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Factors influencing the recovery of molybdenum during the hypochlorite leaching of low grade molybdenite-copper.. Mounsey, Diana Mary 1979-03-05

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FACTORS INFLUENCING THE RECOVERY OF MOLYBDENUM DURING THE HYPOCHLORITE LEACHING OF LOW GRADE MOLYBDENITE-COPPER CONCENTRATES by DIANA MARY MOUNSEY .Sc.(Eng.), A.R.S.M., Imperial College of Science and Technology (University of London), 1977 A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF APPLIED SCIENCE in THE FACULTY OF GRADUATE STUDIES (Department of Metallurgical Engineering) We accept this thesis as conforming to the required standard THE UNIVERSITY OF BRITISH COLUMBIA September 1979 Ccj Diana Mary Mounsey, 1979 In presenting this thesis in partial fulfilment of the requirements for an advanced degree at the University of British Columbia, I agree that the Library shall make it freely available for reference and study. I further agree that permission for extensive copying of this thesis for scholarly purposes may be granted by the Head of my Department or by his representatives. It is understood that copying or publication of this thesis for financial gain shall not be allowed without my written permission. Department of M The University of British Columbia 2075 Wesbrook Place Vancouver, Canada V6T 1W5 Date »** QcUeT '777 (ii) ABSTRACT An investigation into the factors influencing the extraction of molybdenum from copper sulphide molybdenite concentrates in alkaline hypochlorite solutions has been carried out. Sodium carbonate buffer components were found to act as complexing agents for the copper, enabling it to exist in solution as a cupri-carbonate species several orders of magnitude more soluble than the thermodynamically stable copper phase at pH 9.0. Completely selective molybdenum extraction is therefore only possible from solutions contain ing no carbonate. The removal of the carbonate produced adverse side effects for the sodium hypochlorite lixiviant as a result of copper hydroxide formation on the mineral surface which led to rapid decomposition of the hypochlor ite by heterogeneous catalysis. This hydroxide salt was found to consist at least partially of tri-valent copper and it was confirmed that sodium hypochlorite has a sufficiently high redox potential at pH 9.0 to enable the oxidation of copper11 > copper11"1" to take place. Experiment also showed that a tri-valent copper carbonate compound precipitated from leaching solutions containing copper after a suitable nucleation period. It was thus esta blished that the copper present in carbonate containing hypochlorite solu-3-tions exists as the tri-valent complex CuCCO^)^ ; that the induction time to precipitation is a function of the total carbonate content of the (iii) system, and that the solid precipitate is also an active catalyst for hypochlorite decomposition. The leaching of molybdenite - copper concentrates in acid solu tions was studied, and results showed that such a process would not be feasible due to the formation and precipitation of copper molybdate, CuMoO^. This compound is very insoluble, but its formation was found to be suppressed by the presence of sodium bicarbonate at pH values greater than 6.0. Several other insoluble molybdate compounds were found to be capable of forming in both acid and alkaline solutions as a result of hypochlorite dissolution of impurity elements contained in the copper sulphide ores. Potential-pH diagrams constructed to show the thermodynamic stability of the molybdate salts of copper, iron, calcium, zinc, lead and cadmium are included. Calcium was isolated as being the most detrimental impurity ele ment in this respect, due to its common occurrence in copper porphyry ores as well as its solubility in chloride containing solutions. A limited study was carried out to show that this solubility increased in proportion to the chloride content and hence the hypochlorite concentra tion of leaching solutions. Calcium carbonate was found to precipitate in preference to calcium molybdate, but the use of silicates and phos phates as calcium suppressants was also considered as a means of avoiding copper - carbonate complexing while maintaining good molybdenum extraction. Xiv) TABLE OF CONTENTS Page CHAPTER ONE : General 1 1.1 Introduction1.2 Current Molybdenum Production 3 1.3 Hydrometallurgical Production of Molybdenum 6 1.4 Literature Review 9 1.4.1 The extraction of molybdenite with sodium hypochlorite solutions 9 1.4.2 The behaviour of copper in i alkaline carbonate solutions 17 1.4.3 Decomposition of sodium hypochlorite in alkaline solution 2 7 1.4.4 Copper111 compounds 3 9 CHAPTER TWO : Experimental 4 7 2.1 Scope of the Present Investigation 4 7 2.2 Materials 49 2.2.1 Natural minerals 49• 2.2.2 Synthetic minerals 5 4 2.2.3 Sodium hypochlorite  4 2.2.4 Chemical reagents 5 5 2.3 Apparatus 5 5 2.4 Experimental procedure 5 7 2.4.1 Leaching experiments 5 7 2.4.2 Copper111 preparation  8 2.5 Analysis 5 8 2.5.1 Chemical analysis 5;8 i) Copper 5'(v) TABLE OF CONTENTS (Cont'd) Page ii) Molybdenum 59 iii) Calciumiv) Sodium hypochlorite 59 v) Sodium chloratevi) Total chloride 60 2.5.2 Instrumental analysisi) Minerals 60 III ii) Copper 1 CHAPTER THREE : Results and Observations 62 3.1 Sodium Hypochlorite Leaching of Copper Sulphide Minerals 63.1.1 Leaching of ground concentrates in the presence of carbonate 62 3.1.2 Effect of carbonate removal 4 3.1.3 Leaching of synthetic copper sulphides 67 3.1.4 Determination of decomposition products 9 3.1.5 Sodium hypochlorite leaching of massive samples of copper sulphide minerals 72 3.1.6 Variation of total carbonate content during the leaching of ground copper sulphide minerals 75 3.1.7 Effect of varying the hypochlorite concentration 79 3.1.8 Effect of hypochlorite removal 84 3.2 Copper111 87 3.2.1 Analysis of copper111 90 Effect of acidification 90 Evaluation of oxidation state by idometric titration 90 Measurement of oxygen evolu tion using a mercury column 92 (vi) TABLE OF CONTENTS (Cont'd) Page Gas chromatography 94 Instrumental analysis 9i) Magnetic susceptibility 94 ii) Infra-red spectra 95 iii) X-ray diffraction 96 3.2.2 Precipitation of copper11-1- in the presence of carbonate 97 Effect of seeding Effect of acidification 98 I.R. spectra 98 X-ray analysis 9 3.2.3 Hypochlorite decomposition studies 100 3.2.4 Analysis of leaching solutions 100 3.3 Sodium Hypochlorite Oxidation of Molybdenite 102 3.3.1 Sodium hypochlorite leaching of molybdenite and copper sulphide minerals at pH 9.0 104 3.3.2 Sodium hypochlorite leaching of molybdenite and copper sulphide minerals at pH 5.5 114 3.3.3 Copper molybdate 120 Solubility of copper molybdate 127 Catalyzed decomposition of hypochlorite in the presence of copper molybdate 130 X-ray analysis 132 3.3.4 Solubility of calcium in hypochlorite solutions 133 Sodium hypochlorite leaching of calcium sulphate and calcium carbonate minerals 137. Effect of chloride concentration on calcium solubility 140 Removal of calcium from solution 14 0 (vii) TABLE OF CONTENTS (Cont'd) Page CHAPTER FOUR : Discussion 154 4.1.1 Sodium hypochlorite leaching of copper sulphide minerals in the presence of carbonate 154 4.1.2 Removal of carbonate from the system 156 4.1.4 Hypochlorite decomposition products 158 4.1.5 Sodium hypochlorite leaching of massive samples of copper sulphide minerals 160 4.1.6 Variation of the total carbonate content during leaching 161 4.1.7 Effect of varying the hypochlorite concentration 165 4.1.8 Effect of hypochlorite removal 166 4.2.1 Analysis of copper111 170 4.2.2 Precipitation of copper111 in the presence of carbonate 172 4.3 Sodium Hypochlorite Leaching of Molybdenite 173 4.3.1 Sodium hypochlorite leaching of molybdenite and copper sulphide minerals at pH 9.0 174 4.3.2 Sodium hypochlorite leaching of molybdenite and copper sulphide minerals at pH 5.5 174 4.3.3 Copper molybdate 9 4.3.4 Solubility of calcium in hypochlorite solutions 183 Sodium hypochlorite leaching of calcium minerals 196 Effect of chloride concentra tion on calcium solubility 196 Removal of calcium from leaching solutions 200 (viii) TABLE OF CONTENTS (Cont'd) Page CHAPTER FIVE 202 5.1 Conclusions 205.2 Suggestions for Future Work 204 REFERENCES 206 APPENDICES 211 Appendix A Tables of experimental results 211 Appendix B 268 B.l Calculation of surface areas for ground mineral samples 268 B.2 Calculation of available copper for precipitation on mineral surface after copperHI formation 269 B. 3 Hypochlorite redox equilibria at pH 9.0 26Appendix C Equilibria for metal-molybdate Eh-pH diagrams 271 C. l Cu-H20-Mo04C.2 Ca-H20-Mo04 277 C.3 Pb-H20-Mo04 280 C.4 Zn-H20-Mo04 5 C.5 Fe-H20-Mo04 288 C.6 Cd-H20-Mo04 293 (ix) LIST OF TABLES X-ray diffraction pattern for covellite X-ray diffraction pattern for chalcocite X-ray diffraction pattern for Determination of copper"'""'""'" by titration X-ray diffraction pattern for copper11 oxides X-ray diffraction pattern for carbonate Effect of mixing and ageing equi-molecular volumes of Na„MoO. and CuSO„ at pH 5.3 and pH 9.0 X-ray diffraction pattern for copper molybdate 2- 2" 3-Effect of Si03 , C03 , and P04 (sodium salts) on calcium precipitation from OC1 solutions at pH 9.0 Strongest X-ray peaks for various copper molybdate species Comparison of experimental and literature values for solubility of i) CaS04 and ii) CaCO^ in chloride solutions of varying strengths chalcopyrite idometric 111 copper and III copper NaOCl leaching of covellite with carbonate buffers NaOCl leaching of chalcocite with carbonate buffers NaOCl leaching of chalcopyrite with carbonate buffers (x) LIST OF TABLES (Cont'd) IV Effect of separating leaching solution from mineral slurry during agitation of CuFeS2 in NaOCl V Effect of separating leaching solution from mineral slurry during agitation of Cu2S in NaOCl VI NaOCl leaching of covellite in the absence of carbonate buffers VII NaOCl leaching of chalcocite in the absence of carbonate buffers VIII NaOCl leaching of chalcopyrite in the absence of carbonate buffers IX Microprobe examination of massive chalcocite, before and after leaching in NaOCl X Microprobe examination of massive covellite, before and after leaching in NaOCl XI NaOCl leaching of cupric sulphide in the absence of carbonate buffers XII NaOCl leaching of cuprous sulphide in the absence of carbonate buffers XIII NaOCl leaching of covellite with variable [C032~]T at pH 9.0, 35°C XIV NaOCl leaching of chalcocite with variable [C032-]T at pH 9.0, 35°C XV NaOCl leaching of synthetic sopper sulphides with 5 g/1 [C032-]T XVI NaOCl leaching of covellite in the presence of 5 g/1 [C032-]T and 20 g/1 [NaOCl] XVII Agitation of covellite in the presence of carbonate ± NaCl (xi) LIST OF TABLES (Cont'd) Page XVIII NaOCl leaching of covellite in the presence of 10 g/1 [C032~]T and Na2Mo04 XIX Sodium hypochlorite decomposition in the presence of tri-valent copper salts XX NaOCl oxidation of molybdenite at pH 9.0 XXI NaOCl oxidation of 'reagent grade' molybdenum disulphide 228 229 230 231 XXII NaOCl oxidation of molybdenite and covellite ± carbonate buffers 232 XXIII NaOCl oxidation of molybdenite and chalcocite ± carbonate buffers 233 XXIV NaOCl oxidation of molybdenite and chalcopyrite ± carbonate buffers 234 XXV NaOCl oxidation of molybdenum disulphide and chalcopyrite ± carbonate buffers 235 XXVI NaOCl oxidation of molybdenite in the presence of copper sulphate solution 236 XXVII NaOCl oxidation of covellite in the presence of sodium molybdate solution XXVIII NaOCl oxidation of molybdenite at pH 5.5 XXIX NaOCl oxidation of molybdenite and chalcopyrite at pH 5.5 in the absence of carbonate 237 238 239 XXX NaOCl oxidation of molybdenite and covellite at pH 5.5 in the absence of carbonate 240 XXXI NaOCl oxidation of molybdenite and covellite at pH 5.5 in the presence of NaHCO 241 XXXII NaOCl leaching of chalcopyrite at pH 5.5 ± NaHCO„ 242 (xii) LIST OF TABLES (Cont'd) XXXIII NaOCl oxidation of molybdenite and chalcocite at pH 6.5, with NaHCO^ XXXIV NaOCl oxidation of molybdenite and chalcocite at pH 6.5, without NaHC03 XXXV NaOCl oxidation of molybdenite and cupric sulphide at pH 6.0, with NaHCO^ XXXVI NaOCl leaching of molybdenite and cupric sulphide at pH 6.0, without NaHC03 XXXVII NaOCl leaching of molybdenite and chalcopyrite at pH 7.0, with NaHC03 XXXVIII NaOCl leaching of molybdenite and chalcopyrite at pH 7.0, without NaHC03 IXL NaOCl decomposition in the presence of copper molybdate at pH 5.5 XL Determination of the solubility of copper molybdate at pH 5.0-XLI NaOCl oxidation of molybdenite at pH 10.0 XLII NaOCl oxidation of molybdenite and covellite at pH 10.0 in the absence of carbonate buffers XLIII NaOCl leaching of molybdenum disulphide and cuprous sulphide in the absence of carbonate buffers XLIV NaOCl leaching of molybdenum disulphide and cuprous sulphide at pH 9.0, in the presence of carbonate buffers XLV NaOCl oxidation of CuFeS2; Cu2S and CuS at pH 9.0 in the absence of carbonate buffers XLVI NaOCl leaching of molybdenum disulphide and chalcocite in the presence of 0.1 g/1 calcium as CaCl2 solution (xiii) LIST OF TABLES (Cont'd) Page XLVII NaOCl leaching of molybdenite and chalcocite in the presence of CaCl2 and Na2C03/NaHC03 257 XLVIII NaOCl leaching of molybdenite and calcite 258 IL NaOCl leaching of powdered calcite at pH 9.0 259 L NaOCl leaching of Plaster of Paris at 9.0 260 LI Effect of increasing chloride concentration on calcium dissolution from CaSO^-^H-jO 261 LII Effect of increasing chloride concentration on calcium dissolution from CaC03 261 LIII Effect of carbonate on calcium content in hypochlorite solutions 262 LIV NaOCl leaching of covellite and molybdenite with 5 g/1 carbonate 263 LV NaOCl leaching of molybdenite and covellite with 10 g/1 carbonate 264 LVI NaOCl leaching of covellite in the presence of sodium molybdate and 5 g/1 carbonate 265 LVII NaOCl leaching of molybdenum disulphide and chalcocite with 10 g/1 carbonate 266 LVIII NaOCl leaching of molybdenite and chalcopyrite at pH 9.0 in the presence of 1 g/1 Na2Si03 267 (xiv) LIST OF FIGURES Conventional flowsheet for production of molybdenum trioxide Stability relations for copper compounds in the system Cu-H20-02~S-Co2 Potential-pH diagrams for the system Cu-C02-H20, a) and b) Activity ratio diagram and solubility diagram for the Cu-C02~H20 system Experimental leaching apparatus Oxidation of copper sulphide minerals by NaOCl Effect of separating leaching solution from mineral slurry NaOCl consumption during leaching of copper sulphides Effect of carbonate removal on NaOCl decomposition NaOCl decomposition during leaching of synthetic copper sulphides Sodium chlorate production during NaOCl decomposition a) Effect of varied carbonate content on Cu dissolution from covellite ' b) Cu dissolution and NaOCl decomposition for covellite leaching with 10 g/1 [co32"]T Effect of varied carbonate content on NaOCl decomposition (covellite) Effect of varied carbonate content on Cu dissolution from chalcocite (xv) LIST OF FIGURES (Cont'd) Page 15 Effect of varied carbonate content on NaOCl decomposition (chalcocite) 81 16 Cu dissolution and NaOCl decomposition for cupric sulphide leaching (5 g/1 fC032-]T) 82 17 Cu dissolution and NaOCl decomposition for cuprous sulphide leaching (5 g/1 tco32-]T) 83 18 Cu dissolution and NaOCl decomposition for covellite leaching (20 g/1 [NaOCl]) 85 19 Agitation of covellite in solutions of Na2C03/NaHC03 ± NaCl 86 20 Cu dissolution and NaOCl decomposition for covellite with 10 g/1 [C032-] over a period of 6 hours 89 21 NaOCl decomposition in the presence of oxide and carbonate salts of copper111 101 22 NaOCl oxidation of molybdenite at pH 9.0 103 23 NaOCl oxidation of reagent grade molyb denum disulphide at pH 9.0 105 24 NaOCl oxidation of '98%+' molybdenum disulphide at pH 9.0 106 25 NaOCl oxidation of molybdenite and covellite a) in the presence of carbonate buffers 108 b) in the absence of carbonate buffers 109 26 NaOCl oxidation of molybdenite and chal cocite a) in the presence of carbonate buffers 110 b) in the absence of carbonate buffers 111 27 NaOCl oxidation of molybdenite and chal copyrite a) in the presence of carbonate buffers 112 b) in the absence of carbonate buffers 113 (xyi) LIST OF FIGURES (Cont'd) Page 28 NaOCl oxidation of molybdenite at pH 5.5 116 29 NaOCl oxidation of molybdenite and chalcopyrite at pH 5.5 117 30 NaOCl oxidation of molybdenite and covellite at pH 5.5 a) in the absence of bicarbonate 118 b) in the presence of bicarbonate 9 31 NaOCl oxidation of molybdenite and chalcocite at pH 6.5 a) in the presence of bicarbonate 121 b) in the absence of bicarbonate 2 32 NaOCl oxidation of molybdenite and cupric sulphide at pH 6.0 a) in the presence of bicarbonate 123 b) in the absence of bicarbonate 4 33 NaOCl oxidation of molybdenite and chalcopyrite at pH 7.0 a) in the presence of bicarbonate 125 b) in the absence of bicarbonate 6 34 Determination of the solubility of copper molybdate at pH 5.0 128 35 NaOCl decomposition in the presence of copper molybdate at pH 5.5 131 36 NaOCl oxidation of molybdenum disulphide and cuprous sulphide at pH 9.0 134 37 NaOCl oxidation of molybdenite and covellite at pH 10.0 a) in the presence of carbonate buffers 135 b) in the absence of carbonate buffers 6 38 NaOCl leaching of chalcopyrite, chalcocite and covellite at pH 9.0 138 39 NaOCl leaching of powdered calcite and Plaster of Paris at pH 9.0 139-40 NaOCl leaching of molybdenite and calcite in the absence of carbonate 141 (xvii) LIST OF FIGURES (Cont'd) Page 41 Effect of chloride concentration on Ca dissolution from CaS04-%H20 142 42 Calcium sulphate solubility as a function of chloride content 143 43 Effect of chloride concentration on Ca dissolution from CaCO^ ^44 44 Calcium carbonate solubility as a function of chloride content 145 45 Effect of carbonate on calcium in hypochlorite solutions 147 46 NaOCl leaching of covellite and molyb denite at pH 9.0 with 5 g/1 [CO 2~] and 10 g/1 [C032"]T 148 47 NaOCl leaching of chalcocite and molyb denum disulphide with 10 g/1 [C032-]T 14 9 48 NaOCl leaching of covellite in the presence of NaMoO^ and 5 g/1 carbonate 151 49 NaOCl oxidation of molybdenum disulphide and chalcopyrite in the presence of Na2Si03 15 2 50 Mercury column for measurement of gas evolution from Cu111 9 2 51 Potential-pH diagram for the system Cu-CO -H 0 with 10-1 M [C032~] 17 5 52 Eh-pH diagram for (oxide species) the system Cu--H20--MoO. 4 181 53 Eh-pH diagram for the (hydroxide species) system Cu--H20--MoO,, 4 18 2 54 Eh-pH diagram for the system Ca--H20--Mo0„ 4 18-5 55 Eh-pH diagram for the system Pb--H20--Mo0„ 4 18 7 56 Eh-pH diagram for the system Zn--H20--Mo0„ 4 18 8 (xviii) LIST OF FIGURES (Cont'd) Page 57 Eh-pH diagram for the system Fe-H20-MoO 189 58 Eh-pH diagram for the system Cd-H^-MoO^ 191 (xix) LIST OF PLATES Facing Page I Effect of hypochlorite solutions on massive chalcocite samples 7 3 II Effect of hypochlorite solutions on massive covellite samples 7 4 (xx) Acknowledgement The author is greatly indebted to Dr. I.H. Warren for his generous advice, constant interest and invaluable encouragement through out the period of this study. The assistance provided by members of Faculty and technical staff of the Department of Metallurgy, and the help and companionship of fellow graduate students have been much appreciated. Thanks are also extended to Dr. F. Weinberg for the many musical interludes which provided respite from, if not inspiration for, the rigours of research. Financial support in the form of a CIL fellowship is gratefully acknowledged. - 1 -CHAPTER ONE 1.1 Introduction Molybdenite, Mc^/ is the principal source of the molybdenum metal used commercially. It occurs both in primary molybdenum deposits and in more complex ore bodies associated with other sulphide minerals. Powellite, CaMoO^; Wulfenite, PbMoO^ and hydrated ferric molybdate, FeMo^O^ * 8H2° are t'ie most common secondary molybdenum minerals but they are not considered to be of economic importance at the present time. Molybdenum is one of the less common elements with an average concentration in the earth's crust of 0.001%."'" Its principal use is as an alloying element in iron, steel and certain non ferrous metals. Either alone or in the presence of other elements such as nickel, vanadium, manganese and chromium, molybdenum enhances the strength, toughness and hardenability of these materials and can also improve corrosion resistance. Molybdenum compounds have wide application as lubricants, chemical reagents and as catalysts. They are also used extensively in paint and pigment manufacture. Consumption of molybdenum has increased on a world wide basis at an average rate of about 5% per annum for the past decade, and demand is expected to grow at least at this level to the year 2000 and beyond. Western world production for 1977 amounted to 89,500 metric tons, of 2 which the United States produced 68% and Canada 17%. Current estimates of molybdenum reserves indicate that supplies should be adequate in the foreseeable future. A few high grade deposits containing up to 20% molybdenum as MoS0 in rich quartz veins exist, but over 95% of current - 2 -and previous production has been from low grade porphyry type deposits and it seems likely that this figure will increase as more of the richer ores are worked out. Porphyry ores can be characterized into two types: 1) Molybdenum porphyries; where molybdenite is the only contained mineral of economic importance and averages 0.15 - 0.8% MoS2. 2) Copper-molybdenum porphyries; from which copper is extracted as the primary metal value and molybdenum is recovered as a by-product. The molybdenite content of these ores is much lower and ranges from 0.005 - 0.01% MoS2. Copper mineralization consists of primary and secon dary sulphide ores including chalcopyrite, CuFeS,,; covellite, CuS; chal cocite, Cu2S; bornite, Cu^FeS^; and enargite, Cu^sS^. There is often a leached cap associated with the ore body containing secondary and tertiary oxides such as cuprite, Cu20; malachite Cu2(OH)2C03; tenorite, CuO, and 2 chrysocolla, Cu^H^Si^O^Q(OH)g. The overall copper content averages 0.8%. ' By-product molybdenum recovery from porphyry copper ores accounts for almost 50% of total world production, but the fact that ore grades involved are so low means high recoveries are difficult to obtain in many cases. Separation of molybdenite from copper ores is achieved by selective flotation in almost all cases. The principal aim of production plants treating these ores is to optimize copper recovery and values of about 80% are typically obtained. The overall molybdenum yield is usually much lower and averages 50 - 60%. The recent depressed state of the copper market followed by rapid increases in the price of molybdenum to a current value of $9/lb, has given by-product recovery a more important perspective. In cases where copper ores are relatively lean the two metals are almost equal in terms of end product value. The molybdenum - 3 -not recovered is therefore of economic significance and represents a substantial loss of metal resources. 1.2 Current Molybdenum Production A molybdenite concentrate averaging 90% MoS^ is the usual product of beneficiation of run of mine ore. This is then converted to technical grade molybdenum tri-oxide, MoO^, for further processing into ferro-molybdenum and other compounds. The conventional method of obtaining molybdenum tri-oxide follows a flotation-roasting method as outlined in Figure 1. Molybdenum-bearing copper ores are most often mined by open pit methods, after which the run-of-mine ore is subjected to crushing and wet grinding to give a product yielding 50 - 70% minus 200 mesh material. Beneficiation by flotation is used in all cases, largely due to the good natural flotability of molybdenite, making it the easiest method by which separation from copper can be achieved. Complex copper-molybdenum ores are subjected to a collective rougher flotation to give a rougher concentrate averaging 0.3% Mo and 12.5% Cu with some iron, sulphur and calcium oxide. This concentrate is re-ground and cleaned to give an up graded copper flotation concentrate averaging 25 - 35% Cu. The molyb denite is then separated by a selective flotation stage, carried out either by depressing the copper and allowing the molybdenite to float, or by the opposite technique of depressing the molybdenite and floating off the other constituents of the copper concentrate. The former method is generally preferred, giving an initial molybdenum rougher concentrate - 4 -Run of mine ore L_ CRUSHING 8 GRINDING COLLECTIVE ROUGHER FLOTATION j Rougher concentrate (!2-4%Cu,0-3% MoS2) • REGRINDING, CLEANING i Copper flotation concentrate (25-35%Cu) >D-2%Mol <Q-2% Mo THICKENING Cu THICKENING SELECTIVE ROUGHER FLOTATION FILTRATION Mo I I Mo rough e r c onc• (35% Mo S ) I 2 Cu. concentrate DRYING T CLEANING, REGRINDING Mo flotation c one. ( 80% M o SJ*°'50/o(fu THICKENING Mo LEACH Cu FILTRATION DRYING Mo concentrate ( 90%MoS2) I Mo03 PRODUCTION Figure 1: Conventional flowsheet for the production of molybdenum trioxide. - 5 -which is purified further in a counter-current circuit until a 90% MoS2 product, containing less than 0.5% Cu, is obtained. A low tempera ture roast is sometimes incorporated into the cleaning circuit to effect partial oxidation of copper sulphide minerals, thereby retarding their flotability. With careful use of the correct flotation reagents for the mineral ization present in a specific ore, as well as optimization of pH by the use of modifiers, and suitable choice of temperature, up to 85% of the contained molybdenum can be recovered. However, as noted above, condi tions are usually selected to give maximum copper recovery, and this results in a much lower average level of molybdenum recovery. Loss of molybdenum can occur on account of the following factors:"''^ i) The ore is ground to a size suitable for optimum copper recovery and this is usually coarser than the ideal size for molybdenite. This problem is often accentuated during periods of depressed copper prices when an even coarser grind is used to increase throughput. ii) Molybdenite can be lost from the copper rougher flotation tailings because it often exists as large flakes which are trapped by quartz particles. iii) Finer particles are lost from the overflow of middling thickeners after production of the molybdenum rougher concentrate, and this loss is increased by the presence of flocculants. iv) The presence of varying copper sulphide mineralization in an ore can necessitate the use of different depressants due to differ ences in surface properties, for example between chalcopyrite and covellite. Certain reagents can also inhibit molybdenite flotation. - 6 -Naturally floating materials such as sulphur, talc, graphite and coal can be equally detrimental to molybdenum recovery,, and the presence of molyb denum oxides is also undesirable. v) Porphyry ores often have a non-uniform distribution of molybdenum, resulting in large variations of head values in the feed material to the flotation circuit, which present additional difficulties in maintaining continuous high molybdenum recoveries. In many cases the final molybdenite concentrate is subjected to a sulphuric acid or sodium cyanide leach after re-cleaning and prior to thickening and filtration. This serves to remove any copper remaining in the concentrate. Conversion to molybdenum tri-oxide is by means of an oxidizing roast, usually carried out in "Nichols-Herreshoff" type roasters which enable the temperature to be closely monitored and kept between 600 - 700°C. The final MoO^ product contains 56 - 62% Mo and has maximum allowable 7 sulphur and copper contents of 0.25% and 0.75% respectively. •* 1.3 Hydrometallurgical Production of Molybdenum While existing reserves of molybdenite are expected to be adequate to supply an increased demand of up to 5% per annum for some time to come, production of molybdenum on a day-to-day basis is not always adequate. In 1977 there was a production deficit of some 9 million pounds of molyb-2 denum, which had to be supplied from existing stock piles. It would thus seem desirable to research into ways of increasing current molybdenum recovery. As more of the world's higher grade deposits are worked out, molybdenum bearing copper ores typical of those found in British - 7 -Columbia will assume an even more important role as a source of the metal. For these types of ores, a hydrometallurgical method for molybdenum recovery has the potential of giving higher metal extraction and better utilization of raw materials than can be obtained by conventional methods. A hydrometallurgical route could also eliminate either or both of the preliminary and final roasting stages. This would be preferable from an environmental point of view by preventing the evolution of SO^ into the atmosphere. Molybdenite is generally unreactive chemically and its dissolu tion in aqueous media entails oxidation to the MoVI state. Molybdenum will exist in water in oxidation states of +3 to +6, but only the +6 state is stable over a wide range of potential-pH values. The other valence states can be stabilized by certain complexing agents. Oxidation of molybdenite from Mo"tV to MoVI can be effected by dissolution in hot nitric acid, concentrated sulphuric acid, aqua regia and by oxygen under pressure in aqueous solution; as well as by various oxidizing agents such as acid sodium chlorate and sodium hypochlorite. On the alkaline side of 2- VI neutrality the monomer MoO^ is the usual form of Mo , but at pH values -4 below 7.0, and with a molybdenum content in excess of 10 M, this ion Q is usually polymerized. The molybdate ion readily combines with most metal cations, and with the exception of the alkali metals, an insoluble salt is precipitated. The anion is able to attain rapid equilibrium in aqueous solution which is of considerable advantage when considering a wet method for molybdenum extraction. Several studies have been carried out in the last two decades to investigate possible methods of extracting molybdenum by leaching - 8 -molybdenite. The majority of these suggested the use of nitric acid or a strong oxidant in alkaline solution as a lixiviant, with subsequent recovery of molybdenum by means of an ion-exchange or solvent extraction technique, or by reductive precipitation in acid solutions using metal (mercury, zinc, cadmium or molybdenum) or gaseous (hydrogen, hydrogen sulphide, sulphur dioxide) reductants. Investigations into the use of other reducing agents such as hydrazine and sodium sulphite have recently 9 been carried out in this department. The biggest loss of molybdenum in current processes occurs during the collective rougher flotation stage and amounts to some 36%. Further losses of about 10% occur between this point and the final molybdenite product. The most logical hydrometallurgical treatment should therefore use run of mine ore as the starting material for a leaching circuit, but this would not be economically practical for many porphyry ores which contain less than 0.02% Mo. Most of the processes envisaged to date have thus worked on the basis of a leaching circuit in conjunction with existing flotation methods, using either molybdenum or copper rougher concentrates as feed material. Both methods appear to be feasible provid ed the leach is sufficiently selective for molybdenum over copper and other sulphide minerals. While the treatment of un-beneficiated ores is not generally acceptable, the possibility of extracting molybdenum by in-situ leaching or solution mining methods, could have considerable potential value. The advantages of solution mining in terms of reduced capital and operating costs and a reduction in materials handling, a more efficient utilization of resources and less environmental disturbances, as well as the - 9 -disadvantages associated with handling and storing large volumes of leaching solutions and careful prediction and control of their under ground movements, have been well documented in the literature.10 Low grade copper-molybdenum ores could be very amenable to this type of leach ing provided the lixiviant used gave a sufficiently rapid and selective leach of the contained molybdenum values, and could be economically produced. 1.4 Literature Review 1.4.1 The extraction of molybdenite with  sodium hypochlorite solutions Cox and Schellinger carried out laboratory scale tests in an investigation to determine the feasibility of a leach-ion exchange type process for producing molybdic oxide, in which they used sodium hypochlor ite as a lixiviant for three different grades of material containing 0.015%, 1,05% and 63% MoS2 respectively.11 Dresher, Wadsworth and Fassel had previously studied the kinetics of dissolution of molybdenite in alkaline solutions with potassium hydroxide at temperatures in excess of 100°C and pressures up to 700 psi, and had found that suitable rates of 12 leaching could be obtained with these conditions. Cox et al. concluded that use of hypochlorite reagents could give a much more economical process than one necessitating a high pressure-high temperature leach, and optimum conditions for molybdenum recovery were obtained with a 3% solution of sodium hypochlorite at room temperature. They put forward the following stoichiometric relationship as being operative during the leach: - 10 -7NaOCl + MoS2 + 4e >- MoC>4~ + S2032 + 7NaCl (1) The rate of leaching was found to decrease to zero after about thirty minutes, at which point molybdenum extractions ranging from 91 - 99% had been achieved. Both Cox and Schellinger, and Dresher et al. observed that thiosulphate ions were present in the leaching solutions, but the 2-latter workers concluded that S2C>3 was only an intermediate reaction product and that after complete oxidation, only molybdate and sulphate ions would exist in solution: 9 - 2-2-MoS2 + - °2 + 60H y Mo04 + 2S04 + 3H2° (2) A kinetic study of the oxidation of molybdenite using sodium hypochlor-13 ite solutions was reported by Iordinov and Zelikman in 1961. They used a reagent concentration varying from 15 to 60 g/1 OC1 at tempera tures of 20 - 80°C and their starting material was pure MoS2- A detailed description as to the theory of the different stages involved in molyb denite oxidation from a physico-chemical standpoint is given, and the overall reaction was found to be: MoS2 + 9NaOCl + 6NaOH > Na2MoO + 2Na2S04 + 9NaCl + H20 (3) This is in good agreement with the findings of Dresher et al., with 9 moles of oxidant required per mole of molybdenite. Zelikman also found that thiosulphate was present initially but that its concentration went through a maximum value and had reached zero by the end of the reaction. The oxidation of molybdenite to molybdate was found to be first order, with an activation energy of about 5.25 kcal/mole over the temperature range under study. The rate constant increased both with increased - 11 -hypochlorite content and with increased temperature. Shapiro and Kulenkeva carried out tests using a sodium hypochlorite lixiviant to try and obtain efficient molybdenum recovery from intermediate concentrates containing 2 - 6% Mo produced during the processing of 14 disseminated molybdenite or polysulphide ores. Leaching experiments were done in the presence of sodium carbonate and it was found that a 30 g/1 solution of hypochlorite gave a maximum rate of molybdenite decom position at 50°C: MoS2 + 9NaOCl + 3Na2C03 HNa MoC>4 + 2Na2S04 + 9NaCl + 3CC>2 (4) The leach was stated as being selective for molybdenum because any other metal sulphides present would be converted to insoluble carbonates, as represented by the following equation: . MeS. + 4NaOCl .+ Na2C03 v Na2S°4 + 4NaC1 + MeC03 (5) Bhappu, Reynolds and Stahman made a study of the sodium hypochlorite leaching of molybdenite from three different sources: a high grade con centrate containing 96 - 98% MoS2, low grade ores typical of those found in a waste dump (0.2 - 0.8% MoS2) and medium grade lumps of molybdenite averaging 60%. They calculated the stoichiometric relationship for hypochlorite consumption to be the same for each test sample, indicating negligible consumption by gangue material, and in agreement with the findings of Iordinov and Zelikman, that nine moles of OC1 were required to oxidize each mole of MoS . 2 They also confirmed reports by Choppin and Faulkerberry1^ which showed that provided an excess of hypochlorite is present in solution, the - 12 -end product of sulphide oxidation is entirely sulphate. Doubt was thus expressed concerning Cox's claim that thiosulphate was present as a final product of hypochlorite oxidation, and Bhappu et al. stressed that any thiosulphate or other intermediate oxidation state of sulphur would, in their opinion, exist only for very short periods of time before being oxidized to sulphate. These authors concluded that such a hypochlorite process offered considerable potential for molybdenite leaching, especially for the treat ment of low grade ores, on account of the seemingly low consumption of reagent by gangue material. 17 In a later paper however, Bhappu et al. reviewed a number of oxidizing agents for the hydrometallurgical treatment of molybdenite ores and concluded that while alkaline hypochlorite solutions gave rapid and selective molybdenum extraction, the reagent was expensive to manufacture, unstable and corrosive; and the only economically feasible method for its generation in this type of leaching system would be in an electrolytic re-generation circuit from the sodium chloride solution produced by the oxidation reaction. Other reagents considered included sodium chlorate in acid solutions, manganese dioxide-sulphuric acid and nitric acid. The possibilities of bacterial leaching as well as in-situ leaching were also reviewed, especially in the presence of non-sulphide ores such as ferrimolybdate and limonite. For these ores and for those containing mixed oxides and sulphides either an acid-chlorate or a basic hypochlorite-carbonate lixiviant seemed feasible, with the latter being generally pre ferred because of potential problems with the precipitation of molybdate compounds in acid solutions. - 13 -Zelikman reviewed the possibility of leaching molybdenite with 18 hypochlorite again in 1970, particularly from low grade concentrates which would probably also contain copper and iron sulphides. Although such sulphides would be oxidized by hypochlorite in alkaline solution, the oxidation products would be insoluble hydroxides which inhibit the reac tion and prevent subsequent dissolution, thus giving a selective molybde num leach. It was pointed out in this article that the hydroxides of certain transition metal elements, including copper and iron, are also active catalysts for the decomposition of sodium hypochlorite in alkaline media. Catalysis can either be homogeneous or heterogeneous, and it was suggested the latter involves the formation of unstable hydroxides of higher valence states, especially for copper. Optimum conditions for concentrates containing 23% Mo and 9% Cu were found to exist with a hypo chlorite content of 30 g/1 and 20 - 30 g/1 free alkali. Higher tempera tures were found to increase the rate of molybdenum extraction, but as this also gave an increased rate of catalyzed hypochlorite decomposition, leaching at room temperature was recommended. In all the above studies molybdenite leaching has been carried out on the alkaline side of neutrality, with pH 10.0 being stated as optimum in several cases. Sodium hypochlorite is much more stable in alkaline solutions than it is in acid media. A pH range of 5.0 - 7.0 was used in an electrooxidation technique investigated by the United States Bureau of Mines for extracting molybdenum and rhenium from low grade 19 ores, despite this being the region in which decomposition of hypochlor ite to chlorate takes place most rapidly. The method was based on in-situ hypochlorite generation by electrolysis in brine ore pulp. Extractions - 14 -of 90 - 99% molybdenum were reported and a stoichiometric ratio of MoS^: 90C1 was again found to be operative (equation 3). The required pH value was maintained by semi-continuous addition of sodium carbonate and a temperature of 30 - 40°C was used. The overall power consumption was reported as 9.7 kwh/lb Mo extracted, with a current density of 0.5 amp/in2. The study was extended to determine the extractability of molyb denum from copper-molybdenum flotation concentrates, averaging 3 - 28% Mo 20 and 2 - 15% Cu. A maximum extraction of 97% Mo was achieved with a final pH of 7.0: this value apparently decreased at both lower and higher pH levels. It was stated that a minimum copper solubility of 1 - 3 ppm occurred at pH 7.0, and in alkaline solutions copper was dissolved to the extent of 25 ppm due to formation of the soluble compound (CuX^CO^ where X = CI or OH . Scheiner et al. suggested copper molybdate com pounds could adversely affect molybdenum recovery, but oxidation of copper sulphides by hypochlorite could be prevented by adding sufficient sodium carbonate to the system. 21 22 In two later papers ' process solutions were treated, and the technique was demonstrated with a prototype cell. 89 - 98% molybdenum extraction was obtained from concentrates containing less than 1% Cu and about 36% Mo. No copper was dissolved in the pH range 5.5 - 7.0, but large amounts of chlorate were produced in this region and an SO2 reduction step was therefore incorporated into the process. This pro cedure also served to reduce the pH of molybdate solutions for subsequent molybdenum recovery by solvent extraction. Power consumption for treating flotation concentrates was 13.7 kwh which is 4% higher than the corresponding - 15 -value for higher grade molybdenite concentrates. In the prototype cell demonstration 'off grade' flotation concentrates containing 16 - 35% Mo and 6 ~ 15% Cu as chalcopyrite and/or chalcocite were treated. It was reported that the molybdenum extraction decreased to only 75% in the presence of higher initial copper contents, and copper molybdate compounds were reportedly found in leaching tails. It was claimed that any chalco pyrite present in the feed material was unaffected by electrolysis, but that chalcocite was oxidized by sodium hypochlorite with subsequent formation of soluble copper compounds. These then reacted with molyb date ions in solution to give insoluble copper molybdates. All subsequent trials were therefore performed with low copper/ high molybdenum concentrates'and no information was given as to how, if at all, the copper molybdate problem could be rectified. In conjunction with these studies, Barr and co-workers investigated the effects of chlorate production on molybdenum recovery during electrooxidation, and concluded that by-product sodium chlorate was detrimental to molybdenum recovery as well as representing a power loss from the system. Both these factors are unfavourable economically. To minimize chlorate produc tion the authors recommended a low temperature-low current density leach in which the pH was also reduced to 4.0 - 5.0. The use of 'flow through' 23 24 cells was shown to be preferable to lead dioxide anode batch cells. ' By measuring respective oxidation rates of molybdenite and chalco pyrite in alkaline solutions, Stumpf and Berube concluded that the selective leaching of molybdenum from typical flotation concentrates 25 was unlikely. The unit consumption of oxygen for the two minerals is apparently very similar, and this was taken to mean that both species - 16 -would be simultaneously attacked, thereby preventing selective molybdenum dissolution. These predictions are in direct contrast to the findings of 26 Warren, Ismay and King, who showed that molybdenite could be rapidly and selectively leached in alkaline hypochlorite solutions. The feed material in this case was a copper rougher concentrate containing 12.4% Cu and 0.3% Mo. The use of externally generated hypochlorite was proposed. By consideration of power requirements, sodium chloride consumption and other factors for the production of sodium hypochlorite by this method; and also taking into account the potential reduction in molybdenum losses by the use of a process which eliminated cleaner flotation steps, Warren et al. suggested that dismissal of this type of hypochlorite leaching system on economic grounds was no longer valid; and that it would compare favourably with the electrooxidation techniques outlined above. A study was therefore carried out to determine optimum conditions for molybdenum extraction and these were found to exist at pH 9.0, with an excess of hypochlorite and a temperature range of 30 - 40°C. Selectivity of molyb denum over copper was obtained with a ratio of about 100:1, but the presence of even small amounts of copper in solution at pH 9.0 appeared to contradict solubility data. It was postulated that partially oxidized sulphur species could be complexing the copper and holding it in solution. Excessive hypochlorite consumption over that predicted by the stoichiometry of equation 3 was observed in certain cases, and this was attributed to the formation of surface oxides during the oven-drying of copper sul phide minerals, which subsequently act as catalysts for hypochlorite de composition. The study includes a proposed flowsheet for the overall pro cess . - 17 -1.4.2 The behaviour of copper in alkaline  carbonate solutions Addition of a number of anionic ligands to solutions containing 2+ dissolved copper in the form of the aquo-ion [Cu(H 0) ] leads to the 2 6 formation of complexes. Because the six water molecules are not equi distant from the central copper atom, successive displacment by other ligands can readily occur. Ammonia, NH^, is one such complexing agent in alkaline solution and the fact that cupric hydroxide is readily soluble in ammonia solutions has been used to advantage in a number of processes for copper extraction. Ammonium carbonates are used in this respect, but due to the strong complexing action of the ammonia itself, the fact that carbonate also acts as a complexing ligand for copper is generally 2-overlooked. The [CO^ ] anion forms a stronger complex than do any of the halide ions, for example, in alkaline solutions, but the latter are more often considered to stabilize cupric salts in solution. It was observed in the middle of the last century that the addition of certain copper salts to sodium carbonate solutions produced intensely blue solutions which subsequently precipitated light blue crystals of some basic copper carbonate salt. Pickering published the most detailed 28 early study on this subject and he confirmed that salts obtained by the 27 29 above methods reported by Deville and Reynolds were 'double carbonates' with the typical formula Na^Cu(CO^)^•3H20' Pickering noted that the liquid retained negligible amounts of copper after precipitation and assumed that as the liquid and solid were quite different colours, two different substances were involved. He also observed that the solution was much darker than a copper sulphate solution containing an equivalent - 18 -amount of copper. He postulated that it contained a 'cupri carbonate' species in which copper formed part of the anion. Cuprammmonium com pounds with copper displacing the hydrogen atoms of ammonium salts to give a deep purple-blue solution were cited as an analogous case. 29 30 Reynolds and Wood and Jones had previously shown in independent studies that electrolysis of these blue carbonate solutions liberated copper at both the anode and the cathode. The two most common copper carbonates are the minerals azurite, Cu3(OH)2(C03) , and malachite, Cu2C03(0H)2. Pickering reported isolating several other salts including 5Cu0,2C02; 5Cu0,3C02 and 8CuO,3C02,6H20' by mixing copper sulphate solutions with either sodium carbonate or sodium bi carbonate. He found all these basic carbonates to be insoluble in water and sodium carbonate solutions, but slightly soluble in carbonic acid and sodium bicarbonate. He also observed that introducing copper sul phate solutions into alkaline carbonates containing various proportions of NaHC03:Na2C03 produced a very blue solution with a copper content 2-dependent on the C03 :HC03 ratio (and thus of pH), but which for any mixture was directly proportional to the total carbonate content. It was observed that on allowing these solutions to stand for a couple of days precipitation of either malachite or a 'double salt,' Na2 Cu(C03)2 occurred, leaving an almost colourless solution containing very little copper. No clear explanation as to why decomposition should occur after a certain time was given, but Pickering does suggest that the degree of electropositive and electronegative copper present in a given case determined whether malachite or the double salt formed, and that in fact the double carbonate had two isomerides; one (the - 19 -cupri-carbonate) containing anionic copper and the other cationic copper. A proposed molecular structure for each substance is given. Other observations made by Pickering included: i) A slight increase in temperature produced a small increase in the amount of copper dissolved in a given amount of carbonate, but strong heating of the solution gave black colouration and deposition of a black precipitate, probably copper oxide. ii) Addition of excess sodium hydroxide to a carbonate containing copper solution removes carbon dioxide causing decomposition of the solution and formation of a blue basic carbonate salt. Dilution wtih excess water also decomposed the solution, with a black precipitate appearing in some cases. 31 Appleby and Lane also carried out a study of the double carbonates of sodium and potassium with various transition metals including copper. They prepared samples of the salt Na2Cu(C03)2,3H20 by allowing it to precipi tate from copper carbonate solutions: concentrated cupric acetate was added to a sodium carbonate/bicarbonate mixture (100 g Na2CO3:40 g NaHCO^) at 50°C and a clear deep blue solution free of any precipitate was immediately formed. On standing overnight, crystals of the double salt were obtained from this solution and it was observed that the mother liquor could be reused after separation of the crystals, in which case a slightly better yield was obtained. These authors also observed that both crystals and solutions de composed at higher temperatures giving cupric oxide and sodium carbonate, as a result of loosing water and carbon dioxide. The solution was decomposed by addition of excess water. Appleby and Lane concluded that - 20 -crystal deposition was. not due to the solution being super-saturated in the normal sense because, although precipitation occurred slowly and only after several hours of standing, addition of previously formed crystals did not speed up the nucleation procedure. It was also shown that crystallization time was unaffected by atmospheric carbon dioxide and was constant for solutions in either open or sealed flasks as well as those in a desiccator connected to a Kipp's apparatus for carbon dioxide. The following reasons were thus put forward to explain the precipitation behaviour: i) A very slow rate of crystallization is operative and no stable super-saturated solutions exist. ii) The crystals are precipitated via formation of an intermediate compound which undergoes slow transformation into the final produce. In agreement with the observations of Pickering this study showed that the equilibrium copper content in solution increased as a function of total carbonate concentration. It was also noted that clear solutions could only be obtained in the presence of both sodium bicarbonate and sodium carbonate. Omission of the bicarbonate led to almost instantaneous precipitation of a basic carbonate salt. The fact that more copper was dissolved by the carbonate/bicarbonate mixture than would be obtained by dissolving CuCO^ in water was attributed to the formation of the complex 2- 6-ion CuCCO^)^ , or even CuCCO^)^ in the case of a double salt. The possibility of copper being present as a colloidal substance as well as (or instead of) a complex ion, was also discussed. A number of other studies of these characteristically blue copper carbonate solutions and of the basic salts derived from them have been - 21 -32 made, including a polarographic study by Meites, an investigation into the effects of altering the Na^O^CuSO^ on subsequent precipitation by 33 34 Hsu, and a determination of the solubility of malachite by Scaife. In this last study Scaife observed that malachite solubility in sodium bicarbonate solutions was far greater than would be expected form calcu lations for the cupric ion. It was therefore suggested that ionic association or complex formation occurs between cupric and bicarbonate ions. An analogy was drawn between this case and that of cupric/ acetate ion complexes. The total concentration of copper in solution was suggested as being given by the expression [CuT] = [CuC03] + [CuHC03+] + [Cu2+] (6) which seemed to give good agreement with experimentally determined values for copper dissolution. Garrels and Christ,^ give an Eh-pH diagram for the system Cu-0 -S-CO -H 0 at 25°C which shows that at atmospheric pressure (P_„_ = Z Z Z C^2 -3.5 10 ) malachite has a large zone of stability in a pH range from 7.0 -13.0 and potentials higher than 0 - 0.2 V. An increase in the dissolved carbon dioxide content would extend this region of stability (Figure 2). The solubility constant, K , for the reaction s Cu2(OH)2C03 = 2Cu2+ + C032~ + 20H~ (7) -31.90 is given as K = 10 s 36 De Zoubov et al.. investigated the behaviour of copper in sodium bicarbonate solutions of varying strengths, and found that the corrosion of copper metal was significantly affected by the carbonate content. - 22 -Figure 2: Stability relations for copper compounds in the system Cu-H„0-0„-S-C0„. - 23 -Experiment showed that solutions containing less than 0.0032 M bicarbonate were mildly corrosive towards copper wire; those with 0.01 - 0.032 M were practically non-corrosive and a further increase of bicarbonate above 0.1 M produced strongly corrosive solutions. Electrochemical equil ibrium (Eh-pH) diagrams were drawn to explain these observations. It was found that conditions of. stability of the basic copper carbonate mala chite corresponded to the passivation effect observed in the region of 0.01 = 0.032 M bicarbonate, but that dissolution of malachite to give stable, complex cupri-carbonate ions could occur in solutions with a very low bicarbonate content, or in those containing more than 0.1 M. A series of four diagrams are given in which the total dissolved CO^ content, defined as [H2C03] + [HC03~] + [C032~] -3 increases from 10 M to 1 M in concentration (Figure 3). As this value of dissolved CO^ increases, the region of malachite stability grows to higher pH values causing cuprite and tenorite zones to correspondingly diminish. Azurite, 2CuC03Cu(OH)2, becomes stable in a pH range of ^3 - 7 at [C02] values greater than 10 1 M. The dissolved species which are stabilized by the presence of carbonate include the CuC03 (aq.) molecule, 2-the cupri dihydroxo carbonate ion, CuC03(OH)2 , and the cupri carbonate 2-ion, Cu(C03)2 . The CuC03 (aq.) species exists at the lower end of the pH scale, being stable from 6.5 - 10.0 at low carbonate concentrations, and being pushed downwards to 5.0 - 7.0 at higher levels of [C02]. The cupri-carbonate ion appears in a narrow band of pH (9.0 - 10.5) at -2 C02 levels in excess of 10 M and expands to a much wider band (7.0 -- 24 -Figure 3 (a): Potential - pH diagram . for the system Cu-CC^-H^O with 10_3M [CO =]T. - 25 -Figure 3 (b): Potential - pH diagram for the system Cu-CCL-H-O with 10 ° M [C03=]T. 2 2 - 26 -> 11.5) at the 1 M CO^ level. These soluble carbonate species also become stable at increasingly higher potential values for a given pH as the dissolved CO^ content goes up. At E values above this, Cu^O^ (hydrated) is shown as the stable species. Diagrams showing theoretical circumstances of corrosion, passiva tion and immunity for copper as a function of pH and carbonate content, together with solubility diagrams for azurite, tenorite and malachite, are also given. 37 Schindler and co-workers also studied the solubilities of azurite and malachite, and with the help of thermodynamic data produced by Silman for the Cu-H20-CO^ system, stability diagrams showing the zones of predominance for the various components of this system are given. The occurrence of the uniquely stable carbonate complexes CuCO^Caq.) and 2-Cu(C03)2 is again recognized as causing dissolution of the basic copper carbonate minerals: %Cu (0H)2C0 + JsC02 = CuC03 (aq.) + IjH 0 (8) JjCu2(OH)2C03 + 2HC03~ = Cu(C03) 22" + %C02 + | ^0 (9) for which the following equilibria apply: 2+ 2-Cu + C03 c=^ CuC03 (aq.) log k - - 6.73 2+ 2- . 2-Cu + 2C03 Cu(C03)2 log k = 9.83 It was noted that in air-saturated solutions the majority of the copper - 27 -2+ 2-is present as Cu at pH values below 7.0, and as Cu(C03)2 in more alkaline solutions. 38 Stumm and Morgan give both an activity ratio diagram and a solubility diagram for the predominant solid phases and soluble species of the Cu-J^O-CO^ system. They have been constructed for a total -2 II carbonate content of 10 M for the relevant Cu equilibria, and show 2+ that the important soluble species with increasing pH are Cu , CuCO^ (aq.), 2- II Cu(C03)2 and hydroxo-copper anions (Figure 4). In a paper which considers the corrosion of metallic copper in 39 typical sea water solutions, Bianchi and Longhi give a stability diagram showing predominance areas of the species which can exist as cupric com plexes in solutions containing both chloride and carbonate ions. These - 0 - 2- 2- 4-include CuCl , CuCl2 , CuCl3 , CuCl4 , Cu(C03)2 , Cu(C03)3 and 3- 2-Cu(HC0 ) . It is again apparent that the Cu(C0o) ion is predominant 3 5  2 at pH 9.0, indicating that the carbonate complexes copper more strongly than chloride in alkaline solutions. The fact that copper-chloride complexes are weak in alkaline solutions, even in the presence of large quantities of Cl , is also 40 noted by Van Muylder et al. and it is stated that copper would be precipitated from such solutions as CuO (or Cu(0H)2). 1.4.3 Decomposition of sodium hypochlorite in alkaline solution Sodium hypochlorite, NaOCl, is a strong oxidant in alkaline solutions and has been shown to be an effective lixiviant for the extrac tion of molybdenum from molybdenite ores and concentrates, as outlined - 28 -Figure 4: Solubility of Cu (II) in the system Cu-H 0-CO •[CO-=]T = 10_2M. 2 2 2 a) activity ratio diagram. b) solubility diagram. - 29 -above. There are two potential problems associated with the use of this reagent from a commercial standpoint however: i) It is expensive to generate, ii) It is unstable and undergoes spontaneous decomposition under certain conditions of pH and ionic strength. This decomposition reaction can be catalyzed, especially by the salts of certain transition metals. Most early investigators concluded that the costs incurred in hypochlorite generation would be too great to give an economically feasi ble process for molybdenum extraction. In view of the current strong demand and high price of molybdenum, together with the capabilities of modern electrolytic generators for manufacturing hypochlorite, this type of hydrometallurgical process would now seem to compare favourably with existing technology. In a process using externally generated hypochlorite, however, any decomposition of the reagent during leaching is detrimental and would necessitate generation of more than the stoichiometric hypo chlorite required for the oxidation of molybdenite to molybdate. It is therefore important to appreciate the conditions under which decomposi tion takes place, and the extent to which it occurs. Several investigations have reported the results of studies carried out to determine the exact nature of hypochlorite decomposition under varying conditions of pH, temperature, ionic strength etc. It is generally agreed that the products of decomposition are oxygen, chloride and chlorate in varying proportions. The uncatalyzed reaction can be represented as follows: - 30 -HC10 + 20C1~ > cl03~ + 2C1~ + H+ (10) 2HC10 + OCl" >• C103~ + 2C1~ + . 2H+ (11) HCIO + OCl~ —-y 2C1~ + 02 + H2 (12) Equations 10 and 11 both represent decomposition to chlorate. The former is generally applicable in alkaline solutions while the latter 41 more often occurs in acidic media. D'ans and Fruend suggested that in very pure hypochlorite solutions chlorate can be the sole decomposition product, but this is more generally accompanied by simultaneous decomposi-42 tion to oxygen, as shown by Lister. It is this oxygen forming reaction which is usually thought of as being catalysed by certain transition metal hydroxides and oxides. While detailed consideration of the many mechanisms which have been put forward to explain this behaviour has hot been undertaken, a review of some of the relevant work is included as it introduces concepts which are applicable to the present study: 43 Bell measured the velocity of oxygen evolution from solutions of calcium hypochlorite, Ca(0Cl)2, in the presence of a number of salts in both acid and alkaline solution. This velocity increased rapidly in solutions containing Co(N03)2 and NiSO^ and was slightly increased by the presence of BaCl , CaSO , AgNO , HgCl , FeCl , K Cr O and FeSO . 4tC *i O -C- £ I 4 LiCl, Na2C03, KCN, Na2S and KN02 were reported to retard oxygen evolution. 44 Hofman and Ritter in a study of the stability and redox poten tials of hypochlorites, concluded that decomposition was catalyzed to approximately the same degree by CoO, NiO and IrO, while the oxides of manganese, uranium, bismuth, palladium, tellurium and vanadium had no - 31 -effect. A detailed study of the rate of decomposition of sodium hypochlor ite solutions in the presence of cobalt and nickel peroxides, CoO^ and NiO^, and oxides of copper, iron and cobalt, was made by Chirnoaga in 45 1923. He reported that oxygen evolution during decomposition was less than first order and could be represented by the general equation = v1/n (13) dt 1 where c = concentration of hypochlorite at time t, and k^, n are constants. Chirnoaga postulated that OCl ions were adsorbed onto the surface of catalyst particles. Decomposition of these ions would then determine the reaction velocity. The adsorption step would be faster than the sub sequent decomposition, and thus the surface layer would be kept in ad sorption equilibrium with the bulk of the solution. Experiment showed that oxygen evolution increased with time and with concentration (and hence surface area) of the catalyst. The oxides under study were found to decrease in order of catalytic activity: Ni > Co > Cu > Fe, and equation 13 was found to be more applicable for the catalysts with higher activities. Chirnoaga noted that a mixed CoO/NiO catalyst was more active than either oxide alone. This increased catalytic activity of mixed oxides was also' reported 46 by Lewis. He suggested less active catalysts could have a promoting influence when combined with a more active one. Fe^O^ in the presence of CuO for example, gave a more effective catalytic action than that produced by either catalyst alone. Lewis suggested a more simple equation to describe the rate of - 32 -decomposition: dc dt k (14) representing a linear relationship between evolved oxygen and time. He found the rate to be constant for a wide range of initial hypochlorite concentrations, and that it was directly proportional to catalyst concen tration. The mechanism was considered to be one that involved the forma tion and subsequent decomposition of a hypochlorite-catalyst complex in a continuous cycle. Such a process would occur at a constant rate pro vided the active centres on the catalytic surface were entirely covered by reactant. Any loss of these due to coagulation, dehydration etc. would lead to decreased catalytic action. Lewis therefore suggested that the 'promoter' effect occurred as a result of the second substance preserving the active centres on the catalyst itself, thus producing a more efficient (and prolonged) reaction. The first reported study that hypochlorite could be catalytically decomposed to chlorate as well as to oxygen was given in a Russian paper 47 of Glikman and Dain in 1941. They showed that cobalt hydroxide accelerated the reaction but gave no details of the pH region in which this effect was observed. 48 Ayres and Booth showed that both oxygen and chlorate are formed as decomposition products in a"catalysed reaction. Oxidative decomposition was more predominant in less strongly alkaline solutions and overall decomposition reached a maximum around pH = 9.0. It was also observed (15) - 33 -that where the initial pH had a value below 10, decomposition was accom panied by acidification of the solutions, giving a final value of 3 - 4. Chlorine evolution was also observed in such cases. Equations represent ing this observation were cited as: 2HC10 v 02 + 2C1~ + 2H+ (16) 3HC10 y cl03~ + 2C1~ + 3H+ HC10 + Cl~ + H+ y CI + H20 (18) The catalyst used in this study was ixi'dium oxide and Ayres and Booth proposed an overall mechanism involving the formation of an intermediate active complex allowing the production of molecular oxygen which can then recombine with more hypochlorite to give chlorate: C10~ + c~ == x~ + Cl~ (19) x~ = c~ + 0 (20) where c = catalyst and x = active complex C10~ + O y C102~ (21) CIO" + O y C103 (22) Lister, on the other hand, concluded after carrying out an extensive study of both uncatalysed and catalysed reactions of sodium hypochlorite 42 49 decomposition, that chlorative decomposition was not catalysed. ' In the former study he confirmed the findings of Foerster et al.^° that the uncatalysed reaction produces mainly chlorate, is second order and occurs in a two stage process: - 34 -2NaOCl y NaC102 + NaCl (23) NaOCl + NaC102 >• NaC103 + NaCl (24) The first, chlorite forming step is the slower of the two. Lister found that a small part of the decomposition reaction produced oxygen: NaOCl • NaCl + ^02 (25) and that this was a first order reaction. He noted that it could not be concluded with certainty that absolutely no catalysts were present during his experiments. 49 In his study of the catalysed decomposition reaction Lister used manganese, iron, cobalt, nickel and copper oxides and determined that none of these increased the rate of formation of chlorate from hypochlorite, but that the oxygen reaction was catalysed to a different extent by each metal oxide. For cobalt and nickel this was found to be a zero order reaction and for copper it was almost first order. The previously made proposal that this heterogeneous catalysis involves oxi dation of the metal to a higher oxide with subsequent loss of oxygen, followed by reoxidation, was considered by Lister to be quite feasible. The fact that iron and manganese showed only very weak catalytic activity could then be explained by the stability of the higher oxidation states of these metals, in contrast to the cases of cobalt, nickel and copper all of which exist as 'higher oxides' in relatively unstable conditions. This mechanism would not result in increased chlorate production and this was also in agreement with experimental observations. However, more oxygen was evolved than could be accounted for by this type of oxidation-reduction - 35 -cycle so an alternative mechanism was suggested. This still involved the formation of higher valence oxides, but it was proposed that this oxide formed an adsorbed metal-hypochlorite complex on the catalytic surface, and subsequently decomposed: 2M0 + C10~ • M203 + Cl~ (26) M203 + C10~ —M203C10~ (ads.) (27) M203C10~ • 2M0 + Cl~ + 02 (28) Prokopchik^1 dismissed this mechanism of Lister as being experi mentally unfounded. He concluded from his own work that higher oxides were formed as intermediate compounds and that these were actually the catalysts for hypochlorite decomposition. It was found that both oxida tive and chlorative decomposition were catalysed, with the latter occurring as a result of partial hypochlorite oxidation to chlorate during the decomposition of the higher valence oxide. Thus in his study of both homogeneous and heterogeneous hypochlorite decomposition in the presence of copper, Prokopchik determined that an important role was played by copper"'""'"1 compounds. When in solution these existed as anionic cuprates, for example, NaCu(OH)4, while solid compounds such as Cu2C>3 produced heterogeneous catalysis. The tri-valent cuprate was found to be unstable in solution and its formation in the presence of hypochlorite was followed by rapid decom position to bi-valent cuprite: - - 2-4Cu(OH)4 + 40H > 4Cu(OH)4 + °2 + H2° ^29) - 36 -with a reaction rate given by -d (Cu(OH)4~) = k2 (Cu(OH)4~)•(OH_) (30) dt k2 (Cu(OH)4"), since hydroxyl ion concentration is always greater than that of cuprate anions. The decomposition of hypochlorite occurred simultaneously: 2Cu(OH)42~ + Clef + H20 > 2Cu(OH)4~ + Cl~ + OH~ (31) - - 2-2Cu(OH)4 + 20H > 2Cu(OH)4 + + JjO (32) C10~ v==a Cl~ + J5O ( 33) The rate of reaction was found to be first order for hypochlorite concen trations up to 0.1 M, above which it deviated towards zero order. In the case of heterogeneous decomposition it was observed that the action of sodium and calcium hypochlorites on blue cupric hydroxide caused rapid transformation to a brownish-black compound accompanied by catalysed decomposition of the hypochlorite solution. The nature of the catalyst was considered, and Prokopchik concluded that while a copper peroxide compound could be produced by this method, the observed decompo sition data suggested formation of cuprates at high pH values, and the tri-valent hydroxide at lower levels of alkalinity, as being more likely. Formation of a red compound was noted at the end of some experiments, and this was taken to be a deactivated cuprate which no longer had catalytic properties. - 37 -A certain amount of confusion appears to exist concerning the compound responsible for hypochlorite decomposition at pH levels below 11.0. Prokopchik states firstly, as noted above, that cuprates cannot exist at low alkalinity and that oxidation of bi-valent to tri-valent hydroxide thus occurs. He later states, however, that it is impossible to obtain a tri-valent hydroxide, and that in fact no oxidation of bi valent copper will occur below pH 11.5, so that dehydrated cupric hydrox ide is responsible for catalytic action at lower pHs. Since it is emphasized that no hypochlorite decomposition can occur in the absence of copper1"'""'", these statements appear to be self-contradictory. In the proposed mechanism for heterogeneous decomposition, however, both the hydroxide and cuprate are considered: i) at lower alkalinities: Cu(0H)„ + OH -> Cu(OH) + e (34) CIO + 2Cu(0H)2 + H20 •>• CI + 2Cu(0H) (35) 4Cu(0H) (36) (37) ii) higher alkalinity (pH > 12.0): +Cu(0H) 4 + e (38) *C1 + 2Cu(0H) (39) (40) (41) - 38 -34-41 -Consideration of the redox potential for reactions ^.confirmed: that oxidation to the tri-valent state by hypochlorite solutions was thermodynamically possible. A further ambiguity exists regarding the chlorate forming reaction: the above equations suggest this occurs only at higher pH values, and Prokopchik in fact states that chlorative decomposition becomes more pro nounced as alkalinity increases, and may even predominate. Considerable discussion previously however had led to the conclusion that in strong alkaline solutions hypochlorite decomposes to oxygen, while at pH 11.0 and below copper hydroxide significantly accelerated the decomposition to chlorate. At pH 9.0 the mole ratio of ClO^O,, attained a value of 3. Graphical representation based on experimental data is given to illustrate this point. By consideration of the catalytic action of cobalt, iron, and nickel under similar conditions, Prokopchik concluded that catalytic activity decreased in the order Co > Ni > Cu > Fe. The generalized mechanism for any metal Me was given as: CIO 4 Me (OH) -> CI + Me (OH) . (42) n n+1 Me (OH) . y Me (OH) + ho^ (43) n+1 n 2 Me (OH) . + CIO »• C10o + Me (OH) (44) n+1 3 n - 39 -1.4.4 Copper compounds Copper exists in the tri-valent state as a result of 4s and 3d electrons being removed from the copper atom to give a Is2, 2s2 2p6, 3s2 3p^ 3d^ configuration. This is iso-electronic with the nickel1"1" structure, and tri-valent copper compounds are generally diamagnetic as a result of all electrons being paired. Although this is a relatively rare oxidation state for copper, its existence has been acknowledged since the middle of the last century, and several workers have studied the. formation and stability of copper1"'"''" compounds: Crum, in 1845, obtained what he described as a cuprate by the action of an acidic solution of "bleaching powder" on copper nitrate. This occurred as a red powder and was assigned the formula H^Cu^O^ corres-55 ponding to the hydrated form of the sesquioxide, C^O^. . Moser used hydrogen peroxide to oxidize copper hydroxide and obtained a compound approaching the formula CuC^-f^O. Tests with chlorine 56 and bromine produced no oxidation. 57 Scagliarni and Torelli produced an amaranth red compound by the action of potassium persulphate, K S 0 , on copper hydroxide solutions. Z Z 8 The compound evolved oxygen when treated with dilute H^SO^, decolourised KMnO^, and did not contain a peroxide group. It was thus taken to be tri-valent copper oxide, Cu^O^. The earliest report of oxidation of copper1"'" by hypochlorite solu tions is one by Muller and Spitzer60 in which it was shown that tri-valent copper could be produced in strongly basic solutions (>3N) of copper hydroxide by oxidation with hypochlorite. Hypobromite, chlorine and bromine produced similar effects. The colour of the resultant solution - 40 -apparently varied from red and violet to brown and black, dependent on the degree of alkalinity. The same authors carried out electrochemical tests to produce tri-valent copper oxide, and suggested its formation to be caused by i) oxidation of metal ions contained within the copper anode, or ii) oxidation of metal ions contained in the electrolyte solution, ^with subsequent oxide precipitation. 59 Aldridge and Appleby in 1922 did a number of experiments to investigate the peroxidic compounds of copper. Hydrogen peroxide was added to copper containing carbonate solutions and a yellow-brown com pound was rapidly precipitated. Analysis showed this precipitate to contain more oxygen than would correspond to C^O^, but insufficient to give CuO^. It was assumed, however, that a peroxide had been produced, but that its separation in a pure state was not possible. Lepore reported obtaining the compound Cu^0„ (or Cu'Cu.O.) by treat-fa 4 5 4 ing Fehling's solution with hydrogen peroxide.60 He also claimed that the addition of copper nitrate solutions to barium hydroxide gave the compound Cu.O -3H O or Cu (OH) . 2 3 2 2 6 More recent studies regarding the existence of alkaline and alkaline earth metal cuprates have been made by Scholder and Voelskow,61 Prokopchik,62 and Magee and Wood.63 Crystalline barium cuprate, Ba(CuO^)2*H2° was obtained by the addition of potassium hypobromite solutions to cupric hydroxide in the presence of barium salts such as BaCl2 or BaCO^ in alkaline solution, in experiments carried out by Scholder in 1951. The sodium salt, NaCu02, was produced in a similar manner by the action of hypobromite, NaOBr, and sodium bromate, NaBrO^. - 41 -Calcium and strontium cuprates could not be obtained by this method. Calcium cuprate was produced by Prokopchik who used calcium hypo chlorite to oxidize cupric chloride, CuCl^, at a pH value of 12.2. A red-purple precipitate was observed, having a Cu:0 ratio of 1:0.5. Similar oxidation was effected by the use of an equivalent amount of sodium hypochlorite (0.2 M). Magee and Wood carried out an investigation of sodium cuprate1"1""1" stability in alkaline solutions of varying strength. They concluded that the salt was very unstable in all cases and estimated that the CuO^ ion had a half life of only 25 seconds. Copper111 was thus assumed to exist in solution only when complexed. The fact that Scholder and Voelskow found tri-valent copper solutions to be stable in the presence of excess base and hypobromite was suggested to be caused by a hypo-bromite-copper"'""'""'" complex. A number of workers have, in fact, shown that copper1"1"1 can be complexed and hence stabilized either in solution or as a solid compound, by the presence of certain complexing ions. Perio-dates and tellurates are the commonest such groups: 64 III Malatesta, in 1941 obtained a stable periodate copper complex with both chemical oxidation of cupric hydroxide using persulphate, and electrolytic oxidation of copper, followed by addition of periodite. The resultant compound was diamagnetic and was analysed chemically to be tri-65 valent. Malaprade also stabilized the tri-valent state by the presence of potassium periodate, KIO^. He produced firstly a cuprate, KCuCOH)^ by oxidizing copper hydroxide with a KOH/K S 0 mixture, and then added 2 2 o periodate to give a stable complex compound. 66 Lister made essentially the same compound, Na_Cu (10,.) . • 16H„0, as 7 6 2 2 - 42 -well as the corresponding tellurate, Na^CuCTeOg), using sodium hypo chlorite as oxidant. A cupric chloride/sodium hydroxide mixture was added to a 1.5 M NaOCl solution and the blue precipitate which formed initially, turned rapidly black, giving what Lister presumed to be either a tri-valent oxide, Cu^O^, or hydroxide, CuCOH)^. Addition of an acidified sodium periodate solution, NaH 10^, gave a brown compound 3 6 after standing for a short time. Both this precipitate and the complex tellurate were found to contain tri-valent copper by iodide titration and by use of their oxidation equivalent weights. It was observed that oxygen was evolved on acidification of each precipitate and in some cases a scarlet colour was briefly visible. This was therefore suggested to be the colour o'f the Cu3+ ion. To determine the amount of copper present in these compounds in an uncomplexed state, Lister measured oxygen evolution as a result of NaOCl decomposition both before and after addition of the complexing agent. In the former case he found the reaction to be catalyzed by copper, presumably as the Cu(0H)4 ion. The rate of reaction was first order with respect to copper. After addition of either periodate or tellurate anions the rate of oxygen evolution was found to drop consider ably, and to correspond to a rate representing uncatalyzed hypochlorite decomposition. It was therefore assumed that when virtually all the copper existed in a complex state it was no longer an active catalyst for the reaction. Jensovsky prepared complex cuprates by anodic oxidation of copper in alkaline solutions containing periodate and tellurate ions. The maximum yield was obtained at pH 10.2, and increased with temperature. - 43 -The compounds were found to be diamagnetic, and thermal decomposition curves showed loss of water caused decomposition of the copper1"1""1" complex. On account of their relative stability and strong oxidative powers, complexes of tri-valent copper with periodates and tellurates can be used for the oxidimetric determination of a number of inorganic and organic compounds.6^ Titrations using standard solutions of copper1"1""1" are carried out in alkaline media and applications include the determina tion of antimony, calcium, thallium, cyanide and thiosulphate salts. Most other copper1"1"'"1" complexes which have been isolated are organic in nature. These include salts made from carborane ions, Cu(B^QH^QCH)^ , a potassium bis (birueto) cuprate1"1"'1' prepared by peroxodisulphate oxidation of [Cu(-NHC0N-)], and a number of peptide complexes.^2,69 Meyerstein^° carried out a study of the chemical properties of a tri-valent copper 1aquo1 complex by radiolytic means. He reported that Cu111 exists in neutral solution as a CuOH2+ (aq.) or Cu(OH)2+ (aq.) ion and decomposes by the mechanism 2+ 2+ 2CuOH v 2Cu + H202 (45) Such ions were made by the reaction Cu2+ + OH ——+ Cu111 (46) in radiolyzed solutions, followed by a tri-valent copper oxidation of water: Cu3+ + H20 >- CuOH2+ + H+ (47) It was stated that Cu111 will be formed by any oxidant which can lower - 44 -the redox potential of the couple Cu111 - Cu^+ in the presence of various ligands. The potential for the reaction 2+ 3+ Cu > Cu + e (48) 75 is quoted by Latimer to be 1.8 V in alkaline medium at 25°C. 71 Shams El Din and co-workers reported the formation of the sequioxide Cu^O^ by alternate anodic/cathodic polarization of a copper electrode in 0.1 N sodium hydroxide solutions, at oxygen evolution poten tials, i.e. at 0.76 V. The reaction involved was thought to be: 2Cu(OH)2 + 20H • Cu 0 + 3H20 + 2e (49) In an extensive study to investigate the formation of the sesquioxide by 72 various wet and dry chemical methods and electrochemically, Delhez calculated the redox potential for Cu3+ formation as 2.3 V. This was derived from the equation: Cu(0H)3 + e~ + H+ Cu(OH)2 + H20 (50) for which E = 1.57 - 0.054 pH (pH < 13.9). 2+ 3+ Approximate Eh-pH diagrams for the Cu /Cu oxidation-reduction system are given, and this study was later used as a basis for the extension of 73 existing diagrams of the Cu~H20 system given by Pourbaix and co-workers 74 to take account of tri-valent copper and its derivatives. Delhez concluded from his work that the only feasible method of preparing Cu was by oxidation of cupric hydroxide in alkaline solution using either hypochlorite or peroxodisulphate ions as oxidant. While - 45 -electrochemical methods may produce Cu^O^ in some cases, the yield is usually insufficient to carry out meaningful analysis. He thus suggested that all previous workers who had used oxidants including hydrogen perox ide, peroxomonosulphuric acid, chlorine or bromine, had actually produced either copper peroxide, CuC>2 or cuprous oxide, Cu-jO, which in many cases had been confused with the sesquioxide. Using a method of potassium persulphate oxidation of cupric sul phate in alkaline solutions (pH 12.0 - 13.0) similar to that outlined by 57 Scagliarni and Torelli Delhez produced a red amaranth precipitate which contained, but was not exclusively, tri-vaient copper oxide. It was therefore presumed to be a mixture of Cu(OH)2 and O^O^/Cu(OH) , with a ratio of Cu*1:Cu111 dependent on the rates of the relevant reactions: 2Cu(OH)„ + 20H~ + S_0o2_ ^2Cu(OH)_ + 2S02" (51) Z Z o 6 4 4Cu(OH)3 >• 4Cu(OH)2 + 02 + H20 (52) The reaction forming copper111 was catalyzed by the presence of Ba2+ ions, in agreement with the previous findings of Scholder and Voelskow.^ It was also found that the compound existed in an unidentified state of hydration, fluctuating between Cu and Cu(OH)3, apparently at random. For the case of Cu11 oxidation with hypohalogen solutions, Delhez found the results of previous studies to be confusing. Several workers gave strong evidence to show Cu111 had been produced, but it was often difficult to ascertain whether it was Cu,^^, or a tri-valent cuprate. Scholder and Voelskow, for example, had found that oxidation of a bi valent cuprate with hypobromite produced a copper11"1" compound which they assumed to be a cuprate. Delhez concluded it was probably Cu,^^. In - 46 -less alkaline solutions, using cupric hydroxide as a starting material they obtained a different product, having a maximum Cu:available ratio of 1:0.3, but in his own work Delhez never produced more than 35% Cu111 by any method. Consideration of the potentials of 01 , OBr and 0C1 ions showed that only the latter has a sufficiently high value to give a good yield of Cu111. Barium ions also catalyzed this hypochlorite reaction and a mechanism similar to that given by Prokopchik (equations 2+ 34 - 37) was thought to be operative. The catalytic action of Ba ions was due to an increase in the redox potential of the oxidizing solution in which they were present. CHAPTER TWO Experimental 2.1 Scope of the Present Investigation A number of studies have been made to investigate the leaching characteristics of molybdenite in alkaline solutions of sodium hypo chlorite, as outlined above; but very little work has been done to deter mine the oxidation behaviour of copper sulphide minerals when exposed to similar solutions. As one of the principal objectives of finding a hydrometallurgical process suitable for treating molybdenite concentrates is to enhance by-product molybdenum recovery from porphyry ores, it would seem to be of prime importance to gain an understanding of any reactions occurring between copper and the lixiviant, and hence to de lineate any steps necessary to prevent copper dissolution. This would ensure optimum conditions for a selective molybdenum leach, and at the same time prevent copper losses, both of which are fundamental to an economically feasible process. 7 6 Previous work in this department was carried out by Ismay, to research the possibilities of extracting molybdenum from copper-molybdenum sulphide rougher concentrates, by means of a sodium hypochlor ite leach. A proposed process flowsheet was drawn up on the basis of successful molybdenite leaching; and although molybdenum recoveries were high, it was found that some copper was also extracted by the hypochlorite. The object of the present work at its inception was thus to deter mine the extent of the reaction between various copper sulphide minerals - 48 -and sodium hypochlorite under similar conditions to those found by Ismay to be optimum for molybdenum extraction, with a view either to preventing copper from entering solution, or to finding a suitable method for its removal: 1) Initial experiments revealed that the dissolution of copper in hypochlorite solutions at pH 9.0 only occurred in the presence of carbonate buffer reagents but that removal of the carbonate had unde sirable side effects. 2) This led to an investigation of the kinetics of copper catalysed hypochlorite decomposition reactions. The nature of decomposi tion products was also considered. 3) The mechanism of reaction between sodium hypochlorite and copper was studied on a general basis, with and without the presence of carbonate, to increase an understanding of the effects observed during leaching. 4) As the study progressed it became apparent that the presence of copper sulphide minerals adversely affected molybdenum extraction at pH 9.0. A number of experiments were thus carried out at pH levels of 5.5 - 6.0, which had previously been suggested to be optimum for molybdenum recovery in an investigation by the U.S. Bureau of Mines. 5) Further work in alkaline solutions was done using synthetic minerals and this highlighted the fact that certain elements normally present as impurities were detrimental to high levels of molybdenum extraction. A thermodynamic study of the stability of various molybdate compounds in aqueous solution was thus carried out and several Eh-pH diagrams were constructed. - 49 -6) It was then necessary to determine the leaching characteris tics of certain calcium minerals in hypochlorite and other chloride-containing solutions. 7) Consideration of all the above factors led finally to establishing the necessary conditions for maximum molybdenum extraction while minimizing hypochlorite decomposition and copper dissolution. 2.2 Materials 2.2.1 Natural minerals Chalcocite, Cu^S, and covellite, CuS, were obtained as massive samples from Butte, Montana. Chalcopyrite, CuFeS2, was purchased in massive form from Ward's Natural Science Establishemnt, and obtained as a ground concentrate from the Phoenix mine, British Columbia. Large pieces of ore were crushed in rod and ball mills and wet ground to pass a-200 mesh sieve in preparation for leaching as powdered samples. Small pieces of broken ore were polished on a 'Texamat' 5tf diamond polishing wheel for surface examination before and after leaching. In some cases samples were mounted in epoxy resin (Epon 828/DETA) prior to polishing. Molybdenite, MoS2, was obtained as a high grade concentrate from Alice Arm, British Columbia, ground to pass a -325 mesh sieve. Calcite, CaCO^ was obtained in powdered form from the Micro-ionized company (Texas) Ltd. Mineral Analysis Wet chemical analysis gave the following results: - 50 -i) Covellite: Element Cu Fe Ca Zn,Ni This indicates a CuS content of 88.3%. ii) Chalcocite: Element Cu Fe Ca Pb Zn This gives a Cu2S content of 81.0%. iii) Chalcopyrite: Element Cu Fe Ca Zn,Ni Weight % 58.80 3.00 0.16 Trace (<0.05%) Weight % 64.50 4.80 0.12 0.06 Trace Weight % 31.6 29.20 1.51 Trace This gives a CuFeS^ content of 93.3%. Qualitative X-ray analysis using the Scanning Electron Microscope showed that the covellite and chalcocite samples also contained significant amounts of silicon, and traces of sodium and potassium. X-ray diffraction studies gave good agreement with literature values. Silica, Si02 and - 51 -pyrite, FeS2 were the only other major peaks obtained (Tables 1-3). (All diffraction studies reported for copper k radiation) I/I0 - species identified 1 I/I0 - CuS o d A CuS FeS2 sio2 dA (Reported) 8.11 7. 8.18 7 4.26 36 3.33 100 3.28 9 3.29 14 3.21 27 3.22 28 3.04 64 3.05 . 67 2.80 83 2.81 100 2.72 100 2.72 56 2.70 84 2.31 10 2.32 10 2.04 13 2.04 7 1.90 28 1.90 25 1.89 68 1.89 75 • 1.73 33 1.74 34 1.63 100 1.63 3 1.55 31 1.55 37 1.46 6 1.46 6 1.35 8 1.35 7 1.27 6 1.28 1.09 9 10 Tabl=e 1: X-ray Diffraction Pattern for Covellite Sample - 52 -I/I0 - species identified I/I0 - Cu2S o dA Cu2S sio2 O dA (Reported) 4. 25 35 3.77 10 3.35 100 3.60 3.39 3.31 3.26 3.05 10 30 10 20 20 2.86 30 2.88 20 2.74 15 2.73 10 2.63 12 2.67 2.54 2.47 10 10 20 2.39 60 2.40 70 2.28 21 2.20 20 1.96 79 1.97 80 1.87 100 1.87 100 1.82 17 1.69 40 1.70 36 1.64 20 Table 2: X-ray Diffraction Pattern for Chalcocite:Sample - 53 -o o dA ::i/I0 - species identified dA I/I0 CuFeS2 CuFeS2 SiC) 2 (Reported) 3.35 100 3.03 100 3.03 100 1.90 50 1.87 40 1.87 78 1.86 80 1.59 49 1.59 60 1.57 19 1.57 20 1.32 10 1.23 26 1.20 30 1.10 43 1.08 60 Table 3: X-ray Diffraction Pattern for Chalcopyrite Sample iv)r Qualitative spectrographic analysis of the molybdenite showed it to have the following analysis: Element Weight % Element Weight % Al 0.10 Mg 0.01 Bi 0.05 Mn 0.001 Ca 0.50 Mo 58.52 Cr 0.01 S 39.02 Cu 0.05 Si 1.00 Fe 0.50 Sn 0.20 Pb 0.01 Ti 0.03 - 54 -Traces (<0.001%)of gold, silver and strontium were also detected. 2.2.2 Synthetic minerals i) Cupric sulphide analyzing 99.9% CuS was purchased from the Rocky Mountain Research Co., Denver, Colorado. ii) Cuprous sulphide, analyzing 99% + Cu2S was obtained from Matheson, Coleman & Bell, East Rutherford, N.J. Small amounts of these materials were crushed in a pestle and mortar, and screened to give 100% -200 mesh material. iii) Reagent grade molybdenum disulphide from BDH Chemicals, Poole England was used. This analyzed as: Element Weight % Mo 53.4 S 35.6 Si =10 iv) Molybdenum disulphide analyzing >98% MoS2 was purchased from the Venton Company, Danvers, Mass. 2.2.3 Sodium hypochlorite The source of hypochlorite used in all experiments was a household bleach sold under the trade name 'JAVEX,' and containing 50 - 60 g/1 in the as-purchased condition. Suitable dilutions with water were made to this stock solution to obtain the desired concentration for leaching experiments. - 55 -2.2.4 Chemical reagents All other chemicals used were reagent grade. 2.3 Apparatus The majority of leaching experiments were carried out in a 1 litre glass vessel fitted with 4 Teflon baffles at equi-distant points around its inner surface. Solution agitation was effected using a turbine type titanium stirrer coated with a thin layer of 'MICROSTOP' paint. Agitator and vessel dimensions together conformed to a 'standard tank' configura tion. The stirrer was mechanically powered with a Fisher-Dyna mix motor, and the agitation rate was measured for each experiment using a Tecklock chronometric hand tachometer. The leaching vessel was placed in an open water bath which contained a Micro-set thermo-regulator manufactured by the Precision Scientific Company. After setting a desired temperature, this unit was capable of maintaining it to ±0.1°C. The pH of leaching solutions was monitored during each run by plac ing electrodes inside the glass vessel. In initial experiments a Beckman Expandomatic pH meter was used; this was later replaced with a Chemtrix pH controller (Horizon Ecology Co., Chicago, 111.) In combina tion with an ASCO 2-way teflon coated solenoid valve, this controlled the pH automatically -to within ±0.2 pH units. The experimental apparatus is outlined schematically in Figure 5. - 56 -JO PH CONTROLLER A Water bath B Glass leaching vessel •F Thermometer G pH electrodes C Ti mechanical stirrer D Fisher Dyna Mix E Thermo-regulator Figure 5: Experimental leaching apparatus H 2 way solenoid valve controlling addition of buffering solution from I burette, to leaching vessel via J capillary tube - 57 -2.4 Experimental Procedure 2.4.1 Leaching experiments Experiments leaching copper minerals, and copper sulphides and molybdenite together, were carried out in the following way: A required amount of JAVEX reagent was added to 700-800 ml distilled water and the resultant pH (=12.5) was lowered to the required value by addition of 1 N HC1. In cases where carbonate buffering was employed, Na2C03 and NaHC03 were added in ratios calculated from pka values to buffer the solution at the necessary pH level. The solution was made up to 1 litre, transferred to the leaching vessel and placed in the water bath. Slow agitation was imparted until the solution had attained the temperature of the surrounding water, then 'blank' samples were taken for hypochlorite and metal analyses. Agitation was stopped while mineral samples were introduced: where copper sulphides and molybdenite were leached simultaneously they were added as two separate powders, and in some cases addition of one sample was not made until the reaction with the other had proceeded for several minutes. The agitation was then increased. Samples were taken at timed intervals by solution with drawal through a pipette, and filtered using either filter paper or Gooch filter crucibles. In experiments where no carbonate buffering was used the required pH was maintained by dropwise addition of 2 N NaOH from a burette, or, in later experiments by using the pH controller. - 58 -2.4.2 Copper preparation Samples of copper"'""'""'" were made by the action of sodium hypochlorite on cupric chloride using two methods: i) A sodium hypochlorite solution of the required concentration was made up as outlined above, the pH adjusted to 9.0, and heated to 35°C. 8 g of CuCl2-2H20 (equivalent to 3 g/1 Cu) were then added, and the pH maintained by NaOH addition. ii) 8 g of CuCl2-2H20 were dissolved in water and the pH raised to 9.0 with NaOH, to give a blue precipitate of cupric hydroxide. A measured volume of stock hypochlorite solution was then added. After an initial rapid rise in pH this dropped, and had to be maintained at a value of 9.0 as before. After about 5 minutes agitation the solution was completely black, and after a period of brief standing to allow some settling of the con tained precipitate this was filtered, collected and dried in a desiccator. Samples made in the presence of sodium carbonate required a longer period of agitation to give complete precipitation of the copper, but the method used was essentially the same (method (i) was found to be preferable). 2.5 Analysis 2.5.1 Chemical analysis i) Copper: The amount of copper contained in leaching solutions was determined using a Perkin Elmer 306 spectrophotometer with an air-acetylene flame and a wavelength of 324.7 ran, after suitable dilutions with - 59 -distilled water, and by comparison with a calibration curve from known standards. To determine any effect caused by the presence of large amounts of chloride in these solutions, test analyses were done using standard solutions swamped with excess NaCl, but no differences were detected. Copper in ore samples was determined by atomic absorption analysis after dissolution in a warm mixture of concentrated hydrochloric and nitric acids, and subsequent digestion with bromine, Br^. ii) Molybdenum: molybdenum in leaching solutions was deter mined by atomic absorption spectrophotometry with a nitrous oxide-acety lene flame at a wavelength of 313.3 nm. Dilutions prior to analysis were made with a solution containing 10% aluminum chloride and 5% ammonium chloride. 2+ iii) Calcium: [Ca ] in solution was similarly determined at a wavelength of 422.7 nm using an air-acetylene flame. iv) Sodium Hypochlorite: the hypochlorite concentration of each batch of 'JAVEX' used was accurately determined by potassium iodide-77 sodium thiosulphate titration as outlined in ASTM D2022-64. A small sample was diluted with water, then added to an acidified potassium iodide solution and titrated against standardized thiosulphate. Starch indicator solution was added and the disappearance of its characteristic blue colour marked the end point. Sodium hypochlorite concentrations of samples taken during leach ing runs were determined by the same method, usually undiluted. v) Sodium chlorate: the [ClO^ ] content of leaching solutions 77 was determined by a related method, also outlined in ASTM D2022-64. - 60 -The chlorate was reduced by addition of sodium bromide, NaBr, and concen trated hydrochloric acid. After dilution and addition of potassium iodide, the released iodine was titrated against standard thiosulphate as before. This gives the total concentration of (OC1 + ClO^) so the actual chlorate content was determined from the difference between this value and that for hypochlorite in the same sample. vi) Chloride: the total chloride content of hypochlorite solutions both before and after leaching was found: a) by silver nitrate/potassium thiocyanate titration, as 77 outlined in ASTM D2022-64. All the hypochlorite and any chlorate present are reduced to chloride by sodium bisulphate in the presence of dilute HNO^. The total chloride content is then determined using a standard Volhard titration. b) potentiometrically: this method is essentially the same as a) above, but the endpoint of the silver nitrate titration was more accurately determined by immersion of a platinum electrode into the solution, and making a graphical plot of potential vs. [AgNO^]-2.5.2 Instrumental analysis i) Minerals. Qualitative analysis for powdered samples and massive mineral pieces was carried out using the X-ray energy analyzing facility of the scanning electron microscope. Quantitative X-ray analy sis for the ground concentrates was obtained with a Philips X-ray diffrac-tometer, and for surface analysis of massive samples the electron probe micro-analyzer was used. - 61 -ii) Copper a) Infra-red spectrophotometry: samples of the black compound were prepared in mull form with fluorolube oil, and run through a Perkin Elmer 621 grating infra-red spectrophotometer. b) Magnetic susceptibility: a small sample was weighed on a 'Gouy' magnetic balance, both in and out of a magnetic field. More detailed measurements of the gram magnetic susceptibility were also made. c) Ultra-violet visible spectrophotometry: leaching solu tions were examined in the UV/visible regions using a Perkin Elmer ultra violet spectrophotometer with a path length of 2 - 10 nm. d) X-ray diffraction patterns were obtained using the diffractometer. - 62 -CHAPTER THREE  Results and Observations 3.1 Sodium Hypochlorite Leaching of Copper Sulphide Minerals Experiments were carried out to determine the oxidation behaviour of certain copper sulphide minerals in solutions of alkaline sodium hypochlorite. The selected minerals were covellite, CuS, chalcocite, Cu^S, and chalcopyrite, CuFeS^. These are three of the most commonly occurring copper ores associated with molybdenum bearing porphyries. 3.1.1 Leaching of ground concentrates  in the presence of carbonate 10 g samples of each mineral, previously ground to pass a -200 mesh sieve were agitated in solutions containing 7-8 g/1 hypochlorite at a temperature of 35°C and a pH value of 9.0, buffered with a sodium 2-carbonate/bicarbonate mixture. The ratio of HCO^ :CO^ was 14:1 and 10.7 g/1 total carbonate were used initially. A small amount of copper was found to be rapidly dissolved in each case. Small variations were observed between the three minerals, but the copper content of solution averaged 100 - 130 ppm and represented 2 - 3% of the copper introduced into the system (Figure 6, Tables I - III). Copper concentration did not maintain a constant value with respect to time, and after attaining a maximum initial value it fluctuated in an apparently random manner. Leaching solution separated from the mineral Time (minutes) Figure 6: Oxidation.of copper sulphide minerals by NaOCl. - 64 -slurry in the early stages of agitation maintained an almost constant copper content (Figure 7, Tables IV, V). All solutions showed deep blue colouration: significantly more so than standard copper sulphate solutions containing an equivalent amount of copper. Solutions left to stand over night were observed to deposit a black precipitate which, after a further time period of 1 - 3 days, was accompanied by loss of the characteristic blue colouration of the solution. Subsequent analysis revealed these clear solutions to contain negligible amounts of copper. Heating to 50 - 60°C also produced a black precipitate and this was found to be copper oxide by analysis using the Debye Scherrer powder method for X-ray diffraction. Insufficient of the material put down from filtered leach ing samples could be collected to carry out meaningful analysis. Dilution of samples with water also produced a black precipitate in some cases. Sodium hypochlorite consumption was found to vary slightly for each of the three minerals: equal samples of CuFeS^, CuS and Cu^S con sumed 10%, 27% and 18% of the initial hypochlorite present respectively. One mole of copper thus consumed 0.28 M, 0.4 M and"0.22 M OC1 in the respective minerals (Figure 8, Tables I - III). 3.1.2 Effect of carbonate removal Experiments were carried out in which no carbonate/bicarbonate reagents were added to the leaching solution, to determine the effect, if any, on copper solubility. The pH was maintained instead with a sodium hydroxide buffer, added dropwise from a burette, to counteract the observed drop in pH. All other conditions were identical to those of Time (minutes) Figure 7: Effect of separating leaching solution from mineral slurry. - 66 -i--6.0J 4-OJ • chalcocite • covellite . A chalcopyrite (pH 9.0,35°C) 2.0J 1 30 r~ 90 150 Figure 8: Time (minutes) NaOCl Consumption-during copper sulphide leaching. - 67 -previous runs. The analytical results showed this carbonate removal from the system to have a very significant effect: i) Less than 1 ppm copper was dissolved from a 10 g sample of each mineral. ii) The solution pH was found to drop rapidly as agitation commenced, and this was accompanied by sodium hypochlorite decomposition. The actual rate of decomposition was a function of the particular copper mineral being leached: CuS produced 50% decomposition in 9 minutes CuFeS^ produced 50% decomposition in 12.5 minutes Cu^S produced 50% decomposition in 82 minutes. Total decomposition occurred in all three cases, but the rate of decomposi tion was not linear (Figure 9, Tables VI - VIII). 3.1.3 Leaching of synthetic copper sulphides Decomposition was presumed to be caused by catalytic action of copper at the mineral surface. To determine whether the observed differ ences in the rate of reaction were attributable to copper alone, samples of cuprous and cupric sulphide powder were leached under identical con ditions to those used for the natural minerals. In the presence of carbonate results were almost identical to those of covellite and chalco cite with respect to both copper dissolution and hypochlorite consumption. In the absence of carbonate no copper was detected in solution and the - 68 -Time (minutes) Figure 9: Effect of carbonate removal on NaOCl decomposition. - 69 -rate of hypochlorite decomposition was constant in each case (Tables XI, XII; Figure 10). This was linear and gave 50% of the initial 0C1 content remaining after 52 and 54 minutes for C^S and CuS respectively. 3.1.4 Determination of decomposition  products Leaching solutions were analysed for sodium chlorate content to determine whether this was a decomposition product, as given by the equation: 30C1~ * C103~ + 2C1~ (53) or whether oxygen was the only product: 20C1~ v 02 + 2C1~ (54) It was found that some chlorate was produced in all three cases, but that the results for the different minerals were very variable: covellite and chalcopyrite both showed a decrease in the percentage chlorate present in solution as hypochlorite decomposition progressed, being 13.7% and 36.15% respectively after 5 minutes, and decreasing to 9.57% after 45 minutes for covellite and 20.22% after 90 minutes for chalcopyrite (Figure 11, Tables VI - VIII). Chalcocite on the other hand, which produced a much slower rate of catalyzed decomposition, showed the reverse trend in that 8.7% of the decomposition product was chlorate after 5 minutes, and this increased to over 23% after 150 minutes leaching. These general patterns were observed in several different runs, although the absolute values varied slightly in each experiment. - 70 -8.0 150 . Time (minutes) Figure 10: NaOCl decomposition during leaching of synthetic copper Sulphides. - 71 -Figure 11: Time (minutes) NaC103 production during NaOCl decomposition. - 72 -3.1.5 Sodium hypochlorite leaching of massive samples of copper sulphide minerals Small pieces of chalcocite, covellite and chalcopyrite were polished on one face and immersed in solutions of sodium hypochlorite at pH 9.0 for varying amounts of time. Slow agitation was maintained in most cases and the experiments were carried out with and without the presence of carbonate. Mineral surfaces were examined before and after leaching using the scanning electron microscope, the electron microprobe and optical microscopy; and to this end specific areas were marked so that quantitative analysis could be carried out: i) Chalcocite. Immersion in carbonate solutions at pH 9.0 for periods of up to 3 hours produced little difference to the polished surface, although some tarnishing was observed. Approximately 0.003 g/1 copper were detected in solution (3 ppm) from a total surface area of 2 about 6 cm . Electron microprobe examination of marked areas revealed that in most cases both the copper and sulphur content at the surface had decreased, with an overall increase in the Cu:S ratio. In a few small areas an increased sulphur content was observed giving a decreased Cu:S ratio (Table IX). Samples immersed into carbonate containing hypochlorite solutions for longer periods, in excess of fifteen hours, began to form dark green areas on the surface. These areas grew in magnitude with time and after one week the mineral was entirely covered with a green deposit, which analysis showed to have a significantly increased copper content. The deposit was not water soluble, and resembled malachite, Cu (OH) CO .' - 72(b) -PLATE 1 1) Cu^S, unleached 2) after 3 hours in hypochlorite/carbonate solution 3) after 24 hours in hypochlorite/carbonate solution 4) after 3 weeks in hypochlorite/carbonate solution 5) after 12 hours in hypochlorite solution (no carbonate) 6) after 3 days in hypochlorite solution (no carbonate) 7) after 1 week in hypochlorite solution (no carbonate) Effect of NaOCl ± Na„ C0„/NaHC0_ on massive chalcocite samples at pH 9.0, 35°C. - 73 -(Plate I). After 3 weeks immersion in hypochlorite solution, small black patches were observed on the green coating covering the chalcocite sur face . Agitation of solutions containing 7-8 g/1 hypochlorite, but no carbonate, produced similar tarnishing of the chalcocite surface, and after about 12 hours a powdery green deposit formed in certain areas. Within 3 days the entire surface was green, and after 1 week a black deposit covered most of the sample. ii) Covellite. CuS was observed to have a more rapid reaction than Cu^S when exposed to both carbonate-containing solutions and hypo chlorite alone. In the former case the polished surface lost its shine within 10 minutes and white patchy areas were visible. After 4 hours agitation the surface was covered by a white deposit. Probe analysis showed areas with different leaching characteristics to be present: a) areas with an increased Cu:S ratio caused by a large decrease in sulphur content and a relatively small decrease in copper, b) areas with a significantly increased sulphur content and hence a decreased Cu:S ratio. Subsequent washing in carbon tetra chloride, CCl^, removed this 1 excess'sulphur and suggested it was an elemental sulphur deposit (Table X). After exposure to carbonate solutions for longer periods, the entire covellite surface was covered by a dark green deposit similar to that formed on the chalcocite. Gradual dissolution of copper occurred to give a deep blue solution containing up to 100 ppm Cu. After 2-3 weeks standing this solution lost its characteristic blue colour within a 24 hour period and the mineral surface was simultaneously coated with - 73(a) -- 73(b) -PLATE II CuS, unleached. after 10 minutes in hypochlorite/carbonate solution, after 15 hours in hypochlorite/carbonate solution, after 1 week in hypochlorite/carbonate solution, after 3 weeks in hypochlorite/carbonate solution, after 5 minutes in hypochlorite solution (no carbonate), after 15 hours in hypochlorite solution (no carbonate). after 1 week in hypochlorite solution (no carbonate). ect of NaOCl ± Na„C0_/NaHC0„ on massive covellite samples pH 9.0, 35°C. - 74 -a very black deposit (Plate II). This phenomenon was found to be quite reproducible, although the time to precipitation varied slightly for different samples. Covellite samples were rapidly tarnished by sodium hypochlorite solutions containing no carbonate, with the appearance of brown patches within 10 minutes. A powdery green deposit developed within 12 hours and covered the entire surface after about 15 hours,exposure. This substance was found to have a significantly increased copper count (Table X) and was similar in appearance to cupric hydroxide. After 1 week a dark brown/black deposit had developed on top of the green. It was 'flaky,' and easily removable, unlike the precipitate produced by carbonate solutions. The copper contents of both black deposits were similar to those of the green layers which they superceded. iii) Chalcopyrite. Immersion into 'unbuffered' hypochlorite solutions produced a tarnishing of the CuFeS,, surface, and several brown patches were observed to form. The entire sample was covered by a brown coating after 2 weeks standing. Samples exposed to carbonate solutions underwent little change: some tarnishing of the polished surface was observed and several dark brown patches appeared after long standing. SEM examination showed that both treatments produced a similar surface coating. Probe analysis indicated a slight increase in copper and a corresponding decrease in sulphur at the surface. No green material appeared on samples left in either solution for periods of up to 4 weeks. - 75 -3.1.6 Variation of total carbonate content  during the leaching of ground copper  sulphide minerals The surface studies reported above, together with the observed dependence of copper dissolution and hypochlorite decomposition on the presence of Na2C03/NaHC03 in the leaching system, indicated that these reagents played a significantly more important role than had been initially supposed when a carbonate mixture was added as a buffer component. Several experiments were thus carried out in which the total carbonate content of 2- 2-the system was varied from 0.55 g/1 [CO^ ] to 10.0 g/1 [CO^ ] during the leaching of covellite. This was found to have a rather remarkable effect on the behaviour 2-of both copper and hypochlorite. As [CO^ ] increased, the amount of copper initially taken into solution increased in an almost linear manner 2-and varied from 6 ppm with 0.55 g/1 [CO^ ] to 125 ppm with 10 g/1 2-[CO^ ] as shown in Figure 12 a,b and Table XIII. This value of dissolved copper represented a peak concentration which occurred within the first five minutes of leaching and then dropped to a slightly lower level. The second value was maintained for a finite period of time and was apparently a function of the amount of added carbonate. It was followed by a further rapid drop in the copper concentration, accompanied by equally rapid decomposition of the hypochlorite, which had previously maintained a constant value (Figure 13, Table XIII). In experiments 2-using 5 g/1 [CO^ ]. and above, the copper content remained at a second 'plateau' level after hypochlorite decomposition, but those cases con-2-taining 0.5, 1.0 and 2.0 g/1 [CO^ ] showed zero copper in solution after this point. 2.0 120J 80J E cl ol a a o A 10.7 g/1 - 7.5 g/1 [co3 ]T L1.5 • o (pH 5.0 g/1 2.0 g/1 1.0 g/1 0.55 g/1 9.0, 35°C) 404 .1.0 _ o E CO O CD a a o U0.5 Time (minutes) Figure 12(a): Effect of varied carbonate content on Cu dissolution from Covellite. o 1 ! 1 1 I 1 • 1 I 40 120 200 280 360 TIME minutes Figure 12(b): Cu dissolution and NaOCl decomposition for Covellite leaching with 10 g/1 [00^ ]T - QL -- 79 -It was observed that the rate of sodium hypochlorite decomposition was constant, and not therefore a function of dissolved copper or initial carbonate concentration. In all cases the rate was almost identical to that produced during the leaching of covellite in the absence of carbonate. To determine whether or not similar effects occurred with the other copper sulphide minerals, identical runs were carried out using chalcocite and synthetic cupric and cuprous sulphide powders. All three samples showed the dissolved copper content to pass through a rapidly attained maximum value and then to maintain a constant value for an amount of time dependent on the carbonate content of the system. This was again followed by rapid loss of copper and simultaneous hypochlorite decomposi tion (Figures 14 - 17, Tables XIV, XV). The necessary induction period to copper precipitation was almost identical to the analogous case for covellite being 10, 25 and 65 minutes for 1.0, 2.0 and 5.0 g/1 carbonate respectively. The amounts of copper dissolved from each sample were also very similar, but the observed rates of hypochlorite decomposition were not constant. They were however consistent with the rates obtained in each case during leaching without carbonate. 3.1.7 Effect of varying the hypochlorite  concentration To determine whether an increase of hypochlorite strength had any effect on copper dissolution, a covellite sample was leached in a solution containing 20 g/1 [OC1 ], rather than 7-8 g/1 concentration used previously, and in the presence of 5 g/1 carbonate. Results showed that the copper extracted had increased from 60 cL Q. <x> o. CL o O. 20-A-A -1— 40 80 • 5 g/1 [C03"]T A 2 g/1 O 1 g/i (pH 9.0, 35°C) —I 120 Q 160 Figure 14: TIME minutes Effect of varied carbonate content on Cu dissolution from Chalcocite. .0.75 -0.50 CO _ o E CO O X (-—1 CP Q. Q. O .O. -0.25 CO o - 83 -- 84 -0.045 g/1 to 0.1 g/1, and that this copper was held in solution for a longer period before both it and the hypochlorite concentration rapidly decreased. The rate of 0C1 decomposition was consistent with that ob tained in all previous runs using covellite, and the copper remaining in solution after decomposition was the same as in all other runs using 5 g/1 carbonate (Figure 18, Table XVI). 3.1.8 Effect of hypochlorite removal Samples of covellite were agitated in sodium carbonate solutions in the absence of hypochlorite to gain a better insight into the precise role played by the oxidant. Copper was dissolved to a level of 0.06 g/1 in the presence of 10 g/1 total carbonate. This gradually decreased to zero after 70 minutes agitation in cases where sodium chloride was also present, or after 30 minutes in 'carbonate only' solutions (Figure 19, Table XVII). Filtered samples taken for analysis were observed to precipitate a green powder after standing overnight. These solutions had lost their blue colouration and contained negligible amounts of copper. As noted earlier, both in the present study and in previous work carried out by Ismay filtered solutions from experiments leaching copper sulphide minerals in hypochlorite deposited a black precipitate after standing for a short time. Ismay postulated that complexing between copper and a partially oxidized sulphur species could occur, with slow transformation into sulphate in the presence of hypochlorite, causing precipitation of copper from solution as copper oxide, CuO. - 85 -TIME minutes Figure 18: Cu dissolution and NaOCl decomposition from Covellite leaching with 20 g/1 [NaOclJ and 5 g/1 [C0v=]T - 86 -1 I I I I 1 40 80 120 TIME minutes Figure 19: Agitation of Covellite in Na2C03/NaHC03 solution ± NaCl. - 87 -To test this hypothesis, copper sulphate was agitated in a solu tion containing 7 g/1 [0C1 ] and 10 g/1 Na CO /NaHCO at pH 9.0, in sufficient amounts to give 0.1 g/1 copper in solution. If less than this dissolved it could be taken as an indication that the carbonate could not hold this much copper in solution in the presence of fully 2-oxidized sulphur, i.e. as SO^ It was found however that all the copper was taken into solution, giving the same deep blue colour observed in leaching solutions. Further experiment showed a maximum of 0.127 g/1 Cu to be dissolved under these conditions. Removal of hypochlorite from this system enabled only 0 .065 g/1 of the copper to be dissolved giving a less intensely blue solution. Samples left to stand again produced a solid precipitate: black in the former case, and green in the latter where no hypochlorite was present. It was also observed that this precipitation effect could be "seeded," and addition of a small amount of this green precipitate to freshly made copper - carbonate solutions produced complete precipita tion of the dissolved copper, as more green powder, within one hour. ^ HI 3.2 Copper From the above observations it is evident that hypochlorite has a direct effect on the solubility of copper in sodium carbonate solutions, and on the nature and colour of substances precipitated from them. The resemblances between copper sulphide leaching and copper sulphate dissolu tion and the subsequent behaviour of solutions with and without hypo chlorite were found to be significant: - 88 -i) The copper dissolved by a given 0C1 /[CO ] solution is constant whether the starting material is a copper sulphide mineral, or copper sulphate reagent. ii) A lesser amount of copper is held in carbonate solutions containing no hypochlorite. This amount is identical to that found during copper sulphide leaching after hypochlorite decomposition and loss of some of the initially dissolved copper (Figures 12b, 19). iii) The copper carbonate complex decomposes after a certain time period. Copper is thus precipitated as a solid compound which is black in hypochlorite containing solutions and green otherwise. These factors led to the conclusion that oxidation of copper sul phides in the system under study occurs in a "two stage" process. After due consideration of the redox potentials involved, together with consul tation of relevant literature, particularly studies of Lister, Prokopchik and Delhez^'^2' ^2 it was postulated that the observed leaching behaviour involves oxidation of copper to a tri-valent state. Thus the initial, higher copper content of the leaching solution is caused by a copper***-carbonate complex which subsequently decomposes precipitating a black copper*"*""1" solid and causes simultaneous catalytic decomposition of the hypochlorite. The same cycle is then repeated for a copper**-carbonate complex. This also decomposes after a certain time, precipitating the remaining copper as a green powder which is presumably malachite, Cu2(OH)2C03 (Figure 20, Table XVII). - 90 -3.2.1 Analysis of copper A black compound similar in appearance to that precipitated from filtered leaching solutions was produced by the action of sodium hypochlor ite on copper hydroxide, as outlined in Section 2.4.2. The following tests were then carried out (initially on samples made in the absence of carbonate). Effect of acidification A small amount of the substance was added to a solution of 5 N hydrochloric acid. This produced rapid dissolution of the powder giving a pale blue solution, accompanied by vigorous gas evolution. Passage of this gas through a calcium hydroxide solution ("limewater test") showed no cloudiness whatsoever, indicating that no carbon dioxide was present. Evaluation of oxidation state  by idometric titration The amount of excess oxygen contained in the sample was determined from the formula 2Cu 5 l2 = 2Na2S203 => 1 mlN. Na„S 0_ 5 0.0635 g Cu 2 2 3 using the following method: 3+ 2+ Iodide will reduce both Cu and Cu to the +1 oxidation state - 91 -with formation of a Cul precipitate. Thus a difference in the quantity of thiosulphate titrant required in cases a) and b) below allows the 3+ amount of Cu to be calculated: a) A weighed sample of the compound was added to an acidified potassium iodide solution and the liberated iodine titrated against standardized sodium thiosulphate. b) An identical sample was added to dilute sulphuric acid, 2 g of crystalline KI were dissolved and the solution was then titrated against Na2S203 as before. Table 4: Determination of Copper111 by Idometric Titration Sample (g) Na S„0, (0.1 N) 223 (ml) Cu (g) Cu o, "o a) i) 0.5 47.5 0 .302 60.3 2) 0.5 47.5 0 .303 60.6 3) 0.5 47.5 0 .302 60.3 Excess S2°32" Cu (g) °2 (g) °2 Moles b) 1) 0.5 90.47 42.97 0.302 0 .034 0.0021 2) 0.5 91.53 44.03 0.302 0 .035 0.0022 3) 0.5 89.50 42.00 0.302 0 .034 0.0021 Av.=0.00214 M - 92 -Therefore, 0.0021 M of oxygen are present as Cu , presumably CU2°3" Cu associated with this is 0.002 x 2/3 M = 0.00143 M or 0.091 g. Therefore Cu^ = 0.034 g 0 + 0.091 g Cu = 0.125 g total, equivalent to 25% of the sample. For comparison, the same test was done using a standard CuO sample and no difference was observed between cases a) and b). Measurement of oxygen evolution 0 — 1 ' 1 ' " " using a mercury column Figure 50: Apparatus for Measuring Gas Evolution - 93 -The gas evolved on acidification of a weighed sample of the black substance was measured using a mercury column connected to a mercury reservoir in a device resembling a Lunge apparatus outlined above (Figure 50). The mercury column was attached to a flask containing dilute nitric acid and the specimen sample. Evolution of gas on dissolu tion of each sample caused a depression of the mercury level. This was assumed to be oxygen and its volume determined by reduction to STP. Sample : 2 g Depression of. Hg column : 23.4 cm Barometric pressure : 755 mm Hg Temperature : 18°C Reducing to STP: 755 273 VG = 234 x x —- = 21.56 cm Oxygen 1 mole gas at STP occupies 22.4 1 Therefore, 215.6 mm 215.6 _ g_QQgg M oxygen 22400 Copper associated with this as Cu^O: 2/3 x 0.0096 = 0.064 M Cu or 0.4064 g Therefore, percentage of sample which is Cu^O^ = 0.154 g 0 + 0.406 g Cu = 0.56 g or 28% of sample. This is in reasonable agreement with the result obtained titrametrically. Atomic absorption analysis showed the amount of copper contained in a solution obtained by acid dissolution of the black compound to vary between 60 - 64% copper. This is insufficient to account for 70 - 75% of the sample being present as pure cupric oxide, CuO, and it is likely - 94 -that some water is also present, either as adsorbed H^O or as a hydroxy group. Gas chromatography Further confirmation of the presence of available oxygen in the sample was obtained using the gas chromatograph. A sample of gas evolved by acid dissolution of the black compound was passed through the chroma tograph and peaks obtained had a N2:°2 rat^° °^ 71:46 or 1.54:1. This compares to a ratio of 78:40, or 1.95:1, for a standard air sample. The former case thus showed an oxygen enrichment of about 22%. Replications of the above analyses produced variations in the amount of excess oxygen present, but the value was always between 20 -30%. The variations appeared to be fairly random, but two distinct trends were observed: i) An increase in pH during the preparation of the sample gave slightly more copper11"1".. No change in the nature of the substance was observed however, and even at pH values of 12 - 14 no more than 30% available oxygen was obtained. ii) An increase in the time period between precipitation of the black solid to the end of filtration seemed to decrease the percentage yxeld of copper Instrumental analysis i) Magnetic susceptibility III . The above results indicate the presence of copper in the - 95 -substance produced by hypochlorite oxidation of cupric salts, but it was evident that the compound did not consist entirely of a tri-valent copper oxide. Tests to determine its magnetic susceptibility were 3 therefore carried out: as noted previously the d configuration of copper1"1""1" salts produces diamagnetic substances. Initial weighing of samples on a magnetic balance both in and out of a field however, indicated that the compound was paramagnetic. More detailed measurements of the gram magnetic susceptibility were then made, and a valuefor lO^Xg equal to 6.93±0.03 cgs units was obtained. Using suitable corrections for the diamagnetism of copper and oxygen and a value for total copper content of the sample from atomic absorption analysis gave: 6 3 •— 1 10 X Cu = 753 cm mol or magnetic moment, yeff' = 1.33 B.M. Considering that Cu11 compounds normally have 10^ x Cu values of 3 -1 over 1500 cm mol and a ueff. of =1.98 M, it can be assumed that this XXX 11 7 8 compound contains 30 - 40% copper and 60 - 70% copper ii) Infra-red spectra To try and determine the nature of the compound more precisely, and especially whether the copper was present as an oxide or a hydroxide, a sample prepared in mull form with Fluorolube oil was run in the infra red spectrophotometer. The results obtained were far from being consistent: in some cases a distinct absorbance peak was obtained in the 2700 - 3000 cm 1 wavenumber - 96 -band, indicating the presence of a hydroxy, OH group, while other samples gave no peaks at all. Drying the compound at 110°C overnight and re running through the I.R. spectrum caused the disappearance of a previous ly obtained hydroxide peak in about 80% of all samples. iii) X-ray diffraction Small samples were analyzed in the X-ray diffractometer, and the resultant diffraction pattern was compared with one obtained from a standard CuO sample: Cu < CuO Sample CuO Sample (Reported) 0 dA I/Ix dA I/I* o dA Vii • 2.77 9 2.75 12 2.75 12 2.54 100 2.52 100 2.52 100 2. 34 96 2.32 96 2.32 96 1.87 28 1.87 25 1.87 25 1.71 9 1.71 9 1.71 8 1.59 14 1.58 14 1.58 14 1.51 19 1.51 21 1.51 20 1.41 21 1.42 14 1.42 12 1.38 17 1.41 16 1.41 15 1.38 18 1.38 19 1.30 6 1.304 7 1.27 6 1.19 2 Table 5: X-ray Diffraction Patterns for Copper and Copper Oxides - 97 -These results are strongly indicative that the substance consists pri marily of CuO. 3.2.2 Precipitation of copper1"1""1" in the  presence of carbonate Samples made by the action of sodium hypochlorite on cupric hydroxide at pH 9.0 in solutions which also contained sodium carbonate and bicarbonate salts were very similar to those made in the absence of carbonate, but there was a 'nucleation period' before any precipitation occurred. This was a similar effect to that observed during copper sul phide leaching, before hypochlorite decomposition: [co32-]T Time to precipitation 0.5 g/1 30 seconds 2.0 g/1 5 minutes 5.0 g/1 165 minutes 10.0 g/1 12 hours Effect of seeding Addition of a small amount of the filtered precipitate to a car-bonate-hypochlorite solution containing dissolved copper caused instan taneous blackening of the blue solution with subsequent precipitation of more black material. This occurred in all cases, regardless of the initial carbonate concentration. - 98 - Effect of acidification Addition of dilute nitric acid to the sample produced a clear blue solution. Rapid gas evolution was observed and a limewater test indicated a certain amount of carbon dioxide, CO^, was present. Some excess oxygen was also detected by passage of a gas sample through the chromatograph and comparison of the observed retention times with those of a standard air sample. I.R. spectra Analysis of the infra-red spectra produced by samples of this black precipitate gave a consistent peak in the 950 - 1000 cm * wave-2-number band. This indicates the presence of an HCO^ or CO^ group. Peaks in the 2500 - 3000 cm 1 band, corresponding to a hydroxy group, were again found to be rather random in appearance. - 99 - X-ray analysis X-ray diffraction studies gave the following diffraction pattern: o dA I/Il - species Malachite identified CuO dA I/ID - malachite (Reported) 7.43 16 7.41 11 5.99 55 5.02 64 5.06 75 3.68 80 4.07 13 3.25 uniden- :ified 3.69 85 3.02 2.98 2.86 100 2.81 100 2.81 2.52 100 2.52 45 2.51 41 2.32 96 2.47 2.46 35 2.43 20 2.32 17 1.99 10 1.99 11 1.87 25 Table 6: X-ray Diffraction Pattern for Copper Carbonate - 100 -It is thus apparent that the substance made by the oxidation of cupric hydroxide in the presence of carbonate is a mixture of copper*"'""'" carbonate (essentially an oxidized form of malachite) and copper*1 oxide, and that it is this compound which is precipitated from leaching solutions on long standing. A similar black compound was produced by the action of undiluted 'JAVEX' (50 - 60 g/1 0C1 ) on malachite, but solutions containing only 7-8 g/1 gave no visible change to the green powder. 3.2.3 Hypochlorite decomposition studies 2 g samples of each precipitate, i.e. the oxide/hydroxide and carbonate, were agitated in hypochlorite solutions at pH 9.0 and 35°C, to determine the rate of hypochlorite decomposition. Graphical plots of the results suggested this rate to be very similar to that observed during covellite leaching, and considerably more rapid than that produced by the analogous case of cupric oxide agitation in hypochlorite (Figure 21, Table XIX). 3.2.4 Analysis of leaching solutions It was postulated that if copper*** exists in hypochlorite solu tions as a relatively stable carbonate complex, then the spectrum obtained from an ultra-violet-visible spectrophotometric scan should be different from that produced by a copper** carbonate complex, the colour of which is much less intensely blue. However, the only detectable difference between the two solutions - 101 -- 102 -was a strong hypochlorite peak in the UV region of the darker sample. No sensible peaks were obtained in the visible region whatsoever, despite the blue colouration of the samples. A cupric choride sample significantly more concentrated in copper than the carbonate solutions was then run as a standard, and a very broad peak appeared in the 440 - 650 ym region. It was thus assumed that the carbonate complexes were too dilute to give -3 any meaningful peaks ( [Cu] - 1.6 x 10 M). 3.3 Sodium Hypochlorite Oxidation  of Molybdenite 0.3 g samples of natural molybdenite, MoS2, were leached in solutions containing 7—8 g/1 0C1 ±: carbonate at pH 9.0 and 35°C. A value of 0.3 g/1 was chosen because it represents the typical concentra tion of molybdenum found in copper rougher concentrates after flotation. Extractions of the order of 93 - 94% Mo were obtained within 5 minutes of the commencement of leaching. Hypochlorite consumption amounted to about 12% of the initial value (Figure 22, Table XX), corres--2 ponding to the use of 1.69 x 10 moles 0C1 to oxidize 0.17 g/1 Mo (1.77 x 10 3 moles) from MoS2 > NaMo04. These results therefore confirmed the findings of Ismay, that hypochlorite is a rapid and efficient lixiviant for molybdenite oxidation; but the fact that greater than 94% molybdenum extraction was not obtained was not entirely satisfactory. The amount of molybdenite sample used was thus increased to produce sufficient residue after leaching to enable X-ray analysis to be carried out: Figure 22: NaOCl oxidation of molybdenite at pH 9.0. - 104 -i) X-ray fluorescence indicated the following elements to be present in decreasing order of magnitude: Fe > Pb •> Cu,Mo > Zn ii) X-ray diffraction studies identified two lead molybdate species, Pb_MoOr and Pb MoO. as being present. Both of these were 2. b 4 assumed to be insoluble, since re-leaching of about 0.2 g of residue produced no detectable molybdenum in solution. Spectroscopic analysis confirmed that the concentrate did not contain 100% MoS^ (Section (iv)), and that at least 3% of impuri ties were present. The level of extraction obtained therefore represented more than the apparent 94%. Samples of reagent grade molybdenum disulphide were subsequently used, but gave no better results (Figure 23). This material was shown by X-ray analysis (SEM) to contain significant amounts of silicon. Leaching tests were finally carried out with material guaranteed to con tain over 98% MoS^. Extractions of nearly 98% were then obtained (Figure 24). Hypochlorite consumption was slightly higher than for the oxidation of an equivalent amount of the natural mineral (Table XXI). The presence or otherwise of carbonate buffer reagents was found to have no effect on molybdenum extraction. 3.3.1 Sodium hypochlorite leaching of molybdenite and copper sulphide  minerals at pH 9.0 In order to determine the selectivity of molybdenum extraction from molybdenite-copper sulphide mixtures, samples of both minerals were Figure 23: NaOCl oxidation of reagent grade-molybdenum disulphide. - 106 -1 1 1 1 1 To 0.0 12.0 24.0 36.0 48.0 60.0 TIME WINS! Figure 24: NaOCl oxidation of 98% molybdenum disulphide. - 107 -leached simultaneously. In the presence of carbonate, 89% Mo extraction was obtained, accompanied by dissolution of 2 - 3% of the copper (-0.12 g/1). Hypochlorite consumption in the presence of all three copper minerals was virtually the same as for oxidation of molybdenite alone (Figures 25a, 26a, 27a, Tables XXII - XXIV). Removal of carbonate from the system resulted in negligible amounts of copper being dissolved, but also produced a significant decrease in the level of molybdenum extraction: 74% Mo dissolution was obtained in the presence of covellite, 77% in the presence of chalcocite and 79% with chalcopyrite (Figures 25b, 26b, 27b, Tables XXII - XXIV). To confirm that this was a real effect, and not caused by the small amount of M0S2 sample used giving unrepresentative results, 10 g of chalcopyrite were leached with 5 g 'reagent grade' M0S2 under other wise identical conditions. A 2% loss of molybdenum occurred in the presence of carbonate and a 4 - 5% loss in the absence of carbonate (Table XXV). Scheiner et al. stated that copper molybdate compounds were found to be detrimental to good molybdenum recovery in the U.S. Bureau of 20/22 Mines process. Experiments were therefore carried out to see if copper molybdate formation could be induced, although in the present study it appeared that the greater molybdenum loss was occurring from solutions containing no dissolved copper: i) Copper sulphate solution was added dropwise from a burette during molybdenite leaching. ii) Sodium molybdate solution was added to a covellite leach under similar conditions. - 108 -Figure 25 (a) : NaOCl oxidation of molybdenite and covellite in the presence of carbonate buffer. - 109 -Figure 25 (b): NaOCl oxidation of molybdenite and covellite in the absence of carbonate buffer. - 110 -a.Q.O 12.0 24.0 36.0 48.0 60.<& TIMECMINSJ Figure 26 (a): NaOCl oxidation of molybdenite and chalcocite in the presence of carbonate buffer. - Ill -o O.O-in pi-I CO in. r-in OJ in CO 12.0 24.0 36.0 48.0 + MO EXTRACTION x NAOCL CONCENTRATION 12.0 24.0 36.0 TIME(MINS) 48.0 60 in a en in 60.0 Figure 26 (b): NaOCl oxidation of molybdenite.and chalcocite in the absence.of carbonate buffer. - 112 -o a. 0.0 in £-1 in. in to x UJ *< 6D in in 00 C3 in in OJ" 12.0 _1 24.0 36.0 48.0 _l + M0 EXTRACTION x NflOCL CONCENTRATION 60 .na in .a a_ ID LU o -z. 63 O "1—I o CE z. in C3 0.0 Figure 27 (a) 12.0 24.0 36. TIME(MINS) 48.0 60.0 NaOCl oxidation of molybdenite and chalcopyrite in the presence of carbonate buffer. - 113 -Figure 27 (b): NaOCl oxidation of molybdenite and chalcopyrite in the absence of carbonate buffer. - 114 -In both runs copper and molybdenum contents were monitored with respect to time. In the former case no copper was detected in solution, but significant amounts of molybdenum were removed. The presence of copper was also detected on the molybdenite surface after leaching. In the second case with no carbonate present and hence no copper dissolu tion, nearly 20% of the molybdenum added to the system was lost from solution after 60 minutes agitation (Tables XXVI, XXVII). 3.3.2 Leaching of molybdenite and copper  sulphides at pH 5.5 A pH value of 9.0 had been maintained in all experiments reported to date, largely because this was previously found to be beneficial by Ismay for two reasons: i) pH 9.0 is the point at which a maximum rate of molybdenite oxidation occurs in sodium hypochlorite solutions. ii) It is coincidently the pH value at which chlorate production is minimized during hypochlorite decomposition. In view of the adverse effects apparently caused by introduction of copper minerals into the system however, the leaching of molybdenite on the acid side of neutrality was investigated. A pH value of 5.5 was chosen because this is below the region in which chlorate production is most rapid (6.0 - 7.0) . In most cases molybdenite was leached for 5-10 minutes before any copper samples were added to the system. Sodium bicarbonate was added one gram at a time, after copper sulphide addition, and this maintained the desired pH value initially. - 115 -After 10 g of reagent had been added, and in cases where no carbonate was used, sodium hydroxide was used as a buffer reagent, as before. The results clearly showed that an even greater loss of molybdenum occurred under these conditions. Leaching of molybdenite alone gave over 90% recovery, but the introduction of covellite, chalcocite or chalcopyrite reduced the level of molybdenum in solution to about 10% (Figures 28 -30a, Tables XVIII - XXX). The amount of copper dissolved was also much lower than that found leaching CuFeS^ or CuS alone at pH 5.5. Addition of sodium bicarbonate to the system increased this value slightly (Table XXXII), but did nothing to prevent the precipitation of molybdenum (Figures 28 - 30b). It was thus presumed that copper molyb date readily formed at this pH level. Furthermore this substance acted as a catalyst for hypochlorite decomposition as shown by the very rapid loss of OC1 accompanying molybdenum precipitation. The consumption of the lixiviant was higher in experiments leaching only molybdenite at -2 pH 5.5 than for the corresponding case at pH 9.0 (8.3 x 10 M 0C1 for -3 5.72 x 10 M Mo which is higher than the stoichiometric requirement of 9 moles OC1 per mole MoS2). A number of runs were carried out at slightly higher pH values: 6.0, 6.5 and 7.0, to determine the effect on molybdenum recovery. Con siderable losses of both molybdenum and [OC1 ] again occurred in the presence of copper minerals with no carbonate in the system, but recovery was significantly increased on addition of NaHCO^.' A chalcocite/ molybdenite leach at pH 6.5 with bicarbonate in the system gave 89% Mo extraction and the copper and hypochlorite concentrations in solution maintained a constant value for the duration of the experiment. In an -116-Figure 28: NaOCl oxidation of molybdenite at pH.5.5. - 117 -TIME minutes Figure 29: NaOCl leaching of Chalcopyrite/Mplybdenite at pH 5.5. - 118 -°0.0 12.0 24.0 35.0 48.0 60.& 0.0 12.0 24.0 36.0 48.0 60.0 TIME(MINS) Figure 30 (a): NaOCl oxidation of molybdenite and covellite at pH 5.5 in the absence of NaHC0_. - 119 --0.0 12.0 24.0 36.0 48.0 B0.& TIME TMINS1 Figure 30 (b): NaOCl oxidation of molybdenite'and covellite at pH 5.5 in the presence NaHCO-. - 120 -identical run carried out in the absence of bicarbonate, introduction of the chalcocite sample reduced the molybdenum in solution from 80% to 15% within 30 minutes (Figure 31a,b, Tables XXXIII, XXXIV). The copper dissolved initially also decreased with respect to time, and there was a rapid loss of hypochlorite. The fact that copper was entirely responsible for molybdenum precipitation was confirmed by a similar experiment using synthetic CuS at pH 6.0. Results showed the same molybdenum losses occurred without bicarbonate in the system, and that this was prevented by leaching with NaHC03 present (Figure 32a,b, Tables XXXV, XXXVI). Molybdenum depression lessened slightly as the pH was further increased: a CuFeS^/MoS^ leach at pH 7.0 with no carbonate gave 30% Mo extraction. Subsequent addition of sodium bicarbonate in a similar run increased this to over 95%. The hypochlorite decomposed more rapidly in the latter case (Figure 33a,b, Table XXXVII, XXXVIII). 3.3.3 Copper molybdate To determine more about the conditions of formation of this sub stance, especially with respect to pH, experiments were carried out in both acid and alkaline solutions in which equi-molecular solutions of sodium molybdate, Na^MoO^ and copper sulphate, CuSO^ were mixed. At pH 9.0 a fine blue-green precipitate was observed while at pH 5.3 a darker green gelatinous precipitate was visible. Analysis of each solution both before and after mixing revealed that: i) at pH 5.3 copper and molybdenum were lost from solution in a stoichiometric ratio of 1:1. - 121 -Figure 31 (a): NaOCl oxidation of molybdenite and chalcocite at pH 6.5 in the presence of NaHCO-.. - 122 -°-0.0 12.0 24.0 36.0 48.0 . 60.$ TIMF(MINS) Figure 31 (b): NaOCl oxidation of molybdenite-and chalcocite at pH 6.5 in the absence of NaHCO-. - 123 -°-0.0 12.0 24.0 36.0 48.0 60.& §J 1 1 1 1 UCN 0.0 12.0 24.0 36.0 48.0 60.0 TIME(MINS) Figure 32 (a): NaOCl oxidation of molybdenite and cupric sulphide at pH 6.0 in the presence of bicarbonate. - 124 Figure 32 (b): NaOCl oxidation of molybdenite and cupric sulphide at pH 6.0 in the absence of bicarbonate. -125 -Figure 33 (a): NaOCl oxidation of molybdenite and chalcopyrite at pH 7.0 in the presence of NaHC0„. - 126 -o 0.0 12.0 24.0 in in _ r-in ID 36.0 _l 48.0 MO EXTRRCTION v NAOCL CONCENTRATION * X r 24.0 36. TIME(MINS) o 48.0 60 .a in a a en z. O i—i h-a: •I— ID 21 LU CJ O O CE a in a 60,0 Figure 33 (b): NaOCl oxidation of molybdenite and chalcopyrite at pH 7.0 in the absence of NaHC03. - 127 -ii) at pH 9.0 all the copper precipitated and a small fraction of the molybdenum was also lost from solution. Each precipitate was left to stand overnight and then heated over a steam bath at 60°C for 1 hour. The solutions were then reanalyzed. i) At pH 5.3: On standing, slightly more copper and molybdenum had dropped out of solution, again in an approximate Cu:Mo ratio of 1:1. ii) pH 9.0: Conversely at the higher pH, aging had a negligible effect on the copper content of solution (which remained virtually zero) but the molybdenum content increased after standing overnight, and again on heating (Table 7). Solubility of copper molybdate Addition of NajyioO„ solution to CuSO. of an equal concentration 2 4 4 -2 (10 M) was made dropwise from a burette at pH 5.4. It was observed that about 0.8 g/1 molybdenum could be added before any precipitation occurred. Copper was also removed from solution at this point, after which the molybdenum content-, remained constant even with the addition of more molybdate solution: 840 ppm Mo = 0.00875 M 0.00875 M CuMo04 = 1.4 g/1 (Table XL, Figure 34) A further test of solubility was carried out by adding some of the filtered, green CuMoO^ precipitate to 100 ml water and leaving it to stand for several weeks, during which time the Cu and Mo contents of the solution were periodically monitored. Maximum values of 0.75 g/1 Mo and 0.49 g/1 Cu were detected. This corresponds to a copper molybdate 128 -Molybdenum added (moles x 102) Figure 34: Solubility of Copper Molybdate at pH 5.0, 35°C. - 129 -PH Condition [Mo] M [Cu] M Mo ppt M Cu ppt M 5.3 Before Mixing 0.00979 0.00913 — — After Mixing 0.00875 0.00813 0.00104 0.00100 24 Hours Stand 0.00792 0.00740 0.00188 0.00173 30 Min at 60°C 0.00396 0.00252 0.00583 0.00661 9.0 Before Mixing 0.00979 0.00913 — — After Mixing 0.007192 1.57xl0~5 0.00260 0.00911 24 Hours Stand 0.00792 5.12xl0~5 0.00187 0.00908 30 Min at 60°C 0.00849 7.08xl0~5 0.00130 0.00906 Table 7: Effect of Mixing and AgeingEqui-molecular Volumes of Na^MoO  and CuSO,, at pH 5.3 and pH 9.0 - 130 -solubility of 1.24 g/1, or 7.83 x 10 M which is in reasonable agreement with the above value. To confirm the effects observed during leaching, sodium bicarbonate was added to one or other solution prior to mixing. At pH 5.3 this had no effect on either copper or molybdenum precipitation. -3 At pH 9.0 1.57 x 10 moles Cu (100 ppm) were held in solution and only slight lowering of the initial Mo content was observed. Catalyzed decomposition of hypochlorite  in the presence of copper molybdate A 2 g sample of copper molybdate was agitated in a solution containing 7.0 g/1 hypochlorite at pH 5.0 and 35°C. Samples were taken at timed intervals and analysis showed that 50% of the initial hypo chlorite decomposed within 35 minutes and that the rate was almost linear. About 4% of this decomposition was to sodium chlorate, the rest presumably being to oxygen and chloride (Figure 35, Table IXL). No detailed investigation of the effects of varying surface area, pH, temperature etc. during copper molybdate catalysis was undertaken. - 131 -(pH 5.0, 35°C) I 10 ~l— 30 50 TIME minutes gure 35: Effect of Copper Molybdate on NaOCl decomposition at pH 5.0, 35°C. - 132 - X-ray analysis X-ray diffraction studies of a copper molybdate sample gave the following diffraction pattern: . This Study Reported 0 dA o dA I/I0 - 2CuMo04'Cu(OH)2 7.155 10 7.01 20 4.392 20 4.34 35 4.207 50 4.15 55 3.517 100 3.50 100 2.979 10 2.96 20 2.794 18 2.76 20 2.699 8 2.72 20 2.531 36 2 .67 40 2.416 22 2.50 25 2. 313 12 2.40 20 1.728 10 2.29 25 Table 8: x-ray Diffraction Pattern for Copper Molybdate - 133 -3.3.4 Solubility of calcium in  hypochlorite solutions The fact that copper and molybdenum apparently do not combine at pH 9.0 to form a stoichiometric CuMoO^ compound, together with the obser vations made during leaching that solutions containing copper as a soluble carbonate species were not those in which molybdenum extraction was ad versely affected, raised the question that other elements contained in the copper sulphide minerals as impurities could be precipitating insolu ble molybdate salts. An Eh-pH diagram for the Cu-^O-MoO^ was constructed (Figure 51) and this confirmed that copper molybdate exists as a stable compound at pH levels below about 9.2. Further evidence that copper was not primarily responsible for the molybdenum losses at pH 9.0 was then obtained by i) leaching reagent grade molybdenum disulphide and synthetic cuprous or cupric sulphides together and obtaining >95% molybdenum extraction whether or not carbonate (and hence copper) was present in solution (Figure 36); and ii) leaching covellite and molybdenite at pH 10.0, which should be well outside the zone of stability of CuMoO^. In the presence of carbonate -91% of the molybdenum was recovered, but this dropped to only 78% in a leach with no carbonate (Figure 37, Tables XLI, XLII). X-ray and chemical analysis of the ores indicated that iron, zinc, calcium and lead were the major metallic impurities. Leaching in hypo chlorite solutions produced no dissolution of lead or iron, and only negligible amounts of zinc. However, 0.067 g/1, 0.058 g/1 and 0.004'g/1 calcium were detected in solution during the leaching of 10 g samples of Figure 36: NaOCl oxidation of molybdenum disulphide and cuprous sulphide at pH 9.0 ± carbonate. Figure 37 (a): NaOCl oxidation of molybdenite and covellite at pH 10.0 in the presence of carbonate buffers. Figure 37 (b): NaOCl oxidation of molybdenite and covellite at pH 10.0 in the absence of carbonate buffers. - 137 -of chalcopyrite, covellite and chalcocite respectively (Figure 38, Table XLV). It is well known that calcium molybdate and calcium car bonate are both insoluble salts. The fact that calcium will remove molybdenum from hypochlorite solutions at pH 9.0, except in the presence of carbonate reagents, was illustrated by adding a calcium chloride solution to a molybdenite -4 leach. Loss of 1.25 x 10 moles of both calcium and molybdenum was observed when no carbonate was added, but virtually all the calcium was removed and 98% molybdenum maintained in solutions which also contained sodium carbonate/bicarbonate buffers (Tables XLVI, XLVII). Sodium hypochlorite leaching of  calcium sulphate/calcium  carbonate minerals The most commonly occurring calcium minerals are calcite, CaCO^, and gypsum, CaSo4-2H20. With the assumption that either or both of these form the principal source of calcium impurity in the copper sul phide minerals under study, samples of each mineral were leached under identical conditions to those used for copper/molybdenum: 2+ 10 g of CaCO^ containing 4 g calcium yielded =0.016 g Ca in solution, amounting to 0.4% extraction. 10 g of CaSO -IjH 0 ("Plaster of Paris") gave 2 g/1 dissolved calcium which is equivalent to 64.5% extraction (Figure 39, Tables IL, L). That the small amount of calcium dissolved by the action of a 7 g/1 solution of sodium hypochlorite on calcite was sufficient to ad versely affect molybdenum recovery was confirmed by a combined calcite/ molybdenite leach. This gave only 80% molybdenum extraction. The same - 138 -J • 1 r 1 r r 10 30 50 TIME minutes Figure 38: Calcium dissolution during NaOCl leaching of Copper Sulphide Minerals. - 140 -experiment in the presence of carbonate produced over 95% extraction (Figure 40, Table XLVIII). Effect of chloride concentration  on calcium solubility Calcium solubility increases in the presence of chloride solutions. A number of experiments were carried out to determine whether the solu bility observed during the leaching of calcium minerals in hypochlorite solutions prepared from a commercial bleach, was in line with that expected from sodium chloride solutions of equivalent strength. The total chloride content was varied from zero to 2 M and it was found: i) The amount of calcium dissolved from 10 g samples of both calcite and gypsum increased as a function of the total chloride content . (Figures 41, 43, Tables LI, LII). The absolute amounts were higher in the latter case. ii) Identical results were obtained whether the chloride was present as NaCl only, NaOCl made from 'JAVEX', or as a mixture of both. iii) The results were in good agreement with values reported in the literature for calcium solubility in chloride solutions (Figures 42, 44) . iv) The rate of extraction was fast in all cases and there was a slight increase as the chloride content was increased. Removal of calcium from solution It is evident that leaching molybdenite in the presence of soluble calcium salts is undesirable. It was therefore necessary to find a method - 141 -Figure 40: NaOCl oxidation of molybdenite and calcite at pH 9.0, in the absence of carbonate. - 142 -40 80 TIME minutes Figure 41: Effect of Chloride Concentration on Ca dissolution from CaS04- 1/2 H20. - 143 -moles Sodium Chloride 0.5 CD CO E CO E O CO .O. 1.64 0.84 1.5 _i • Experimental • Literature (pH 9.0, 35°C) 40 —i— 80 120 Figure 42: [Sodium Chloride] 9rams/litre Calcium Sulphate Solubility as a Function of Chloride Content. L5.0 CO jD O E •2.5 o x E O co .O. - 144 -TIME minutes Figure 43: Effect of Chloride Concentration on Ca dissolution from CaC0„. - 14 5 -Figure 44: Calcium Carbonate Solubility as a Function of Chloride Concentration. - 146 -2+ of removing Ca from the leaching solution, without detriment to the molybdenum. Despite the fact that calcite is slightly soluble in hypo chlorite solutions, experiment showed that the presence of excess carbonate in the system suppressed formation of CaMoO^. Further tests 2-indicated that a minimum of 3.19 M CO^ per mole Ca were required to ensure complete precipitation as CaCO^ (Figure 45, Table LIII). In practice however, it appeared that the quantity of carbonate required by the system may be dictated by factors other than for calcium suppression. It was noted in Section 3.1.6 that low carbonate concen trations produced a loss of hypochlorite more quickly than higher con centrations during the leaching of Cu^S, CuS and CuFeS2. Experiments in which molybdenite and copper sulphides were leached together showed 2-this nucleation time decreased further for a given [CO^ ]^ concentration. This was apparently due to the presence of molybdenite. CuS/MoS2 runs at the 5 g/1 and 10 g/1 carbonate level showed hypochlorite decomposition to occur after 45 and 110 minutes respectively, representing a decrease of 30 - 50% over the nucleation time for CuS alone. Copper dissolution was also slightly lower (Figure 46, Tables LIV, LV). A similar effect was observed with Cu2S/MoS2 and 10 g/1 carbonate (Figure 47, Table LVII). 2-At the 5 g/1 [CO^ ] level it was also found that a gradual loss of hypochlorite occurred continuously from the beginning of agitation and before the rapid loss caused by copper111 catalysis. Leaching CuS in the presence of sodium molybdate confirmed this trend and showed loss of both copper and molybdenum from solution in a 1:1 ratio suggesting that some CuMoO. forms even in the presence of carbonate. This then causes 4 catalyzed hypochlorite decomposition before Cu111 precipitation moles of Carbonate added Figure 45: Effect of Carbonate Addition on Calcium Precipitation from Solution. 120. E Q. cL CD a CL o .o. A - NaOCl --: • Cu. • 10 g/1 [co3=] A 5 g/1 [C03=] (pH .9.0, 35°C) A>.: N 80 TIME minutes 120 160 -6 8 CD CO 14 E ^ CO O) O O CO V2 Z Figure 46: Copper dissolution and NaOCl Decomposition during Covellite/Molybdenite leaching at pH 9.0, with varied [C03~JT> 4^ 03 160 8 T 1— 1 1 1 ' 1 I 40 80 120 160 TIME minutes Figure 47: Copper Dissolution and NaOCl Decomposition during Chalcocite/Molybdenite Leaching at pH 9.0 (10 g/1 [C03=] ). - 150 -(Figure 48, Table LVI). A similar run with double the amount of carbon ate produced no loss of copper or molybdenum and no hypochlorite decom position for over 4 hours (Figure 12b, Table XVIII). Leaching in the presence of larger amounts of carbonate would thus extend the period of maximum molybdenum extraction prior to copper111 precipitation and 0C1 decomposition, as well as preventing CuMoO^ and CaMoO^ formation. Silicates and phosphates also have insoluble calcium salts. Various beaker experiments were done to investigate the effectiveness of these reagents in precipitating calcium, as an alternative to carbonate, thereby preventing copper dissolution. Sodium meta silicate, Na^iO^-SH^O was found to be most efficient in this respect (Table 9). A molybdenite/chalcopyrite leach was subsequently carried out in 2-the presence of 1 g/1 SiO^ • Almost 99% Mo recovery was obtained (Figure 49, Table LVII). No copper was detected in solution and the hypochlorite maintained a constant value for the duration of leaching. - 151 -- 152 -Figure 49: NaOCl oxidation of molybdenite and chalcopyrite in the presence of Na„SiO„. - 153 -Additive [Ca] in soln. Additive [Ca] ppt (g/D (ppm) (Moles) (Moles) Ratio — 94.12 — 2.35xl0~2 [C032-]: 2- 2+ CO^ :Ca^ 0.25 90.31 4.166xl0~3 9.50xl0~4 4.39:1 0.50 53.17 8.33xl0~3 1.02xl0~3 8.17:1 1.00 18.8 1.66xl0~2 1.88xl0~3 8.83:1 [Si032-]: 0.25 20.03 3.28xl0~3 1.85xl0~3 1.77:1 1.00 3.45 1.31xl0~2 2.26xl0~3 5.04:1 [P043-]: 0.25 53.17 3.16xl0~3 1.02xl0~3 3.09:1 1.00 25.20 1.27xl0~2 1.70xl0~3 7.47:1 Table 9: Effect of SiO , CO and P04 (sodium salts) on Calcium  Precipitation at pH 9.0, in Hypochlorite Solutions. - 154 -CHAPTER 4  Discussion 4.1.1 Sodium Hypochlorite Leaching of  Copper Sulphide Minerals in the  Presence of Carbonate Initial tests to investigate the leaching behaviour of covellite, chalcocite and chalcopyrite in sodium hypochlorite solutions at pH 9.0 reaffirmed Ismay1s observation made during leaching experiments with copper rougher concentrates, that a certain amount of copper was rapidly extracted by the hypochlorite and held in solution. The thermodynamic solubility of copper hydroxide which is the 73 -9 stable specxes at pH 9.0, as given by Pourbaix is 8.0 x 10 which is several orders of magnitude less than the observed experimental value -3 of 10 M. It thus seemed likely that complexing of some kind was occurring. Copper, in common with other elements of the first transition series, is capable of forming aqueous complexes with a large number of anions, but these are more limited on the alkaline side of neutrality than in acid solutions. The most common at higher pH levels are ammonia and cyanide, but the only components present in the system under study were chloride (as CI and as 0C1 ) and carbonate, as well as sulphide con tained in the copper ores which would be expected to oxidize to sulphate, 2-SO^ , in hypochlorite solution. To determine if one or more of these species could complex copper under the prescribed conditions, relevant literature was reviewed and it was found: i) Certain cupric chloride complexes are known to exist, but - 155 -39,40 are found predominantly in acid solution. ii) Copper hypochlorite can exist in aqueous solution but the only known method of preparation is with acidic hypochlorite, and the 79 species is relatively unstable. Various references to Cu - 0C1 complexes were found but none contained definite proof that such a complex could be made. iii) Bivalent copper can combine with various intermediate sul phur oxidation species to form a number of salts, such as CuS„0 '4H„0, z 6 2 79 CuS^Og, some of which are soluble in aqueous solution. iv) Cupric carbonate complexes exist, and are relatively stable 27-40 in alkaline solution. Thus carbonate seemed to be the most likely complexing agent at a pH value of 9.0, and the fact that the leaching solutions showed such strong blue colouration for the amounts of copper they contained was 27 30 31 in line with observations made by Deville, Pickering and Appleby ' ' that the alkaline 'cupricarbonates' have a characteristic deep blue colour. Although slight variations in the amount of copper dissolved occurred between the three minerals, they were all of the same magnitude -3 (0.12 - 0.13 g/1 equivalent to 2.0 x 10 M copper). The fact that the same amount of sample was used in each case, implying that different amounts of copper were initially present, suggests that this value represented the maximum amount of copper which is soluble under the given conditions of pH, temperature and carbonate concentration. For a total carbonate content of 10 ^ M, De Zoubov et al. show the solubility of copper11 to be 4.0 x 10 ^ M at pH 9.0, where it exists as - 15'6 -the cupricarbonate ion CuCCO^)^ (Figure 51). This is slightly lower than the value obtained in this study and suggests that carbonate 2-may not be the only complexing agent for copper. Assuming that Cu(C03)2 forms the major complex species, however, and that all the sulphide associated with extracted copper is oxidized to sulphate, the reaction can be described by the equation: CuS + 4NaOCl + 2NaOH + 2HCC>3~ >-2-Cu(C03)2 + 4NaCl + Na2S04 + 2H20 (55) The reasons why such apparently random variations in the copper content occurred throughout a leaching run, or why a black precipitate was deposited from filtered solutions on standing overnight, were not immediately apparent. 4.1.2 Removal of Carbonate from the System Experiment confirmed that the carbonate buffer components were primarily responsible for holding copper in solution: negligible amounts were dissolved when the NaHC03/Na2CC>3 mixture was replaced by sodium hydroxide solution. This would seem to be an important factor when considering a leaching process which requires the selective extraction of molybdenum from porphyry copper ores. The solubility of copper in carbonate solu tions has been overlooked by a number of other workers including Shapiro 14 and Kulenkeva who claimed that in a carbonate system, any metal sul phides other than molybdenum would be precipitated as insoluble carbonates; - 15-7 -and Bhappu and Scheiner et al. ' ' both of whom suggested hypo chlorite leaching in the presence of carbonate would give a selective molybdenum leach. However, the potential advantage of removing the carbonate, and hence any soluble copper, from the system is offset by the observed effects on the hypochlorite. The small amounts of copper dissolved in the carbonate buffered case consumed very little hypochlorite and no further loss of strength was observed for up to four hours leaching. On the other hand, leaching with no carbonate caused rapid decomposition of the hypochlorite, accompanied by a drop in the pH of the solution, necessitating the addition of sodium hydroxide as an external buffer component. Sodium hypochlorite is a strong oxidant and it is likely that the surfaces of the copper minerals are rapidly oxidized on contact with it. At pH 9.0 and with no complexing agent present, the copper precipitates as a hydroxide on the mineral surface, thereby preventing any copper dissolution: CuS + 4NaOCl + 2Na0H >- Cu(OH)2 + NaS04 + 4NaCl (56) The fact that molybdenite and chalcopyrite have approximately 25 equal oxidation rates suggested to Stumpf and Berube that a selective molybdenum leach was not possible, but these authors seem to have over looked the relative solubilities of the two species in alkaline solution. The formation of a surface hydroxide is presumably the cause of the observed hypochlorite decomposition: oxides and hydroxides of various transition metals, including copper, are known to be active catalysts / - 158 -43-51 for the heterogeneous decomposition of hypochlorite solutions. Decomposition can occur either to oxygen and chloride or to chloride and chlorate. The former reaction is generally thought to be the more common in alkaline solutions. It was reported by Prokopchik, however,^ that some chlorate is produced at high pH values, and that in the presence of copper hydroxide catalysts this reaction becomes more predominant as the pH is lowered. Thus at pH 9.0 a ratio of C1C>3:02 of about 3:1 was obtained. Sodium chlorate is undesirable in a hypochlorite system utilizing complete salt recycle for regeneration of the lixiviant. The production of NaClO^ in large amounts would therefore necessitate solu tion treatment to remove it; for example, using S02 reduction, as 20 reported in the U.S. Bureau of Mines process. Chlorate has also been shown to interfere with the subsequent recovery of molybdenum from leach-23 24 ing solutions. ' 4.1.4 Decomposition Products During the Leaching of Copper Sulphide Minerals The production of NaClO^ as a result of NaOCl decomposition in the present study was found to vary from mineral to mineral. The fact that chlorative decomposition increased with respect to time during chalcocite leaching, but decreased in the cases of covellite and chalcopyrite may be associated with their relative catalytic rates, considering that chalco cite produced a much slower rate of decomposition than the other two. In no case was a C10:02 ratio of anything approaching the 3:1 reported by Prokopchik at pH 9.0 found to exist. The most obvious reason for the observed differences in catalytic - 159 -decomposition rate between chalcocite, chalcopyrite and covellite would be a difference in the active surface area, since the reaction is heterogeneous. Calculations for the overall surface area of -200 mesh 2 samples of covellite and chalcocite, give values of 33,300 cm and 2 27,000 cm respectively (Appendix B). The difference is therefore in sufficient to account for the 6:1 ratio observed between the rates of decomposition given by CuS and C^S. Furthermore, the use of synthetic copper sulphides indicated a constant rate of NaOCl decomposition for both cupric and cuprous salts. This could be the actual rate of cataly sis of copper hydroxide, with that observed during covellite and chalco pyrite leaching being caused by some sort of "promoted catalysis." Both these minerals were found to contain traces of nickel and cobalt as impurities. Nickel salts have been shown by several workers to be more active catalysts than the corresponding copper salts, and it is known that negligible amounts of both nickel and cobalt can cause sig-45 49 51 nificant hypochlorite decomposition. ' ' Chirnoaga also reported that the action of 'mixed' catalysts produced a more rapid reaction than either one alone. This was confirmed by Lewis who proposed a mechanism 45 46 for this type of reaction. ' It is therefore suggested that the enhanced rates of decomposition observed in the present study are caused by the combined action of copper and nickel hydroxide catalysts, giving a faster rate than would be produced by copper alone. The presence of carbonate in the system obviously deactivates the mineral surface: complexing a small amount of oxidized copper as 2-the cupri-carbonate ion, Cu(C02)3 , prevents formation of a hydroxide catalyst, and the soluble copper species does not in itself act as a - 160 -catalyst, so that no hypochlorite decomposition occurs. 4.1.5 Sodium Hypochlorite Leaching of Massive  Samples of Copper Sulphide Minerals Experiments using massive samples were carried out to obtain more information, both qualitatively and quantitatively, about the nature of surface reactions occurring during leaching. Covellite was found to have a more rapid reaction with the hypochlorite than chalcocite, whether or not carbonate was present in the system. Cu2S nas been shown to leach in a two stage process in acid solution, with the second 80,81 stage resembling the leaching of CuS or similar mineral. It is possible that even in the limited reaction which occurs on exposure to hypochlorite solutions, there is a difference between the oxidation behaviour of chalcocite and covellite. This would represent a more fundamental cause for the differences in the rates of NaOCl decomposition noted above. The patchy white areas present on the covellite surface after immersion into hypochlorite containing carbonate solutions for short periods probably indicate sulphate formation although why elemental sul phur should be present in certain areas as well is a little puzzling. 16 Choppm and Faulkenberry reported that elemental sulphur could be produced during the oxidation of aqueous sulphide solutions by hypochlor ite whenever the ratio of S:0C1 was less than 1:4. In all other cases 2-sulphate, SO^ , was the only end product. It is thus possible that hypochlorite depletion occurred in localized areas of the mineral surface, enabling small amounts of sulphur to be formed. - 161 -Green deposits produced on both covellite and chalcocite surfaces after longer periods of exposure to hypochlorite solutions are almost certainly malachite, (^^(OH^CO^. This substance forms the stable phase in the Cu-CO^-I^O system at pH 9.0 (Figure 2). Further oxidation of the mineral surface after the saturation point for the cupri-carbonate species in solution has been reached will result in precipitation of a basic copper carbonate. In the absence of carbonate, copper hydroxide forms on the mineral surface as a light green, powdery deposit. Micro-probe analysis confirmed that both deposits had a significantly increased Cu:S surface ratio, due to a large increase in the copper content at the surface. The appearance of a black deposit on top of both green substances coating the original covellite surface, accompanied by a loss of blue colouration from the carbonate containing solutions was taken to be analo gous to the deposition of a black precipitate from filtered solutions obtained during the leaching of CuS and Cu2S powders. Subsequent leach ing of ground material with variable amounts of carbonate in the system, 51 consideration of the findings of Prokopchik and others on the mechanism of hypochlorite decomposition in the presence of copper hydroxide cata lysts, together with extensive analysis of the compounds made by the action of NaOCl on copper salts at pH 9.0 (Section 3.2.1) led to the conclusion that these black substances are tri-valent copper compounds. 4.1.6 Variation of the Total Carbonate Content  During Leaching A lower amount of carbonate in the leaching system caused less - 162 -copper to be dissolved, and it was held in solution for only a finite period of time. This time period, and the copper concentration in solution were found to be proportional to the contained carbonate. The rapid drop in copper concentration which subsequently occurred, accompanied by simultaneous hypochlorite decomposition can be explained by assuming sodium hypochlorite is capable of oxidizing copper to the tri-valent state, and that it is held in solution in this oxidation state by the complexing action of carbonate until such time as a solid copper1"'"'1' compound is able to nucleate and precipitate on the mineral surface. This then becomes an active catalyst for hypochlorite decomposition. In the absence of any carbonate, the hydroxide formed on the mineral surface is therefore a copper*1"1" compound rather than normal cupric hydroxide, Cu(OH)2, and hypochlorite decomposition presumably occurs by the mechanism outlined by Prokopchik (applicable to pH values below 12.0): C10~ + 2Cu(OH)2 + H2<3 > Cl~ + 2Cu(OH)3 (57) 4Cu(OH)3 >• 4Cu(OH)2 + °2 + H2° (58) (4Cu(OH)3 + CIO" >• 4Cu(OH)2 + C103~ + 2H20) (59) The observation that exposure of a small surface area of copper to a hypochlorite solution with minimal agitation, results in formation of copper11 hydroxide before deposition of the black copper*11 hydroxide illustrates step (i) of this mechanism and subsequent hypochlorite decom position would presumably follow a +2/+3 oxidation-reduction cycle. The loss of copper from solution containing 10 g/1 carbonate accompanying hypochlorite decomposition amounted to about 0.06 g/1 Cu. - 163 -If the total surface area of mineral present is taken to be 33,000 cm 2 for covellite and 27,000 cm for chalcocite, then it can be shown that this 60 ppm copper is sufficient to form a mono-molecular layer covering the entire surface (Appendix B). Analysis showed that the compound made in the presence of carbonate was actually a tri-valent copper111 carbonate, rather than the tri-valent oxide/hydroxide made with no carbonate present. The black precipitate observed to be deposited from filtered leach solu tions is therefore a copper11"1" carbonate and is presumably the same substance which coats the mineral surface when precipitated from solutions still in contact with mineral samples. The fact that this compound and the tri-valent hydroxide apparently catalyze the decomposition of hypo chlorite at the same rate is not too surprising; but the reason why chalcocite, covellite and synthetic copper sulphide still produce differ ent rates of decomposition after loss of copper from a carbonate system, when it is assumed that the entire surface is covered by a Cu111 carbon ate layer in each case, is less obvious. It could well be that catalysis is enhanced at certain active sites on the mineral surface and although the entire surface is covered by Cu1"'""'", only certain areas produce cataly sis. Any area in which adsorption can occur more readily than in other surrounding areas will act as a preferential site for catalysis. The redox potential of a sodium hypochlorite solution at pH 9.0 has a value of 1.2 V (Appendix B). This is within the region shown by. Delhez and coworkers to contain tri-valent copper as a stable compound, 74 described as hydrated Cu203- Various other studies have shown that copper1"1" salts can be oxidized to the +3 oxidation state in the presence of hypochlorite solutions, although most of this work was carried out at - 16 4 -pH levels greater than 9.0.^2'^'^2 in fact Prokopchik specifically states that tri-valent copper compounds cannot be made at pH less than 11.5, and the black product so obtained is merely dehydrated copper11 hydroxide. The present study would seem to dispute this view. In their paper acknowledging the existence of cupri-carbonate ions and the role they play in the solubility of malachite, De Zoubov 36 et al. do not consider the possibility that at sufficiently high redox potentials a tri-valent copper - carbonate complex species could exist, rather than the Cu2°3 solid substance shown in the relevant Eh-pH diagrams (Figure 3). It is unlikely that these workers used sufficiently strong oxidizing solutions to produce such a species, but it seems perfectly reasonable to assume that a complexing agent for copper11 is equally capable of holding copper^^^in solution, given that the aqueous medium has a sufficiently high potential which the hypochlorite solution does. Equation 55 could then be modified: CuS + 4NaOCl + 3NaOH + 3HC0 ~ >-3-Cu(C03)3 + Na2S04 + 4NaCl + 3H20 (60) The Cu3+ ion is not generally thought of as being stable in solu tion: its half life has been suggested by Magee and Wood^3 to be of the 66 64 65 order of 25 seconds. However, Lister, Malatesta, Malaprode, Berka et al.^ and others, have all shown that complexing agents such as tellurates and periodates will stabilize copper1"1""1" in solution. Tri-valent cuprates are also known to be soluble and relatively stable ' but these substances occur only at higher pH values. Due to its relevance to the present study, it is interesting to note Lister's observation that any copper present in solution in an - 165 -uncomplexed state, as Cu(OH)^ for example, showed catalytic activity towards sodium hypochlorite; but addition of complexing ions to the system prevented any such decomposition. This is in line with the observation noted above that copper existing in solution as a Cu111 - carbonate com plex did not produce any hypochlorite decomposition, but as soon as any solid material precipitated the decomposition reaction was catalyzed heterogeneously. 4.1.7 Effect of Varying the Hypochlorite  Concentration The fact that more copper could exist in leaching solutions con taining 7-8 g/1 hypochlorite in the presence of carbonate, than that thermodynamically calculated by De Zoubov et al. for an equivalent amount of carbonate and no oxidant, suggests that more copper can exist as a Cu111 complex than for the corresponding Cu11 complex. This was also illustrated in experiments done in the absence of hypochlorite (see below). The increase in soluble copper caused by an increase in the hypochlorite concentration is thus a related phenomenon: the presence of chloride is known to increase the solubility of certain metal salts and in the present study this was found to apply to calcium in relation to another part of the work (Section It seems likely therefore that copper is an analogous case, and that while carbonate is the princi pal complexing agent, chloride increases both the solubility and stability of the complex, giving a longer nucleation time to Cu111 precipitation. - 166 -4.1.8 Effect of Hypochlorite Removal Removal of hypochlorite from a copper sulphide carbonate leach-2-ing system showed quite clearly that less copper existed as the CutCO^)^ species than as CuCCO^)^3 . It was a little surprising that any copper at all was dissolved by agitating covellite in a solution of sodium carbonate/bicarbonate, but this is probably attributable to air-oxidation of the copper during agitation. It is significant that the dissolved copper content of these solutions containing no hypochlorite, was identical to that contained by the leaching solutions after decomposition of the hypochlorite. It was ascertained that at the higher values of total carbonate (5.0 g/1; 7.5 g/1 and 10.0 g/1) a second 'plateau' level corresponding to the Cu11 carbonate complex existed before this also decomposed, leaving no copper in solution. At lower values of contained carbonate the amounts of copper and time periods involved are too low to observe a distinct "3-stage" curve. The reason why a small drop in the copper content of the solution was so consistently observed in all experiments carried out in the presence of carbonate within the first five minutes of leach ing is still not known precisely. It could be a result of super-saturation by the very rapid oxidative action of hypochlorite on the mineral surface, necessitating precipitation of a small amount of copper to retain equilibrium solubility. While it is clear that the black precipitate put down from hypo chlorite containing solutions is an oxidized form of the malachite precip itated from copper*"1" carbonate solutions, the appearance of either substance, and the related precipitation of copper during leaching is most peculiar. - 167 -Ismay's hypothesis that the substance could be a partially oxidized sulphur species was dismissed by showing that the same amount of copper was dissolved in the presence of carbonate and hypochlorite whether the starting material was a copper sulphide or copper sulphate. It is now clear, however, that copper*'1*"'" was also produced in this study during the leaching of chalcopyrite with an unspecified, but presumably small, amount of carbonate buffer at pH 9.0, because the curves obtained were very similar to those shown in Figures 11 and 13 for covellite and chalcocite leaching. Observations made during copper sulphate agitation in carbonate solutions bare a strong resemblance to those in similar 28 31 experiments carried out by Pickering and Appleby and Lane. These workers found that mixing Na^CO^/NaHCO^ solutions with copper sulphate or copper acetate produced a deep blue solution which subsequently precipi tated either a greeen powder, CuCOH^CO^, or well defined blue crystals assumed to be a sodio-cupric double salt, Na^Cu^O^)^. The green powder observed in the present work is almost certainly malachite. Appleby noted that precipitation of the double salt could not be accelerated by "seeding," but in this case it was found that addition of previously precipitated powder, either as copper*"''"'" to hypochlorite solutions, or as malachite to copper** carbonate, gave loss of copper and further precipitation within a very short time. This would seem to confirm that the observed precipitation can . only occur after nucleation of a small amount of solid has taken place, and that this is the slow step of the process. Filtered solutions take longer to put down a precipitate than the corresponding leaching solutions in contact with a mineral because the mineral surface can provide - 168 -favourable sites on which nuclei can form. Nucleation generally occurs when the super-saturation of the separating phase has reached a certain value such that the activation barrier has been surmounted. The free energy of formation of a nucleus consists of two terms: 3 AG. = ~ V° M In- + 4T7r Z : 3V a0 where a/ao is the concentration ratio or degree of supersaturation V = molecular volume 3 = ihterfacial energy r = radius The second term represents the work necessary to create a new surface, i.e. for the formation of a new phase, while the first takes account of the energy involved in making new bonds. This assumes the nuclei to be spherical. As the degree of supersaturation increases, then AG_. will decrease, and this usually results in smaller nuclei forming by a decrease in radius, r. The term describing the rate of formation of crystal nuclei, J, can be expressed: A G_ J = A - —a-exp kT where A represents the efficiency of ionic or molecular collisions. Homogeneous nucleation is a very slow process which rarely occurs in practice; various foreign particles can act as 'catalysts' for nucleation - 169 -by reducing the energy barrier presented by AG, thus inducing heterogen eous nucleation. The induction period for nucleation may be defined by the expression k = tc*"-1* where n = number of ions required to form a cluster of critical size c = concentration of solution k = constant n is effectively reduced during heterogeneous nucleation by the incorpor ation of foreign particles into a nucleation site. The induction period involved in the precipitation of copper*"'""'" from leaching solutions in the presence of varying amounts of carbonate under otherwise identical conditions, follows an almost linear relation ship with carbonate concentration. After nucleation and precipitation of copper111 has occurred leaving a solution containing a copper*1 - carbonate complex, the same cycle is repeated with precipitation of malachite after another induction period for the nucleation of copper** crystals. Appleby's suggestion that delayed precipitation occurs due to a very slow rate of crystalliza-31 tion, thus appears to be reasonable. His observation that seeding did not enhance precipitation of the double salt, Na2Cu(C03)2, may be an indication that solutions decomposing to give malachite, Cu(OH) (CO^) , are more super-saturated. The reason for the observed variations in copper concentration of - 170 -some leaching solutions (Section 3.1.1), is probably precipitation of XXX 28 some copper prior to analysis. Pickering noted that excess water could induce precipitation of the solid compound. It was found in this study that filtered solutions sometimes turned black on dilution with water. Subsequent analysis would therefore indicate a lower copper con tent than was actually present on initial removal of the sample. 4.2.1 Analysis of Copper111 Because insufficient of the precipitated material could be collected to carry out any meaningful analysis, similar material was manufactured in much larger quantities by the action of sodium hypo chlorite on copper1"1" salts at pH 9.0. The black compound so formed was assumed to be the same as that produced by Lister^ under almost identical conditions, which he subsequently complexed with periodate or tellurate to give a tri-valent copper salt. He presumedthe initial material was a copper"'""'""1" hydroxide. Similarly in the present study the gas evolved on acidification of the black precipitate was taken to be oxygen, pro duced as a result of the reduction of Cu1""''"" v Cu*1: 2Cu(OH)3 + 4HC1 y 2CuCl2 + 5^0 + ho^ 3+ - 2+ Cu + e > Cu (61) The presence of oxygen was proven by iodometric titration and gas chromatography, arid its volume measured by the use of a mercury column. Compounds made in the absence of carbonate did not contain carbon dioxide. - 171 -The results consistently showed, however, that insufficient oxygen was present for the black powder to be 100% copper111, and the various analytical methods indicated that 25 - 30% available oxygen existed. While the percentage of copper1"'""'" did increase slightly with pH, there was no significant change in appearance of the compound when prepared in more strongly alkaline solutions. A number of other workers have stated that the sesquioxide, Cu^O^, occurs as a red compound, and that its formation is catalyzed by the presence of barium and calcium ions. In the present study, addition of these ions to solutions con taining 0.1 M hypochlorite made little difference to the final product, and no red material was ever detected. It is therefore proposed that the black compound produced is a mixed copper oxide or hydroxide, consisting partially of copper in the +3 state, and partially in the +2 state, with the latter actually repre senting the majority. The fact that it is black would support this 72 view: this is a characteristic of mixed oxides. Delhez also reported obtaining a mixture of copper""""1""1" and copper1"'", with the tri-valent state again amounting to no more than 30% of the total, but his compound was exclusively red. It should be noted however that he was using solutions significantly more concentrated in hypochlorite than 0.1 M, at a higher pH level and with a greater ionic strength than in this case. It is quite likely that he obtained a mixture of the two different oxides: 30% Cu2°3 + 70% CuO which appeared red, while the black compound des cribed here is actually a mixed valence state compound consisting of Cu^O^ and CuO together. Assuming a mixture of 75% copper"'""'" and 25% III 2+ 3+ copper the ratio Cu :Cu would be 3:1. - 172 -Cu-O, + 3CuO "-Cu^Cv (62) 2 3  6 This gives a compound having an oxidation state of 2.4. By further analogy with Delhez's work, and from the results obtained by infra-red spectrophotometry, it appears that the compound fluctuates between an oxide and a hydroxide. Thus it could consist of Cu(OH)3 and Cu(OH)2 giving a compound such as Cu^(OH) with an oxidation state of 2.5. This would have a copper content of about 60% which fits the experimental analysis more closely. Magnetic susceptibility tests further confirmed that the compound is only partially Cu111, while X-ray diffraction suggested it to be cupric oxide. No reported X-ray diffraction data for copper1"'""'" compounds was found. It is possible that Cu2°3 1S amorphous, in which case no diffraction peaks would be detected. 4.2.2 Precipitation of copper11"1" in the  presence of carbonate X-ray tests carried out on compounds made in hypochlorite contain ing carbonate solutions showed peaks corresponding to both malachite and cupric oxide, confirming it to be a mixture of a carbonate and an oxide. There were also some unidentified peaks which could represent the difference between Cu1""" carbonate and the Cu""*"'"'"" salt. Infra-red studies and limewater tests suggested further that this compound was at least partially a carbonate. The nucleation times involved in its preparation are not linear with respect to carbonate concentrations. The conditions of temperature, agitaiton etc. were not maintained perfectly constant during preparation of each sample however, as they were in the leaching experiments. It was again shown that 'seeding' would induce precipitation of the compound. Both oxide/hydroxide and carbonate samples were found to catalyze the hypochlorite decomposition reaction at approximately the same rate, when added in equal amounts to a 0.1 M solution. This is to be expected, because both compounds contain approximately the same amount of copper11"1", and it has been shown by Lister,-3^ Prokopchik^4 and Delhez^2 that Cu"""1""" is the active catalyst for decomposition. When none is present initially oxidation of Cu""""1" must take place before the reaction occurs, thereby 2+ 3+ setting up a Cu /Cu oxidation-reduction cycle. This is presumably why cupric oxide showed a slower rate of decomposition under identical conditions. 4.3 Sodium Hypochlorite Leaching of  Molybdenite, MoS,, Leaching molybdenite mineral in sodium hypochlorite solutions at 2- . pH 9.0 did not initially yield 100% molybdenum as MoO^ in solution. Examination of leach residues revealed the presence of iron, copper, zinc and silica, as well as lead molybdate which is insoluble. Thus the assumption of 100% MoS^ was a false one, and spectroscopic analysis of unleached samples revealed at least a 3% impurity content. Subsequent leaching with "98% +" molybdenum disulphide produced extractions in excess of 97%, confirming Ismay's findings that molybdenite leaches rapidly and completely at pH 9.0, 35°C and in solutions containing about 0.1 M OC1 . Hypochlorite consumption was found to correspond to the - 174 -previously reported stoichiometric MoS2 + 9NaOCl + 6NaOH :—* Na2Mo04 + 9NaCl + 2Na2S04 + 3H20 (63) 4.3.1 Sodium hypochlorite leaching of  molybdenite and copper sulphide  minerals at pH 9.0 It has already been reported in conjunction with the leaching behaviour of copper sulphide minerals, that carbonate plays a significant role in this system besides that of a buffering reagent. In addition to acting as a complexing agent for copper and hence a stabilizer of the hypochlorite however, it was also found to affect molybdenum extraction. When excluded from the system, the level of molybdenum extraction dropped from 98% to only 76 - 84%. The actual depression was a function of the particular mineral being leached. It was postulated that copper molybdate compounds were interfering with molybdenum extraction, but it was rather surprising that the case in which no copper was present in solution showed the greatest molybdenum loss. Oxidation of molybdenite in the presence of CuS04 solution suggest ed further that precipitation of copper molybdate was occurring. This did not explain why molybdenum was also precipitated from a covellite/ sodium molybdate leach, in which no copper had been dissolved. 4.3.2 Leaching of molybdenite and copper  sulphides at pH 5.5 Experiments were carried out in acidic solutions to see if an - 175 -- 176 -improved molybdenum yield could be obtained, even if this incurred production of more sodium chlorate than at pH 9.0. -2 Copper exists in solution to a maximum concentration of 10 M 2+ as the Cu ion at pH 5.5. This fact proved very detrimental to molyb-2+ denum extraction. Cu ions can combine with the molybdate species 2-MoO^ giving insoluble CuMoO^. The rapid loss of 90% of the molybdenum in slution on introduction of a copper sulphide mineral can be interpreted as precipitation of this molybdate salt. The presence or otherwise of carbonates in the system did not have the effect of preventing molybdenum losses that had been observed in alkaline solutions. The only detectable difference between a carbonate and non-carbonate leach in the experiments at pH 5.5 was a slightly higher copper content in the former case. The fact that such a small increase in pH, a value of 6.0, resulted in the carbonate being able to suppress formation of copper molybdate, can be explained by careful 36 consideration of the Eh-pH diagrams given by De Zoubov et al. showing the existence of the stable copper carbonate compounds and complex ions for the Cu-H^O-CO^ system (Figure 3). The Cu(C03) aq. molecule exists as a soluble species at pH levels between 5.0 and 7.0 dependent on the total carbonate content of the solution. At the level used here (10 g/1 HCO^ E 1.6 x 10 1 M tc03]T) it is stable at pH values above 5.4; below 2+ this the copper exists exclusively as the Cu ion (Figure 51). Thus 2+ leaching at pH 5.5 puts copper at the stability boundary between Cu / CujCO^) aq. The experimental results indicated that copper is only com-plexed by carbonate to a very small degree, shown by the small increase in dissolved copper, and that this association is insufficient to prevent - 177 -copper molybdate formation. An increase in pH to 6.0 puts copper into the region where Cu(CC>3) aq. is stable in solution, thereby preventing a reaction with the molybdate. Leaching in the absence of carbonate, 2+ 2-however, enables Cu and MoO^ to combine as before. The observed rapid decomposition of the hypochlorite was presumed to be caused by copper molybdate catalysis, suggesting again that this cannot form in carbonate solutions above pH 5.5 - 6.0 where no excess hypochlorite decomposition was observed. As the pH of the leaching system is further increased, the 2+ -5 solubility of the Cu ion decreases to 1.6 x 10 M. This is still sufficient to give a 60% loss of the molybdenum, coupled with an even greater loss of hypochlorite i) due to CuMoO^ catalysis and ii) because this is the region in which hypochlorite most readily decomposes to chlorate. These findings would seem to question the feasibility of the U.S. Bureau of Mines electrooxidation process at the suggested pH levels of 20 21 23 5.0-7.0. ' ' The statement made in conjunction with this investi gation that a minimum of copper solubility occurs at pH 7.0 is in itself a little dubious since this is shown to occur for CuO and Cu(0H)2 at a 73 value of 9.0, and is not greatly modified by the presence of chlorides. Secondly, no mention in the literature of the soluble compounds Cu(OH)2C03 or CuC^CO^ which Scheiner et al. claim as being responsible for copper solubility in alkaline solutions, has been found; and the cupri-carbonate 2-ion, Cu(C03)2 appears to be the only species which can impart significant solubility to copper at pH values above 9.0. Scheiner mentions the compound copper molybdate as being detrimental to molybdenum extraction - 178 -and reports detecting it in leaching tails, but does not appear to have studied the conditions for its formation. In this electrooxidation process the pH is allowed to rise during leaching by the addition of sodium carbonate on a semi-continuous basis. In the light of present findings any copper molybdate which did form, would be in the initial stages of a run when the pH is still below 6.0. The most surprising conclusion of the U.S. Bureau of Mines study, however, is the fact that chalcopyrite is unaffected by hypochlorite, but that chalcocite is oxidized to a certain extent giving soluble copper compounds which are capable of combining with the molybdate ions to form insoluble copper molybdate. In the present study no difference in the leaching behaviour of the dif ferent copper sulphide minerals exposed to hypochlorite solutions has been detected, except in the catalyzed rate of hypochlorite decomposition. Furthermore, and in contrast to Scheiner's statement, it is in the presence of the soluble copper compound Cu(C03)2 aq. that copper molybdate formation 2+ is suppressed. Only when copper exists in solution as the cupric ion Cu , in the absence of carbonate or at pH <6.0 can a reaction with molybdate ions occur. This should hold true for any copper mineral introduced into the system. This implies that leaching copper - molybdenum concentrates on the acid side of neutrality is only feasible in the presence of carbonate, and in a limited pH region of =6.0 - 7.5. This would not give any improve ment in molybdenum selectivity over that at pH 9.0 in the presence of carbonate, and would have the added disadvantage of producing large amounts of sodium chlorate. Barr and coworkers have shown that this represents a source of power loss in the electrooxidation process and that - 179 -it also interferes with subsequent molybdenum recovery. ' Their suggestion of leaching at pH values of 4.0 - 5.0 to overcome the chlorate problem, is hardly practical on account of the stability of copper molyb date in .this region. 4.3.3 Copper molybdate The simple molybdate, CuMoO^, is not well documented in the literature. It is known to be insoluble, in common with most other 82 transition metal molybdate salts. Zelikman quotes this solubility as 0.017 M which is slightly higher than that found in the present work of 0.00783 M. X-ray data for a number of copper molybdate species has been pub lished. The one corresponding closely to diffraction patterns obtained in this study (Table 8), is the hydrated species Cu3(MoO^)2(OH)2 which occurs in nature as the mineral Lindgrenite. Data for other :copper molybdates are shown in Table 10: Species dA (1) dA (2) dA (3) CuMoO,, 4 3.72 3.36 2.71 CuMoO. 4 3.05 3.30 3.53 Cu2Mo05 3.54 3.45 3.32 Cu3Mo2°9 3.44 2.63 3.39 CU4-xMO3012 3.42 3.31 2.65 CU4-xM°3°12 3.42 2.64 2.76 CU6M°4°15 3.38 2.89 2.63 ( This study 3.52 4.20 2.53 ) ( Cu3(Mo04)2(OH)2 3.50 4.15 2.67 ) Table 10: Strongest X-ray Peaks for Various Copper Molybdate Species - 18 0 -The catalytic effects on hypochlorite decomposition observed during leaching in acid solutions were confirmed by agitating copper molybdate in sodium hypochlorite solutions. This is a further reason for preventing 'in-situ' formation of the compound even in small amounts during molybdenite leaching. An Eh-pH diagram for the Cu-H^O-MoO^ system was constructed to confirm the thermodynamic stability of copper molybdate in acid hypochlor ite solutions. In Figure 52 the solid species considered are Cu, C^O, CuO, CuMo04 and Cu203, and in Figure 53, Cu, Cu20, Cu(OH)2, CuMo04 and Cu203. In both diagrams the dissolved species have been taken as Cu+, 2+ 3+ — 2— 2— — Cu , Cu , Cu02H , Cu02 and Mo04 , HMo04 . The diagrams show the equilibria between stable substances when the activities of copper con taining ions in solution are 10 ^ M, and of molybdenum species are 10 1 M. Considering the solid species to be oxides, Figure 52 shows that copper molybdate has a region of stability from pH -0.3 to 8.65, with an oxidation potential varying from =:0.03 to 2.4 V. Thus at pH 9.0 there should be no interference with molybdenum extraction from CuMo04 forma tion. An increase in copper concentration would decrease the pH limit for CuMo04 stability, while decreasing molybdenum activity would increase it (and vice versa). In Figure 53, in which the hydroxides are shown as the stable phase, the zone of copper molybdate stability is extended slightly to a pH value of 9.18. As the hydroxide is the meta-stable phase and would form prior to the oxide, it is apparent from this diagram that CuMo04 formation could pose a problem to a leach at pH 9.0. It also shows clearly why a molybdenite - copper sulphide leach would not be feasible - 181 -Figure 52: Eh-pH diagram for the system Cu-H20 - Mo04 (oxide species). (Activities of copper ions 10_6M and activities of molybdenum ions 10_1M). \ - 182 -Figure 53:. Eh-pH diagram for the system Cu-H20 - MQO^ (hydroxide system). (Activities of copper ions 10~6M and activities of molybdenum ions 10_1M). - 18.3 -in the pH region from 5.0 - 7.0, and illustrates that copper formation would not be encountered in acid solutions due to the higher redox poten tial of the +3 state. Although copper molybdate could therefore exist at pH 9.0 from a thermodynamic point of view, the fact that no copper was dissolved in the absence of carbonate, together with results obtained from leaching with synthetic copper sulphides and at pH 10.0 (Section 3.3.4), confirmed that copper itself was not the principal cause of poor molybdenum recovery at this pH value. In addition to this, aging of copper containing molybdate solutions showed that at pH 5.5 CuMoO^ readily formed as a stoichiometric compound while at pH 9.5 a co-precipitation phenomenon occurred. Any copper added to a solution at this pH will precipitate as Cu(0H)2, and the fact that some molybdenum dropped out of solution initially but reappeared on standing or heating, is strongly indicative of the molybdenum co-precipi tating with the copper. A similar effect would explain why molybdenum was lost from a leaching solution to which copper sulphate was added at pH 9.0, i.e. due to co-precipitation with the copper as Cu(OH)2. 4.3.4 Solubility of calcium in  hypochlorite solutions A combination of these various factors brought to light the fact that other elements, contained in the copper sulphide minerals under study as impurities, could be capable of forming insoluble molybdates and thus be detrimental to good molybdenum recovery. In order to - 18 4 -elucidate which one or more elements were responsible in the present case, as well as to determine if any other likely gangue elements could cause similar problems with other ore samples, the following factors were taken into consideration: i) whether an element can combine with the molybdate ion and form an insoluble species which is thermodynamically stable at pH 9.0. ii) whether hypochlorite oxidation leads to the element existing as a stable cation in solutions at pH 9.0. iii) whether the element is likely to be found as an impurity in copper - molybdenum porphyries, and associated minerals. To determine the regions of stability of the simple molybdates of elements known to be contained in the copper sulphide minerals under study, Eh-pH diagrams were constructed for the following systems: Ca-H20-Mo04 Fe-H20-Mo04 "Pb-H20-Mo04 Cd-H 0-MoO„ 2 4 Zn-H20-Mo04 shown as Figures 54-58 respectively. The activity of the molybdate ion has been taken as 10 ^ M in all cases, and that of the metal containing -6 -4 species is {10 } except for lead, where {10 } was found to more appro priate. Figure 54: calcium molybdate: The diagram drawn up for calcium 2+ species and their interaction with molybdate, shows that the Ca ion 2+ has a much larger zone of stability than the corresponding region for Cu resulting in a calcium molybdate compound with a wide range of stability. - 18 6 -When considering the solid species CaO, calcium molybdate is stable up to a pH of 19.14, while for the hydroxide system the molybdate would extend to a pH .of. 14..0. Figure 55: lead molybdate: PbMoO^ is stable up to a pH value 2-of 14.2 before dissolution into plumbite ions, HPbO^ , occurs. The redox potential varies between 0.3 V at this pH, to 2.0 Vat pH 0. Thus at a value of 9.0, the potential at which PbMoO^ decomposes to give PbO^ is 0.92 V and this molybdate should not be stable in hypochlorite solutions with a predicted value of 1.2 V below pH values of 13.5. Figure 56: zinc molybdate: It is seen that zinc molybdate has a large zone of stability which exists at all potentials more positive than 0.4 V and up to a pH value of 18.6. The white, amorphous zinc hydroxide was considered as the stable oxide phase in setting up the diagram, and it is apparent that the molybdate is more stable thermo-2+ - 2-dynamically than Zn , Zn(OH)2, HZnC>2 and ZnC>2 until dissolution of this final species occurs at pH 18.0. Figure 57: iron molybdate: Iron is of major interest with respect to its interaction with molybdate, being present in large amounts in chalcopyrite, CuFeS2 and bornite, Cu^FeS^, as well as a common impurity in chalcocite and covellite. It is unlikely however that ferrous iron could exist in hypochlorite solutions, thus FeMoO^ should not pose a problem. This is confirmed in the Eh-pH diagram; FeMoO^ is stable up to a pH of 10.8, but the maximum potential it can withstand is 1.2 V, at pH 2.4, and at pH 9.0 decomposition to Fe(0H)3 would occur at less than 0.5 V. 92 Zelikman mentions theexistence of a ferric molybdate compound, Figure 55: Eh-pH diagram for the system Pb-H„0-MoO^ (Activities of lead ions 10~^M, activities of molybdenum ions 10_^M). PH Figure 56: Eh-pH diagram for the system Zn-H^O-MoO^. (Activities of zinc ions 10~^M, and activities of molybdenum ions ^Q-IJM) • ' Figure 57: Eh-pH diagram for the system Fe-H„0 MoO^ (Activities of ferric and ferrous ions 10"%, ana activities of molybdenum ions 10_1M). - 19 0 -Fe^(MoO^) and notes that it will precipitate from aqueous solution under certain conditions. The mineral ferrimolybdite, Fe^(MoO^) -SH^O, is quite commonly found in nature but is not a source of molybdenum metal. No thermodynamic data for such an iron11"'" molybdate compound could be found and it has not been included in the Eh-pH diagram shown. Figure 58: cadmium molybdate: Cadmium and zinc have several chemical similarities and are found in close association in many ore bodies. The present study also shows that their respective molybdates form under similar conditions: CdMoO^ is stable up to pH 12.5 and above a potential of 0.4 V (pH 0.0)- -0.9 V (pH 14.0). The presence of 2+ Cd ions in solution is thus undesirable during molybdenite leaching. 2-The stability boundary between the simple molybdate ion MoO^ and polymolybdate species, described as HMoO^ , occurs at pH 6.0. Above this value, combination with metal cations will result in precipitation of a simple molybdate salt such as CuMoO^. In more acidic solutions it is possible that a complex molybdate species will form the stable species. 2-On the other hand, with a sufficiently large concentration of MoO^ in equilibrium with HMoO^ , a simple molybdate may still precipitate. Experimental results for molybdenite - copper sulphide leaching carried out at pH 5.5 suggests a stoichiometric Cu:Mo ratio of 1:1, indicating formation of the simple molybdate CuMoO^. The exact nature of molyb denum species at low pH values is not known. MoO^ is the stable species at higher potentials below pH 3.0. MoC>2 precipitates at a potential of 0.4 V (pH 0.0), decreasing to -0.1 V at pH 6.0 and -0.55 at pH 10.0. Between these two species molybdenum probably exists in the +5 oxidation 9 state. - 191 -Figure 58: Eh-pH diagram for the system Cd-H^O-MoO^. (Activities of cadmium ions 10~^M and activities of molybdenum ions 10_1M) - 19 2 -In conjunction .with these thermodynamic considerations for the stability of various molybdates, the following data was accumulated for the solubility of a number of molybdate compounds: Compound Solubility (g Mo04/1) Source PbMo04 1.2 x lo"6 83 SrMoCK 4 3.0 x lo"4 83 CaMoO„ 4 a) 5.0 x 10~4; b) 1.3 x 10~3 a) 83; b) 82 BaMo04 5.5 x lo"4 83 CdMoO„ 4 6.7 x io"4 83 FeMoO„ 4 7.6 x lo"4 82 CuMoO. 4 3.8 x io"3 82 Ag2Mo04 3.86 x 10 3 83 ZnMoO. 4 a) 3.70 x 10 ; b) 5.0 x 10 2 a) 82; b) 83 Of the above elements barium, strontium and silver were not con sidered because it was not felt they would be present in sufficiently large quantities to cause molybdenum losses. From the remainder it appears from the Eh-pH diagrams that lead and iron are both insoluble in hypochlorite solutions as discussed above. This was borne out experimentally: samples of galena, PbS and pyrite, FeS2 were leached under identical conditions to those used for copper/ molybdenum minerals and no lead or iron whatsoever was detected in solution. Copper molybdate can be described as a borderline case with respect - 193 -to leaching at pH 9.0 since the thermodynamic boundary occurs at pH 8.65 for the oxide and 9.2 for the hydroxide'; but the fact that both these salts are insoluble should prevent any molybdate formation. As noted previously, the presence of carbonate imparts a certain degree of solubility to copper in alkaline solution, and experimental results suggested that a copper - carbonate association occurred in preference to CuMoO^ formation. The relevant equilibria are: Cu2+ + 2C032~ v==^ Cu(C03)22~ log k = 9.83 2+ 2-Cu + Mo04 v=^ CuMo04 log k = -2.027 It was observed, however, that leaching in the presence of small amounts of carbonate did result in some copper molybdate formation (Section, suggesting that less carbonate produced a weaker Cu-C03 bond. Zinc, cadmium and calcium are all undesirable elements in this system on account of their ability to exist as soluble species at pH 2-9.0 and hence to combine with Mo04 and precipitate. It has been shown that sodium hypochlorite is a feasible lixiviant for the extraction of 84 zinc from sphalerite ores. However, it is not likely that either zinc or cadmium would be present in more than trace amounts in most copper -molybdenum ore bodies, and analysis certainly showed that insufficient amounts were present to account for the large losses of molybdenum observed during leaching experiments. By a process of elimination in consideration of all the above factors it was concluded that calcium was the element primarily respons ible for the observed poor molybdenum recoveries. This would readily - 19.4 -explain the loss of molybdenum from combined copper sulphide molybdenite experiments in the absence of carbonate. Calcium was present as an impuri ty in all three mineral samples used; it can exist as a soluble species at pH 9.0, and has a very insoluble molybdate salt. In the presence of carbonate, however, the calcium will precipitate preferentially as CaCO^ so that negligible interference with molybdate occurs. The fact that copper is also dissolved in the latter case is coincidental. 2+ 2- ^ Ca + MoO. log k = -7.38 4 S 2+ 2 — Ca + CO., log k = -8.55 3 S Stoichiometric amounts of calcium and molybdenum to give a 1:1 ratio for CaMoO^ were precipitated from a solution containing molybdenite and calcium chloride, except in the presence of carbonate buffers when all the calcium was precipitated with no effect on the molybdenum. The fact that extraction in the presence of carbonate and copper sulphide minerals was slightly below the levels obtained for leaching molybdenite alone, could be attributed to one or more of the following factors: i) formation of small amounts of copper molybdate, even in 2-the presence of soluble Cu(C0.j)2 , ii) co-precipitation of the molybdenum with calcium carbonate, iii) slight solubility of CaCO^ due to the presence of chloride, enabling CaMoO^ to form in small amounts. The implications of calcium molybdate precipitation are unfavour able for an alkaline leaching process in which the potential feed material is a copper - rougher concentrate, which would unavoidably contain a certain amount of gangue material as well as the copper - molybdenum - 195 -metal values. The problems of such gangue materials being soluble in hypo chlorite solution and capable of forming insoluble molybdates would pose an even greater threat to an 'in-situ' leaching process where good selec-tivity of metal values is a necessity. It is interesting to note Bhappu's comments in this respect: he states that either a basic carbonate -hypochlorite or an acid chlorate leach would be capable of dissolving molybdenum from mixed sulphide - oxide ores, but that an alkali leach would be preferable on account of the tendency of ferric iron, calcium and aluminum to precipitate as insoluble molybdates in acid solution. Presumably he was also relying on the presence of carbonate to suppress molybdate formation in alkali solution. The major calcium minerals likely to be associated with copper -molybdenum porphyry ores were considered to be calcite and gypsum, and a brief review of the geology of porphyry type deposits confirmed this: Chalcopyrite, molybdenite and pyrite are the most commonly occurring minerals in such ore bodies. Other sulphides found in association with these three include bornite, chalcocite, galena, sphalerite, covellite, pyrrhotite, bismuthinite and enargite. The most abundant gangue material is quartz but magnetite, tourmaline, hematite and fluorite are also common. Calcite, CaCO^, and gypsum, CaSO^'211^0, are found in association with either quartz or tourmaline distributions. In most ore deposits they occur in the youngest rocks present, and represent late stage veining which is therefore usually post-ore. Calcite, together with zeolite can intergrow with the sulphide minerals and is also found, with gypsum in veins and fracture fillings in the quartz. Fibrous gypsum can also be present as vugs in the main late stage quartz - sulphide veins. The distribution of the copper and molybdenum mineralization can - 19 6 -vary considerably between different ore bodies. In some cases both metals are contained within the central core of a deposit, while in others the copper especially is found in the outer zones, and therefore in associa tion with quartz, sericite, pyrite etc. The molybdenite is often unsystem-atically distributed throughout the ore body and can cross cut the areas in which calcite and gypsum are found. Of the molybdenum-bearing porphyry copper ores mined in British Columbia, deposits at Bethlehem Copper, Lornex, Valley Copper, Ox Lake, Brenda and the 'J.A.' ore body are all known to contain calcite and/or gypsum in close proximity to the metal sulphides. In the so-called molybdenum porphyries, calcium containing minerals are less common, but in the Boss Mountain deposit and at Endako, significant amounts of calcite are found in association with quartz and clay minerals contained in the ore bodies.4 Sodium hypochlorite leaching of calcium minerals Having established that calcite and gypsum are likely gangue mater ials in low grade copper-molybdenum ores, experiments were carried out in which samples of these minerals were leached in hypochlorite solutions under identical conditions to those used for molybdenite. 2+ It was confirmed that both minerals yield Ca in solution in the presence of hypochlorite, and that the amounts extracted were sufficient to account for the observed molybdenum losses during copper sulphide -molybdenite leaching. Effect of chloride concentration  on calcium solubility Samples of calcite and gypsum leached in the presence of varying - 197 -amounts of total chloride showed that calcium solubility increases as a function of [CI ] and that the results were in good agreement with literature values (Table 11). Stronger hypochlorite solutions can thus be expected to extract more calcium. Identical results were obtained whether 0C1 or an equivalent amount of NaCl alone were used. This did not support the suggestion made in conjunction with the experimental observation that calcium sulphate is soluble in hypochlorite solutions, namely that this was due to calcium complexing agents contained in a commercial bleach. Furthermore, an assurance was obtained from the manu facturers of JAVEX that the product, when sold, consists exclusively of 89 sodium hypochlorite and water. The reason for the increased solubility of calcium in chloride solutions is a change in activity coefficient due to an increased ionic strength. The thermodynamic solubility product for calcium sulphate is derived from the equation: 2+ 2-CaS04*xH20 ===== Ca + S04 + xH^O . - 198 -i) Calcium sulphate Chloride Content [Ca] Dissolved [Ca] Dissolved (Moles Cl~) (Experiment) g/1 (Reported) g/1 Source 0.0 0.712 0.666 85 0.825 86 0.2 1.326 1.068 85 1.224 87 0.5 1.528 1.532 87 1.442 88 0.7 1.688 1.496 88 1.0 1.860 1.833 88 1.800 87 1.5 2.012 1.959 88 2.0 2.024 2.116 88 2.180 87 ii) Calcium carbonate Chloride Content [Ca] Dissolved [Ca] Dissolved Source (Moles CI ) (Experiment) g/1 (Reported) g/1 0.0 0.016 0.022 85 0.014 86 0.25 0.016 0.048 85 0.60 0.049 0.056 85 0.70 0.060 0.056 85 1.37 0.055 0.055 85 Table 11: Comparison of Experimental and Literature Values for Solubility of i) CaSC~>4 and ii) CaCCy in Chloride Solutions  of Varying Strengths - 199 -for which k sp = tmCa2+] [mS042"] [6Ca2+] ^SO^ [aH20]X where m„ , m„ „ are iholal concentrations Ca S0„ 4 &n ' &cr\ are activity coefficients 4 x = number of moles of water a = water activity in the solution. 2+ 2-Substituting for activity coefficients of Ca and S04 from the Debye Huckel expression: -A z log 6 1 1 + B/l where ionic strength I = 4 m CaSo4 = m NaCl z = charge on ion i A,B are constants then log k = log k +2 - x log aTT „ SP SP 1+ B/F H2° Assuming that m 2+ = m 2-Ca. bu 4 I3* then, log S = log S° + 4A j- - — log an 1+B12 2 H2 where S = molal solubility fo CaSC>4 at any ionic strength, I and S° is the theoretical solubility when 1=0. It is apparent therefore that S increases with an increase in m, nll NaCl and hence of ionic strength, I. - 200 - Removal of calcium from  leaching solutions The minimum carbonate:calcium ratio for effective precipitation of CaCo^ and prevention of calcium molybdate formation was found to be 2- 2+ 3.19 M [CO^ ]: 1 M [Ca ]• For the amount of calcium contained in the covellite mineral under study, 0.058 g/1, this gave a minimum carbonate requirement of 0.55 g/1. However, leaching with this value of total 2- III [CO^ ] at pH 9.0 gave almost instantaneous precipitation of copper and rapid decomposition of the sodium hypochlorite, as noted earlier. Increasing the carbonate above the stoichiometric requirement lengthened the induction period to precipitation of copper1"1"''', but gave more copper in solution initially. The fact that the presence of molybdenite seeming ly decreased this induction period is probably due to the lower hypo chlorite content of solution as a result of the oxidation of MoS2 to Na^oO^, rather than any effect of the molybdenum itself. The small degree of copper molybdate formation which occurred at lower carbonate levels has been pointed out previously to be detrimental to molybdenum recovery and an undesirable catalyst for hypochlorite decomposition. There would seem to be two alternatives therefore for the extraction of molybdenum from molybdenite - copper sulphide concentrates: i) to leach in the presence of large amounts of carbonate buffer reagents and hence to precipitate calcium, as well as other soluble im purities such as zinc and cadmium, as insoluble carbonate salts and to prevent any copper molybdate precipitation. At the same time a small percentage of the copper would be dissolved. This would enable all the molybdenum to be extracted before nucleation of copper1"'"''" and decomposition - 201 -of the hypochlorite lixiviant could occur, but would probably necessitate a copper - molybdenum separation step prior to molybdenum recovery from the solution. ii) to leach with no carbonate in the system, and thereby to ob tain complete selectivity of molybdenum over copper, but to decompose all the hypochlorite reagent as a result of heterogeneous copper catalysis. This would necessitate constant regeneration of the leaching solution. In case (ii) an alternative suppressant for calcium would have to be found, and experiment suggested either phosphates or silicates may be feasible. Leaching chalcopyrite and molybdenite together in the presence of silicate in fact showed that very good molybdenum recoveries were obtained, coupled with no copper dissolution and an apparent stabiliza tion of the hypochlorite. Copper meta-silicate, CuSiO^, is known to exist, and such a compound could form at the mineral surface in preference to Cu(0H)2 and/or Cu(0H)3 and thus prevent rapid catalytic decomposition of the hypochlorite. While the use of silicate appears to be a very viable alternative for leaching at pH 9.0, it should be noted: i) Silicates and polymolybdates can combine in acid solutions and problems could be incurred during subsequent molybdenum recovery if silicate was present in the leaching solutions. ii) Silicate reagents are likely to be more expensive than carbonates. The most favourable process thus seems to be one incorporating leaching in the presence of sodium carbonate/bicarbonate at pH 9.0, pro vided a suitable method for copper - molybdenum separation after leaching can be found. - 202 -CHAPTER 5 5.1 Conclusions 1) The selective extraction of molybdenum from copper sulphide -molybdenite concentrates can be rapidly and completely obtained by leach ing with sodium hypochlorite solutions at pH 9.0 and a temperature of 35°C, in the absence of carbonate buffer reagents. 2) Under these conditions the surfaces of the copper minerals act as catalysts for the heterogeneous decomposition of sodium hypochlor ite due to the formation of an oxide or hydroxide salt of tri-valent copper on the mineral surface. 3) Alkaline carbonate solutions are capable of dissolving small amounts of copper as copper - carbonate complexes which are considerably 2+ more soluble than the Cu ion alone at pH values above 7.0. It is suggested that the presence of hypochlorite in such solutions enables the copper to exist as a tri-valent cupri-carbonate species, and that precipitation of a solid carbonate compound containing up to 30% of the copper in the +3 oxidation state occurs after a suitable nucleation period has passed. This time period and the original copper content of the solution are a function of the total carbonate concentration in the system. 4) Soluble impurity elements which have insoluble molybdate salts are detrimental to good molybdenum extraction. Calcium is suggested as being the principal source of potential molybdenum losses in the - 20 3 -proposed copper sulphide - molybdenite leaching process. The presence of carbonate in the system would prevent calcium molybdate formation by the alternative precipitation of calcium carbonate, and a similar mechanism would apply to zinc and cadmium impurities. The use of sili cates would provide an alternative means of calcium precipitation with out the coincidental side effects of copper complexing, but may prove detrimental to subsequent molybdenum recovery. 6) Lead and iron do not exist as soluble species in hypochlorite solutions at pH 9.0 and will not interfere with molybdenum extraction. Copper molybdate is of minor importance at pH 9.0 due to the low 2+ solubility of the Cu ion in alkaline solution, but in acid solution precipitation of 90% of the dissolved molybdenum can be attributed to copper molybdate formation. This compound is a catalyst for hypochlorite decomposition and has a solubility of about 1.25 g/1 as CuMoO^. 7) Leaching in acid solutions would be feasible only in the presence of carbonate and above a pH value of 6.0, because this is the 2+ region in which the cupric ion, Cu , is complexed by carbonate anions, and copper molybdate formation is therefore prevented. The pH region 6.0 - 7.0 has previously been shown to be the most favourable for sodium chlorate formation. 8) Leaching in alkaline hypochlorite solutions in the presence 2-of sufficiently large amounts of carbonate (25 g/1 [CO^ ] for example) would give good molybdenum extraction with approximate stoichiometric consumption of the lixiviant; but would entail simultaneous dissolution of a certain amount of copper. This copper would probably require separation - 204 -from the molybdenum before the recovery stage. To minimize sodium hypochlorite decomposition such a separation step should be made before III . • . copper precipitation can occur. 5.2 Suggestions for Future Work Further work could be usefully done in the following areas relating to the present study: i) A more detailed study of the surface reactions occurring between hypochlorite and different copper sulphide minerals would deter mine more precisely why such variations were observed in the rates of catalyzed decomposition. ii) Further analysis of the carbonate containing hypochlorite solutions, to provide a more conclusive indication of the soluble copper species present. Polarographic techniques would probably give different half wave potentials for copper carbonate complexes in the +2 and +3 oxi dation states. iii) Analysis of solid compounds thought to contain tri-valent copper proved to be somewhat difficult due to a lack of published data for standard compounds. If pure samples of Cu^O^ or (""^(CO^)^ could be manufactured, X-ray diffraction patterns, gram magnetic susceptibilities, solubility data etc. could be usefully obtained. iv) Leaching was found to be most successful in alkaline solutions. It would therefore be logical to research into ways of recovering molyb denum from leach solutions without lowering pH. The ideal method would - 205 -be one of solvent extraction or ion exchange operating at pH 9.0. This would give copper - molybdenum separation and molybdenum recovery in a single stage and would prevent unnecessary chlorate formation by maintain ing an alkaline system throughout the process. v) Use of typical copper rougher concentrates•as feed material would illustrate the feasibility of this type of process more clearly; and show whether or not the suggested conditions are optimum for good molybdenum recovery. Flotation studies on the leaching tails could also be usefully pursued. - 206 -REFERENCES 1. 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Chem., 1_, No. 12, 1956, p. 2778. - 211 APPENDIX A Tables of Experimental Results Table I: NaOCl leaching of covellite in the presence of carbonate buffers (Figures 6,8) Temp. 35°C,- pH 9.0, Agitation 748 rpm, [NaOCl] 7.68 g/1, CuS 10 g, NaHC03 14.15 g, Na2C°3 1-01 9-Time [Cu] [NaOCl] (mins) ppm g/1 5 133 5.64 10 133 15 125 20 127 5.60 30 127 45 127 60 121 5.62 90 126 120 128 150 130 5.64 - 212 -Table II: NaOCl leaching of chalcocite in the presence  of carbonate buffers (Figures 6,8) Temp 35°C, Agitation 720 rpm, [NaOCl] 7.8 g/1, pH 9.0, Cu2S 10 g, NaHC03 14.15 g, Na2C03 1.01 g Time [Cu] [NaOCl] (mins) ppm g/1 5 122.5 7.08 10 122 15 115 6.07 30 116 45 118 6.10 60 118 90 113 6.05 120 118 150 115 6.07 ) / - 213 -Table III: NaOCl leaching of chalcopyrite in the presence of  carbonate buffers (Figures 6,8) Temp 35°C, Agitation 803 rpm, [NaOCl] 8.13 g'/l, pH 9.0, CuFeS2 10 g, NaHC03 14.15 g, Na2C03 1.01 g. Time [Cu] [NaOCl] '(mins) ppm g/1 2 105 5 104 7.56 10 96 7.44 20 99 7.36 30 96 7.29 50 101 7.24 60 106 7.24 80 106 100 99 7.22 120 95 150 101 180 104 7.20 210 106 - 214 -Table IV: Effect of separating leaching solution from mineral slurry during agitation of CuFeS^ in NaOCl (Figure 7) Temp 35°C, Agitation 760 rpm, [NaOCl] 8.0 g/1, pH 9.0, CuFeS2 10 g, NaHC03 14.15 g, Na2C03 1.01 g. Time [Cu] slurry [Cu] filtered soln (mins) ppm ppm 5 118 129 10 111 129 15 108 129 20 105 130 30 103 128 45 109 130 65 105 129 100 104 130 140 110 130 180 117 131 - 215 -Table V: Effect of separating leaching solution from mineral slurry  during agitation of Cu^S in NaOCl (Figure 7). Temp 35°C, Agitation 782 rpm, NaOCl 7.45 g/1, pH 9.0, Cu2S 10 g, NaHC03 14.15 g, Na2C03 1.01 g. Time [Cu] slurry [Cu] filtered soln (mins) ppm ppm 5 112 123 10 115 123 15 106 123 20 110 122 30 108 122 45 113 123 70 109 123 90 110 123 145 106 123 -216 -Table VI: NaOCl leaching of covellite in the absence of carbonate  buffers (Figures 9,11) Temp 35°C, Agitation 784 rpm, pH 9.0, [NaOCl] 7.61 g/1, CuS 10 g. Time [Cu] [NaOCl] [NaC103] NaC103 (mins) ppm g/1 g/1 % 0 0.05 7.61 0.179 5 0.29 5.93 0.304 13.70 10 0.22 2.70 0.429 20 0.40 0.39 0.490 9.40 30 0.30 0.08 0.537 9.57 45 0.20 — 0.540 - 217 -Table VII: NaOCl leaching of chalcocite in the absence of carbonate buffers (Figure s 9,11) Temp 35°C, Agitation 780 rpm, pH 9.0, [NaOCl] 7.69 g/1, Cu2S 10 g. Time (mins) [Cu] ppm [NaOCl] g/i [NaCIO ] g/i NaCIO % 0 0.13 7 .69 0.304 5 0.19 6.98 0.340 8.73 10 6.60 0.358 9.21 20 0.20 6.30 0.429 17.70 30 5.63 0.501 21.22 45 0.25 5.03 0.633 22.81 60 4.53 0.716 20.20 80 0.21 3.91 100 3.30 0.859 23.42 120 0.19 2.78 0.913 23.03 150 2.18 0.922 23.31 180 0.23 1.56 - 218 -Table VIII: NaOCl leaching of chalcopyrite in the absence of  carbonate buffers (Figures 9.11) Temp 35°C, Agitation 785 rpm, pH 9.0, [NaOCl] 7.73 g/1, CuFeS 10 g. Time [Cu] [NaOCl] [NaC103] NaCIO (mins) ppm g/1 g/1 % 0 0.03 7.73 0.358 5 0.09 5.70 0.751 35.0 10 0.08 4.54 0.805 25.9 20 0.11 2.78 0.966 22.8 30 0.10 1.73 1.048 21.4 45 0.09 0.64 1.199 20.7 60 0.05 0.23 1.180 80 0.11 0.11 1.191 20.2 100 0.10 - 219 -Table IX: Microprobe examination of massive chalcocite before and after leaching in NaOCl. Temp 35°C, Agitation 729 rpm, pH 9.0, [NaOCl] 7.06 g/1 2 Polished Cu„S surface 3.2 cm Area S Count Before Cu Count Before Cu/S Ratio S Count After Cu Count After Cu/S Ratio 7287 81631 11.18 4873 71025 14.58 7963 74964 9.41 4429 70794 15.98 5699 71766 12.59 4463 70688 15.84 1 4698 73564 15.66 4555 71092 15.61 5136 75063 14.19 4517 71194 15.76 5449 75061 13.33 4496 71196 15.84 5528 75144 13.59 4508 70874 15.72 7283 82548 11.33 7867 79133 10.06 7094 82268 11.60 8000 77862 9.73 2 7232 82722 11.44 8072 78918 9.78 7243 82370 11.37 8030 79544 9.91 7220 82314 10.95 8459. 76290 8.88 - 220 -Table X: Microprobe examination of massive covellite, before and  after leaching in NaOCl Temp 34.9°C, Agitation 730 rpm, pH 9.0, [NaOCl] 8.1 g/1, Polished 2 CuS surface 4.5 cm . Area S Count Cu Count Cu/S S Count Cu Count Cu/S Before Before Ratio After After Ratio 14161 67660 4.78 10338 67705 6.55 14132 67131 4.75 10434 67189 6.44 1 14408 66725 5.63 10264 66974 6.53 14356 66850 4.66 9733 66380 6.82 14516 66570 4.59 9593 66683 6.95 14191 66695 4.70 29880 32991 1.10 14020 66912 4.77 28731 33094 1.15 2 13932 66553 4.78 28605 52541 1.84 14111 66321 4.69 29641 50942 1.72 13400 64360 4.80 13352 64230 4.81 After CuCl„ Wash A 12751 64821 4.71 13748 64491 4.69 3 13704 66909 4.88 00541 79590 147.11 (Green Deposit) 13826 66894 4.84 00670 83579 124.74 - 221 -Table XI: NaOCl leaching of cupric sulphide, CuS in the absence  of carbonate (Figure 10) Temp 35°C, Agitation 821 rpm, pH 9.0, [NaOCl] 7.67 g/1 CuS 10.0 g. Time [Cu] [NaOCl] (mins) ppm g/1 3 0.15 7.52 10 0.17 7.48 20 0.19 6.17 35 0.20 5.04 80 0.20 3.01 100 0.20 1.65 - 222 -Table XII: NaOCl leaching of cuprous sulphide, Cu2S> -*-n tne absence  of carbonate (Figure 10) Temp 35°C, Agitation 806 rpm, pH 9.0, [NaOCl] 7.14 g/1, Cu2S 10.0 g. Time [Cu] [NaOCl] (mins) ppm g/i 0 0.02 7.14 5 0.10 6.84 25 0.10 5.49 45 0.10 4.04 55 0.08 2.73 85 0.10 1.63 100 0.10 0.45 Time 0.55 g/1 [CO ] 1.0 g/1 [C03] 2.0 g/1 [C03] 5.0 g/1 [C03] 7.5 g/1 [C03] 10.7 g/1 [C03] (mins) [Cu] NaOCl [Cu] NaOCl [Cu] NaOCl [Cu] NaOCl [Cu] NaOCl [Cu] NaOCl ppm g/i ppm g/i ppm g/i ppm g/i ppm g/i ppm g/i 0 7.82 0.19 7.67 7.66 0.2 7.44 7.45 0.9 7.85 5 5.6 6.94 14.5 6.84 21.5 7.05 46.4 6.43 84.0 6.69 119.2 7.07 10 5.0 5.48 3.0 6.24 12.0 6.69 44.1 76.8 119.3 6.31 20 1.3 3.38 3.9 4.29 12.0 6.59 40.3 6.43 77.4 6.64 114.3 30 3.8 1.13 3.3 1.32 13.0 6.39 38.0 77.8 105.8 45 2.5 — 3.5 0.38 7.5 4.96 37.6 6.47 80.0 113.7 60 2.5 — 3.4 -- 4.5 0.52 29.0 5.53 77.2 6.62 117.2 6.32 90 1.5 0.09 20.2 0.25 75.0 6.69 116.5 120 17.3 — 67.2 6.65 116.2 6.31 140 41.8 4.71 107. 3 180 42.1 — 110.1 200 110.8 240 116.1 6.30 255 91.0 5.00 280 64.0 2.00 320 64.0 — Table XIII: NaOCl leaching of covellite with variable [C03 ]T at pH 9.0, 35°C (Figures 12,13) - 224 -Time 1.0 g/1 lco32-] 2.0 g/1 [co32-] 5.0 g/1 [co32-] (mins) [Cu] ppm [NaOCl] g/i [Cu] ppm [NaOCl] g/i [Cu] ppm [NaOCl] g/i 0 8.12 7.89 8.06 5 2.93 7.67 8.5 7.52 10 1.00 7.52 8.0 7.52 40.8 7.59 20 1.55 7.14 8.0 30 0.85 6.92 6.3 7.52 7.52 45 0.33 6.39 38.3 60 6.00 6.3 6.47 35.2 7.52 90 0.33 5.45 4.8 5.65 23.5 7.00 120 0.28 4.49 2.8 4.86 22.5 5.94 160 0.25 3.81 22.5 4.90 190 0.30 3.08 2.3 3.30 22.0 4.35 240 2.2 0.45 300 3.0 — Table XIV: NaOCl leaching of chalcocite with variable [CO., at pH 9.0 , 35°C (Figures 14,15) I - 225 -Table XV: NaOCl leaching of synthetic copper sulphides with 5 g/1 [CO.,""* ] (Figures 16,17) pH i > ).0, Temp 35 °C, Agitation 821 rpm CuS Cu2S Time (mins) [Cu] ppm [NaOCl] g/i [Cu] ppm [NaOCl] g/i 0 7.67 7.14 5 46.6 7.52 44.8 15 44.8 43.8 7.07 30 44.0 7.52 43.8 45 40.0 7.48 43.7 6.84 60 27.5 6.17 30.0 5.49 80 27.5 5.04 23.8 4.04 100 24.6 3.01 10.0 2.73 120 11.0 1.65 7.1 1.63 140 6.7 5.4 0.45 - 226 -Table XVI: NaOCl leach of covellite in the presence of 5 g/1 [CO ] and 20 g/1 [NaOCl] (Figure 18) Temp 35°C, pH 9.0, Agitation 800 rpm, CuS 10 g, NaHCO 6.5165 g, Na2C03 0.4720 g. Time [Cu] [NaOCl] (mins) ppm g/1 0 19.17 2 106 18.04 10 117 25 109 40 103 16.54 60 97 155 94 13.53 195 49 5.86 205 28 3.35 225 22 0.75 265 22 - 227 -Table XVII: Agitation of covellite in the presence of NaHCO^/Na^CO^ ± NaCl (Figure 19) Temp 35°C, Agitation 817 rpm, pH 9.0, CuS lOg, NaHC03 14.15 g, Na2CC>3 1.01 g. Time Carbonate only Carbonate and 14.0 g/1 NaCl (mins) [Cu] ppm [Cu] ppm 10 62.5 64.5 20 62.3 64.5 35 60.0 64.0 60 28.5 53.0 75 11.9 41.5 90 9.5 25.0 120 16.- 228 -Table XVIII: NaOCl leaching of covellite in the presence of  10 g/1 [C032~], and Na2MoQ4 (Figure 20) Temp 35°C, [NaOCl] 7.82 g/1, Agitation 803 rpm, pH 9.0, NaHC03 14.15 g, Na2C03 1.01 g, CuS 10 g, Na^MoO„ 2 4 1.538 g = 0.6 g/1 Mo. Time [Cu] [NaOCl] [Mo] Mo % (mins) ppm g/i g/i 0 7.82 0.66 100 5 136 6.76 0.63 95.5 20 133 60 133 6.54 0.63 95.5 120 131 6.54 180 133 6.52 0.65 98.5 255 132 6.50 270 118 4.96 0.65 98.5 290 85 1.65 305 66 0.60 0.61 92.0 360 64 — 410 20 — 0.65 98.5 440 9 - 229 -Table XIX: Sodium hypochlorite decomposition in the presence of tri-valent copper salts (Figure 21) Temp 35 copper °C, Agitation 700 -samples: 2 g. • 760 rpm, pH 9.0, Time (mins) Ill copper oxide [NaOCl] g/1 III copper carbonate [NaOCl] g/1 cupric oxide, CuO [NaOCl] g/1 0 7.74 7.67 7.82 5 5.94 6.75 7.80 10 3.99 5.40 7.37 20 1.80 3.50 7.22 30 0.38 2.05 6.58 45 — 0.20 6.35 60 — 6.00 90 5.45 - 230 -Table XX: NaOCl oxidation of molybdenite (Figure 22) Temp 35°C, pH 9.0, Agitation 780 rpm, [NaOCl] 7.7 g/1, MoS2 0.3 g. Time [Mo] Mo extracted [NaOCl] (mins) ppm % g/1 0 — — 7.70 5 168 93.3 6.83 10 168 93.3 15 169 93.8 20 170 94.4 6.82 30 170 94.4 45 170 94.4 60 169 93.8 6.82 - 231 -Table XXI: NaOCl oxidation of 'reagent grade' molybdenum  disulphide (Figures23,24) Temp 35°C, Agitation 803 rpm, pH 9.0, MoS2 1.0 g Reagent grade MoS^ 98% + MoS„ 2 Time (mins) [Mo] g/i Mo extrn. "O [NaOCl] g/i [Mo] g/i Mo extrn. % [NaOCl] g/i 0 8.54 8.72 5 0 .535 89.3 5.62 0.375 62.5 6.91 10 0 .535 89.3 0.531 88.5 6.05 20 0 .535 89.3 5.60 0.587 97.9 5.68 30 0 .536 89.3 0.587 97.9 45 0 .536 89.3 5.59 0.584 97.9 5.00 60 0 .536 89.3 5.59 0.587 97.9 4.65 - 232 -Table XXII: NaOCl oxidation of molybdenite and covellite, with  and without carbonate buffers (Figure 25) Temp 35°C, Agitation 800 rpm, pH 9.0, MoS2 0.3 g CuS 10.0 g (10 g/1 total carbonate) MoS 2, CuS + NaHCO^/Na 2C°3 MoS2, CuS + NaOH Time [Mo] Mo extr [Cu] [NaOCl] [Mo] Mo extr. [Cu] [NaOCl] (mins) ppm % ppm g/i ppm % ppm g/i 0 — — — 7.37 — — — 7.43 5 160 88.9 110 132 73.3 0.25 4.44 10 160 88.9 108 4.89 132 73.3 0.30 3.01 20 160 88.9 106 133 73.8 0.27 1.13 30 160 88.9 104 4.66 133 73.8 0.23 0.38 45 160 88.9 107 133 73.8 0.32 — 60 160 88.9 103 4.04 133 73.8 0.27 — - 233 -Table XXIII: NaOCl oxidation of molybdenite and chalcocite, with and without carbonate buffers (Figur e 26) Temp 35°C, Agitation 785 rpm pH 9 .0, MoS2 0. 3 g, Cu2S 10.0 g (10 g/1 total carbonate) MoS2, Cu2S + NaHC03/Na 2C°3 MoS2, Cu2S + NaOH Time [Mo] Mo extr. [Cu] [NaOCl] [Mo] Mo extr. [Cu] [NaOCl] (mins) ppm % ppm g/i ppm % ppm g/i 0 — — 7.89 — — — 7.35 5 159 88.5 96.0 139 77.2 0.13 6.50 10 169 89.2 93.0 4.61 139 77.2 0.13 6.35 20 160 89.0 94.0 139 77.2 0.25 6.02 30 162 89.9 93.0 4.53 139 77.2 0.25 5.57 45 162 89.9 90.0 139 77.2 0.25 4.95 60 162 90.0 89.0 4.44 139 77.2 0.23 3.60 - 234 -Table XXIV: NaOCl oxidation of molybdenite and chalcopyrite with  and without carbonate buffers (Figure 27) Temp 34.9°C, Agitation 792 rpm, pH 9.0, MoS2 0.3 g, CuFeS 10.0 g (10 g/1 total carbonate). MoS2, CuFeS2 + NaHC03/Na 2C°3 MoS2 CuFeS2 + NaOH Time [Mo] Mo extr. [Cu] [NaOCl] [Mo] Mo extr. [Cu] [NaOCl] (mins) ppm % ppm g/i ppm Q. "6 ppm g/i 0 — — 7.50 — — — 7.13 5 154 86.1 106 5.32 143 79.4 0.53 4.51 10 156 86.6 100 144 80.0 0.41 4.17 20 156 86.6 98 5.30 142 78.8 0.32 2.82 30 158 87.8 98 144 80.0 0.32 1.97 45 158 87.8 94 5.25 144 80.0 0.27 0.94 60 158 87.8 99 5.20 144 80.0 0.11 — - 235 -Table XXV: NaOCl oxidation of molybdenum disulphide and chalcopyrite  with and without carbonate buffers Temp 35°C, Agitation 780 rpm, pH 9.0, CuFeS2 10 g, MoS0 5 g (reagent grade) MoS^ only MoS2 , CuFeS2 + NaOH MoS2, CuFeS2 + [co32-] Time (mins) [Mo] g/i Mo extr. 0. ~o [NaOCl] g/i [Mo] g/i Mo extr. % [NaOCl] g/i [Mo] g/i Mo extr. % [NaOCl] g/i 0 24.55 — — 20.32 — — 20.21 5 2.67 89.0 11.32 2.54 84.7 8.32 2.64 87.5 7.37 10 2.66 88.7 2.54 84.7 4.62 2.64 87.5 20 2.68 89.3 11.30 2.54 84.7 2.83 2.62 87.5 7.28 30 2.67 89.0 2.54 84.7 1.56 2.62 87.3 45 2.67 89.0 11.27 2.54 84.7 0.34 2.62 87.3 7.26 60 2.67 89.0 11.25 2.54 84.7 — 2.64 87.5 - 236 -Table XXVI: NaOCl oxidation of molybdenite in the presence of  copper sulphate solution Temp 35.1°C, NaOCl 7.5 g/1, Agitation 792 rpm, pH 9.0, MoS2 0.3 g, CuS04 0.4 g/1 = 0.1 g/1 Cu Time Cu added [Cu] in soln [Mo] Mo extrd. (mins) ppm ppm ppm % 5 — — 168 93.3 10 12.8 0.05 142 79.1 20 25.6 0.10 136 75.3 30 38.4 0.08 135 75.2 45 67.5 0.08 134 74.8 60 100 0.11 132 73.5 - 237 -Table XXVII: NaOCl oxidation of covellite in the presence of sodium  molybdate solution Temp 35.1°C, Agitation 821 rpm, pH 9.1, CuS 10.0 g, Na2Mo04 2.15g= 0.1 g Mo, [NaOCl] 7.63 g/1 Time Mo added [Mo] in soln Mo in soln [Cu] (mins) ppm ppm % ppm 10 11.04 9.79 89.0 0.23 20 25.55 22.48 88.0 0.31 30 39.23 33.38 85.1 0.31 45 51.80 44.55 86.0 0.32 60 75.45 62.62 83.0 0.33 90 100.00 81.00 81.0 0.33 - 238 -Table XXVIII: NaOCl oxidation of molybdenite at pH 5.5 (Figure 28) Temp 35°C, Agitation 880 rpm, pH 5.4, [NaOCl] 6.92 g/1, Mo S2 1.0 g Time [Mo] Mo extracted [NaOCl] (mins) g/1 % g/1 0 — — 6.92 5 0.550 91.6 10 0.563 93.8 2.71 20 0.548 91.3 30 0.563 93.8 2.53 45 0.564 94.0 60 0.563 93.8 2.21 - 239 -Table XXIX: NaOCl oxidation of molybdenite and chalcopyrite at pH 5.5 (Figure 29) Temp 35°C, Agitation 803 rpm, pH 5.55, [NaOCl] 6.54 g/i CuFeS2 10 g , MoS2 0. 3 g. MoS2 , CuFeS2 + NaOH MoS2, CuFeS2 + 10 g NaHC03 Time [Mo] ppm Mo extr. % [Cu] ppm [Mo] ppm Mo extr. % [Cu] ppm 5 142 78.8 — 146 82.2 — (CuFeS2 added) 10 4.1 2.27 12.9 0.2 0.10 37.5 20 4.5 2.50 8.7 0.3 0.16 33.7 30 4.6 2.56 8.1 1.7 0.94 27.2 45 4.5 2.50 8.1 1.6 0.88 25.0 60 4.6 2.56 8.4 1.2 0.66 25.0 - 240 -Table XXX: NaOCl oxidation of molybdenite and covellite at pH 5.5  (no carbonate)(Figure 30 a) Temp 35°C, Agitation 802 rpm, pH 5.45, [NaOCl] 6.17 g/1 CuS 10.0 g, MoS2 0.3 g. Time [Mo] Mo extracted [Cu] [NaOCl] (mins) ppm % ppm g/1 5 141 78.3 — 5.55 (CuS added) 10 5.8 3.2 18.0 1.54 20 6.2 3.4 15.2 0.03 30 6.6 3.6 13.5 45 6.6 3.6 12.9 60 6.7 3.7 12.8 90 6.2 3.4 12.4 - 241 -Table XXXI: NaOCl oxidation of molybdenite and covellite at pH 5.5, with sodium bicarbonate (Figure 30 b) Temp 34.9°C, Agitation 800 rpm, pH 5.50, [NaOCl] 6.92 g/1 CuS 10 g, Mo S 0.3 g, NaHCO 10.0 g. Time [Mo] Mo extracted [Cu] [NaOCl] (mins) ppm % ppm g/1 5 153 85.0 — 6.92 (CuS added) 10 2.4 1.33 36.0 0.15 20 2.4 1.33 35.0 30 2.5 1.38 32.0 45 3.0 1.66 30.0 60 2.4 1.33 30.0 - 242 -Table XXXII: NaOCl leaching of chalcopyrite at pH 5.5 with and without bicarbonate in the system Temp 35°C, Agitation 805 rpm, pH 5.40, [NaOCl] 7.2 g/1, CuFeS2 10 g, NaHC03 9.28 g. CuFeS2 only CuFeS2 + HCO3-Time [Cu] [NaOCl] [Cu] [NaOCl] (mins) ppm g/i ppm g/i 5 36.5 5.7 48.8 5.64 10 34.2 4.6 46.2 4.52 20 30.6 3.4 45.3 3.30 30 28.5 1.7 44.8 3.30 45 25.0 0.5 43.6 3.28 60 24.1 — 42.5 3.25 90 24.3 — - 243 -Table XXXIII: NaOCl leaching of molybdenite and chalcocite at pH 6.5,  with bicarbonate in the system (Figure 31 a) Temp 35°C, Agitation 880 rpm, pH 6.5, [NaOCl] 6.92 g/1 Cu2S 10 g, MoS2 1.0 g, NaHC03 10 g. Time [Mo] Mo extracted [Cu] [NaOCl] (mins) g/1 % ppm g/1 0 — — 6.92 10 0.531 88.3 34.7 2.71 20 0.531 88.3 29.5 2.11 30 0.529 88.1 25.0 2.11 45 0.529 88.1 25.0 2.10 60 0.529 88.1 25.0 2.06 - 244 -Table XXXIV: NaOCl leaching of molybdenite and chalcocite at  pH 6.5, with no carbonate present (Figure 31 b) Temp 35°C, Agitation 880 rpm, pH 6.5, [NaOCl]7.52 g/1, Mo S2 l.Og, Cu2S 10 g Time [Mo] Mo extracted [Cu] [NaOCl] (mins) g/1 % ppm g/1 10 0.515 85.8 32.6 2.65 20 0.472 78.6 30.0 2.34 30 0.465 77.5 29.5 1.97 45 0.387 64.5 24.3 1.03 60 0.280 46.6 18.7 0.20 90 0.253 42.2 16.2 - 245 -Table XXXV: NaOCl leaching of molybdenite and cupric sulphide at  pH 6.0 in the presence of carbonate (Figure 32) Temp 35°C, Agitation 803 rpm, pH 6.5, [NaOCl] 7.04 g/1, CuS 10.0 g, MoS2 1.0 g, NaHC03 10.0 g Time [Mo] Mo extracted [Cu] [NaOCl] (mins) g/1 % ppm g/1 10 0.539 89.8 39.5 5.04 20 0.531 88.5 24.5 2.86 30 0.549 91.5 25.5 45 0.560 93.3 25.5 2.63 60 0.560 93.3 24.9 100 0.560 93.3 25.0 2.56 - 246 -Table XXXVI: NaOCl leaching of molybdenite and cupric sulphide at  pH 6.0, with no carbonate present Temp 35°C, Agitation 857 rpm, pH 6.5, [NaOCl] 6.39 g/1, CuS 10.0 g, MoS2 1.0 g. Time [Mo] Mo extracted [Cu] [NaOCl] (mins) g/1 % ppm g/1 10 0.473 78.8 32.7 3.39 20 0.455 75.8 37.2 2.63 30 0.441 73.5 17.2 2.18 45 0.288 48.0 11.8 2.03 60 0.093 15.5 4.5 1.28 90 0.061 10.2 0.3 0.38 - 247 -Table XXXVII: NaOCl leaching of molybdenite and chalcopyrite at  pH 7.0 with bicarbonate present (Figure 33 a) Temp 35°C, Agitation 785 rpm, pH 7.0, [NaOCl] 6.5 g/1, CuFeS2 10 g, MoS2 1.1 g, NaHC03 10 g Time [Mo] Mo extracted [Cu] [NaOCl] (mins) g/1 % ppm g/1 5 0.62 93.9 56.6 2.63 10 0.64 96.9 58.2 0.23 20 0.65 98.5 52.5 — 30 0.64 96.9 48.5 — 50 0.65 98.5 42.5 — 70 0.65 98.5 41.3 - 248 -Table XXXVIII: NaOCl leaching of molybdenite and chalcopyrite at pH  7.0, no carbonate (Figure 33 b) Temp 35°C, Agitation 785 rpm, pH 7.0, [NaOCl] 6.54 g/1, CuFeS2 10 g, MoS2 1.0 g. Time [Mo] Mo extracted [Cu] [NaOCl] (mins) g'/l % ppm g/1 5 0.173 28.75 12.5 10 0.180 30.00 8.2 3.01 20 0.185 30.83 3.0 2.11 30 0.190 31.66 1.9 60 0.195 32.50 1.3 1.96 120 0.195 32.50 0.7 205 0.203 33.75 0.5 1.94 - 249 -Table IXL: NaOCl decomposition in the presence of copper molybdate  at pH 5.5 (Figure 35) Temp 35°C, Agitation 792 rpm, pH 5.40, [NaOCl] 6.84 g/1, Cu Mo04 2 g. Time [NaOCl] [NaCIO ] (mins) g/1 g/1 0 6.84 — 10 6.39 0.109 20 5.30 0.785 30 4.14 1.430 45 2.31 1.480 - 250 -Table XL: Determination of the solubility of copper molybdate at  pH 5.0 (Figure 34) pH 5.0 Na2Mo04-2H20 4.13 g = 0.096 g/1 Mo CuSO4-5H20 6.22 g 5 0.064 g/1 Cu Mo added Mo in soln Cu in soln Mo pptd Cu pptd (Moles)xlO2 (Moles)xlO2 (Moles)xlO2 (Moles)xlO (Moles)xlO 0.2 0.188 0.889 0.12 0.24 0.4 0.350 0.884 0.50 0.29 0.6 0.604 0.883 -- 0.30 0.8 0.792 0.860 0.08 0.53 1.0 0.881 0.795 1.20 1.18 1.2 0.880 0.595 3.20 3.18 1.4 0.881 0.385 5.20 5.28 - 251 -Table XLI: Sodium hypochlorite leaching of molybdenite and covellite  pH 10.0, with carbonate (Figure 37 a) Temp 35.1°C, Agitation 789 rpm, pH 10.0, [NaOCl] 7.54 g/1, MoS2 0.3 g, CuS 10.0 g, Na2C03 5.37 g, NaHC03 7.24 g. Time [Mo] Mo extrn. [NaOCl] (mins) ppm % g/1 5 148 82.2 6.63 10 150 83.3 20 154 85.5 6.43 30 158 87.8 45 160 88.9 6.25 60 164 91.1 6.23 - 252 -Table XLII: Sodium hypochlorite leaching of molybdenite . and covellite at pH 10.0 with no carbonate (Figure 37 b) Temp 35°C, Agitation 790 rpm, pH 10.0, [NaOCl] 7.61 g/1, CuS 10 g, MoS 0.3 g. Time [Mo] Mo extrn. [Cu] [NaOCl] (mins) ppm % ppm g/i 5 142 78.9 0.5 6.01 10 142 78.9 0.7 3.10 20 142 78.9 0.5 0.60 30 141 78.3 0.5 0.05 45 143 79.4 0.4 — 60 143 79.4 0.3 - 253 -Table XLIII: NaOCl leaching of molybdenum disulphide and cuprous  sulphide at pH 9.0, with no carbonate (Figure 36) Temp 35°C, Agitation 808 rpm, pH 9.0, [NaOCl] 33.32 g/1, MoS2 5 g, Cu2S 10 g. Time [Mo] Mo extrn. [NaOCl] (mins) g/1 % g/1 5 2.66 88.67 10 2.65 88.30 19.36 20 2.66 88.67 30 2.66 89.0 10.53 45 2.68 89.0 70 2.68 89.0 4.14 - 254 -Table XLIV: NaOCl leaching of molybdenum disulphide and cuprous sulphide at pH 9.0, with 10 g/1 carbonate (Figure 36) Temp 35°C, Agitation 803 rpm, pH 9.0, [NaOCl] 33.08 g/1, MoS2 5 g, Cu2S 10 g. Time [Mo] Mo extrn. [NaOCl] (mins) g/1 % g/1 5 2.68 89.01 20.14 10 2.68 89.01 20 2.70 90.00 19.36 30 2.72 90.67 45 2.72 90.67 16.16 60 2.72 90.67 13.90 - 255 -Table XLV: NaOCl leaching of chalcopyrite, chalcocite and covellite  at pH 9.0 with no carbonate (Figure 38) Temp 35°C, Agitation 799 rpm, pH 9.0, [NaOCl] 7.74 g/1, Cu samples 10 g-CuFeS2 CU2 S CuS Time [Ca] [Cu] [Ca] [Cu] [Ca] [Cu] (mins) ppm ppm ppm ppm ppm ppm 5 65.0 0.68 3.48 0.07 8.0 0.13 10 65.3 0.75 3.55 0.12 9.2 0.15 20 66.2 0.52 5.35 0.15 9.5 0.09 30 66.8 0.33 4.03 0.08 8.1 0.05 40 66.8 0.32 4.10 0.04 8.5 0.12 50 66.6 0.30 4.18 0.07 8.8 0.17 60 66.8 0.52 4.20 0.13 9.8 0.13 - 256 -Table XLVI: NaOCl leaching of MoS^ and Cu^S in the presence of 0.1 g/1  calcium as CaCl^ solution Temp 35°C, Agitation 821 rpm, pH 9.0, [NaOCl] 8.0 g/1, MoS2 1 g (reagent grade), Cu2S 10 g, CaCl2 0.3703 g = 0.1 g Ca. Time [Ca] [Mo] Mo extr. Mo pptd. Ca pptd. (mins) ppm g/1 % (x 10 M) (x 104M) 0 97.25 — — — — 10 92.50 0.522 87.2 1.56 1.24 20 91.25 0.522 87.2 1.56 1.50 30 92.50 0.525 87.5 1.25 1.24 45 92.50 0.525 87.5 1.25 1.24 60 92.50 0.525 87.5 1.25 1.24 - 257 -Table XLVII: NaOCl leaching of molybdenite and chalcocite in the  presence of CaCl^, and Na^CO^/NaHCO^ Temp 34.9°C, Agitation 820 rpm, pH 9.0, [NaOCl] 7.89 g/1, MoS2 1.0 g, CU2S 10 g, CaCl2 0.3703 g, NaHCt>3 14 .15 g, Na2C03 1.01 g-Time [Mo] Mo extr. [Ca] [mins) g/i % ppm 5 0.487 81.5 2.15 10 0.538 89.2 1.90 30 0.537 89.0 1.90 60 0.538 89.2 1.90 100 0.540 89.9 1.90 135 0.542 90.0 1.90 - 258 -Table XLVIII: NaOCl leaching of molybdenite and calcite with no  carbonate (Figure 40) Temp 34.9°C, Agitation 805 rpm, pH 9.0, [NaOCl] 9.5 g/1, MoS2 1 g, CaC03 10 g. Time [Mo] Mo extrn. [Ca] (mins) g/1 % ppm 10 0.495 82.5 9.00 20 0.495 82.5 10.12 30 0.495 82.5 10.13 60 0.500 83.3 10.25 120 0.505 83.4 10.36 - 259 -Table IL: NaOCl leaching of powdered calcite, CaCO^, at pH 9.0  (Figure 39) Temp 35°C, Agitation 780 rpm, pH 9.0, [NaOCl] 6.69 g/1 CaCO., 10 g = 4 g Ca Time [Ca] Ca extr. (mins) ppm % 0 0.25 5 13.35 0.33 10 13.40 0.34 20 14.50 0.36 30 15.25 0.38 50 15.50 0.39 70 16.00 0.40 - 260 -.: NaOCl leaching of plaster of paris, CaSO^*%Hn0 at pH 9.0 (Figure 39) 4 2 Time [Ca] Ca extrn. [NaOCl] (mins) g/i Q, "O 5 1.81 60.2 7.89 10 1.87 62.1 20 1.89 63.0 7.89 30 1.89 63.0 50 1.90 64.5 7.94 70 1.90 64.5 Temp 35°C, Agitation 805 rpm, pH 9.0, [NaOCl] 8.01 g/1, CaSO^H^O 10 g LI: i) CaSo4-J5H20 Time 0 g/1 NaCl 12 g/1 NaCl 28 g/1 NaCl 40 g/1 NaCl 58.5 g/1 NaCl 86.8 g/1 NaCl 117 g/1 NaCl (mins) [Ca] g/1 [Ca] g/1 [Ca] g/1 [Ca] g/1 [Ca] g/1 [Ca] g/1 [Ca] g/1 10 0.435 0.960 1.394 1.860 2.06 20 1.548 1.684 1.864 2.012 2.08 30 0.659 1.314 40 1.688 2.012 2.10 50 1.528 1.868 60 0.712 1.326 1.686 2.10 90 0.712 1.326 1.528 1.690 1.860 2.012 LII: ii) CaC03 Time (mins) 0 g/1 NaCl [Ca] ppm 13 g/1 NaCl [Ca] ppm 20 g/1 NaCl [Ca] ppm 35 g/1 NaCl [Ca] ppm 40 g/1 NaCl [Ca] ppm 80 g/1 NaCl [Ca] ppm 10 6.5 12.25 16.80 21.50 43.20 35.75 20 14.50 18.30 33.50 55.05 30 15.5 15.25 34.75 48.75 56.80 45 16.0 15.50 49.00 58.60 60 16.5 16.00 35.0 49.00 60.30 55.00 Tables LI, LII: Effect of increasing chloride concentration on calcium dissolution from i) calcium sulphate ii) calcium carbonate (Figure 41-44) pH 9.0, Temp 35°C, Agitation 780 - 860 rpm - 262 -Table LIII: Effect of carbonate on calcium content in hypochlorite  solutions (Figure 45) pH 9.0, Temp 35°C, NaOCl 8.0 g/1 [C032 ]T added [CO^-] [Ca] [Ca] [Ca] pptd. g/1 (moles) ppm (moles) (moles) 0 — 1.750 0.0438 2.838 0.047 0.330 0.0082 0.0356 5.00 0.083 0.110 0.0027 0.0411 10.00 0.167 0.020 3.8xl0~4 0.0434 25.00 0.417 0.005 9.3xl0~5 0.0437 - 263 -Table LIV: NaOCl leaching of covellite and molybdenite with 5 g/1  carbonate (Figure 46) Temp 35°C, Agitation 821 rpm, pH 9.0, [NaOCl] 8.77 g/1, CuS 10 g, MoS2 1 g, [C032"]T 5 g/1 Time [NaOCl] [Cu] [Mo] Mo extrn. (mins) g/1 ppm g/1 %' -10 2.71 14.3 0.510 85.0 30 2.71 14.0 0.510 85.0 45 2.11 10.0 0.516 86.0 65 0.75 6.5 0.510 85.0 85 0.23 4.0 0.546 91.0 120 — 5.0 0.540 90.0 - 264 -Table LV: NaOCl leaching of molybdenite and covellite in the presence  of 10 g/1 carbonate (Figure 46) Temp 35°C, Agitation 803 rpm, pH 9.0, [NaOCl] 7.5 g/1 2-MoS2 1.0 g, CuS 10 g, [C03 ] 10.7 g/1. Time [Cu] [Mo] Mo extrn. [NaOCl] (mins) ppm g/1 % g/1 5 109 0.485 80.8 2.33 10 89 0.490 81.7 2.20 40 99 0.535 89.1 2.18 60 98 0.542 90.3 2.18 120 98 0.542 90.3 2.02 150 79 0.543 90.5 0.25 180 63 0.542 90.5 - 265 -Table LVI: NaOCl leaching of covellite and sodium molybdate with  5 g/1 carbonate (Figure 48) Temp 35°C, Agitation 785 rpm, pH 9.0, CuS 10 g, Na2Mo04 0.5 g = 0.35 g [Mo], NaHC03 0.46 g, Na2C03 6.52 g, [NaOCl] 7.5 g/1. Time [Cu] [Mo] Mo in soln [NaOCl] (mins) ppm g/1 % g/1 10 39.0 0.341 97.4 6.32 30 37.0 0.323 92.3 5.68 60 27.5 0.315 90.0 1.95 70 19.5 0.305 87.1 1.65 85 8.4 0.295 84.3 0.60 - 266 -Table LVII: NaOCl leach of molybdenum disulphide and chalcocite with  10 g/1 carbonate (Figure 47) Temp 35°C, Agitation 803 rpm, pH 9.0, MoS 1.0 g, 2-Cu2S 10 g, [C03 ]T 10.68 g/1. Time [Cu] [Mo] Mo extrn. [NaOCl] (mins) ppm g/1 % g/1 0 — — — 7.06 10 96.0 0.403 67.16 3.61 30 93.0 0.555 92.50 2.71 60 90.0 0.570 95.00 2.48 120 89.0 0.570 95.00 2.36 150 66.0 0.570 95.00 1.55 180 64.0 0.570 95.00 0.60 240 16.0 0.570 95.00 - 267 -Table LVIII: NaOCl oxidation of molybdenum disulphide and chalcopyrite with 1 g/1 Si0o2~ added to the system, pH 9.0 (Figure 49T Temp 35°C, Agitation 803 rpm, [NaOCl] 8.4 g/1, MoS2 1 g, CuFeS2 10 g, Na2Si03 3.29 g, pH 9.0. Time [Cu] [Mo] Mo extrn. [NaOCl] (mins) ppm g/1 % g/1 10 0.25 0.579 96.5 5.04 20 0.13 0.579 96.5 30 0.17 0.583 97.2 5.04 60 0.22 0.590 98.4 90 0.15 0.593 98.9 5.04 - 268 -APPENDIX B B.l Calculation of surface areas for ground  mineral samples The specific surface (surface per unit weight)of a screened material is given by the expression: n = x2 k-l k-l S*l = 6hk . *2 " XI *2 A (k-l) k k AVxx X2 " Xl where x^ = unit crystal size x^ = size of -mesh particle h = shape factor k = constant A = specific gravity For copper sulphide minerals ground to pass a -200 mesh screen: x^ = 5 x 10 cm x2 = 0.0074 cm h =1.75 (i.e. assume irregularly shaped particles have 1.75 x the surface area of a cube for same screen size) k = 1.02 A = 5.6 - 5.8 (Cu2S) = 4.5 - 4.6 (CuS) Cu^S 6 x 1.75 x 1.02 - 0.0074(0-°2) - (5 x 10~8)0"02 2 -1 UT021 ~ , -8,1.02 cm g 2 5.65 x 0.02 0.0074v ' - (5 x 10 ) = 94.779 x 28.632 cm2g-1 2 -1 = 2713.71 cm g - 269 -6 x 1.75 x 1.02 • 0.0074(°-02) - (5 x 10 8)°-02 2 -1 CUS ~ 4.60 x 0.02 0.0074(1-°2) - (5 x IO"8)1'02 " ? 116.41 x 28.632 cm2g-1 = 3333.05 cmV1 2 10 g of chalcocite will have a total surface area of 27,137 cm and 10 g of covellite will have a total surface area of 33,330 2 cm . 90 (Error of calculation can be 30 - 50%) B.2 Calculation of available copper for precipitation  on mineral surface after copper111 formation Radius of copper molecule = copper precipitated = Avagadro's number = Number of Cu molecules Total area of Cu This compares with a surface area of 2.71 x 10' cm" for a 10 g 4 2 sample of chalcocite, and 3.33 x 10 cm for a 10 g sample of covellite. The small amount of copper precipitated as Cu111 is thus sufficient to form a mono-molecular layer on the mineral surface. o -a 1A = 10 cm 60 ppm =0.06 g/1 23 6 x 10 0.06 , ln23 = ,- x 6 x 10 63. D = 5.660 x 1021 = 5.669 x 1021 x 10~8 = 5.669 x 1015 cm2. B.3 Hypochlorite redox equilibria at pH 9.0 1) HC10 + H + e = JgCl^aq) + H20 E = 1.59 + 0.0591 {log (HC10) - Jglog (Cl2) - pH} - 270 -2) *SCl2(aq) + e Cl" E = 1.40 + 0.0591 {%log (Cl2) - log (Cl )} (1) + (2) 3) HC10 + H+ + 2e = Cl~ + H20 E = 1.49 + 0,°^91 {log (HC10) - log (Cl ) - pH} 4) HC10 = HCIO H+ + CIO" E° = 1.49 (2) - (4) + (1) : 5) CIO + 2H+ + 2e = Cl" + H20 E = 1.738 + °'0^91 (log (CIO ) - log (Cl ) - 2pH> at pH 9.0 E = 1.738 + {0-05291) • (-18) = 1.738 - 0.532 E = 1.206 V - 271 -APPENDIX C Reactions and equilibria pertaining to potential- pH diagrams for various metal molybdate systems shown in Figures 52-58 C.1 Cu-H^O-MoO^ system C.l.l Substances considered. Oxidation Number Species AG° (kcal) Source of Data 0 Cu° — +1 Cu20 -34.98 73 +1 Cu+ +12.00 73 +2 CuO -30.40 73 +2 Cu(OH)2 -85.30 73 +2 CuMo04 -202.70 91 +2 ^ 2+ Cu +15.53 73 +2 2-Cu02 -43.50 73 +2 Cu02H -61.42 73 +3 CU2°3 -41.50 74 +3 CuO - -26.87 74 +3 r, 3 + Cu +72.50 74 (+4 Mo03 -120.00) 73 (+6 HMoO„~ 4 -161.95) 73 (+6 MoO 2~ 4 -213.60) 73 +6 CuMoO. 4 -205.42 35 Activity of copper containing species in solutions has been taken as 10 ^ M, and that of molybdenum containing species as 10 1 M. 272 C.1.2 Two dissolved substances + 2+ 1) Cu ->Cu + e 2+ E = 0.153 + 0.0591 log ^ (Cu+) E = 0.153 2) Cu+ + 2H20 >-HCu02 + 3H+ + e" E = 1.733 - 0.1773 pH + 0.0591 log (HCu°2-i-(Cu ) E = 1.733 - 0.1773 pH + 2- „ + 3) Cu + 2H20 >Cu°2 + + e E = 2.510 - 0.2364 pH + 0.0591 log -^^-2—-(Cu ) E = 2.510 - 0.2364 pH 4) Cu2+ + 2H20 >flCu02 + 3H+ log (HCU°2-^- = -2.72 + 3pH; pH = 8.91 (Cu +) 2- + 5) HCu02 > Cu°2 + H 2-log -^^2 ]- = -39.88 + 4 pH; pH = 13.15 (HCu02~) ^,2+ 3+ 6) Cu • Cu + e 3+ E = 2.475 + 0.0591 log U ; E = 2.475 (Cu ) 2-7) Cu02 Cu02 + e (CuO ) E = 0.721 + 0.0591 log ; E = 0.721 (Cii02 ) - 273 -8) HCu02 >Cu02 + 6 E = 1.498 - 0.0591 pH + 0.0591 log E = 1.498 - 0.0591 pH (Cu02 ) (HCu02~) 9) Cu2+ + 2H20 =*CuO + 4H+ + e" (CuO ) E = 3.078 - 0.2364 pH + 0.0591 log (Cu2+) E = 3.078 - 0.2364 pH 3+ - + 10) Cu + 2H.0 >-Cu0o + 4H (Cu02 ) log * — = -10.28 + 4 pH; pH = 2.57 (Cu3+) C.1.3 One solid substance and one dissolved substance + 2+ 11) Cu20 + 2H >-2Cu + H20 + 2e 2+ E = 0.203 + 0.0591 pH + 0.0591 log (Cu ) E = -0.152 + 0.0591 pH 12) Cu20 + 3H20 + 4H+ + 2e »- 2HCu02~ + 4H+ + 2e~ E = 1.783 - 0.1182 pH + 0.0591 log (HCu02~) E = 1.428 - 0.118 pH 2+ 13) Cu ->Cu + 2e 2+ E = 0.337 + 0.0295 log (Cu ); E = 0.337 274 .2+ + 14) a) Cu + H20 CuO + 2H 2+ log (Cu ) = 7.89 - 2 pH; pH = 6.93 2+ + b) Cu + 2H20 > Cu(OH)2 + 2H 2+ log (Cu ) = 9.071 - 2 pH; pH = 7.54 15) a) CuO + H20 y HCu02~ + H+ log (HCu02~) = -18.83 + pH; pH = 12.83 b) Cu(OH)2 > Cu02 + 2H+ log (Cu02 ) = -30.80 + 2 pH; pH = 12.4 2- + 16) CuO + H20 Cu02 + 2H log (CuO 2 ) = -31.98 + pH; pH = 12.99 2+ + -17) Cu + 3H20 >• Cu203 + 6H + 2e E = 1.783 - 0.1182 pH + 0.0591 log (HCu02 ) E = 1.428 - 0.118 pH 18) a) CuO + H20 • Cu02_ + 2H+ + 2e_ E = 2.609 - 0.1182 pH + 0.0591 log (Cu02~) E = 2.254 - 0.118 pH b) Cu(0H)2 > Cu02~ + 2H+ + 2e" E = 1,267 - 0.118 pH - 0'°291 log (Cu02~) E = 1.149 - 0.118 pH - 275 -'3+ + 19) Cu + 3H20 y Cu203 + 6H 3+ log (Cu ) = -6.09 +3 pH; pH = -0.3 20) Cu2°3 + H2° 2Cu02~ + 2H+ log (Cu02~) = -16.31 + pH; pH = 10.31 2- 2- + 21) Cu02 + Mo04 + 4H y CuMo04 + 2H20 2- 2-4 pH = 67.18 - 0.0591 log ((Cu02 ) • (Mo04 )'); pH = 10.815 2- + 22) Cu02H + Mo04 + 3H y CuMo04 + 2H20 2-3 pH = 49.06 - 0.0591 log ((CuC^H )•(Mo04 )); pH = 9.99 3+ 2-23) CuMo04 y Cu + Mo04 + e E = 3.026 + 0.0591 log ((Cu3+)•(MoC^2-)) E = 2.671 2- + -24) CuMoO„ + 2H O y CuO„ + MoO„ + 4H + e 4 2 2 4 2-E = 3.626 - 0.236 pH + 0.0591 log ((Cu02 )•(MoO ) E = 3.271 - 0.236 pH C.1.4 Two solid substances 25) 2Cu + H20 y Cu20 + 2H+ + 2e~ 0.471 - 0.0591 pH 26) a) Cu + H20 y CuO + 2H+ + 2e~ E = 0.570 - 0.0591 pH - 276 -b) Cu + 2H20 y Cu(OH)2 + 2H+ + 2e~ E = 0.609 - 0.0591 pH 27) a) Cu20 + H20 y 2CuO + 2H+ + 2e~ E = 0.669 - 0.0591 pH b) Cu20 + 3H20 y 2Cu(OH)2 + 2H+ + 2e~ E = 0.741 - 0.0591 pH 2- + 28) Cu20 + 2Mo04 + 2H • 2CuMo04 + H20 + 2e E = 0.35 + 0.0591 pH - 0.029 log (MoO 2~) E = 0.379 + 0.0591 pH 2- + 29) a) CuO + Mo04 + 2H > CuMo04 + H20 2 pH = 17.23 - 0.0591 log (MoO 2~) pH = 8.62 2- + b) Cu(OH)2 + Mo04 + 2H CuMo04 + 2H20 2 pH = 18.46 - 0.0591 log (MoO^2-) pH = 9.25 30) a) CuO + H2Q > Cu2°3 + 2H+ + 2e" E = 1.648 - 0.0591 pH b) 2Cu(OH)2 y Cu203 + H20 + 2H* + 2e~ E = 1.572 - 0.0591 pH - 277 -31) 2CuMo04 + 3H20 2- + Cu„0^ + 2MO0,, + 6H + 2e 2 3 4 E = 2.658 - 0.0197 pH - 0.0197 log (MoC>4 ) 2.677 - 0.0197 pH 32) Cu + MoO 2- 4* CuMoO„ + 2e 4 E = 0.0589 - 0.0591 log (Mo04 ) E = 0.118 C. 2 Ca-H20-Mo04 system C.2.1 Substances considered Oxidation Number Species AG (kcal) Source of Data -2 0 +2 +2 +2 +2 +4 ( +4 ( +6 ( +6 +6 CaH 2 Ca CaO Ca(OH) CaMoO„ n 2+ Ca Ca02 Mo02 MoO„ 3 HMoO ~ 4 MoO.2 ~ -35.80 -144.40 -157.64 -345.80 -132.18 -143.00 -120.00 ) -161.95 ) -213.60 ) -205.42 73 73 73 91 73 73 73 73 73 35 Activity of calcium containing species in solution has been taken as 10 ^ M and the activity of molybdenum containing species as IO-1 M. 278 -C.2.2 One solid substance and one dissolved substance 2+ + 1) a) Ca + H20 y CaO + 2H log (Ca2+) = 32.63 - 2pH; pH = 19.315 b) Ca2+ + H20 • Ca(OH)2 + 2H+ 2+ log (Ca ) = 22.91 - 2 pH; pH = 14.45 2) CaH„ > Ca2+ + 2H+ + 4e~ 2 E = -1.045 - 0.0295 pH - 0.0148 log (Ca2+) E =. -1.134 - 0.0295 pH 2+ 3) Ca y Ca + 2e 2+ E = -2.866 - 0.0295 log (Ca ); E = -3.043 C.2.3 Two solid substances 4) CaH2 y Ca + 2H+ + 2e" E = 0.776 - 0.0591 pH 5) a) CaH2 + H20 y CaO + 4H+ + 4e~ E = -0.563 - 0.0591 pH b) CaH2 + 2H20 y Ca(OH)2 + 4H+ + 4e -0.706 - 0.0591 pH 6) a) Ca + H20 y CaO + 2H+ + 2e E = -1.902 - 0.0591 pH - 279 -b)r Ca + 2H20 >-Ca(OH)2 + 2H+ + 2e~ E = -2.189 -. 0.0591 pH 2-7) Ca + MoO. > CaMoO„ + 2e 4 4 0.0591 2-E = -3.044 - - log (Mo04 ) ; E = -3.0.14 2- + 8) a) CaO + Mo04 + 2H >• CaMo04 + H20 2 pH = 38.36 + log (Mo042~) pH = 19.15 2- + b) Ca(OH)2 + Mo04 + 2H • CaMo04 + 2H20 2-2 pH = 28.99 + log (Mo04 ) pH = 14.00 2- + 9) CaH„ + MoO. y CaMoO„ + 2H + 4e 2 4 4 E = -1.134 - 0.0295 pH - 0.0148 log (Mo04 E = -1.192 - 0.0295 pH 2- + 10) CaMoO. + 2H_0 y CaO„ + MoO„ + 4H 4 2 2 4 2-4 pH = 81.09 - log (Mo04 ) pH = 20.52 11) a) CaO + H20 y Ca°2 + 2H+ + 2e_ E = 1.260 - 0.0591 pH b) Ca(OH)2 + H20 y Ca02 + 2H+ + 2e~ E = 1.547 - 0.0591 pH 2-. - 280 -C.3 Pb-H^O-MoO^ system C.3.1 Substances considered Oxidation „ . AG° Source Species , Number (kcal) of Data 0 Pb — 73 +2 PbO -45.25 91 +2 PbMo04 -205.42 73 +2 Pb2+ -5.81 73 +2 HPb02~ -81.00 73 +2.67 Pb,0„ -147.6 73 3 4 +3 Pb2°3 -98.42 73 +4 Pb02 -52.34 73 +4 Pb4+ +72.30 73 +4 PbO 2~ -66.34 73 2-+4 PbO, -67.42 73 4 ( +4 Mo02 -120.00 ) 73 ( +6 Mo03 -161.95 ) 73 ( +6 HMo04- -213.60 ) 73 +6 MoO.2- -205.42 35 -4 Activity of lead containing species in solution taken as 10 M, and of molybdenum containing species as 10 M. Stable bi-valent oxide considered to be anhydrous plumbous oxide, PbO. 281 C.3.2 Two dissolved substances 2+ - + 1) Pb + 2H20 y HPb02 + 3H log (HPb°2 ) = -28.02 + 3 pH; pH = 9.34 (Pb2+) 4+ 2- + 2) Pb + 3H20 y Pb03 + 6H (PbO 2"~) log — = -23.06 + 6 pH; pH V= 3.84 (Pb4+) 2- 4- + 3) Pb03 + H20 y Pb04 + 2H (PbO 4") log r-j- = -40.87 + 2 pH; pH = 20.44 (PbO ) 2+ 4+ 4) Pb y Pb + 2e 4+ E = 1.694 + 0.0295- log ( } (Pb2+) E° = 1.694 2+ 2- + -5) Pb + 3H O y PbO + 6H + 2e 2-(PbO ) E = 2.375 - 0.177 pH = 0.0295 log -(Pb2+) 2- + -6) HPbO + HO y PbO + 3H + 2e Z Z j> (Pb03 ")' E = 1.547 - 0.0886 pH + 0.0295 log _ (HPb02 ) E = 1.547 - 0.0886 pH - 282 C.3.3 Two solid substances 7) Pb + H20 y PbO + 2H+ + 2e" E = 0.248 - 0.0591 pH 8) 3PbO + H20 y Pb304 + 2H+ + 2e" E = 0.972 - 0.0591 pH 9) 2pb3°4 + H2° 3Pb2°3 + 2H+ + 2e~ E = 1.228 - 0.0591 pH 10) Pk304 + 2H20 > 3Pb02 + 2H+ + 2e~ E = 1.127 - 0.0591 pH 11) pb203 + H20 — -> 2Pb02 + 2H+ + 2e~ E = 1.093 - 0.0591 pH 2-12) Pb + MoO, y PbMoO, + 2e 4 4 E = -0.672 + log (Mo042~) E = -0.613 2- + -13) PbMo04 + 2H20 y Pb02 + Mo04 + 4H + 2e E = 1.99 - 0.1182 pH + 0.0591 log (MoO^") E = 1.931 - 0.1182 pH 2- + 14) 3PbMo04 + 4H20 y Pb304 + 3Mo04 + 8H + 2e E = 3.724 - 0.236 pH + 3(0.0591) log (MoO^" E = 3.547 - 0.236 pH - 283 C.3.4 One solid substance and one dissolved substance 2+ + 15) Pb + H20 y PbO + 2H 2+ log (Pb ) = 12.65 - 2 pH; pH = 8.325 16) PbO + H20 y HPb02~ + H+ log (HPb02 ) = -15.36 + pH; pH = 11.36 4+ + 17) Pb + 2H20 > Pb02 + 4H log (Pb4+) = -8.26 - 4 pH; pH = -1.065 2- + 18) Pb02 + H20 -> Pb03 + 2H 2-log (PbO. ) = -31.22 + 2 pH; pH = 13.66 2+ 19) Pb • Pb + 2e E = -0.126 + 0.0295 log (Pb2+) E = -0.244 20) Pb + 2H20 y HPb02 + 3H+ + 2e" E = 0.702 - 0.0886 pH + 0.0295 log (HPb02 ) E = 0.584 - 0.0886 pH 2+ + -21) 3Pb + 4H20 y Pb304 + 8H + 2e E = 2.094 - 0.2364 pH - 0.0886 log (Pb2+) E = 2.448 - 0.2364 pH 22) 3HPb02~ + H+ —~-y Pb^ + 2H20 + 2e~ E = -0.390 ••+ 0.295 pH - 0.0886 log (HPbO^) - 284 -E = -0.0356 + 0.0295 pH 23) Pb2+ + 2H20 — -> Pb02 + 4H+ + 2e~ E = 1.449 - 0.1182 pH - 0.0295 log (Pb2+) E = 1.567 - 0.1182 pH 24) HPb02~ + Mo042~ + 3H+ • PbMo04 + 2H2<3 log (HPb02~) = -46.33 + 3pH; pH = 14.11 4+ 25) PbMo04 y Pb + 4e E = 3.347 - 0.0147 log (Pb4+) E = 3.288 2- 2- + -26) PbMo04 + 3H20 >- Pb03 + Mo04 + 6H • + 2e E = 2.915 - 0.177 pH + 0.0295 log (PbO 2~) E = 2.797 - 0.177 pH - 285 -C. 4 Zn-H^O-MoO^ system C.4.1 Substances considered Oxidation Number Species AG° Source (kcal) of Data 0 +2 +2 +2 +2 +2 +2 ( +4 ( +6 ( +6 +6 Zn Zn(OH)„ ZnMoO„ * 2+ Zn HZn02~ ZnO 2~ 2+ ZnOH Mo02 Mo03 HMoO.~ ' 4 MoO2~ -76.94 -294.40 -35.18 -110.90 -93.03 -78.70 -120.00 ) -161.95 ) -213.60 ) -205.42 73 91 73 73 73 73 73 73 73 73 35 Activity of zinc containing species taken to be 10 ^ M, and of molybdenum containing species as 10 -1 M. Stable oxide considered to be zinc hydroxide Zn(OH)2 (orthorhombic) C.4.2 Two dissolved substances 1) Zn2+ + H20 • ZnOH+ + H+ (7nOH~'~ ^ log = -9.67 + pH; pH = 9.67 (Zn+) + - + 2) ZnOH + H20 > HZn02 + 2H (HZn02~) log — = -17.97 + 2 pH; pH = 8.98 (ZnOH ) - 286 -2+ - + 3) Zn + 2H20 y HZnC>2 + 3H (HZn02~) log — = -27.63 + 3 pH; pH = 9.21 (Zn +) 2- + 4) HZnO y ZnO + H 2-(Zn02 ) log = -13.11 + pH; pH = 13.11 (HZn02 C.4.3 Two solid species 5) Zn + 2H20 >- Zn (OH) + 2H+ + 2e E = -0.439 - 0.0591 pH 2- + 6) Zn(OH)2 + Mo04 + 2H • ZnMo04 + 2H20 2-2 pH = 91.84 + log (Mo04 ) pH = 45.42 2-7) Zn + Mo04 + 2e >- ZnMo04 2-E = -1.929 - 0.0295 log (Mo04 ) E = -1.899 C.4.4 One solid substance and one dissolved substance 8) Zn(OH)2 + H20 *• HZn02~ + H+ log (HZn02 ) = -16.68 + pH; pH = 10.68 - 287 -2- + 9) HZnC- + MoO. + 3H y ZnMoO„ + 2H„0 2 4 4 2 2-3 pH = 66.94 - log ((HZn02 )•(Mo04 )) pH = 20.31 2- 2- + 10) ZnO„ + MoO, + 4H y ZnMoO, + 2H~0 2 4 4 2 2-4 pH = 80.04 - log ((Zn02 )•(MoO )) pH = 18.51 2+ 11) Zn y Zn + 2e E = -0.763 + 0.0295 log (Zn2+); E = -0.94 2- + -12) Zn + H20 y Zn02 + 4H + 2e E = 0.441 - 0.1182 pH + 0.0295 log (Zn022~) E = 0.264 - 0.1182 pH 2+ + 13) Zn + 2H20 y Zn(OH)2 + 2H 2+ log (Zn ) = 10.96 - 2 pH; pH = 8.48 14) Zn + 2H20 y HZn02~ + 3H+ + 2e~ E = 0.054 - 0.0886 pH + 0.0295 log (HZn02~) E = -0.123 - 0.0886 pH - 288 -C. 5 Fe-H^O-MoO^ system C.5.1 Substances considered Oxidation „ . AG° Source Species Number (kcal) of Data 0 Fe +2 Fe(OH) -58.88 73 2+ +2 Fe -20.30 73 +2 HFe02~ -90.63 73 +2 FeMoO, -235.30 91 4 +3 Fe(OH). -161.93 73 3+ +3 Fe -2.53 73 +3 Fe(OH) + -106.202+ +3 FeOH -55.91 73 ( +4 Mo02 -120.00 ) .73 ( +6 Mo03 -161.95 ) 73 ( +6 HMoO -213.60 ) 73 +6 MoO„ ~ -205.42 35 V +6 FeO„ -111.69 73 4 Activity of iron containing species in solution taken as 10 ^ M, and of molybdenum species as 10 M. Ferrous and ferric hydroxides, Fe(OH)2 and Fe(OH) have been used as the stable oxides. C.5.2 Two dissolved species 2+ - + 1) Fe + 2H20 ——* HFe02 + 3H < (HFeO ) log = -31.58 + 3 pH; pH = 10.53 (Fe2+) - 289 -3+ 2+ + 2) Fe + H20 FeOH + H 2+ log (Fe°" ) =•- -2.43 + pH; pH = 2.43 (FeJ ) 3) FeOH2+ + H20 —>• Fe(OH)2+ + H+ log (Fe(°H)2 } = -4.69 + pH; pH = 4.69 (Fe(OH)2+) .,2+ 3+ 4) Fe Fe + e (Fe3+) E = 0.771 + 0.0591 log (Fe^ ) E = 0.771 5) Fe2+ + H20 >: Fe(OH)2+ + H+ + e" (Fe (OH) 2+. E = 0.914 - 0.0591 pH + 0.0591 log (Fe2+) E = 0.914 - 0.059 pH 6) Fe2+ + 2H90 ^-r* Fe (OH) + + 2H+ + e~ 2 + (Fe(OH) ) E = 1.91 - 0.1182 pH + 0.0591 log (Fe2+) E = 1.91 - 0.1182 pH 7) HFeO + H+ —Fe (OH) + + e~ + (Fe(OH)2 ) E = -0.675 + 0.0591 pH + 0.0591 log :  (Fe02H~) E = -0.675 + 0.0591 pH 2- + -8) HFe02 + 2H20 —Fe04 + 5H + e (Fe042~) E = 1.001 - 0.0738 pH + 0.0148 log (HFe02 E = 1.001 - 0.0738 pH - 290 -3+ 2- + -9)' Fe + 4H20 —Fe04 + 8H + 3e (FeO ") E = 1.700 - 0.1580 pH + 0.0197 log -(Fe3+) E = 1.700 - 0.1580 pH 10) Fe(OH)2+ + 3H20 ——> Fe042~ + 7H+ + 3e~ (FeO 2~) E = 1.652 - 0.1379 pH + 0.0197 log (Fe(OH)2+) E = 1.652 - 0.1379 pH + 2- + -11) Fe(OH)2 + 2H20 —y Fe04 + 6H + 3e E = 1.559 - 0.1182 pH + 0.0197 log (Fe°4 ) 1.559 - 0.1182 pH (Fe(OH)2+) C.5.3 Two solid substances 12) Fe + 2H20 y Fe(OH)2 + 2H+ + 2e~ E = -0.047 - 0.0591 pH 13) Fe + 3H20 y Fe (OH) + 2H+ + 2e~ E = 0.059 - 0.0591 pH 14) Fe(OH)2 + H20 >-Fe(OH)3 + 2H+ + 2e~ E = 0.271 - 0.0591 pH 2-15) Fe + MoO, + 2e >• FeMoO, 4 4 2-E = -0.648 - 0.0295 log (Mo04 ) E = -0.6185 2- + 16) Fe(OH)2 + Mo04 + 2H >- FeMo04 + 2H20 291 -2-2 pH = 61.804 + log (MoO ) 4 pH = 30.402 17) FeMo04 + 3H20 y Fe (OH) 3 + 3H+ + MoO 2~ + e~ E = 1.637 - 0.1773 pH + 0.0591 log (MoO 2~) E = 1.578 - 0.1773 pH C.5.4 One solid substance and one dissolved substance 18) Fe2+ + 2H20 y Fe(OH)2 + 2H+ log (Fe2+) = 13.29 - 2pH pH = 9.645 19) Fe(OH)2 y HFe02~ + H+ log (HFe02 ) = -18.30 + pH pH • = 12.30 20) Fe3+ + 3H 0 y Fe (OH) + 3H+ 3+ log (Fe ) = 4.84 - 3 pH pH = 3.613 2+ + 21) Fe (OH) + 2H O y Fe (OH) + 2H log (Fe(OH)2+) = 2.41 - 2 pH pH = 4.205 22) Fe(OH)2+ + H20 • >-Fe(OH)3 + H+ - 292 log (Fe(OH)2+) = -2.28 - pH pH = 4.28 2+ 23) Fe y Fe + 2e E = -0.440 + 0.0295 log (Fe2+) E = -0.617 24) Fe + 2H20 ^ HFe02 + 3H+ + 2e~ 2+ E = 0.493 - 0.0886pH + 0.0295 log (Fe ) E = 0.316 - 0.0886 pH 3+ 25) Fe Bj Fe + 3e E = -0.037 + 0.0197 log (Fe3+); E = -0.1552 2+ + -26) Fe + 3H20 • Fe(OH) + 3H + e E = 1.057 - 0.1773 pH - 0.0591 log (Fe2+) E = 1.4116 - 0.1773 pH 27) HFe02 + H20 -y 2Fe(OH)3 + e -0.810 - 0.0591 log (HFe02 ); E = -0.4356 2- + 28) FeMo04 + 2H20 y HFe02 + Mo04 + 3H 3 pH = 38.46 + log ((HFe02~)-(Mo042~)) pH = 10.82 3+ 2-29) FeMoO. 4- Fe + MoO, + e 4 4 2-E = 1.186 + 0.0591 log (Mo04 ); E = 1.127 - 293 -r 2-H 2 —• "I-30.) FeMoO + HO y Fe (OH) + MoO + H + e E = 1.326 - 0.0591 pH + 0.0591 log ((FeOH2+) (Mo042*)) E = 0.971 - 0.0591 pH + + - 2-31) FeMo04 + 2H20 y Fe(OH)2 + 2H + e + Mo04 E = 1.599 - 0.1182 pH + 0.0591 log ((FeOH+)• (Mo042-)) E = 1.244 - 0.1182 pH C.6 Cd-H2Q-MoQ4 system C.6.1 Substances considered Oxidation „ . AG° Source Species ,, _ „ , Number (kcal) of Data 0 Cd — 73 2 Cd(OH)2 -56.44 73 2 CdMoO„ ' -283.30 91 2+ 4 2 Cd -18.58 73 2 HCdO ~ -86.50 73 6 MoO„2- -205.42 35 4 Activity of cadmium containing species in solution taken as 10 6 M, and that of molybdenum species as 10 1 M. 'Inactive' cadmium hydroxide, Cd(OH)2, has been used as the oxide species. - 294 C.6.2 Two dissolved species 1) Cd2+ + 2H20 HCd02 + 3H+ (HCdO ) log — = -33.34 - 3 pH; pH = 11.11 (CcT ) C.6.3 Two solid species 2) Cd + H20 >- Cd(OH)2 + 2H+ + 2e~ E = 0.05 - 0.0591 pH 2- + 3) Cd(OH)2 + Mo04 + 2H • CdMo04 + 2H20 2-2 pH = 65.73 + log (MoC>4 ) pH = 32.365 2-4) Cd + MoO, + 2e > CdMoO, 4 4 2-E = -0.713 - 0.0295 log (Mo04 ); E = -0.684 C.6.4 One solid substance and one dissolved substance 5) Cd2+ + 2H20 >-Cd(OH)2 + 2H+ log (Cd2+) = 13.81 - 2 pH; pH = 9.905 6) Cd(OH)2 > HCd02 + H+ log (HCd02 ) = -19.54 + pH; pH = 13.54 7) Cd y Cd2+ + 2e~ E = -0.403 + 0.0295 log (Cd2+); E = -0.508 - 295 -8) Cd + 2H20 y HCd02~ + 3H+ + 2e~ E = 0.583 - 0.0886 pH + 0.0295 log (HCd02 E = 0.406 - 0.0886 pH 2- + 9) CdMo04 + 2H20 • y HCd02 + MoC>4 + 3H log (HCd02~) - 3 pH = -43.69 pH = 12.564 C.8 Molybdenum species (all systems) C.8.1 Substances considered Oxidation „ . AG° Source Species „ ,\ _ „ ^ Number (kcal) of Data 0 Mo +3 Mo3+ -13.80 73 +4 Mo02 -120.00 73 +6 Mo03 -161.95 73 +6 MoO,2" -205.42 73 4 +6 HMoO ~ -213.60 73 4 Activity of dissolved species taken as 10 "*" M. Hexa-valent oxide considered to be molybdic trioxide, MoO^ C.8.2 Two dissolved substances 2- + 1) HMoO. y MoO. + H 4 4 2-, (MoO ) log — = -6.00 + pH; pH = 6.0 (HMoO ) 296 C.8.3 One solid substance and one dissolved substance 2) Mo03 + H20 K HMo04~ + H+ log (HMo04 ) = -3.70 + pH; pH = 2.70 3+ 3) Mo y Mo + 3e E = -0.200 + 0.0197 log (Mo3+); E = -0.2197 2- + 4) Mo + 4H20 y Mo04 + 8H + 6e 2-E = 0.154 - 0.0788 pH + 0.0098 log (Mo04 E = 0.144 - 0.0788 pH 3+ + -5) Mo + 2H20 y Mo02 + 4H + e E = 0.311 - 0.2364 pH - 0.0591 log (Mo3+) E = 0.370 - 0.2364 pH 3+ + -6) Mo + 3H20 y Mo03 + 6H + 3e E = 0.317 - 0.1182 pH - 0.0197 log (Mo3+) E = 0.337 - 0.1182 pH 7) Mo02 + 2H20 y HMo04 + 3H+ + 2e~ 0.429 - 0.0886 pH + 0.0295 log (HMo04 ) E = 0.399 - 0.0886 pH 2- + 8) Mo02 + 2H20 y Mo04 + 4H + 2e 2-E = 0.606 - 0.1182 pH + 0.0295 log (Mo04 ) E = 0.577 - 0.1182 pH - 297 -C.8.4 Two solid substances 9) Mo + 2H20 > Mo02 + 4H+ + 4e~ E = -0.072 - 0.0591 pH 10) Mo02 + H20 y Mo03 + 2H+ + 2e E = 0.320 - 0.0591 pH C 


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