3 E f f e c t of Adding Sodium S u l f i t e at 130°C An experiment was c a r r i e d out taking 0.13 moles/1 of sodium s u l f i t e f o r 0.115 moles/1 of i r o n at 130°C. Only 0.16 g/1 of i r o n went i n s o l u t i o n a f t e r two hours, for an introduced amount of 8.3 g/1. A X-ray d i f f r a c t i o n pattern of the residue did not give any clear i n -formation as to i t s composition. However, the i n i t i a l FeO structure had disappeared, suggesting that FeO had been dissolved and p r e c i p i t a t e d i n a d i f f e r e n t form. 3. Reductive Leaching of Goethite at 130°C As mentioned i n A a l l the runs were c a r r i e d out using aqueous ammoniacal ammonium carbamate solutions containing 18 moles/l NH^ and 3 moles/1 C O 2 . Unless otherwise mentioned, 2 g of goethite ore s i z i n g 65-150 mesh and wet screened were used f o r a l l the runs. The volume of 100 g/l Fe d issolved 7 5 U 50 h 2 5 50 _L 100 Tem perature ( °C) 150 Figure 9 : E f f e c t of temperature on the s o l u b i l i t y of FeO i n the system N H 3 - C 0 2 -H 2 0 . g /I Fe dissolved ( N H 4 ) 2 S O 4 I g ) Figure 1 0 : E f f e c t of adding sulfate on the s o l u b i l i t y of FeO at 130 0 C, 50 the s o l u t i o n was 600 cc. Reducing agents such as formaldehyde (HCOH) methanol (CH^OH) and formic acid (HC00H) gave no ferrous i r o n i n s o l u t i o n . S u l f i t e i n the presence of a ca t a l y s t and hydrazine alone yielded i r o n i n s o l u t i o n . 3.1 Reductive Leaching of Goethite Using Sodium S u l f i t e i n Presence of  Anthraquinone Catalyst Compounds Although sodium s u l f i t e i s known as a powerful reducing agent no reduction of goethite occurred when sulfite alone was added. However, copper ions or anthraquinone acted as c a t a l y s t s and a l i m i t e d r e a c t i o n took place. 3.1.1. E f f e c t of Sodium S u l f i t e Concentration Two grams of goethite were leached with 1, 2, 4 and 6 g of sodium s u l f i t e and a constant 0.1 g anthraquinone i n 600 ml of s o l u t i o n . As seen i n Figure 11, the amount of goethite leached i n a given, time increased with increasing s u l f i t e up to 4 g of s u l f i t e then s u l f i t e had no further e f f e c t . The parabolic shape of the d i s s o l u t i o n curves suggested that leaching and p r e c i p i t a t i o n phenomena might be competing. The more s u l f i t e was added, up to about 4 g, the f a s t e r goethite was i n i t i a l l y reduced. However, the more complexed i r o n and s u l f i t e ions were i n s o l u t i o n the fa s t e r i r o n appeared to be p r e c i p i t a t e d out of s o l u t i o n . 3.1.2. E f f e c t of AnthraquinoneConcentration The e f f e c t of adding 0.1, 0.2, 0.3 and 0.4 g of anthraquinone to 600 cc of s o l u t i o n was studied with a constant 6 g amount of sodium s u l f i t e , D i s s o l u t i o n rates increased with increasing anthraquinone mg Fe dissol ved 100 2 00 3 0 0 Time ( min.) F i g u r e I I : E f f e c t o f s u l f i t e c o n c e n t r a t i o n on t h e d i s s o l u t i o n o f g o e t h i t e a t c o n s t a n t a n t h r a q u i n o n e c o n c e n t r a t i o n . 52 concentration (Figure 12) but not proportional/. The log^log p l o t between the amount of i r o n dissolved and the amount of anthraquinone at the same leaching time led to the following r e l a t i o n (Figure 13). [Fe complexed] = k [anthraquinone]^*^ (8) 3.1.3. E f f e c t of the Nature of the Catalyst A run using 0.4 g of anthraquinone and a run using 0.6 g of sodium 3 sulfonate anthraquinone were ca r r i e d out i n the presence of 4 g of sodium s u l f i t e and 600 cc of s o l u t i o n . The leaching curves are p l o t t e d i n Figure 14. Although the same mole concentration of c a t a l y s t was used i n both runs the i n i t i a l d i s s o l u t i o n rate when 3 sulfonate anthraquinone was used was s u b s t a n t i a l l y l a r g e r than when anthraquinone was used. This suggested that anthraquinone dissolved slower than sodium {$ sulfonate anthraquinone. 3.1.4. E f f e c t of Leaching the Residue The residue of a previous run was washed with water and leached under the same conditions with 4 g of sodium s u l f i t e and 0.6 g of sodium 3 sulfonate anthraquinone i n 600 ml of s o l u t i o n . This reproduced the same leaching curve. (Figure 15). I t was concluded that f i n e p a r t i c l e s within the goethite were not the reason of the large i n i t i a l d i s s o l u t i o n rates of goethite. 3.1.5. E f f e c t of Doubling the Amount of Goethite A run was c a r r i e d out taking 4 g of goethite, 0.6 g of sodium 3 sulfonate anthraquinone and 4 g of sodium s u l f i t e i n 600 cc of s o l u t i o n . o 0 .1 g anthraquinone I i i i 1 1 1— 0 100 2 0 0 300 Ti m e (min . ) Figure 12 : E f f e c t of anthraquinone concentration on the d i s s o l u t i o n of goethite at constant s u l f i t e concentration and at 130 ° C. -0.5 Log ( mg Fe) - 1 . 0 o 9 0 minutes • 150 A 3 0 0 00 - 0 . 7 5 - 0 . 5 0 Log ( g anthraquinone) Figure 1 J : Log-log plot'between anthraquinone concentration and the amount of iron dissolved at 130 °C. -p-200 -mg Fe ±- u dissolve d A anthraquinone 10 0 A Na? sulfonate anthraquinone 0 • i l l 0 100 2 00 300 400 5 00 Time (min . ) Figure Ik : E f f e c t of the s o l u b i l i t y of the c a t a l y s t . Figure 15 : Leaching of goethite with s u l f i t e at 130 °C. E f f e c t of adding s u l f a t e , of doubling the i n i t i a l amount of goethite and of r e d i s s o l v i n g the residue. The rate at which i r o n appeared i n s o l u t i o n was doubled f o r a same amount of i r o n i n s o l u t i o n (Figure 15). This suggested that the rate of the leaching process was surface c o n t r o l l e d . 3.1.6. E f f e c t of Adding Sodium Sulfate at the Start Adding s u l f a t e had no e f f e c t on the d i s s o l u t i o n rate of goethite (Figure 15). 3.1.7. E f f e c t of Adding Sodium S u l f i t e i n the Middle of the Run In order to check i f s u l f i t e had decomposed by a side r e a c t i o n a f t e r a long leaching period the same amount of sodium s u l f i t e as at the s t a r t of the run was added i n the middle of the run. (Figure 16). Only a s l i g h t l y increased rate was observed. This confirmed that s u l f i t e was not decomposed during the run, 3.1.8. E f f e c t of Adding Sodium g Sulfonate Anthraquinone i n the Middle of the Run Sodium g sulfonate anthraquinone could have been reduced so f a r by s u l f i t e as to lose i t s c a t a l y t i c power. Therefore the same amount of sodium g sulfonate anthraquinone was added i n the middle of the run as at the beginning. Figure 16 shows that the d i s s o l u t i o n rate was s l i g h t l y increased which confirmed that the c a t a l y s t was s t i l l a c t i v e . The f a c t that leaching of washed residue led to the same d i s -s o l u t i o n curve, and that s u l f i t e and the c a t a l y s t remained unaltered during the leaching of goethite suggests that the surface of the goethite p a r t i c l e s changed" i n nature during the leaching. 200 m g Fe dissolved 100 A, : 6g N o 2 S 0 3 , 0. 6 g • + 0.6 g a nth. B, B, onthro quinone = A + 6g Na 2 S0 3 100 200 3 0 0 4 0 0 Tim© (min.) 5 0 0 Figure 16 : E f f e c t of adding s u l f i t e or sodium sulfonate anthraquinone i n the middle of the run oh the leaching of goethite with s u l f i t e at 13o °C 00 59 3.1, Reductive leaching of Goethite Using Hydrazine Sulfate at 130°C 3.2.1. E f f e c t of Hydrazine Sulfate Concentration The e f f e c t of 1, 2 and 4 g of hydrazine s u l f a t e on the leaching of goethite was investigated (Figure 17). A parabolic type of leaching was observed. The d i s s o l u t i o n rate of goethite decreased to a low value although only 10 to 20% of the i r o n had been dissolved. A log-log p l o t between the i r o n concentration i n s o l u t i o n and the hydrazine s u l f a t e concentration l e d to the following r e l a t i o n at the same leaching time. (Figure 18). 2 [Fe complexed] = k [N^.H^O^] 3 (9) 3.3. Reductive Leaching of Goethite Using Hydrazine Hydrate at 84°C A s t r a i g h t l i n e r e l a t i o n s h i p was obtained between the amount of ir o n dissolved and time when hydrazine hydrate was used (Figure 19). This suggested that the HSO^ r a d i c a l of hydrazine s u l f a t e was associated i n some way with the decreasing rates of leaching observed when using that compound. It should be mentioned at t h i s point that although a large amount of goethite with a large excess of hydrazine hydrate were sealed i n the autoclave the goethite d i s s o l u t i o n rate was very low at 8^°C. Indeed, only 0.72 g of i r o n from 18 g dissolved i n s o l u t i o n a f t e r a four hour leaching time which was equivalent to 4% of the added goethite. At 35°C only 0.1% of the i r o n was leached from goethite a f t e r one hour when hydrazine hydrate was used. The very small d i s s o l u t i o n rate of goethite at low temper-ature w i l l be of.great i n t e r e s t i n the d i s s o l u t i o n of i r o n containing manganese acide.with hydrazine hydrate. 150 A • i g 2g N^ 4 H 2 S0 4 II mg Fe 0 4 g u dissolv ed 1 00 5 0 0 J 1 • 1 I 1 1 I I , 1 0 10 0 2 0 0 3 0 0 Time ( min.) Figure 17 : E f f e c t of hydrazine sulfate concentration on the leaching of goethite at 130 °C. 1.25 L o g (mg Fe ) 1.00} 0.75 0.50 A 2 7 0 minutes A 150 " „ 9 0 • 0.25 0.0 0.3 0.6 Log(g N2H4.H2S04) 0.9 Figure 1 8 : Log-log pl o t between hydrazine concentration and the amount of i r o n dissolved at 1 3 0 °C. Constant time p l o t s . 63 C. REDUCTIVE LEACHING OF PYROLUSITE Five grams of p y r o l u s i t e wet screened to 65 - 150 mesh were used i n 600 ml of s o l u t i o n i n a l l the runs unless otherwise indicated. The temperature at which the runs were c a r r i e d out w i l l be indicated. As always the s o l u t i o n contained 18 moles/1 NH^ and 3 moles/1 C0 2. 1. S o l u b i l i t y of MnO 1.1. E f f e c t of Time at Constant Temperature In each of the runs the f i n a l temperature was reached as quickly as pos s i b l e . Then equilibrium was allowed to e s t a b l i s h . For each run 15 g of MnO were added to 600 ml of s o l u t i o n ( i . e . 19.35 gm/1). Since MnO was r e a d i l y soluble above 40°C a l l MnO was dissolved during the heating. However, the s o l u b i l i t y of MnO decreased with temperature leading to quick p r e c i p i t a t i o n of MnCO^ between 80°C and the f i n a l temperature. As seen i n Figure 20 some time i s needed to achieve a stable equilibrium of soluble manganese at the f i n a l temperatures. 1.2. E f f e c t of Temperature The s o l u b i l i t y of MnO decreased with increasing temperature (Figure 21). In a closed system the s t a b i l i t y region of the manganese complex became narrower with increasing temperature so that carbonate was p r e c i p i t a t e d out of s o l u t i o n (manganese carbonate with the rhodochrosite structure was i d e n t i f i e d by X-ray analysis and by CO^ a n a l y s i s ) , For each mole of MnCO^ p r e c i p i t a t e d one mole of NH^ C0 2 NH 2 i s used. Experiments were c a r r i e d out with d i f f e r e n t i n i t i a l amounts of synthetic MnO. At low temperature (less than 50°C) a l l the MnO was usually soluble. The maximum s o l u b i l i t y was found to be 70 to 75 g/1 of manganese, Initial conditions: O 0.5 M/l MnO © 1.27 M/l MnO 0 50 100 150 , T©mperoture (° C ) Figure 21 : S o l u b i l i t y of MnO with temperature. E f f e c t of the i n i t i a l amount of MnO, of adding s u l f i t e or s u l f a t e . 66 By comparing d i s s o l u t i o n curves I and XI i n Figure 21 i t was found that the s o l u b i l i t y of manganese at elevated temperature depended on the i n i t i a l amount of MnO which was added. This d i f f e r e n c e could be ex-plained by the following considerations! 1) In run I 0.5 moles/1 of synthetic MnO were added to the carbamate so l u t i o n . P r e c i p i t a t i o n of manganese carbonate occurred when the tem-perature at which a maximum of 0.5 moles/1 of MnO were soluble was ex-ceeded (between 95 and 100°C, Figure 21). Then by going to higher temperatures a maximum of 0.5 moles/1 of MnCO^ could be p r e c i p i t a t e d corresponding to a destruction of 0.5 moles/1 of ammonium carbamate. 2) In run II 1.27 moles/1 of synthetic MnO were used. P r e c i p i t a t i o n of MnCO^ occurred at a lower temperature than i n run I (between 60 and 65°C). Therefore, the s o l u b i l i t y of MnO at elevated temperature was smaller when the i n i t i a l amount of MnO introduced i n s o l u t i o n was larger due to ammon-ium carbamate depletion. 1.3 E f f e c t of Adding Sulfate The s u l f a t e content of the leaching l i q u o r had no e f f e c t on the s o l u b i l i t y of MnO (Figure 21). 1.4 E f f e c t of Adding S u l f i t e A run was carried out using 1.33 moles/1 of MnO and 1.2 moles/1 of sodium s u l f i t e (Figure 21). S u l f i t e had no influence on the s o l u -b i l i t y of MnO below 50°C although more than 1 mole/1 of MnO had been dissolved i n s o l u t i o n . Above 50°C the MnO s o l u b i l i t y decreased f a s t . An X-ray d i f f r a c t i o n pattern of the residue showed that both hydrated ammonium manganese s u l f i t e (NH.)- Mn(SO_)„. H O and manganese carbonate were present. Both p r e c i p i t a t i o n of ( N H ^ Mn ( S Q 3 ) 2 . H 20 and Mn C0 3 lowered the ammonium carbamate content of the s o l u t i o n , the former by diminishing the ammonia content of the liquor, the l a t t e r by decreasing the carbon dioxide content of the s o l u t i o n . I t i s not known i f both compounds co p r e c i p i t a t e or p r e c i p i t a t e separately. However, since no s u l f i t e was l e f t i n s o l u t i o n at 130°C i t i s assumed that (NH^) 2 Mn(S0 3) 2. H 20 p r e c i p i t a t e d f i r s t inducing c o p r e c l p i t a t i o n of MnC03 by loss of NH 3 and water. This was probably the reason of the f a s t p r e c i p i t a t i o n of manganese out of s o l u t i o n above 55°C and the very low f i n a l s o l u b i l i t y of MnO at 130°C, e.g. 0.1875 g/1. 2. Reductive leaching of P y r o l u s i t e In order to f i n d a low-cost, e f f e c t i v e reducing agent, formal-dehyde, methanol, formic acid and NaN02 were tested. None gave manganous ions i n s o l u t i o n . S u l f i t e and hydrazine yielded manganese i n s o l u t i o n . 2.1 Reductive Leaching of P y r o l u s i t e Using Sodium S u l f i t e The leaching curves A, B, C and D are reported i n Figure 22 and were obtained by varying temperature and the amounts of s u l f i t e and p y r o l u s i t e i n s o l u t i o n . In each of the runs a f i n a l equilibrium was obtained and the amount of manganese remaining i n s o l u t i o n depended on the operating conditions mentioned above. The shape of the d i f f e r e n t curves suggested that s u l f i t e acted i n two ways. I t reduced p y r o l u s i t e to MnO more r e a d i l y with increasing temperature and p y r o l u s i t e concentration (curves A and B). I t pre--c i p i t a t e d manganese from s o l u t i o n p r o p o r t i o n a l l y to the s u l f i t e and manganese content of the s o l u t i o n . The p r e c i p i t a t e was i d e n t i f i e d as hydrated ammonium manganese s u l f i t e (NH 4) 2.Mn(S0 3) 2.H 20 by X-ray g/l Mn dissolved T i m e t min.) Figure 22 ': E f f e c t of s u l f i t e concentration on the reductive leaching of p y r o l u s i t e . analysis of the residue (curves C and D), With the help of curves A and B some t h e o r e t i c a l model could be presented by assuming that the reduction r e a c t i o n of p y r o l u s i t e was: Mn0 2 + S0 3 -> MnO + SO^ (10) before any p r e c i p i t a t i o n was taking place (this l s probably true f o r curves A and B since almost a l l the s u l f i t e had been used i n the re-duction of p y r o l u s i t e ) . The rate of d i s s o l u t i o n could then be expressed by: = k. [SO " J (11) dt and having: rso = i = rso = i " [ M n 0 ] 3 L b ° 3 J i n i t i a l d [MnO] dt = k l [ t l ~ ( M n 0 ) ] ( 1 2 ) By i n t e g r a t i o n : [MnO] = K. (1- e k l t ) (13) aq i A p l o t of log (K^ - Mn complexed) versus time must be s t r a i g h t l i n e of slope k^. Relation (13) appeared to be followed for runs A and B. (Figure 23). In runs C and D the i n i t i a l d i s s o l u t i o n rate was too high to be measured and only p r e c i p i t a t i o n curves were obtained. 2.2. Reductive Leaching of P y r o l u s i t e Using Hydrazine Hydrate as a  Reducting Agent 2.2.1. Leaching of P y r o l u s i t e with a Large Excess of Hydrazine Hydrate at 85°C 0.9 0.8 _ Log(K -Mn) ; ; Log(l£?M/l) 0.7 ( A ) ; 4 g Na 2 S0 3 ( B ) ; 8 g 0.6 ^ V ^ A ) 0.5 ^ ^ ^ ^ 0.4 i i i i i i i 1 .2 1.3 1.4 1.5 1.6 1.7 1.8 1.9 L o g ( t i m e ) Figure 23 : Log-log pl o t between ( K x- Mncomplex ) and time at 130 °C. A run with 1.67 moles/1 of 64% hydrazine i n water and 0.61 moles/1 of unsized p y r o l u s i t e ore was c a r r i e d out i n order to check the pressure change with the progress of the reaction. If p y r o l u s i t e was reduced by hydrazine according to the following stoichiometric r e -a c t i o n : 2Mn02 .+ N 2H 4 ->- 2Mn0 + 2H 20 + N 2* (14) one must detect a proportional increase i n the pressure due to N 2 evolution; i t could be assumed that the s o l u b i l i t y of N 2 i n s o l u t i o n was small. The d i s s o l u t i o n of p y r o l u s i t e was so f a s t that v i r t u a l l y no p y r o l u s i t e was l e f t when the desired temperature was reached. Hence the increase i n the t o t a l pressure during heating compared to the equilibrium pressure obtained i n absence of reduction at corresponding temperatures was equal to the N 2 p a r t i a l pressure. (Figure 24). Knowing the tem-perature and the t o t a l gas phase volume at each time, the amount of nitrogen produced could be computed and hence a t h e o r e t i c a l leaching; curve according to r e l a t i o n (14) could be c a l c u l a t e d . (Figure 25). The measured amounts of manganese i n s o l u t i o n were pl o t t e d on the same f i g u r e . A deviation of 10% was noted, that i s more hydrazine was used than ex-pected from reaction (14). Indeed, some NH^ could have been formed or side reductions could have taken place. . • • i 2.2.2. E f f e c t of Temperature Because of the f a s t d i s s o l u t i o n of p y r o l u s i t e even at low tem-peratures precautions were taken to check and possibly to avoid d i f f u s i o n control of the d i s s o l u t i o n reaction. p 0 50 BOO 150 Tomperature (°C) Figure 2k : Pressure v a r i a t i o n i n the d i s s o l u t i o n of p y r o l u s i t e with hydrazine hydrate as a reducing agent. 40 g Mn dissolved 30 20 10 5 0 -A A m@asur©< A calculate) 100 Ti me ( min.) 150 Figure 25 : Dissolution of p y r o l u s i t e using hydrazine hydrate as a reducing agent. Increasing the s t i r r i n g rate had no e f f e c t on the d i s s o l u t i o n of py r o l u s i t e ( F i g u r e 26). Two s e r i e s of runs were investigated. The f i r s t series were ca r r i e d out using 5 g of p y r o l u s i t e s i z i n g 50 - 70 mesh. I t was ex-pected that the large p a r t i c l e s i z e and the narrow p a r t i c l e s i z e d i s -t r i b u t i o n would decrease the d i s s o l u t i o n rate of p y r o l u s i t e and enable more accurate measurement of the i n i t i a l leaching rates. The second serie s were ca r r i e d out using 5 g of p y r o l u s i t e s i z i n g 65-150 mesh. This i s the s i z e which was generally used f o r the leaching experiments. In t h i s second case however, the temperature had to be maintained below 50°C to enable measurement of the i n i t i a l leaching rates. The d i s s o l u t i o n curves are p l o t t e d i n Figures 27 and 28 fo r both ser i e s r e s p e c t i v e l y . The reaction rate depended on the hydra zine concentration and on the amount of undissolved p y r o l u s i t e . However, even i f the hydrazine concentration decreased during the runs, s t r a i g h t portions could be obtained i n the early stages of the leach. Rates were calculated by taking s t r a i g h t l i n e s through the f i r s t three points of the leaching curves. I t was then possible to use Arrhenius pl o t s f o r the c a l c u l a t i o n of apparent a c t i v a t i o n energies, as shown i n Figure 29. An apparent a c t i v a t i o n energy of 16.7. K cal/mole f o r the f i r s t series runs was i n good agreement with the value of 17.4 K cal/mole obtained f o r the second seri e s runs. A f i n a l value of 17.0 K cal/mole - I K cal/mole was deduced from the Arrhenius p l o t s . 2.2.3. E f f e c t of Hydrazine Hydrate Concentration on the Reductive  Leaching of P y r o l u s i t e Leaching curves for 0.005, 0.044, 0.087, 0.175 and 0.262 moles per l i t e r of hydrazine hydrate are given i n Figure 30. The s t r a i g h t Time ( min.) Figure 26 : E f f e c t of s t i r r i n g on the d i s s o l u t i o n of p y r o l u s i t e with hydrazine at 130 °C. 0 I I I ! I 1 L _ 0 50 100 150 Time ( min.) Figure 27 : E f f e c t of temperature on the leaching, of p y r o l u s i t e ( 50-70 mesh .) with hydrazine hydrate . Mn dissolved r@lusit© 65 -150 mesh 100 300 Figure 28 200 Time ( min.) E f f e c t of temperature on the leaching of puro l u s i t e ( 65-I5O mesh ) with hydrazine hydrate. I 1 1 1 I I I I 2.9 3 0 31 3-2 3-3 3-4 3-5 , 0 0 Q • T (° K ) Figure 29 : Arrhenius plot for leaching of p y r o l u s i t e with hydrazine hydrate. oo 0 50 100 150 Tim© ( min .) Figure 30 : E f f e c t of hydrazine concentration on the d i s s o l u t i o n of pyrolusite at 30 °C. l i n e r e l a t i o n s h i p was followed up to a c e r t a i n amount of manganese i n s o l u t i o n and was used f o r c a l c u l a t i o n of the i n i t i a l r ates. If i t was assumed that reaction (14) represented the leaching reaction, the following rate equation can be deduced: " - * 2 [N 2H 4]^ (15) dt *" 2 The a c t i v i t y of p y r o l u s i t e probably i s concerned with i t s surface area and the configuration of active s i t e s . By p l o t t i n g log (rate) versus log [N^H^] a s t r a i g h t l i n e was obtained, the slope of which was equal to 0.56 (Figure 31). This meant that r e l a t i o n (15) was e s s e n t i a l l y correct as far as hydrazine concentration was concerned. The t h e o r e t i c a l slope of 0.50 i s 10% d i f f e r e n t of the measured value of 0.56. I t i s again observed that a small excess of hydrazine i s needed over the t h e o r e t i c a l amount predicted by equation (14). Rate constant may involve other parameters besides temperature such as surface area and s i z e of the p y r o l u s i t e c r y s t a l s . 2. 2. 4. E f f e c t of Doubling the I n i t i a l Amount of P y r o l u s i t e In one of the runs,10 g of p y r o l u s i t e was added instead of 5 g. The hydrazine concentration was of 0.175 moles/1 i n both runs. As seen i n Figure 32,twice as much time was needed to obtain a same amountof manganese i n s o l u t i o n f or 5 g of p y r o l u s i t e than f o r 10 g of p y r o l u s i t e . I t may be concluded that the rate determining step i n d i s s o l u t i o n of p y r o l u s i t e occurred on the surface of the mineral. 2 h Log(rote) L og (mg/min) Log( N 2 H 4 ) Log ( g) Figure 31 : Log-log plot between hydrazine concentration and the d i s s o l u t i o n rate of p y r o l u s i t e at 30 °C. Figure 32 : E f f e c t of doubling the i n i t i a l amount of pyr o l u s i t e on the d i s s o l u t i o n of p y r o l u s i t e with hydrazine at 3 0 °C. D. SOLUBILITY OF AMMONIUM PHOSPHATE IN THE NH„ - CO„ - H n0 SYSTEM ; : j 2 2 1. E f f e c t of Temperature Phosphate s o l u b i l i t y increased r a p i d l y with increasing temper-ature (curve A of Figure 33). Below 35°C the s o l u b i l i t y of phosphate was small. This was i n agreement with the c r y s t a l l i z a t i o n of fibrous phosphate out of s o l u t i o n on cooling. Usually manganese ores found i n nature contain 0-1% phosphate. If 100 g/1 of p y r o l u s i t e were added to the leaching s o l u t i o n there would be a maximum of 1 g/1 of phosphate present. Therefore, curve A was obtained s t a r t i n g with no more than 2.26 g/1 of phosphorus i n the form of ammonium phosphate NH^ HPO^. This l a s t was probably easier dissolved than would have been phosphates out of a r e a l ore. Curve A has a S shape denoting that the s o l u b i l i t y of phosphate increased exponentially with temperature. 2. E f f e c t of Adding FeO The maximum s o l u b i l i t y curve of phosphorus i n the presence of complexed i r o n i s pl o t t e d i n Figure 33 and the v a r i a t i o n of i r o n i n s o l -ution i n Figure 34. The s o l u b i l i t y product of i r o n phosphate was reached when the temperature exceeded 60°C. However, phosphate continued to be dissolved while i r o n phosphate was p r e c i p i t a t i n g , f o r the i r o n con-tent i n s o l u t i o n decreased continuously above 60°C (Figure 34). About 2 g/1 of i r o n was p r e c i p i t a t e d out taking with i t 1.07 g/1 of phosphorus out of s o l u t i o n . This showed that 1 mole of i r o n p r e c i p i t a t e d 1 mole of phosphate, probably under the form of FePO^ or FePO^.H^O. X-ray d i f f r a c t i o n pattern of the residue did not reveal c l e a r l y the presence of FePO. but did not exclude i t either (low peaks were present). oo 3. E f f e c t of Adding MnO The maximum s o l u b i l i t y curve of phosphate i n the presence of complexed divalent manganese i s plo t t e d i n Figure 33 and the amount of manganese i n s o l u t i o n i s seen i n Figure 34. The s o l u b i l i t y pro-duct of manganese phosphate had a much lower value than that of i r o n phosphate, since already at 35°C a manganous phosphate p r e c i p i t a t e d out Above 80°C the s o l u b i l i t y of phosphate was maintained below 0.01 g/1 (hardly measureable). The manganese consumption was low (Figure 34); only 0.34 g/1 of phosphate had been p r e c i p i t a t e d by 0.61 g/1 of manganese. Hence, manganese not only p r e c i p i t a t e d out a l l the phos-phate i n s o l u t i o n but also prevented further phosphate d i s s o l u t i o n . I t i s suggested that an in s o l u b l e manganese species coated the phosphate p a r t i c l e s . In the presence of enough complexed manganese i n s o l u t i o n the s o l u b i l i t y of phosphate was very small. Above 35° C the s o l u b i l i t y product of manganese phosphate was very small and below 35°C the s o l u b i l i t y of phosphate i t s e l f was small independently of the amount of manganese present i n s o l u t i o n . I t should be noted that phosphate presence i n s o l u t i o n i s undesired i n the production of manganese compounds. E. RECOVERY OF MANGANESE FROM SOLUTION WITH OXYGEN Several runs were c a r r i e d out under d i f f e r e n t oxygen pres-sures. The p r e c i p i t a t e which was obtained was a very f i n e black powder which could not be i d e n t i f i e d by X-ray d i f f r a c t i o n . A thermal d i f f e r e n t i a l analysis did not give any c h a r a c t e r i s t i c endo or exothermic peak but a continuous endothermic transformation. Hence, the p r e c i p i -tate was considered to be i n the amorphous st a t e . Two samples were 87 heat treated under vacuum f o r 24 hours, one at 200°C and one at 400°C. The T.D.A, analysis showed endothermic regions at 300°C, 650°C and 1100°C, These regions probably correspond to the endothermic trans-formations. 3Mn0 2 - M n 0 1 < 5 5 M n 0 1 . 5 5 * a M n 2 ° 3 a Mn 20 3 + M n 3 ° 4 and i t could be assumed that amorphous Mn02 was p r e c i p i t a t e d by 0^ at 30-40°C. 1. E f f e c t of the Oxygen Pressure on the P r e c i p i t a t i o n Rate of Complexed  Manganese out of Solution The e f f e c t of 80, 110, 220 and 475 p s i of oxygen pressure on the p r e c i p i t a t i o n rate of complexed manganese out of s o l u t i o n was i n -vestigated. (Figure 35). The p r e c i p i t a t i o n rate'of manganese dioxide increased r a p i d l y with oxygen pressure. At constant oxygen pressure the rate decreased as the manganese concentration i n s o l u t i o n decreased. Assuming that the p r e c i p i t a t i o n r eaction can be expressed by the following equation: k 3 Mn complex, v + x 0„ •+ Mn0_ + products (16) U q \ 2 (gas) Z(s) i n which x and the products are unknown; a rate equation can be deduced: d[Mn complex. ,.] x = k 3 [Mn complex]^ ( f 0 2 ) (17) At constant oxygen pressure, expression (17) can be integrated between an i n i t i a l amount of complexed manganese i n s o l u t i o n and some smaller amount Figure 35 : E f f e c t of oxygen pressure on the p r e c i p i t a t i o n of manganese from solution at ± 35 °C• ob oo 89 reached a f t e r a time t : [Mn Complex] ,1 = e~ J (18) [Mn complex]. r i n i t i a l with k* = k 3 (|»02) X A logarithmic p l o t of the l e f t member of expression (18) versus time should be a st r a i g h t l i n e f o r each oxygen pressure and the slope of the l i n e s should then be equal to the k^ values. Indeed, s t r a i g h t l i n e s were "obtained (Figure 36) and the time r e l a t i o n s h i p of manganese i n s o l u t i o n (eq 18) could be deduced. The order of the reaction i n oxygen could be calculated from a p l o t of l o g ( k 3 ) versus log ( f C ^ * This p l o t was a st r a i g h t l i n e of slope equal to x, the order of the reaction i n oxygen. A re-l a t i v e l y good s t r a i g h t l i n e was obtained (Figure 37). The calculated order i n oxygen pressure was 1.6. The products are unknown but are probably composed by R^O, CO^, NH^ and N 2« The mechanism of t h e 1 r e -action i s also unknown at t h i s point. 2. E f f e c t of the S t i r r i n g Rate on the P r e c i p i t a t i o n Rate For one of the runs the s t i r r i n g rate was m u l t i p l i e d by three, e.g. from 900 R.P.M. to 2,700 R.P.M. This only resulted i n a 10% r i s e of the p r e c i p i t a t i o n rate of manganase out of so l u t i o n by oxygen. The transfe r of oxygen from the gas phase to the l i q u i d phase depended meanly on the oxygen pressure (Figure 38). L o g ( Mn*)/(MnJb 3 -47 5 psi — A ^ 2 20 psi - 7 - / . 110 psi * i 8 0 psi _ _ _ j g > — 50 100 Ti me ( min.) 850 Figure 36 : Log of manganese concentration i n so l u t i o n versus time f o r d i f f e r e n t oxygen pressures. o Figure 38 : E f f e c t of s t i r r i n g on the p r e c i p i t a t i o n of manganese from s o l u t i o n at constant oxygen pressure. 93 DISCUSSION AND CONCLUSIONS A. DISSOLUTION OF FeO AND MnO Both FeO and MnO were soluble i n the NH^ - CO - H^ O system with 18 moles/1 NH^ and 3 moles/1 CO^. The divalent i r o n complex formed by d i s s o l u t i o n i n the system was found to be more stable with increasing temperature than the analogous divalent manganese com-plex. The s o l u b i l i t y of FeO remained constant at 1.2 moles/1 from 20°C to 130°C while the s o l u b i l i t y of MnO remained constant at 1.2 moles/1 only from 20°C to 50-55°C and then decreased almost to zero at 130°C. This was due to the breakdown of the manganous complex into i n s o l u b l e MnCO^ ( i n the rhodochrosite form). This d i f f e r i n g behaviour of i r o n and manganese i s only apparent i n a closed system of course, for i t was shown by W e l s h ^ that i n an open system the ferrous complex also broke down by p r e c i p i t a t i n g FeCO^ (in the s i d e r i t e form) when the ammonia was able to leave the system and at a tem-perature of 80°C. The nature of the complexes i s unknown, but Dean and W e l s h 4 ^ ' ^ postulated the following: 0 Mn(NH V - °- C- N H2 3 ^ - + 0 NH. 4 and Fe(NH_) ^ c ' m 2 3 ?v - + 0 NH 4 where y > x > 1. In t h i s study d i f f e r e n t complexes were assumed. It i s im-94 portant to note that amino acids e x i s t as dipolar (zwitter) ions i n the c r y s t a l l i n e state"'' 7. Such ions may form chelate rings with metals. NH, (CH 9) £ 00H = +NH„ (CH„) COO" z z n j z n It i s suggested that f o r n = 0, a weak zwitter carbamate ion might e x i s t i n s o l u t i o n : +NH 3 C00~ which might then coordinate on d i v a l e n t metals and form pseudo-square planar complexes (extra coordination of NH 3 i s probable): H2 N 0 / X / \ 0 = C Mn C = 0 \ / \ / 0 N However, such a complex might not be acceptable f o r stereo-chemical reasons and must therefore be considered as"hypothetic. .;. ( i . e . the p o s s i b i l i t y f o r the carbamate ion to form a r i n g around the manganous ion) Tervalent manganese and i r o n cannot be coordinated i n the. same way with carbamate ions and at the same time respect s t a b i l i t y conditions. Tetravalent manganese could be coordinated with four carbamate ions but since carbamate ions are not l i k e l y to be strongdipolar ions such coordination probably does not occur. Some experimental r e s u l t s were i n favor of the above mentioned complex; at 40°C the concentration of ammonium carbamate i n s o l u t i o n was 2.7 moles/1. Both MnO and FeO dissolved quickly and t h e i r maximum s o l u b i l i t y was found to be 1.2 moles/1 even when an excess of MnO and FeO was used. Therefore i t seemed possible that more than 1 mble of carbamate reacted with 1 mole of MnO or FeO during d i s s o l u t i o n . 58 It should also be noted that Bernard i s o l a t e d manganous carbamate i n a pure NH^-CX^ system, avoiding hydrolysis of the complex by water, He mentioned the following complex: Mn (C0 2 N H 2 ) 2 » 4NH3 B. REDUCTIVE DISSOLUTION OF GOETHITE As goethite does not di s s o l v e i n the carbamate s o l u t i o n i t was f i r s t reduced to y i e l d FeO which was r e a d i l y soluble at elevated temperature. The d i s s o l u t i o n was c a r r i e d out at 130°C since higher tem-4 peratures might a f f e c t the structure of goethite . Various reducing agents were used. I, Sodium S u l f i t e Sodium s u l f i t e did reduce goethite but only with the help of I j c a t a l y s t s such as Cu , anthraquinone, or sodium 3 sulfonate anthraquinone. Experiments showed that: 1) The rate of d i s s o l u t i o n of goethite decreased quickly with time, even with excess s u l f i t e . 2) At constant c a t a l y s t concentration the leaching rate of goethite f i r s t increased with increasing s u l f i t e concentration then decreased when excess s u l f i t e was used. 3) At constant sodium s u l f i t e concentration, increasing amounts of 96 cat a l y s t increased the goethite leaching rate. 4) Sulfate was without e f f e c t on the leaching rate of goethite. This suggested that two reactions were competing: a) d i s s o l u t i o n of goethite (proportional to the concentration of s u l f i t e ) b) p r e c i p i t a t i o n out of s o l u t i o n of complexed i r o n (proportional to both s u l f i t e and i r o n concentration i n solution) II Hydrazine Sulfate No ca t a l y s t was needed when hydrazine s u l f a t e was used as a reducing agent of goethite. I t was concluded from the experiments that the rate at which i r o n appeared i n s o l u t i o n decreased quickly with time. This fast decrease i n the di s s o l u t i o n s of goethite might be associated with the presence of some s u l f i t e i n s o l u t i o n . I I I Hydrazine Hydrate 1) The d i s s o l u t i o n of goethite with time followed a s t r a i g h t l i n e r e l a t i o n s h i p when hydrazine hydrate was used instead of hydrazine s u l f a t e . 2) At 80°C,4% of the i r o n was leached a f t e r four hours while at 35°C only 0.4% of the i r o n was dissolved i n the same time. This was useful f o r at low temperature:1;the goethite remained almost i n -soluble whereas p y r o l u s i t e did not. This i s discussed i n the next section. C. REDUCTIVE DISSOLUTION OF PYROLUSITE Py r o l u s i t e was more r a p i d l y reduced to MnO than goethite to FeO by both s u l f i t e and hydrazine. 97 I. Sodium S u l f i t e Leaching of p y r o l u s i t e using s u l f i t e as a reducing agent could be ca r r i e d out without the presence of a c a t a l y s t i n s o l u t i o n . j Generally the d i s s o l u t i o n rates f o r p y r o l u s i t e were one order of mag-nitude higher than f o r goethite. I t was concluded from the experiments that: 1) Increasing amounts of s u l f i t e increased the i n i t i a l d i s s o l u t i o n G rates of p y r o l u s i t e . 2) When excess s u l f i t e was used, only p r e c i p i t a t i o n curves were ob-tained. The p r e c i p i t a t e was i d e n t i f i e d as hydrated ammonium manganese s u l f i t e and might be formed by rea c t i o n : H 20 + Mn complex ( v + 2 (NH^SC^ -> (NH^MnCSO ) H 0 (aq) 3) For low s u l f i t e concentration the d i s s o l u t i o n rate of p y r o l u s i t e appeared to be f i r s t order i n s u l f i t e concentration. As i n the case of goethite leaching both reductive d i s s o l u t i o n and p r e c i p i t a t i o n reactions were competing. No more than 6 g/1 of manganese could be maintained i n s o l u t i o n at 84°C. II Hydrazine Hydrate As hydrazine s u l f a t e might have l e d to s i m i l a r problems f o r py r o l u s i t e as f o r goethite, hydrazine hydrate was used. At 35°C, 90% of the manganese was leached i n one hour when a 50% excess of hydrazine 24 hydrate was used. Hydrazine , N^H^, may be thought of as derived from ammonia by replacement of a hydrogen atom by the -NH^ group. I t might therefore be expected to be a base, but somewhat weaker than NH^, which i s the case. I t i s a b i f u n c t i o n a l base: 98 N 2H 4 + H 20 = N 2 H 5 + + 0H~ K 2 5p c - 8.5 x l ( f 7 (aq) (aq) (aq) N 2 H 5 + + H 2 0 - + O H - a q ) K 2 5 0 c = O x I O " 1 6 (aq) (aq) Aqueous hydrazine i s a powerful reducing agent i n basic s o l u t i o n ; one reaction which i s quantitative with some oxidants i s : 40H~ + N 2H 4 -> N 2 + 4H 20 + 4e E° = 1.16V 25 26 27 However, NH^ and HN^ are also obtained under various conditions ' ' The r e s u l t s of the reductive d i s s o l u t i o n of p y r o l u s i t e i n -dicated that: 1) The o v e r a l l reaction was given by: ZMnO., v + N. H. + «NH.CO_NH. 2(s) '2 4, N 4 2 2, N (aq) (aq) 2(Mn complex), . + 2H„0 + N 0 ( a q ) 2 (aq) 2 ( q ) 2) The apparent a c t i v a t i o n energy f o r p y r o l u s i t e reductive d i s -s o l u t i o n i n s o l u t i o n was 17.0 K cal/mole - I K cal/mole below 50°C. Above 50°C the a c t i v a t i o n energy dropped to a much lower value. 3) Above a concentration of 1.3 g/1 of hydrazine the reac t i o n order was 0.56 i n hydrazine concentration. Below t h i s concentration the reaction order increased towards 1.0 4) Doubling the amount of p y r o l u s i t e and using the same hydrazine concentration doubled the rate at which manganese appeared i n sol u t i o n . Therefore, the rate determining step of the d i s s o l u t i o n 99 of p y r o l u s i t e occurred probably at the surface of the mineral. 5) Manganous oxide dissolved much quicker than p y r o l u s i t e under the same leaching conditions, 6) Increasing the s t i r r i n g rate had no e f f e c t on the leaching rate of p y r o l u s i t e In the following discussion an attempt i s made to r e l a t e the above mentioned experimental r e s u l t s and some conclusions a r r i v e d at 25 38 39 59 by Morgan > Brenet ' and Vetter with a general mechanism which i s operative during the reductive d i s s o l u t i o n of p y r o l u s i t e i n aqueous ammoniacal ammonium carbamate s o l u t i o n . It i s considered that the f i r s t step of the d i s s o l u t i o n of pysolusite involves hydration of the oxide surface i n the folowing manner: K l ^ V J + H2°(aq) = CMnO > I.H 0 (1) s s Morgan and Stumm showed that the zero point of charge of c o l l o i d a l manganese dioxide occurred at pH 2.8 with respect to net OH and H + ions bound. With increasing pH, the charge becomes progress-i v e l y more negative. The leaching experiments i n the present work were ca r r i e d out at pH 9-10. The potential-determining r o l e of OH ions at pH.9 may be v i s u a l i z e d as a d i s s o c i a t i o n of H + ions from surface OH groups: K2 ( Mn0 2)J.H 20 = (Mn0 2) x|.OH- + H + (2) s s In t h i s condition the oxide surface w i l l probably have an o v e r a l l negative charge due to the hydroxy1 groups on the surface. 100 Hydrazine i s known to react with water by giving hydrazinium cations i n s o l u t i o n as follows: K* + N 2H 4 + H 2 ° = N 2 H 5 + ° H ( 3 ) (aq) (aq) (aq) (aq) The o v e r a l l negative charge of the p y r o l u s i t e surface f a c i l i t a t e s adsorption of hydrazinium cations. This may be written: (Mn0 2) x| .OH" + N2H+ = 4 (MnO^| .0H.H5 N 2 (4) s (aq) s Subsequently a redox reaction occurs on the surface: k l (Mn0 2) x|OH.H 5N 2 + (Mn0 2) x_ 2 (MnO)2|.H20 s s (5) + 2H 0 + N * 1 (aq) ^(g) Step (5) may involve a progressive reduction of the p y r o l u s i t e 38 as mentioned by Brenet , who studied the reduction of y - Mn02 with hydrazine i n sodium hydroxide buffered s o l u t i o n at pH 9. Divalent manganese was not soluble i n such a s o l u t i o n so that progression of the reaction with d i l a t i o n of the y - Mn0 2 unit c e l l was observed according to re a c t i o n : Mn02 + f- N H -> MnO _ . (OH) + | N 39 However, Brenet showed that progression of the reduction i n s o l i d phase was possible because of the non-stoichiometry of y - Mn0 2 which contained OH groups i n the l a t t i c e . Since p y r o l u s i t e i s the 0 form of manganese 101 dioxide such a progression was not observed and only a surface r e a c t i o n took place. Indeed, B - Mn0 2 i s a stoichiometric compound and the rate of the hydrogen reduction f e l l to zero once the surface was covered with insoluble MnO.OH or MnO.H20. In the NH3-CO2-H20 system |MnO.H20 reacted with the leach s l i q u o r and so new s i t e s of the p y r o l u s i t e surface could be activated by water and more hydrazinium cations continued to be adsorbed a f t e r de-sorption of the r e s u l t i n g manganese complex: k 2 (Mn02) 2|(MnO) 2.H 0 + 2NH4 C0 2 NH + X s 2 (aq) (Mn0 2) x_ 2 [Mn(NH 0)„ (C0„ NH 0)„]„ + 2Ho0 (6) *3'2 v 2 2'2J2 2 k 3 (Mn0 2) x_ 2|[Mn(NH 3) 2 (C0 2 N H 2 ) 2 ] 2 -8 (Mn0 2) x_ 2| + 2 Mn(NH 3) 2 (C0 2 N H ^ (7) s (aq) Equilibrium for reactions (1) and (2) i s established f a s t 25 according to Morgan . It was observed that MnO dissolved much f a s t e r than p y r o l u s i t e . Although i n those runs very f i n e powder was used f o r MnO and 65-150 mesh grains f o r p y r o l u s i t e i t i s assumed that reactions (6) and (7) are f a s t . I f reaction (7) was slow and rate-determining i t would be expected Hwt" the d i s s o l u t i o n rate of p y r o l u s i t e on hydrazine decreased f a s t e r with increasing hydrazine concentration than i t was observed within the experi-mental range of conditions studied. 102 Since a r e l a t i v e l y high a c t i v a t i o n energy was measured for the d i s s o l u t i o n of p y r o l u s i t e the rate-determing step appeared to be a chemical r e a c t i o n . If the adsorption of hydrazinium cations (4) was rate-deter-mining, the rate of d i s s o l u t i o n of p y r o l u s i t e should follow the following r e l a t i o n : [N 2 E 5 + ] H — = A [ N ^ ] ^ + B (8) (rate) where A and B are constants f o r a steady state. This was not at a l l the case. This only leaves reaction (5) to be considered f o r the rate-determining step. The rate equation f o r the d i s s o l u t i o n of manganese v i a t h i s mechanism may be written as: d[Mn complex] I i = k/* [(Mn0 2)J.OH.H N ]"* (9) dt 1 s Using equilibrium reactions (1), (2) and (3), rate equation (9) becomes: d[Mn complex]( a q) dt (0H~) (H ) [( M n 0 2 ) x ) ] l N 2 H 4 ] h _s (aq) (10) The product (OH )(H +) i s constant and i f i t i s assumed that the surface area of p y r o l u s i t e i s not modified i n the early stages of the dis-s o l u t i o n i t may be seen that the rate of d i s s o l u t i o n of p y r o l u s i t e i s d i r e c t l y proportional to [®2^^2 (aq)' This was also observed i n the ex-periments . I t should be noted that the d i s s o l u t i o n rate of p y r o l u s i t e might be pH dependant i f the adsorption of hydroxyl ions on i t s surface was not l i n e a r l y proportional to the pH of the s o l u t i o n . This was not detected i n the experiments since the pH remained constant at 9.5-10 during the d i s s o l u t i o n of p y r o l u s i t e . From the o v e r a l l r e a c t i o n i t i s deduced that: [N 2H 4] = [N 2H 4] - | [Mn complex] } (aq) (aq) i n i t i a l (aq) one mole of hydrazine reacting with p y r o l u s i t e corresponding to two moles of manganese complexed i n s o l u t i o n . Expression (10) can then be integrated when i t i s assumed that the product k^.K^.K2.K^.K4 [(Mn0 2) x|] i s constant for a given temperature s This assumes that the active surface area of p y r o l u s i t e i s not a l t e r e d during the reductive d i s s o l u t i o n . The f i n a l expression a f t e r i n t e g r a t i o n i s : [Mn complex}£ a (^ = B - (B - Dt) 2 = Dt (2B 2 - Dt) with: t = time (min.) B = 2[N 2H 4] (moles/1) ( a q ) i n i t i a l k-.K..K^.K-.K. 1 1 2 3 4 (12) D = (OH") (H +) (moles/1)^ (min.) A = [(Mn0„) | ] , concentration of exposed s i t e s on p y r o l u s i t e s k^ = rate constant f o r the rate-determining r e a c t i o n . molesv , . v —2 ( q—) (min.) 104 = equilibrium constant f o r the hydration of the py r o l u s i t e surface - equilibrium constant forthe net OH bound on the p y r o l u s i t e surface Kg = equilibrium constant of the reaction of hydrazine with water Kj, = equilibrium constant f o r the sorption of hydra-zimium cations on the p y r o l u s i t e a c t i v e s i t e s Equation (13) shows that the d i s s o l u t i o n of p y r o l u s i t e follows a parabolic r e l a t i o n s h i p with time, which i s observed i n the ex-periments. At the s t a r t of the runs, B =2[N 2H 4]^ n^ t^ a^ i s large com-pared with the manganese complexed i n s o l u t i o n and the rate of d i s s o l u t i o n i s almost constant. This accounts f o r the s t r a i g h t portions of the d i s -s o l u t i o n curves: \, [Mn complex] ^  ^ = 2B 2 D.t (t small) It was observed that f o r very low concentrations of hydrazine i n s o l u t i o n the order of the rate-determining step approached one. This might mean that the rate-determining r e a c t i o n (5) became: k4 (Mn02) |0H.H N 2 (Mn02) (MnO) | ^ O . H ^ + H ^ (14) s x-1 s The remaining diazene (N 2H 2) could then act i n two possible ways: 1) react with a neighbouring Mn0 2 s i t e as follows: k 5 (Mn0 2) x|H 20.N 2H 2 -> (MnO ). |MnO.H20 + H 20 + x-1 33 2) desorp into s o l u t i o n giving disproportionation according to r e a c t i o n : 2N 2H 2 + H 20 -> N 2 H 5 + + 0H~ + (15) (aq) (aq) Xaq) (aq) (gas) 33 35 Reaction (15) i s f a s t as mentioned by Cahn and Huang I t should be noted that the stoichiometry of the o v e r a l l r e a c t i o n i s not al t e r e d by r e a c t i o n (15) since hydrazinium cations are regenerated i n s o l u t i o n and only N 2 i s obtained as a by-product. The much slower rate of d i s s o l u t i o n observed for goethite than for p y r o l u s i t e i n the presence of hydrazine hydrate may be explained on the basis of d i f f e r e n t surface behaviour of the oxides i n acid-base'sol-utions. Indeed, the zero point of charge (z.p.c.) for i r o n (111) oxide was measured at pH 8.5 by Parks and De B r u y n ^ t h i s value i s i n contrast with the much lower pH value of 2.8 measured by Morgan and Stumm f o r the z.p.c. of manganese (IV) oxide. In the present work, the pH of the leach-ing s o l u t i o n remained constant at 9.5-10 during the d i s s o l u t i o n of goethite and p y r o l u s i t e . In t h i s s i t u a t i o n the surface of p y r o l u s i t e would show a larger OH concentration per unit area compared to goethite, for i t was shown by the above investigators that the net OH bound on the oxide surface increased r a p i d l y with increasing pH above the pH corresponding to the z.p.c. of the oxide. As a consequence of t h i s d i f f e r e n c e i n the negative charge concentration of the two mineral surfaces, the adsorption of the hydrazinium cations would also be expected to d i f f e r . In the present experiments p y r o l u s i t e dissolved approximately one thousand times 106 fa s t e r than goethite and t h i s may correspond to the differences i n magnitude of the negative charges on the minerals. D, DISSOLUTION OF PHOSPHATE Ammonium phosphate was soluble i n small quantity at 30°C, but i t s s o l u b i l i t y increased quickly with increasing temperature. Manganese ores are very often associated with phosphates i n nature. Phosphate i s undesired i n the manufacturing of manganese products, e.g. manganese used i n the deoxydation of high q u a l i t y s t e e l s . When FeO was added i n s o l u t i o n p r e c i p i t a t i o n of an i r o n phosphate compound started above 50°C. When MnO was added i n solution, the phosphate content,of the carbamate s o l u t i o n was maintained below 0.01% of the i n i t i a l amount added, at a l l temperatures. The manganous carbamate p r e c i p i t a t e d phosphate out of s o l u t i o n already above 35°C and also prevented more phosphate to be dissolved probably by coating the phosphate p a r t i c l e s with an insolu b l e species. Below 35°C the phosphate i t s e l f exhibited low s o l u b i l i t y i n the carbamate s o l u t i o n . This was demonstrated when c r y s t a l l i z a t i o n of fibrous ammonium phosphate occurred a f t e r cooling of the s o l u t i o n . In the previous sections i t was shown that goethite dissolved slowly with hydrazine f o r temperatures between 30-40°C. On the other hand, p y r o l u s i t e leached at a high rate with hydrazine over the same temperature range. Since i n presence of complexed manganese also the phosphate content of the carbamate s o l u t i o n was kept low, i t was concluded that the leaching of manganese from i r o n and phosphate containing man-ganese dioxides with hydrazine hydrate might be successful f o r temperatures below 40°C E. RECOVERY OF MANGANESE WITH OXYGEN The l a s t part of t h i s work was concerned with the recovery of 51-53 maganese from s o l u t i o n . Previous workers investigated the leaching of MnO i n carbamate solutions at room temperature i n an open system. When the s o l u t i o n became saturated with the manganous carbamate complex t h i s l a s t was p r e c i p i t a t e d from s o l u t i o n under the form of rhodochrosite by heat and loss of ammonia. In t h i s work, rhodochrosite could be ob-tained without l o s s of ammonia and carbon dioxide by increasing the temperature above 50°C i n absence of oxygen. V i r u a l l y a l l the manganese had p r e c i p i t a t e d at 130°C. However, i n view of making y - MnO^ s u i t a b l e f o r battery c e l l s , i t was then necessary to r e - d i s s o l v e the manganous carbonate i n a c i d i c s o l u t i o n before an e l e c t r o l y s i s could be achieved. It was proposed i n t h i s work to p r e c i p i t a t e MnO^ from s o l u t i o n by breaking down the manganese complex with an oxygen over-pressure at 30-40°C (temperature of the leaching experiment) and keeping the ammonia arid carbon dioxide content of the s o l u t i o n constant. From the experiments i t was concluded that: 1) The rate equation f o r manganese p r e c i p i t a t i o n out of s o l u t i o n was: d[Mn c o m p l e x ] d [ M n 0 2 ] dt dt 3 2 k. ( f 0 2 ) [Mn complex]^ ^ I t i s seen that low temperature, large oxygen over-pressure and high manganese content of the s o l u t i o n accelerate the p r e c i p i t a t i o n process.. 2) The nature of the p r e c i p i t a t e could be i d e n t i f i e d as MnO^ only a f t e r a 24 hour heat treatment of the product under vacuum at 200°C. The product was therefore thought to p r e c i p i t a t e i n the amorphous state. 3) Increasing the s t i r r i n g rate had only a small e f f e c t on the rate of p r e c i p i t a t i o n of manganese. This suggested that the tr a n s f e r of oxygen from the gas phase to the l i q u i d phase was not a r a t e - l i m i t i n g step. F. APPLICATIONS AND ENTENSIONS OF THE PRESENT WORK The r e s u l t s obtained i n the present work suggest a p o t e n t i a l a l t e r n a t i v e to the Dean Leute process f o r the treatment of low-grade manganese ores containing high proportions of i r o n and phosphate. The reductive roast might be replaced by a reduction of the ore into the processing s o l u t i o n using hydrazine hydrate as a reducing agent. A sealed v e s s e l would be necessary so that the reduction step could be completed i n neutral atmosphere at moderate temperature and pressure. For temperatures below 40°C the i r o n i n s o l u t i o n could be maintained at a very low value i n comparison to the dissolved manganese. Any phosphate i n the o r i g i n a l ore would not appear i n the f i n a l s o l u t i o n . The use of a closed v e s s e l might enable the recovery of very pure MnCOg since by heating up to 130°C v i r t u a l l y a l l the manganese i s pr e c i p i t a t e d out of s o l u t i o n while i r o n would remain soluble (even up to 70 g/1) A l t e r n a t i v e l y , r e l a t i v e l y pure manganese dioxide would be ob-tained by passing oxygen through the saturated manganese s o l u t i o n . Further studies on p r e c i p i t a t i o n i n these solutions might enable the production.of battery grade MnO , A d i r e c t e l e c t r o l y s i s of the manganese carbamate containing s o l u t i o n should also be investigated. This work may be extended with success to the leaching of manganese sea nodules. Processing of the sea nodules i n the ammonium carbamate s o l u t i o n with hydrazine as a reducing agent would enable the extraction of manganese while n i c k e l , copper and cobalt would be ex-pected to be reduced to the m e t a l l i c state and remain behind i n the residue. The p r a c t i c a l a p p l i c a t i o n of the above outlined p o s s i b i l i t e s almost c e r t a i n l y l i m i t e d , c u r r e n t l y , by the cost of hydrazine. Future research might be directed towards seeking an a l t e r n a t i v e to hydrazine as a reducing agent or to seeking a l t e r n a t i v e routes f o r production of hydrazine s o l u t i o n s . 110 REFERENCES 1. 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Welsh, J.Y., U.S. Patent 2, 834, 726, (1958). 54. Fortune, W.B. and Mellon, M.G., Ind. Eng. Chem, Anal. Ed., 10, 60, (1938). i e 55. G. Chariot , Masson and C P a r i s , "Nouveau T r a i t e de Chimie Analytique", p.786, (1961). i e 56. G. Chariot , Masson and C P a r i s , "Nouveau T r a i t e de Chimie Analy-tique", p.850, (1961). 57. Nakamoto, K., "Infrared Spectra of Inorganic and Coordination Com-pounds", Wiley, New York, (1963). 58. Bernard, ,M,A. and Borel, M.M., B u l l . Soc. Chim., 6, 2362, (1968). 59. Vetter, K.J. and Jaeger, N., Electrochem. Acta., 11, 401-419, (1966). 60. Parks, G.A. and De Bruyn,P.L. , J . Phys. Chem., 66, 967, (1962). 113 TABLE I E f f e c t of [Na2S0 3] concentration on the leaching of goethite at 130°C (Figure 11) and - 0.17 g/1 anthraquinone [Na„S0 3] % Fe Dissolved % Fe Dissolved (Molarity,M) a f t e r 90 mins. A f t e r 300 mins. 0.013 2.88 6.48 0.026 3.84 8.37 0.053 4.03 8.52 0.079 3.77 8.50 To t a l i r o n added: 0.027 m o l e s / l i t e r 114 TABLE II E f f e c t of [^H^.I^SO^] concentration on the leaching of goethite at 130°C. (Figure 17). [N H..H SO,] % Fe Dissolved % Fe Dissolved (Molarity, M) After90 mins. A f t e r 270 mins. 0.015 3.20 5.80 0.031 4.93 9.80 0.061 7.57 14.74 To t a l i r o n added: 0.027 m o l e s / l i t e r TABLE III E f f e c t of temperature on the s o l u b i l i t y of FeO and MnO. (Figures 9, 21). Temperature Mn Fe (°C) ( g / l i t e r ) ( g / l i t e r ) 21 - 47.5 40 67.0 70.0 50 68.3 -60 70.0 67.2 80 46.8 73.2 100 21.4 68.8 124 1.82 -130 - 69.5 140 — 68.8 116 TABLE IV Log^log pl o t between [S0 o ]. . . . - [Mn complex) and time. 3 i n i t i a l tr * (Figure 23) . Time Log (time) [S0 3 ] l n l t l a l ([S0 3 ] i n i t J a l [Mn complex]) Log (min) 100 x A (M/l) (M/l) 20 30 40 50 1.304 1.478 1,600 1.699 0.053 0.053 0.053 0.053 0.0385 0.0334 0.0304 0.0275 0.585 0.524 0.481 0.440 20 30 40 50 70 1.304 1.478 1.600 1.699 1.845 0.106 0.106 0.106 0.106 0.106 0.0840 0.0769 0.0718 0.0678 0.0635 0.924 0. 886 0.856 0.831 0.802 117 TABLE V Pressure v a r i a t i o n i n the leaching of p y r o l u s i t e with hydrazine hydrate. (Figures 24, 25). Time Temperature To t a l Pressure (min) (°C) (psi) 5.00 30 48.0 20.0 7.50 42 105.0 25.0 12.00 66 260.0 60.0 16.00 84 340.0 105.0 135.00 120 500.0 240.0 Equilibrium Pressure (psi) Time Temperature Overpressure Mn i n Solution Mn i n Solution (min) (°C) (psi) (g/1) (g/1) measured calculated 5,00 30 28 - 4.51 7.50 42 .80 - 12.40 12.00 66 200 - 28.80 16.00 84 235 30.60 32.20 135.00 120 260 31.85 32.30 118 T A B L E V I E f f e c t of [Na 2S0 3] concentration on the leaching of p y r o l u s i t e at 130°C (Figure 22). [Na 2S0 3] (Molarity, M) % M i n Solution A f t e r 30 mins. % Mn i n Solution A f t e r 200 mins. 0.053 0.106 0.630 44.5 63.3 85.7 T o t a l manganese added: 89.0 100.0 69.0 0.043 m o l e s / l i t e r TABLE VII E f f e c t of temperature on the rate of leaching of p y r o l u s i t e (50-70 mesh) i n 0.0874 M N„H..H_0 (Figure 27). Temperature 1000 (°C) T(°K) 61 2.995 50 3.100 43 3.161 38 3.215 30 3.300 Rate Log Rate (gr Mn/ . ) 6  V 6 'mm. 0.04246 -3.155 0.03180 -3.445 0.02076 -3.880 0.01280 -4.350 0.00600 -5.120 TABLE VIII E f f e c t of temperature on the rate of leaching of p y r o l u s i t e (65-150 mesh) i n 0.0874 M N^.E^O (Figure 28). Temperature 1000 Rate Log Rate CC) T(°K) .(gr Mn/ . ) 6 mxn. 47.5 3.120 0.04180 -3.170 40.0 3.195 0.03625 -3.315 35.5 3.240 0.02560 -3.660 30.0 3.300 0.01553 -4.150 22.0 3.400 0.00597 -5.150 121 TABLE IX E f f e c t of [^H^.H 0] concentration on the rate of leaching of p y r o l u s i t e (65-150 mesh) at 30°C. (Figure 30). [N 2H 4.H 20] (Molarity, M) L o g 1 0 [ N 2 H 4 ' H 2 0 ] Rate (gr Mn/ ) mm. L o g 1 Q Rate 0.00437 0.0437 0.0874 0.1748 0.2622 -2.360 -1.360 -1.059 -0.758 -0.582 0.00151 0.01065 0.01553 0.02210 0.02854 -2.822 -1.973 -1.809 -1.656 -1.545 T o t a l manganese introduced: 0.072 m o l e s / l i t e r 

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