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Cobalt precipitation by reduction with sodium borohydride Lu, Jianming 1995

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C O B A L T P R E C I P I T A T I O N B Y R E D U C T I O N W I T H S O D I U M B O R O H Y D R I D E by Jianming L u B . Eng., Northeastern University of Technology, P. R. China, 1983 M . Eng., Shanghai University of Technology, P. R. China, 1990 A THESIS S U B M I T T E D I N P A R T I A L F U L F I L L M E N T O F T H E R E Q U I R E M E N T S F O R T H E D E G R E E OF M A S T E R OF A P P L I E D S C I E N C E in T H E F A C U L T Y OF G R A D U A T E S T U D I E S Department of Metals and Materials Engineering We accept this thesis as conforming to the required standard T H E U N I V E R S I T Y OF B R I T I S H C O L U M B I A December 1995 © Jianming L u , 1995 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 or her representatives. It is understood that copying or publication of this thesis for financial gain shall not be allowed without my written permission. The University of British Columbia Vancouver, Canada Department DE-6 (2/88) 11 A B S T R A C T The reaction of cobalt reduction with borohydride is very complicated. Various authors obtained different reaction stoichiometries and have proposed a number of mechanisms. There are conflicting claims about the removal of cobalt from zinc electrolyte with sodium borohydride. The present research has focused on the stoichiometry of cobalt reduction with borohydride and on the removal of cobalt from zinc sulphate solution. Cobalt reduction with borohydride releases hydrogen ions resulting in a decrease in the pH of the solution. The efficiency of cobalt reduction increased with increasing concentration of NaOH in the reducing solution. The NaOH in the reducing solution neutralized the hydrogen ions released during cobalt reduction. The best reduction efficiency without the precipitation of cobalt hydroxide is one mole of sodium borohydride to reduce one mole of cobalt (II). If the pH was controlled at 4, the maximum reduction efficiency was about 81 % for a ten-minute addition time and the addition of more sodium borohydride did not increase the reduction efficiency. The reduction efficiency increased with increasing pH from nil at pH 2 to 96% at pH 6 and decreased with increasing temperature. X-ray diffraction patterns and T E M patterns of the recovered precipitates showed them to be amorphous. After a two-hour heat-treatment at 500 °C, the X-ray diffraction pattern of the precipitate showed well-defined peaks due to C 0 2 B with the main peak attributable to cobalt. The single crystal T E M pattern obtained was consistent with that of C 0 2 B . The particle size was about 20-100 nm. The atom ratio of Co to B increased with increasing temperature. The reduction of cobalt ions in the absence of interfering species was completed within several seconds. The time for the reduction of most of the cobalt ions decreased with increasing temperature from 24 seconds at 5 °C to less than 2 seconds at 35 °C. Zinc ions had a dramatic inhibitory effect on cobalt reduction. Several tens of iumol/L of zinc ions completely inhibit cobalt reduction with borohydride. The main cause of inhibition is that zinc ions compete with those of cobalt for borohydride ions and zinc borohydride forms and hydrolyzes rapidly. The resulting zinc ions react further with borohydride. Zinc ions catalyze the hydrolysis of borohydride. Zinc ions also have an inhibitory effect on nickel, cadmium and lead reduction. Ill T A B L E O F C O N T E N T S Abstract i i Tables of Contents i i i Lists of Tables vi Lists of Figures v Acknowledgments x i 1. Introduction 1 2. Literature Review 3 2.1 Application of Sodium Borohydride 3 2.1.1 Preparation of Ultrafme Amorphous Materials 4 2.1.2 Recovery of Precious Metals in Hydrometallurgy 5 2.1.3 Pollution Control Application 6 2.1.4 Electroless Plating 7 2.1.5 Analytical Chemistry 7 2.2 Hydrometallurgical Extraction of Zinc 8 2.2.1 Introduction 8 2.2.2 Impurities in Zinc Electrolyte 10 2.2.3 Methods of Cobalt Removal from Zinc Electrolyte 11 2.2.3.1 Cobalt Removal by Chemical Oxidation 11 2.2.3.2 Precipitation Method 11 2.2.3.3 Solvent Extraction Method 12 2.2.3.4 Cobalt Cementation with Zinc Dust 12 2.2.3.5 Cobalt Cementation with Manganese Dust 14 2.2.3.6 Cobalt Removal with Sodium Borohydride 15 2.3 Chemistry of Boron 16 2.3.1 Physical and Chemical Properties of Boron 16 2.3.2 Boron Compounds 17 2.3.2.1 Borides : 17 2.3.2.2 Oxygen Compounds of Boron 18 IV 2.3.2.3 Boron Hydride (Borane) and Tetrahydride 19 2.3.2.4 Metal Borohydride 20 2.4 Aqueous Chemistry of Sodium Borohydride 23 2.4.1 Aqueous Solution of Sodium Borohydride 23 2.4.1.1 Solubility and Stability 23 2.4.1.2 Thermodynamic Properties 24 2.4.1.3 Hydrolysis of of Borohydride Ions 26 2.4.2 Reduction of Metal Ions with Borohydride 28 2.4.3 Aqueous Chemistry and Borohydride Reduction of Cobalt ; 31 2.4.3.1 Aqueous Chemistry of Cobalt 31 2.4.3.2 Borohydride Reduction of Cobalt Ion 33 2.4.3.3 Kinetics of Cobalt Reduction with Borohydride 42 2.5 Summary 42 3. Experimental 44 3.1 Experimental Objectives and Methods 44 3.2 Equipment 44 3.3 Materials 45 3.4 Experimental Procedure 48 3.4.1 Procedure for the Reduction of Cobalt 48 3.4.2 Heat-treatment of the Precipitates 48 3.4.3 X-Ray Diffraction of the Precipitates 49 3.4.4 Scanning Electron Microscopic Analysis of the Precipitates 49 3.4.5 Transmission Electron Microscopic Analysis of the Precipitates 49 4. Results and Discussion 50 4.1 Stoichiometry of Cobalt Reduction with Sodium Borohydride 50 4.2 The Effect of pH on Cobalt Reduction Efficiency with Sodium Borohydride 67 4.3 The Effect of Addition Rate on Cobalt Reduction Efficiency 68 4.4 The Effect of Temperature on Cobalt Reduction Efficiency 69 4.5 Redissolution of Reduced Cobalt 70 4.6 Identification of the Precipitate 72 4.6.1 Morphology and Properties of Precipiates 72 4.6.2 Identification by X-Ray Diffraction 75 4.6.3 T E M Identification of the Precipitate 78 4.6.4 The Effect of Temperature on the Composition of the Precipitate 84 4.6.5 The Effect of Addition Rate on the Composition of the Precipitate 85 4.7 Possible Reaction of Cobalt Reduction with Borohydride 86 4.8 Kinetics of Cobalt Reduction with Borohydride 89 4.8.1 Features of the Kinetics of Cobalt Reduction. 89 4.8.2 The Effect of Temperature on the Reaction Rate 92 4.8.3 The Effect of pH on the Reaction Rate 93 4.8.4 The Autocatalytic Effect of Cobalt Reduction 96 4.9 Cobalt Precipitation from Zinc Sulfate Solution 98 4.9.1 The Inhibitory Effect of Zinc Ions on Cobalt and Nickel Reduction 98 4.9.2 The Effect of Additives on Cobalt Reduction 105 4.9.3 Possible Inhibitory Mechanism of Zinc Ions 106 5. Conclusions 108 6. Recommendations I l l 7. References 112 Appendix 1 Joint Committee Powder Diffraction (JCPDS) Card of Randomly Oriented Cobalt (Cubic) 129 Appendix 2 Joint Committee Powder Diffraction (JCPDS) Card of Randomly Oriented Co 2 B (Tetragonal) 130 Appendix 3 Joint Committee Powder Diffraction (JCPDS) Card of Randomly Oriented Boron (Rhombohedral) 131 Appendix 4 Procedure for Colorimetric Cobalt Analysis 132 L I S T O F T A B L E S Table 2-1 Theoretical weight ratio of reduced metals obtainable from ionic species per amount of sodium borohydride 3 Table 2-2 Impurity levels in various zinc plant electrolytes 10 Table 2-3 Some physical and chemical properties of sodium borohydride 22 Table 2-4 p H o f solutions o f N a B H 4 at 24 °C 24 Table 2-5 Half-lives for the hydrolysis of the borohydride ion in the presence o f transition ions 27 Table 2-6 Effect of additives on the reduction of cobalt 41 Table 4-1 Moles o f sodium borohydride required to reduce one mole o f cobalt ions for different basic solutions and different initial p H values 51 Table 4-2 Calculation of the camera constant ( A u calibration) for T E M 79 Table 4-3 Calculation of the d-spacing of the precipitate 80 Table 4-4 Calculated theoretical angles and measured angles 80 vii L I S T O F F I G U R E S Figure 2-1 Schematic diagram of hydrometallurgical zinc extraction 9 Figure 2-2 Reaction data for the isothermal cementation of cobalt onto zinc in the presence of additional zinc ions at 95 °C 13 Figure 2-3 Co-B phase diagram 18 Figure 2-4 Elements forming metal tetrahydroborate compounds 20 Figure 2-5 The solubility of sodium borohydride in water at different temperatures 23 Figure 2-6 E h-pH diagram for B-H 2 0 system at 25 °C 25 Figure 2-7 Eh-pH diagram for Co-H 20 system at Co concentration of 0.5, 0.05, 0.005, 0.0005Mand25 °C 32 Figure 3-1 Schematic diagram of the experiment setup 46 Figure 3-2 Schematic diagram of sample heat-treatment 47 Figure 4-1 Cobalt reduction fraction, pH and potential vs. time with 35 °C, 10-minute addition time and initial pH 7, the reducing solution being 0.25 M NaBH 4 + 0.05 M NaOH 53 Figure 4-2 Cobalt reduction fraction, pH and potential vs. time with 35 °C, 5-minute addition time and initial pH 7, the reducing solution being 0.25 M NaBFL + 0.05 M NaOH 54 Figure 4-3 Cobalt reduction fraction, pH and potential vs. time with 35 °C, 10-minute addition time and initial pH 5.5, the reducing solution being 0.25 M NaBH4 + 0.05MNaOH 55 Figure 4-4 Cobalt reduction fraction, pH and potential vs. time with 35 °C, 10-minute addition time and initial pH 8, the reducing solution being 0.25 M NaBH4 + 0.05 M NaOH 56 Figure 4-5 Cobalt reduction fraction, pH and potential vs. time with 35 °C, 10-minute addition time and initial pH 5.5, the reducing solution being 0.25 M NaBH4 57 Figure 4-6 Cobalt reduction fraction, pH and potential vs. time with 35 °C, 10-minute addition time and initial pH 7, the reducing solution was 0.25 M NaBH 4 58 Figure 4-7 Cobalt reduction fraction, pH and potential vs. time with 35 °C, 10-minute addition time and initial pH 8, the reducing solution being 0.25 M N a B H 4 59 Figure 4-8 Cobalt reduction fraction, pH and potential vs. time with 35 °C, 10-minute addition time and initial pH 5.5, the reducing solution being 0.25 M N a B H 4 + 0.125 M NaOH 60 Figure 4-9 Cobalt reduction fraction, pH and potential vs. time with 35 °C, 10-minute addition time and initial pH 7, the reducing solution being 0.25 M N a B H 4 + 0.125 M NaOH 61 Figure 4-10 Cobalt reduction fraction, pH and potential vs. time with 35 °C, 10-minute addition time and initial pH 9, the reducing solution being 0.25 M N a B H 4 + 0.125 M NaOH 62 Figure 4-11 Cobalt reduction fraction, pH and potential vs. time with 35 °C, 10-minute addition time and initial pH 5.5, the reducing solution being 0.25 M N a B H 4 + 0.25 M NaOH 63 Figure 4-12 Cobalt reduction fraction, pH and potential vs. time with 35 °C, 10-minute addition time and initial pH 7, the reducing solution being 0.25 M N a B H 4 + 0.25 M NaOH 64 Figure 4-13 Cobalt reduction fraction, pH and potential vs. time with 35 °C, 10-minute addition time and initial pH 8, the reducing solution was 0.25 M N a B H 4 + 0.125 M NaOH 65 Figure 4-14 The reduction fraction vs. time and mol ratio of N a B H 4 to C o 2 + at 35 °C and p H 4 66 Figure 4-15 The effect of pH on cobalt reduction efficiency at 35 °C 67 Figure 4-16 The effect of addition rate of borohydride on cobalt reduction efficiency 68 Figure 4-17 The effect of temperature on cobalt reduction efficiency at pH 4 69 Figure 4-18 The redissolution of reduced cobalt at 35 °C and pH 4 71 Figure 4-19 Surface morphology of the fresh precipitate prepared under the conditions: pH 5.5-7.5, 35 °, 500 rpm stirring rate, mol ratio of N a B H 4 to C o 2 + = 2, 10-minute addition time, initial [Co 2 +] = 30 mg/L and 200 x 73 Figure 4-20 Surface morphology of the fresh precipitate prepared under the conditions: pH 5.5-7.5, 35 °, 500 rpm stirring rate, mol ratio of N a B H 4 to C o 2 + = 2, ix 10-minute addition time, initial [Co 2 +] = 30 mg/L and 40,000 x 73 Figure 4-21 Transmission electron micrograph of the precipitate under the conditions: pH 5.5-7.5, 500 rpm stirring rate, mol ratio of N a B H 4 to C o 2 + = 2, 10-minute addition time, initial [Co 2 +] = 30 mg/L and 50000 x 74 Figure 4-22 Transmission electron micrograph of the precipitate under the conditions: pH 5.5-7.5, 500 rpm stirring rate, mol ratio of N a B H 4 to C o 2 + = 2, 10-minute addition time, initial [Co 2 +] = 30 mg/L, 500 °C heat treatment for two hours and 100,000 x '. 77 Figure 4-23 X-ray diffraction pattern of the fresh precipitate under conditions: pH 5.5-7.5,35 °C, 500 rpm stirring rate, mol ratio of N a B H 4 to C o 2 + = 2, and 10-minute addition time 76 Figure 4-24 X-ray diffraction pattern of the precipitate under the conditions: pH 5.5-7.5, 35 °C, 500 rpm stirring rate, mol ratio of N a B H 4 to C o 2 + = 2, 10-minute addition time and 500 °C heat-treatment for two hours 76 Figure 4-25 X-ray diffraction pattern of the precipitate under the conditions: pH 5.5-7.5, 35 °C, 500 rpm stirring rate, mol ratio of N a B H 4 to C o 2 + = 2, 10-minute addition time, and heating in boiling water for 80 minutes 77 Figure 4-26 X-ray diffraction pattern of the precipitate under the conditions: pH 5.5-7.5, 35 °C, 500 rpm stirring rate, mol ratio of N a B H 4 to C o 2 + = 2, 10-minute addition time, heating in boiling water for 80 minutes and 500 °C heat-treatment for two hours 77 Figure 4-27 X-ray diffraction pattern of the precipitate under the conditions: pH 5.5-7.5, 35 °C, 500 rpm stirring rate, mol ratio of N a B H 4 to C o 2 + = 2, and 10-minute addition time 78 Figure 4-28 T E M pattern of the precipitate under the conditions: pH 5.5-7.5, 35 °C, 500 rpm stirring rate, mol ratio of N a B H 4 to C o 2 + = 2, and 10-minute addition time 81 Figure 4-29 T E M pattern of the precipitate under the conditions: pH 5.5-7.5, 35 °C, 500 rpm stirring rate, mol ratio of N a B H 4 to C o 2 + = 2, 10-minute addition and 500 °C heat-treatment for two hours 82 X Figure 4-30 T E M pattern of gold film 83 Figure 4-31 The effect of temperature on the composition of the precipitate 84 Figure 4-32 The effect of the addition rate on the composition of the precipitate 85 Figure 4-33 pH and potential of the solution vs. time under the conditionsl5 °C, 2500 rpm stirring rate and initial pH 5.5 90 Figure 4-34 pH vs. time for the solutions with different compositions and addition of different additives under the conditions: 15 °C, 2500 rpm stirring rate and initial pH 5.5 91 Figure 4-35 pH vs. time at different temperatures 92 Figure 4-36 pH vs. time at initial pH 6.0 and 15 °C 93 Figure 4-37 pH vs. time at initial pH 7.0 and 15 °C 94 Figure 4-38 pH vs. time at initial pH 7.5 and 15 °C 94 Figure 4-39 pH vs. time at initial pH 7.8 and 15 °C 95 Figure 4-40 pH vs. time at initial pH 7.0 and 35 °C 95 Figure 4-41 pH vs. time at initial pH 7.8 and 35 °C 96 Figure 4-42 pH vs. time data for autocatalytic effect on cobalt reduction 97 Figure 4-43 pH vs. time data demonstrating the autocatalytic effect on cobalt reductio 97 Figure 4-44 The inhibitory effect of zinc ions on cobalt reduction 98 Figure 4-45 The inhibitory effect of zinc ions on nickel reduction 99 Figure 4-46 The effect of the additives on cobalt reduction for 10-minute addition time 100 Figure 4-47 The effect of the additives on cobalt reduction for 5-minute addition time 101 Figure 4-48 The effect of the additives on cobalt reduction for 2-minute addition time 102 Figure 4-49 The effect of the additives on cobalt reduction for 1-minute addition time 103 Figure 4-50 pH vs. time data for 30 mg/L C o 2 + + 0.0001 M Z n 2 + after the addition of sodium borohydride solution 104 XI A C K N O W L E D G E M E N T S I would like to express my sincere appreciation to my supervisor, Dr. D. B. Dreisinger for his thoughtful supervision and constructive discussions. M y stay at U B C was made possible due to his funding support. I am very grateful to Dr. W. C. Cooper for reviewing and editing this thesis, helpful discussions and kind help and care. I would also like to acknowledge Dr. D. Tromans for providing constructive ideas. Dr. B. Wassink's kind help and reading of this thesis are very much appreciated Many thanks to my fellow graduate students and staff of the hydrometallurgy group in various areas, and with whom I have enjoyed working. Finally, I would like to thank my beloved wife, my parents , brothers and sisters for giving me moral support. 1 1. I N T R O D U C T I O N Sodium borohydride discovered by Schlesinger is an efficient water sol5ble reducing agent [1-3]. It has the following features: (1) a low reducing equivalent weight of 4.75 g/mole e", (2) high reducing power (-1.24 V SHE at pH 14) and (3) redox reaction in different media (water and organic solvents, and acidic, neutral and alkaline conditions). It was first used in the paper and textile industries and the military [4]. The commercial importance of sodium borohydride reached its first peak in the fifties in U.S. military research programs [4, 5]. Recently borohydride reduction of metal ions is the basis of several commercial processes including: (1) the preparation of selective catalysts, magnetic materials [6-24], (2) the recovery of precious metals (silver, gold, platinum and palladium group metals) [6, 25-35], (3) the removal and recovery of heavy metals (mercury, lead, copper, nickel, cobalt and silver) from aqueous discharge streams [36-47], and (4) the electroless plating of nickel, cobalt, copper, silver and gold onto metallic or nonmetallic surfaces such as glass, ceramics and plastics [5, 6, 48-63]. Sodium borohydride has also widely used to reduce metal cations for analysis [5, 64-72]. Finally it may be potentially applicable to emerging hydrometallurgical applications. Most of the world's zinc is now produced by the roast-leach-electro winning process because it has the following advantages over the pyrometallurgical method: (1) a better zinc recovery and energy efficiency, (2) a high purity product, (3) efficient labour utilization [73]. The zinc electrowinning process is very sensitive to the presence of impurities in the electrolyte. Some impurities such as antimony, germanium, cobalt, nickel and copper can cause extensive redissolution of the electrodeposited zinc, leading to a decrease in the current efficiency of zinc electrodeposition with time and an impure zinc product [74-83]. Therefore zinc electrolyte purification is a key operation for the economical extraction of zinc by the hydrometallurgical process. Many methods have been proposed [84-113], but they are still not satisfactory because they are expensive, inefficient, and potentially toxic. Sodium borohydride reduction has been proposed as a way to overcome these disadvantages. Cobalt reduction with sodium borohydride has been widely used to make magnetic materials and catalysts and to remove and recover cobalt from waste water. It is considered as a new available technology for treating waste water for cobalt removal [41, 42]. 2 Cobalt reduction with sodium borohydride is very sensitive to reaction conditions including temperature, the method and rate of addition, p H and the presence of other ions such as zinc and copper. The stoichiometry and reaction mechanism of cobalt reduction with sodium borohydride is still unclear. There are various reports with differing results. Polyakov et al [114, 115] reported that 1.1 moles of sodium borohydride can reduce 4 moles of cobalt ions to lower the concentration of cobalt ions in zinc sulphate electrolyte (120 g/1 zinc ions) to 0.01 ppm in the presence of triethanolamine. However, Awadalla et al [47] reported that 2 moles of sodium borohydride can only reduce 1 mole of cobalt ions and that zinc ions have a strong negative effect on cobalt reduction. Cominco staff [116] have reported that borohydride cannot reduce cobalt from zinc sulphate solution. Thus it is necessary that a careful study of cobalt reduction with sodium borohydride be conducted to see i f this might be a feasible process. Consequently the present research was undertaken with the following objectives: 1. To understand the stoichiometry of the reduction of cobalt ions with sodium borohydride and optimize the cobalt reduction efficiency. 2. To study the feasibility o f removing cobalt from zinc sulphate electrolyte with sodium borohydride. It was hoped that the results of this study would help to increase the efficiency o f cobalt reduction with sodium borohydride and cobalt removal and recovery from aqueous solutions. The presentation of this work is divided into four major sections. Chapter 2 deals with a review of the literature, providing a summary of the current ideas about cobalt reduction or its removal from aqueous solution with sodium borohydride. Chapter 3 describes the experimental procedures used in this study and the experimental parameters tested, while Chapter 4 provides a review and discussion of the experimental results. Chapter 5 presents a summary of the work. 3 2. L I T E R A T U R E R E V I E W 2.1 APPLICATIONS OF SODIUM BOROHYDRIDE Sodium borohydride was discovered by Schlesinger in 1942 while looking for volatile uranium compounds [1]. Brown [2] pioneered the chemistry of this compound. The first application of sodium borohydride was its solvolysis in the presence of cobalt salts to produce hydrogen gas [3]. Today, it is used in numerous applications in the paper and textile industries, in stabilizers for organic materials and in catalysts. Sodium borohydride is a very efficient water soluble reducing reagent and has the following features: 1. Sodium borohydride has a low equivalent weight of 4.75 g/mol e" and one mole of sodium borohydride can supply eight moles of electrons. Table 2-1 lists the theoretical weight ratio of reduced metals obtainable from ionic species per amount of sodium borohydride. Table 2-1 Theoretical weight ratio of reduced metals obtainable from ionic species per amount of sodium borohydride Metal ions Weight ratio(g metal/g reductant) Pb 2 + 22 A g + 22 H g 2 + 21 A u i + 14 Ir 4 + 10 P d 2 + 11 Pt 4 + 10 C u 2 + 7 N i - 6 C o 2 + 6 C d 2 + 12 4 2. Sodium borohydride has a high reducing power. Its redox potential is -1.24 V vs. S H E at p H 14 decreasing to -0.48 V at p H 0. 3. Sodium borohydride redox reaction can take place in different media such as water and organic solvents, and under acidic, neutral and alkaline conditions. These properties distinguish it favourably from other reducing agents. Hence sodium borohydride has been widely used to reduce a variety of metal cations to metallic state. These technologies are the basis of several commercial processes and are also important to the analytical chemistry of many elements [4-72]. 2.1.1 P R E P A R A T I O N O F U L T R A F I N E A M O R P H O U S M A T E R I A L S Amorphous materials have usually been prepared by three techniques: l iquid quenching, vapor deposition and solid-state reactions. However, a method o f chemical reduction with borohydride to prepare ultrafme amorphous powders has recently attracted increasing attention because of the unique preparation process and the fine morphology o f the resulting powder. Amorphous materials can used as ferrofluids, catalysts, and magnetic recording materials [6-24]. A s catalysts, amorphous materials were found to be more active and to exhibit different catalytic properties as compared to the corresponding crystalline metals [5, 12]. M u c h attention has focused on nickel and cobalt reduction. The reduction reaction generally takes place as follows: 4BH4~ +2M2+ + 9H20 = M2B + 3H3B03 + \25H2 (2-1) where M is N i or Co. 5 2.1.2 R E C O V E R Y OF PRECIOUS M E T A L S IN H Y D R O M E T A L L U R G Y G O L D Sodium borohydride can be used as an efficient agent for reducing gold and silver in acidic solutions of thiourea, thiocyanate, thiosulfate, chloride, and nitrate [25, 26]. The Au(I) ion is reduced to metallic gold in the form of very fine crystals. The reduction reaction from thiourea solution can be represented as follows: SAu[CS(NH2 )2]+ +BH4~ + 2H20 = SAu + SCS(NH2)2 + B02 +SH+ (2-2) The recovery of gold from cyanide solutions can be carried out effectively i f the cyanide is first oxidized to cyanate. SAu(CNO)3 + 3NaBH4 + 2AOW = 3NaB02 + 2 4 C M T + \SH20 + SAu (2-3) Also it is reported that gold was recovered directly from cyanide and thiourea leach solutions [27]. Carbonate is used as an eluant for gold in a carbon-in-pulp process and the direct and selective recovery of gold from carbonate solutions has been reported [28]. The barren solution is in a condition such that it can be recycled to the upstream process. Gold ions are present on the activated carbon as cyanide complexes and are expected to be in the carbonate eluate also as cyanide complexes. The reaction can be represented as follows: 2Au(CN)2~ + BH4~ + 20H' = 2Au + 4CN~ + B02 + 3H2 (2-4) SILVER Sodium borohydride is also used to recover silver from various sources such as spent photographic fixer solution, electronic scrap , mirrors etc. [5]. In the case of spent photographic fixer, for example, the reaction may be written as : 6 8^g(S2<93)23~ +BHA~ +2H20 = B02- +W+ +\6S2032- +SAg (2-5) Silver is recovered from a carbonate solution having a pH in the alkaline range by the addition of stabilized alkali metal borohydride [28]. P L A T I N U M GROUP M E T A L S Sodium borohydride can be used as an efficient agent to reduce platinum group metals from acidic solutions, such as leach or strip liquors at ambient temperature [5, 29-33]. Compared to the common reductants, the use of sodium borohydride leads to higher recovery, lower cost, and simpler operation. The reaction for PtCl62" is: 2PtCl62' + BH- + 2H20 = 2Pt + BO{ + SHCl + 4CP (2-6) Sodium borohydride is also used to separate iridium [34] and to reduce rhodium in the process of solvation and purification of rhodium [35]. 2.1.3 POLLUTION CONTROL APPLICATION Several industries now employ sodium borohydride to reduce heavy metal cation impurities from their effluent streams to meet stringent environmental requirements [5, 36-47]. These include: 1. The use of sodium borohydride to remove and recover copper, lead, etc. from their electroless plating baths and etchants in printed circuit board manufacturers. 2. The use of sodium borohydride to remove silver and cadmium from plant effluents in photographic film manufacture to meet discharge limits. 7 3. The use sodium borohydride to remove soluble mercury from the effluents in mercury-cell chloralkali plants. 4. The use of sodium borohydride to remove and recover cobalt, copper and lead from other waste water sources. 2.1.4 ELECTROLESS PLATING Sodium borohydride is a highly effective reducing agent for the electroless plating of nickel, cobalt, silver, gold and copper on metallic and nonmetallic surfaces [5, 6, 48-63]. Depending on the bath conditions, either a metal or a metal boride layer may be formed. Borohydride-reduced coatings provide superior hardness, wear resistance, lubricity and uniformity, solderability and brazability and lower labor costs. Nickel electroless plating has been most widely studied and used to prepare amorphous coatings with different properties such as Ni -B , Ni-P, Ni-P-B. 2.1.5 A N A L Y T I C A L CHEMISTRY Sodium borohydride has been widely used to reduce metal cations for analysis [5, 64-72]. The sodium borohydride reduction of nanogram quantities of arsenic, antimony, bismuth, tin, germanium, mercury, tellurium, selenium, and lead to produce votatile hydrides for detection by atomic absorption, gas chromatography and emission spectroscopy has been widely reported [5, 66-70]. Sodium borohydride is also used to reduce silver, copper and nickel in the spectrophotometric determination [64]. 8 2.2 HYDROMETALLURGICAL EXTRACTION OF ZINC 2.2.1 INTRODUCTION In 1880, Luckow [73] applied for the first patent for electrolytic zinc production. Since then, the process has advanced to the point where more than 80 % o f the world's production o f zinc is produced by the roast-leach-electrowinning process (RLE). The RLE process has been adopted in many plants because it offers a number of advantages including good zinc recovery and energy efficiency, a high purity product, efficient labour utilization and relatively minor environmental problems. Basically, the R L E process consists of the following five major steps: 1. Concentrates containing zinc sulphide plus several other impurities are roasted in air at 800 - 975 °C to form acid-soluble zinc oxide, other less soluble oxides, and sulfur dioxide gas. 2. Calcine containing zinc oxide and other oxides is leached in electrolyte containing sulphuric acid. ZnO + H2 SO, = ZnSO, + H20 (2-7) Zinc ferrite and certain other impurities such as silver, bismuth, lead, and barium form insoluble precipitates. Additional leaching steps are often used to dissolve more insoluble zinc ferrite (e.g. ZnFe204). 3. Dissolved iron and other impurities such as arsenic, antimony, germanium, indium, tin, mercury and lead are precipitated from solution as hydroxides when the p H is adjusted to 4.5-5.1 by returning acid and calcine addition. 2Fe3+ + 3ZnO + 3H20 = 2Fe(OH)2 +3Zn2+ (2-8) Copper, nickel, cobalt, aluminum, arsenic and germanium are only partially co-precipitated [74] or absorbed with iron, while cadmium, magnesium and chloride remain in solution. Further treatment is needed before zinc can be efficiently electrowon from solution. 9 4. Dissolved impurities such as cadmium, copper, cobalt and nickel are removed from solution by zinc dust cementation. This step is divided into two stages. The first is the removal of copper, cadmium, nickel, arsenic, thallium and antimony at about 50 °C; the second is the removal o f cobalt at above 70 °C in the presence of "activators", either antimony, antimony-copper or arsenic-copper, for successful cementation. M2++ Zn = M + Zn2+ (2-9) 5. Zinc is recovered from the purified electrolyte by electrowinning onto aluminum cathode sheets. The zinc deposit is stripped off the cathodes, melted and cast into bars or ingots for shipment. The generated sulphuric acid is recycled to the leach step. ZnSO, + H10 = Zn + H,SO,+\/ 2Q2 (2-10) £ o c cu Q. CO Zinc Concentrate I3Z Roasting I Calcine Leaching Calcine Hydrolysis • pH 4.5-5.1 Zn Dust Cold Stage Cementation 50-60 a C Sb(As), Cu Zn Dust Hot Stage Cementation 70-80 2 C Electrowinning -SO2 Fe, Ga, In, Sb "As, Hg, Sn .Cu,Cd,Ge, As.Ni Co Figure 2-1 Schematic Diagram of Hydrometallurgical Zinc Extraction 10 2.2.2 IMPURITIES IN ZINC E L E C T R O L Y T E Electrolyte purification is a key operation for the economical extraction of zinc by the hydrometallurgical method. Impurities in the electrolyte can lead to: 1. A decrease in the current efficiency during electrowinning [75, 76]. 2. Formation of dendrites on the deposited zinc, leading to short circuits and stripping difficulties. 3. A decrease in the quality and purity of the final zinc product. Low levels of cobalt, copper, antimony, iron, cadmium, arsenic, germanium, tin, selenium, silver, bismuth, and barium in the electrolyte can affect the quality of the zinc product and the current efficiency [77, 78]. These impurities can codeposit with the zinc and become incorporated into the cathodic zinc, or can lower the hydrogen overpotential. Ohoyama et al [79] ranked the effects of impurities by the amount of H 2 evolved on the zinc cathode. Cobalt was considered to be one of the more detrimental impurities in the zinc electrolyte. At cobalt concentrations larger than 0.5 mg/L, cobalt alone reduces the zinc electrowinning current efficiency. However, the presence of cobalt can also exacerbate the effects of other impurities on current efficiency [80]. The range of impurities in some zinc plant electrolytes is listed in Table 2-2. Table 2-2 Impurity levels in various zinc plant electrolytes (mg/L). Impurities Cominco [81] National Zinc[82] V M Balen [83] Kidd Creek [80] Co 0.3 0.7 0.23 0.3 Cu 0.2 <0.1 0.2 0.1 Cd 0.3 2.3 5 <0.1 Fe <5.0 <0.1 15 5 Sb 0.03 0.01 0.01 <0.02 As - <0.01 0.01 <0.02 N i <0.05 0.3 0.01 <0.1 11 2.2.3 METHODS OF C O B A L T R E M O V A L F R O M ZINC E L E C T R O L Y T E The seriousness of cobalt impurities was realized as early as 1917, when the Electrolytic Zinc Company of Australasia attempted to treat the Broken Hi l l N . S. W. ores, which were high in cobalt content. Since then a number of methods have been patented and or employed for the removal of cobalt from zinc sulfate electrolytes, depending on the operating conditions in the zinc electrowinning plants. 2.2.3.1 C O B A L T R E M O V A L B Y C H E M I C A L OXIDATION 1. A strong oxidant, lead peroxide or calcium plumbate was used in the early 1920's to oxidize cobalt to its cobaltic state at 40-50 °C followed by precipitation or hydrolysis of Co2(S04)3 [84, 85]. Livingston and Field [86] reported that cobalt could be removed by agitating the electrolyte with lead peroxide or manganese dioxide, followed by precipitation of cobalt: 2CoSOA + Pb02 + H20 = Co203 + H2S04 + PbSO, (2-11) 2CoS04 + Mn02 + H20 = Co2 03 + H2 S04 + MnSO, (2-12) 2. Sodium hypochlorite is used to oxidize cobalt to its cobaltic state [87]. Ammonia has to be added as a neutralizing agent and the temperature is about 70 °C and The residence time necessary was about 2-3 hours. The reaction is: NaOCl + 2Co2+ + 2H20 = Co203 + NaCl + 4H+ (2-13) 2.2.3.2 PRECIPITATION METHODS 1. Clevenger patented two methods in 1918 [88, 89]: a process of cobalt removal during iron purification with the addition of permanganates or manganates and a precipitation process for cobalt as the red cobalt-nitroso-beta-naphthol with nitroso-beta-naphthol or both beta-12 naphthol and sodium nitrite, added together or separately. Alpha-nitroso-beta-naphthol (in the form of sodium naphthenate and sodium nitrite) has been used to precipitate cobalt from zinc sulfate solutions commercially. Zinc electrolytic plants in Australia make use of this technology [90]. The precipitated cobalt-nitrosonaphthate is recovered and oxidized at elevated temperatures to a cobalt oxide. Richard et al [91] reported that when alpha-nitroso-beta-naphthol dissolved in a minimal amount of ethanol is added to the zinc sulfate solution, 97-99 % of the cobalt is precipitated. 2. Cobalt can be removed from zinc sulfate solution with organic nitrogen-containing complexing agents in 100 - 120 % stoichiometric excess. Suitable agents are ethylenediamine, hydroxylamine or urea [92]. The reaction is carried out in the presence of barium oxide or sulphide, in an amount equal to 3 times that of the cobalt present. The compound is added after the complexing agent in the form of a powder [92]. 3. The xanthate method in which cobalt is precipitated as cobalt xanthate was widely used in Russia [93]. Cobalt is precipitated in the form of cobalt xanthate. 2.2.3.3 SOLVENT E X T R A C T I O N M E T H O D Solvent extraction has been studied to remove cobalt from zinc sulfate solution containing 190 g/L Zn and 50 ppm cobalt and the following complexing reagents gave the best results: alpha-nitroso-beta-naphthol, beta-nitroso-alpha-napthol, 1, 2-cyclohexane dione dioxime (nioxime), di-2-pyridyl ketone oxime, and dimethyglyoxime [91]. These reagents in suitable solvents (kerosene plus isodecanol) lowered the cobalt and nickel levels to < 0.1 ppm. Stripping of the complexed reagents from the loaded organic phase was not possible except for nioxime. This method has not been applied commercially. 2.2.3.4 C O B A L T CEMENTATION WITH ZINC DUST Because of the favourable redox characteristics of zinc metal, it is used to electrochemically precipitate less electropositive impurities from zinc electrolyte. Of all the 13 impurities in zinc electrolyte, cobalt has been found to be one of the most troublesome to remove by cementation, despite favourable thermodynamics. The cobalt cementation equilibrium constant is 2><1016, which suggests that cobalt should be completely removed from the solution with zinc dust and no back reaction would be expected [94, 95]. However, cobalt cementation is very slow and unless special measures are taken such as the addition of activators, e.g. arsenic or antimony trioxide, the Co 2 + /Zn reaction is practically useless for the removal of cobalt from the zinc sulfate electrolyte. As shown in Fig. 2-2, the cementation of cobalt with zinc dust is slow even in the absence of zinc ions and the addition of only 1 g/L zinc ions to the solution slows down the rate of cobalt cementation significantly. The pH of the zinc electrolyte is usually controlled in the range of 4-5. Too low a pH (<4) increases the rate of the competing hydrogen reduction reaction and too a high a pH (> 5) can result in the formation of hydroxide or basic zinc sulfate salts and thus passivation of the residue's surface. 1.0 0.9 0.8 0.7 0.6 0.5 H 0.4 0.3'I 0.2 OH 0.0 o • [Zn2+]=lg/l o [Zn2+]=0 — i — 100 300 Time (minute) Figure 2-2 Reaction data for the isothermal cementation of cobalt onto zinc in the presence of additional zinc ions [93] The following are some of the technologies which have been applied to cobalt removal by zinc cementation. 1. Cobalt cementation with zinc dust at 75 °C in the presence of copper and arsenic was patented in Australia in 1919 [96]. The process is believed to take place via precipitation of 14 cobalt arsenide [94, 97-99]. In the copper-arsenic process, typical additions of copper and arsenic to the electrolyte are, respectively, 200 mg/L and 50-200 mg/L and the temperature is above 90 °C [94, 97]. It has been proposed that a galvanic cell is formed between zinc dust and copper particles. Cobalt is deposited on the copper by the galvanic current. Its deposition is improved greatly when it co-deposits with arsenic [100]. 2. In 1946, antimony compounds were found to be substitutes for AS2O3 with the advantage of lower toxicity, but cobalt cementation with copper-antimony activation requires more zinc dust than when copper-arsenic is used and proceeds at a lower rate [98]. Typical additions are 30 mg/L copper as copper sulfate and 1.5 mg/L antimony as antimony tartrate with 4 g/L zinc dust at a temperature above 70 °C [100]. The addition of antimony only, improves the cementation of cobalt to some extent, whereas copper by itself has an almost negligible effect on the rate. The combination of the two elements gives the best results. 3. A method for purification with pretreated zinc dust was patented in 1972 [101]. Atomized zinc dust containing 0.002-5 wt % Sb and 0.05-10 wt % Pb was used at 60-80 °C. A similar method was patented in 1975 in Japan [102]. The zinc dust containing antimony and 0.01-5% graphite was reported to improve the removal of cobalt from zinc sulphate solution [103]. 4. Addition of Sb (V) is more effective than the addition of Sb(III) in promoting the removal of cobalt ions from zinc sulphate solution by zinc cementation. The time of purification and the amount of zinc consumed are reduced [104, 105]. 5. Finely divided antimony powder was used as a substitute for soluble antimony. The addition of only 0.002 mg/L antimony dust could remove all cobalt at 90 °C when no copper was present [105, 106]. 6. Indium [108], tin [109], thallium [110], and mercury [111] compounds were found to accelerate cobalt cementation with zinc dust. 2.2.3.5 C O B A L T CEMENTATION WITH M A N G A N E S E DUST In order to increase the efficiency of cobalt removal from zinc electrolyte, manganese powder has been tested for the cementation of cobalt [112, 113]. Considering the more 15 favourable redox characteristics of manganese, the rate of the Co (H)-Mn reaction should be very high, but the reaction rate is still slow and the antimony and copper compounds have to be added to the solution to accelerate the reaction and prevent the redissolution of reduced cobalt. The cobalt removal rate using manganese powder was improved by a factor of 3 compared to using similarly sized zinc powder. However, fifty times the stoichiometric requirement for manganese powder was needed, and there was a strong tendency for the cemented cobalt to redissolve. So manganese powder has not been used commercially for the purification of zinc electrolyte. 2.2.3.6 C O B A L T R E M O V A L WITH SODIUM BOROHYDRIDE The removal of cobalt from zinc electrolyte has been studied and practiced for about 80 years. Now most modern plants in the world employ the zinc cementation process with arsenic-copper or antimony processes to remove cobalt from the zinc electrolyte. However, these zinc dust purification processes are still not satisfactory. The main problem is that these processes have to use a very large excess of zinc dust, a high temperature and a long purification time (more than 2 hours). Thus they are an expensive component of the zinc hydrometallurgy process. In order to realize a more economic extraction of zinc by hydrometallurgical processes, zinc production plants have been looking for a better method to replace the present processes. Sodium borohydride has been studied to remove cobalt from waste water and implemented as the best available technology for the treatment of such waste water. Polyakov [114, 115] reported the use of sodium borohydride to remove cobalt from zinc sulfate solution to 0.01 ppm. This purification process could have many advantages over the present process for example, the high reduction efficiency, the low purification temperature, the short reaction time, and the simplification of the purification process and equipment. However, Cominco Ltd.'s staff [116] reported that cobalt ions cannot be reduced from zinc sulfate solutions. In view of the conflicting reports it was deemed worthwhile to study carefully the possibility of cobalt removal with sodium borohydride from zinc sulfate solution. 16 2.3 CHEMISTRY OF B O R O N 2.3.1 P H Y S I C A L A N D C H E M I C A L PROPERTIES OF B O R O N Boron is a unique and exciting element. It is the only non-metal in Group HI of the periodic table and shows many similarities to its neighbor, carbon and its diagonal relative silicon. Thus, like carbon, and silicon, it shows a marked propensity to form covalent, molecular compounds, but it differs sharply from them in having one less valence electron than the number of valence orbitals, a situation referred to as "electron deficiency". This has a dominant effect on its chemistry. The chemistry of boron is discussed in a number of books [117-123]. Boron has two stable naturally occurring isotopes 1 0 B and U B . The physical properties of elemental boron are greatly affected by its complex polymorphism and contamination by irremovable impurities. Boron is an extremely hard refractory solid of high melting point, low density, and very low electrical conductivity. Crystalline forms are dark red in transmitted light and powdered forms are black. The most stable (B -rhombohedral) modification has a melting point of 2180 °C, a boiling point of 3650 °C, and a density of 2.35 g cm"3 [119]. The chemical reactivity of boron itself obviously depends markedly on the purity, crystallinity, state of subdivision, and temperature. Crystalline boron is extremely inert chemically. It is unaffected by boiling HC1 or HF, only slowly oxidized by hot, concentrated nitric acid when finely powdered, and either not attacked or only very slowly attacked by many other hot concentrated oxidizing agents. Finely divided amorphous boron can be pyrophoric in air, whereas massive crystalline boron oxidizes only slowly at 800 °C. Boron does not react directly with hydrogen or the inert gases, but reacts with all other non-metals (except germanium and tellurium) under suitable conditions. It also reacts readily and directly with all metals at elevated temperatures to form borides. Boron reacts with a number of oxides at elevated temperatures to form B2O3 or its derivatives. Boiling water reacts slowly with powdered boron of moderate purity, and the reaction can become violent at red heat. At 830 °C, a stream of water passed over boron at 10~3 to 10"6 atm produces H3BO3. A powdered mixture of boron and B 2 0 3 is stable at 900 °C in vacuum. 17 2.3.2 B O R O N COMPOUNDS Five types of boron compounds can be distinguished, each having its own chemical characteristics which can be rationalized in terms of the type of bonding involved with each resulting in highly distinctive structures and chemical reactions: 1. Metal borides ranging from M5B to MB66-2. Boron hydrides and their adducts and derivatives including carboranes and polyhedral borane-metal complex. 3. Boron trihalides and their adducts and derivatives. 4. Oxo compounds including polyborate, borosilicates, peroxoborates. 5. Organoboron compounds and B - N compounds. 2.3.2.1 BORIDES The borides comprise a group of over 200 binary compounds which show an amazing diversity of stoichiometries and structural types: M 5 B , M 4 B , M 3 B , M7B3, M 2 B , M B , M B 3 ; M B 6 6 , etc. There are also numerous ternary and more complex phases. Metal-rich borides are extremely hard, chemically inert, involatile, refractory materials with melting points and electrical conductivities which often exceed those of the parent metals. Borides are normally prepared as powders. The ability to withstand attack by molten metals and salts has commended borides or boride-coated metals as high-temperature reactor vessels, vaporizing boats, etc.. There are about eight general methods to prepare borides: 1. Direct synthesis from the elements or boron and a metal hydride at about 1000 °C. 2. Reduction of metal oxides with boron. 3. Co-reduction with hydrogen of a volatile boron compound and of a volatile compound of the metal. 4. Reduction of BCI3 or other halides with a metal. 5. Fused-salt electrolysis of a metal oxide with a source of boric oxide at 700 - 1000 °C. 6. Co-reduction of oxides with carbon at temperatures up to 2000 °C. 7. Reduction of metal oxides (or M + B2O3) with boron carbide. 18 8. Co-reduction of a metal oxide and boric oxide by a metal such as A l and Mg. Cobalt boride can be made by the direct combination of powdered Co and B at 1000 °C , and electrolytic deposition from fused salt media. Cobalt has the following three main types of borides: C04B (orthorhombic), C 0 3 B (orthorhombic), C 0 2 B (tetragonal), CoB. Figure 2-3 is Co-B phase diagram. Cobalt boride is not attacked by HC1, or H2SO4 solutions but is soluble in hot nitric acid solutions. Weight Percent Boron 0 10 20 30 40 50 60 70 80 100 2200 I , 1 i •••>•, '• • i I '• i I ...«•••••• I • 1 6 10 20 30 40 50 60 70 80 90 100 Co Atomic Percent Boron B Figure 2-3 Co-B phase diagram [124]. 2.3.2.2 O X Y G E N COMPOUNDS OF B O R O N These are among the most important compounds of boron, comprising nearly all of the naturally occurring forms of the element. 1. Boron oxide. The principal oxide is boric oxide, B2O3, which is formed when boron is heated in air or oxygen, but is generally prepared by dehydration of boric acid. Boric oxide dissolves in water with considerable evolution of heat to give a solution of boric acid, B(OH) 3 . 2. Boric acids. The principal oxoacid of boron is orthoboric acid B(OH)3. The other boric acids are metabolic acids, H B O 2 , polyboric acid, B303(OFf)4~ etc.. Boric acid is moderately 19 soluble in water with a large negative heat of solution so that the solubility increases with temperature. It is a very weak and exclusively monobasic acid that is believed to act, not as a proton donor, but as a Lewis acid, accepting OH"[l 19]: B(OH)3+ H2Ot> B(OH)~ + i T pk=9.25 (2-14) A t concentrations <0.025 M , essentially only the mononuclear species B ( O H ) 3 is present; but at higher concentrations the acidity increases, and p H measurement is consistent with the formation of polymeric species such as [118]: 3B(OH)3oB,03(OH)A- +H+ +2H20 pk = 6.84 (2-15) When heated, boric acid loses water in stepwise fashion: Heal Heat ,~ . ^ N . B(OH),<^ HB02 oB20, 1V V 5 H70 H20 2.3.2.3 B O R O N H Y D R I D E ( B O R A N E ) A N D T E T R A H Y D R I D E S Boranes are colourless, diamagnetic, molecular compounds of moderate to low thermal stability. The lower members are gases at room temperature but with increasing molecular weight they become volatile liquids or solids. The boranes are all endothermic and their free energy o f formation is also positive; their thermodynamic instability results from the exceptionally strong interatomic bonds in both boron and H 2 rather than the inherent weakness of the B - H bond. Boranes are extremely reactive and several are spontaneously flammable in air. They are extremely versatile chemical reagents but the very diversity of their reactions makes a general classification unduly cumbersome. Diborane (B 2 H6) occupies a special place because all o f the other boranes are prepared from it. It is also one of the most studied and synthetically useful reagents in the whole of 20 chemistry. It is spontaneously flammable in air and has a higher heat of combustion per unit weight of fuel than any other substance except H2. Borane, B H 3 , appears to have a transitory existence during the thermal decomposition of diborane: B 2 H 6 < f a s t > 2BH, (2-17) BH3 reacts rapidly with diborane to give B3H9. The bridge bonds in B2H6 are readily cleaved, even by weak ligands, to give either symmetrical or unsymmetrical cleavage products, e.g., B2H6 + 2H~ -> 2BH4~ (2-18) BH4" itself can be unidentate, bidentate, or tridentate ligand as illustrated by its complexes. 2.3.2.4 M E T A L B O R O H Y D R I D E Borohydride is also called tetrahydroborate. The elements known to form such compounds are shown in Figure 2-4 [123]. LL Be N a M g A l K C a [Ga] Rb Sr [In] [Sn] Cs B a li(iv) Cr [ M n l [Fe] [Co] N i [Cu]* Z n Y Zr(IV) N b Hf(IV) [Ag]* [Cd] L a Sm E u G d Tb D y Ho Er T m Y b L u Th U N p Figure 2-4 Elements forming metal tetrahydroborate compounds. Asterisks indicate compounds stabilized at room temperature by coordination with phosphine ligands. 21 The elements underlined form simple hydroborates, M ( B H 4 ) n , which are stable or decompose only slowly at ambient temperatures; those not underlined have only been isolated with other groups, in addition to the hydroborate group, attached to the metal, while those in brackets form compounds which have been reported to be unstable at room temperature but may be isolated at lower temperatures. The stability of metal hydroborate compounds may be related qualitatively to the electronegativity of the central metal atom. Compounds involving those elements possessing an electronegativity higher than that of boron are unstable or do not exist e.g., Si(BH 4 ) 4 and P(BH 4 ) 3 and attempts to prepare them result in the corresponding hydride being formed, e.g. S i H 4 and PH3. In these compounds the element of higher electronegativity competes for the hydride ion of the hydroborate group, BH3»H~, more effectively than the borane group, B H 3 . In the transition metal series the stability of the compounds decreases across the series as the electronegativity increases. Thus, for example, while compounds of titanium, zirconium and hafnium are known, those of nickel, cobalt or copper only exist at ambient temperatures when ligands are also bonded to the metal atom, essentially reducing its electronegativity (e.g. [Co(NH3)g] (BH4)2»NH3). The stability of metal hydroborates has also been discussed in relation to their percentage ionic character, and those compounds with less ionic character than diborane are expected to be highly unstable [123]. The polarizability of the tetrahydroborate ion has been calculated to be 3.9±0.1A, which is similar to that of the bromide ion (4.16 A) [126]. The tetrahydroborates therefore vary in character from the ionic salt-like compounds of the alkali metals, which are only slowly hydrolyzed, to typically covalent volatile compounds such as those of aluminium, and zirconium, which show similar properties to diborane reacting explosively with air and being rapidly hydrolyzed. Of all of the borohydrides, alkali metal borohydrides are most widely studied and used and their properties are better understood. The major reaction of alkali metal borohydrides can be divided into three types: metathetic reactions, protolysis reactions and substitution reactions. The metathetic reactions are those in which the borohydride group is transferred from one element to another. A metal borohydride reacts with another metal compound (MX) in a suitable solvent precipitating the hydroborate derivative of M , or the metal hydroborate formed is more 22 covalent and often volatile in nature, probably involving direct bonding between the metal and borohydride group in the form of hydrogen bridge bonds, e.g., NaBH4 + RbOH-> RbBH4 + NaOH (2-19) 3NaBH4+AlCl3 -> Al(BH4 )3 + 3NaCl (2-20) 2NaBH4 + ZnCl2 -> Zn(BH4 )2 + INaCl (2-21) Sodium borohydride is the cheapest and most widely produced of all metal borohydrides. It is produced in three forms: powder, pellets and SWS™ (12 % solution of sodium borohydride in caustic soda). Some of physical and chemical properties are listed in Table 2-3. Table 2-3 Some physical and chemical properties of sodium borohydride [122] Melting point (°C) 505 Decomposition temperature (°C) 315 * 1 Density (g/cm ) 1.08 Enthalpy of formation, (kcal/mol ,298 K) -43.8 Entropy (cal/mol., 298 K) -30.19 Structure Face centered cubic a Na-B distance, (A) 3.08 Zinc borohydride is a white, ether-soluble solid which quantitatively decomposes into the elements at > 85 °C [125]. The preparation of zinc borohydride has been reported in ether or tetrahydrofuran by metathetic reactions involving zinc chloride with lithium, sodium or potassium hydroborate, and by the action of dioborane on zinc hydride. It reacts vigorously with water [123]. 23 2.4 AQUEOUS CHEMISTRY OF SODIUM BOROHYDRIDE 2.4.1 AQUEOUS SOLUTION OF SODIUM BOROHYDRIDE 2.4.1.1 SOLUBILITY A N D STABILITY Sodium borohydride forms a dihydrate, NaBH4» 2H2O, which may be crystallized from cold aqueous solution. The stability regions of the two crystalline forms, N a B H 4 and NaBH4»2H20, are shown in Figure 2-5. The curve below 36.4 °C shows the solubility of NaBH4»2H20 while that above 36.4 °C represents the solubility of NaBH 4 . Temperature, °C Figure 2-5 The solubility of sodium borohydride in water at different temperatures [123]. The stability of sodium borohydride in water is dependent on the temperature and pH. The hydrolysis reaction is accelerated by increasing the temperature and lowering the pH. As the borohydride is basic, the higher the concentration, the more stable the solution. The pH of the solutions of sodium borohydride is shown in Table 2-4. The hydrolysis of sodium borohydride causes a rise in pH, and the rate of decomposition therefore decreases. For example, a 0.01 M solution of N a B H 4 has an intitial pH of 9.6 which changes during hydrolysis to 9.9. It is obvious that the addition of sodium hydroxide will stabilize aqueous sodium borohydride. 24 Table 2-4 pH of solutions of NaBFL at 24 UC [5] Concentration of NaBFL, (M) pH 1.000 10.48+0.02 0.100 10.05+0.02 0.010 9.56±0.02 2.4.1.2 THERMODYNAMIC PROPERTIES OF BOROHYDRIDE Sodium borohydride is an ionic salt-like compound which will dissociate into sodium ions and borohydride anions in water [125]. NaBH4 • 2H2Q o Na+ + BH~ + 2H2Q A G ° = -4070 cal K e q u u i b n u m = 966 at 25 °C (2-22) The equilibrium constant is 1.4 x 104 at 25 °C, indicating that most of the sodium borohydride is dissociated. In its physical behaviour, borohydride ion resembles the halide ions much as ammonium ion resembles the alkali cations. Its ionic radius of 2.03 A lies between those of bromide and iodide ions. However, its polarizability of 3.94A is slightly less than that of the bromide ions. Figure 2-6 shows the Eh-pH diagram for the B-H2O system at 25 °C. The following species are considered: B, BH4", HBO2, H3BO3, B(OHV, and B2O3. The species BO2" is often referred to in the literature. However, spectroscopic studies have shown that BO2" is hydrolyzed to B(OH)4_ [121] and so B0 2 " is not considered. Some polymeric species are not considered because under the practical reaction conditions, the concentration is much lower than 0.025 M . At pH < 9.24, the product of the oxidation of borohydride is H3BO3. Its redox reaction and potential are described by the following equations. NaBH4 <=> Na+ + BH4 A G ° = -5660 cal K ^ n t , ™ ™ = 1.4 x 104 at 25 °C (2-23) 25 BH' +mO = H,B03 + 1H+ + Se Eh = -0.481 - 0.0511 pH (2-24) (2-24A) At pH > 9.24, the product of the oxidation of borohydride is B(OH) 4", but for simplicity, B(OH) 4" is often written as B0 2 ~. The redox reaction and potential are expressed by the following equations, BH~ + 80H~ = B(OH)~ + 4H20 + Se Eh = -0.413 - 0.059\pH (2-25) (2-25A) 2.0 1.5 1.0 0.5 - .5 - 1 1 1 1 1 1 -H 20 ~~ " — • _ H3BO3 - - B(OH)4" -— ~ ~ ~ — - ____ H 2 ~~ ~~ ~- — - BH 4 1 1 1 1 1 1 6 8 P H 10 12 14 Figure 2-6 E h -pH diagram for the B - H 2 0 system at 25 °C [127, 128]. The activity of species 0.0001. 26 2.4.1.3 HYDROLYSIS OF BOROHYDRIDE IONS In spite of extensive studies on the hydrolysis of the borohydride ion, the details of the mechanism of the reaction and the intermediates involved, remain obscure. The ion is stable in alkaline solution since its salts are capable of recrystallization from such a medium. Hydrogen evolution becomes more rapid as the pH of the solution is decreased, and several early reports comment on the importance of pH on the hydrolysis reaction [123]. The hydrolysis is described by the following reaction: BH~ +H+ + 3H20 = H3BO, + 4H2 (2-26) The borohydride ion undergoes acid catalyzed hydrolysis. Mesmer and Jolly [129] have studied the kinetics of the hydrolysis over the pH range 3.8-14 and have shown that the rate may be represented by the equation: where k] = 2.18 x 101 1 T exp(-4000/T) mole" min"1 and k 2 = 1.72 * 107 T exp(-10380/T) min"1. They measured the isotope effect in the hydrolysis reaction using heavy water and concluded that the most probable mechanism was one involving the initial formation of BH5 in the rate-determining step, i.e. d[BH4~]/dt k{[H+][BH4-] + k2[BH4~] (2-27) BH4 + H+ slow > BH5 > BH3 + H2 (2-28) BH3 + 3H20 fast > H3B03 + 3H2 (2-29) Kreevoy [130] reported that borohydride undergoes acidic and non-acidic hydrolysis and proposed the following reaction mechanism: 27 H+ +BH~Z>H2BH, (2-30) H20 + BH; Z H2BH3 + OH' (2-31) k-H20 H2BH2 — H 2 + BH2 (2-32) BH2+3H20 v e r y f a s t >3H2+B(OH)3 (2-29A) where k H = 9.9><105 M _ 1 s e c " k _ H 2 0 = 2.2* IO' 7 sec"1 and k . H 2 o /k 2 = 2.3 at 25 °C. Schlesinger first reported that some metal ions act as catalysts for the hydrolysis of borohydride ion [3]. Cobalt and nickel are particularly effective. The action o f these compounds was attributed to the formation of the respective borides, C o 2 B and N i 2 B , which served as the catalysts. The relative effect o f various salts has been studied by Brown and Brown [131]. Their results are tabulated in Table 2-5. Table 2-5 Half-lives for hydrolysis of borohydride ion in the presence of transition ions [131] Metals Compounds Time, min. Iron F e C l 2 38 Cobalt C o C l 2 9 Nicke l N i C l 2 18 Ruthenium R u C l 3 0.3 Rhodium R h C l 3 0.3 Palladium P d C l 2 180 Osmium O s 0 4 18.5 Iridium I r C l 4 28 Platinum H 2 P t C l 6 1 28 Gomez-Lahoz [41] reported that borohydride undergoes cobalt-catalyzed and non-catalyzed hydrolysis. The following kinetic equation was proposed for the hydrolysis: J [ g f 4 " ] = kx[BH4][H+] + k2[BHf][H+][Co2B] (2-33) dt 2.4.2 REDUCTION OF M E T A L IONS WITH BOROHYDRIDE The metal reduction reaction with borohydride is sensitive to the reaction conditions. In neutral and acidic solutions (pH < 9.24), the reduction reaction of divalent metal ions is generally described by the following equation: BH~+ 4M2+ + 3H20 = H3B03 + 4M + 1H+ (2-34) In alkaline solutions (pH > 9.24), the reduction reaction is expressed by the equation: BH~ + 4M2+ + SOH~ = 4M + B(OH)~ + 4H20 (2-35) What usually takes place is a combination of reduction and hydrolysis which depends on reaction conditions, concentration, mixing efficiency and the kinetics of the competing reactions. Therefore, the combined reaction at pH < 9.24 can be expressed as: BH4~ + 3H20 + xM2+ = H3B03 + xM + (4-x)H2 + (2x-\)H+ (2-36) where x is less than 4. For some noble metals such as silver and gold, the reduction products are pure metals and the reduction reaction can be expressed by the equations (2-34) and (2-36) [132-137]. 29 The reaction with more electropositive metals such as Co, Fe, N i , gives highly dispersed black products which contain metal and boron [7, 14, 15, 41, 42, 47, 135-138]. There is some adsorbed hydrogen on the products. They are usually called borides because their composition is similar to that of known borides. The reduction of cobalt with borohydride is generally expressed as: 4BH4~ + 2Co2+ + 9H20 =Co2B + 3B(OH)3 + \2.5H2 (2-37) The formation of the borides correlates with the catalytic effects of Co, N i and Fe. The other metals have little or no catalytic effect and the precipitates contain very little boron. The hydrolysis of borohydride is also catalyzed by Co, N i , and Fe. The stoichiometry of borohydride reduction depends on the experimental conditions, the electronegativity of the metal and catalysis on borohydride, pH, and the form of the metal salt in solution. Silver has a minimal catalytic effect on the hydrolysis of borohydride and the stoichiometry of silver (I) reduction with borohydride in different media (pH 3-10, and different water-soluble compounds) can be almost entirely expressed by equation (2-34). Silver metal is produced [134]. The reduction of arsenic, antimony, and bismuth compounds, such as trihalides or oxides of these elements, e.g. (C 6 H 5 ) 2 AsCl3, SbCl 4 \ K(SbO)(C 4 H 4 0 6 ) , H 3 Sb0 3 , gives their hydrides, e.g. A s H 3 , AS2FJ.6, SbH 3 , Sb 2 H 6 , and B i H 3 , which partially react further with their compounds to produce elemental Sb, As [5, 123, 136, 139-141]. 3BH4~ + 4H3Sb03 + 3H+ -> 4SbH3 + 3H3B03 + 3H20 (2-38) The reaction with zinc ions produces zinc borohydride which decomposes rapidly into hydride and borane. Zinc hydride is stable, but will hydrolyze. In acid solution, zinc enters the solution in the form of zinc ions and no precipitate is formed [136]. Zn2+ + 2BH4- -> Zn(BH4)2 (2-39) 30 Zn(BHA)2 -> ZnH2 + 2BH3 (2-40) ZnH2 + 2H+ -> Zn 2+ (2-41) If the pH is higher than that required for the formation of zinc hydroxide, zinc is precipitated as zinc hydroxide. Khain [132] reported that the reactions of borohydride ions with metal ions must be classified as reactions taking place via bridged active complexes, that is, by an inner-sphere mechanism. The first stage is the replacement of ligands by BH4" ion with a subsequent stage involving a spontaneous redox reaction or electron transfer of the intermediate complex. The role of the bridging ligand is played by the hydride ion, H". The entire process can be represented as series of successive stages: The formation of the intermediate compound, Mz+(aq) + nBH- -> M(BH4) + (aq) (2-42) The formation of a bridging bond, M(BHA)n ^M--+-(H--BH,)n (2-43) The transfer of a hydrogen atom and rupture of the bridging bond, Mz+ - (H~ - BHi)n -> MH;-n + nBH, (2-44) The electron transfer and bond rupture, MHnz-" -> M + (z - n)H+ + (2n - z) / 2H2 The catalytic decomposition of borane, (2-45) 31 BH-. M >B + 3/2H2 (2-46) The hydrolysis of borane, BH3 + 3H20^> B(OH)3 + 3H2 (2-47) The increase in the alkalinity of the medium leads to an increase in the hydrolysis of the metal ions, and more stable hydroxocomplexes or inert polynuclear ions M q(OH) p ( q z" p ) are formed. In this case, initial M(BH 4) n z~ n formation is suppressed, and the reactions with BH 4" become slower (M = Pt, Pd, Rh, Au) or practically stop (M = Pb, Sn, Cu). 2.4.3 AQUEOUS CHEMISTRY AND BOROHYDRIDE REDUCTION OF COBALT 2.4.3.1 AQUEOUS CHEMISTRY OF COBALT The principal oxidation states of cobalt in aqueous solutions are +2 and +3. In the absence of ligands other than water, the solutions of cobalt (II) contain the pink, high-spin, hexa-aquo ion [Co(H20)6]2+[142], which is the thermodynamically most stable state. Cobalt (III) is not thermodynamically stable in the absence of ligands. The redox reaction and potential for the Co(II)/Co couple are: Co 2+ + 2e = Co (2-48) Eh = -0.277 - 0.0591 log(a C o2 + ) (2-48A) 32 Figure 2-7 shows the E h-pH diagram for Co-H 20 system at [Co2+] = 0.5, 0.05, 0.005, 0.0005, 0.00005 M (assuming concentation = activity). From these diagrams, for [Co ] = 0.5, 0.05, 0.005, 0.0005, and 0.00005 M , the initial pH's for the precipitation of Co(OH)2 are 6.6, 7.1, 7.6, 8.1, and 8.6 respectively. Therefore the cobalt reduction reaction should be carried out at a pH below that at which Co(OH)2 will precipitate. Otherwise cobalt metal will not be produced. Furthermore, at pH's below 4.82, 5.32, 5.82, 6.32, 6.82 respectively for [Co2+] = 0.5, 0.05, 0.005. 0.0005, and 0.00005 M , cobalt metal may be oxidized by FT. 2 0 i 1 1 ' 1 1 1 1 0 2 4 6 8 10 12 14 pH Figure 2-7, Eh - pH diagram for the Co-H 20 system at Co activities of 0.5, 0.05, 0.005, 0.0005, and 0.00005 and 25 °C [127, 128]. 33 Cobalt ions in aqueous solutions hydrolyze and form hydroxocomplexes [143-145]. A s in the case of the hydrolysis equilibrium reactions for other metallic ions, the following reactions have been assumed to determine the equilibrium constants for the first ionization step of the hydrolysis of cobaltous ions: Co(aq)2+ + H20 <=> Co(OH)+(aq) + H+ (aq) (2-49) The equilibrium constant K a , = 1.58 x 10"1 0 [140]. Co(OH)+ + H20 o - Co(OH)2(aq) + H\aq) (2-50) The equilibrium constant K a 2 = 1.26 x 10"9 [143]. A t higher concentrations, cobalt ions polymerize and form C o 2 ( O H ) 2 2 + (pk = 9.44) and C o 6 ( O H ) 6 6 + (pk A = 42.55) [146]. For cobalt sulfate solution, there exist the following equilibrium reactions (147-149]; Co2+ + S042~ <=> CoS04(aq) log K =2.69 at 25 °C [149] (2-51) H+ + S042~ <=> HS04~ log K = 2.0 at 25 °C [148] (2-52) 2.4.3.2 B O R O H Y D R I D E R E D U C T I O N OF C O B A L T I O N Borohydride reduction of cobalt ion has been used to prepare magnetic materials and catalysts, and remove cobalt from waste water and therefore the reaction has been studied. It is desirable to establish the mechanism of the reductions in order to determine the best conditions for the various applications. However relatively little systematic work has been done in this area. In the case of metal boride formation, such studies would be very difficult because o f the extreme 34 rapidity with which the boride is generated. The following several reaction mechanisms have been proposed to explain the reaction product. 1. Mochalov first proposed a cobalt reduction mechanism according to his results [135]. The composition of the products obtained by the reaction of cobalt (II) with borohydride was expressed as C 0 2 B with some adsorbed hydrogen also present. However, the product behaved like a mixture of cobalt metal and elemental boron. Unlike the authentic boride, the products had non-stoichiometric compositions which varied with the experimental conditions. A t 0 °C, 25 °C and 75 °C, the ratios of Co to B were 1.84, 2.05, and 2.48 respectively. Authentic C 0 2 B is kinetically inert and dissolves slowly in HC1 solution (1:1) with the evolution o f hydrogen. However, the product of the borohydride reduction of cobalt ions dissolved readily in HC1 solution (1:1) and the amount of evolved hydrogen corresponded to the contents o f free metal and boron. Even boiling water dissolved the product and the same stoichiometric hydrogen was obtained. The precipitate was readily oxidized in air and was pyrophoric. If the product was exposed to water and air, it gradually changed into hydroxide and boric acid. The product was able to cement Pb, Cu , A g , and Pd from solutions of their salts. It readily reduced M n 0 3 ~ , NO3", I, T i 4 + , F e 3 + and was used in analytical chemistry. When covered with water or other liquid, the precipitate was ground with mercury to produce an amalgam containing 8-10 % Co. The mercury was distilled and free cobalt metal precipitated. X-ray investigations of the products produced under many conditions showed that they were amorphous and did not show any spectral features of authentic boride. However, when the product was boiled in water, the clear alpha and beta cobalt phases were apparent from the X-ray diffraction pattern. Mochalov [131] concluded that the precipitate produced by the reduction of cobalt (II) with borohydride was a mixture of highly dispersed cobalt metals and boron. Except for the similarity in composition to real cobalt boride, the product had different properties and might be considered to be a pseudo-boride. 35 The proposed mechanism involves the initial formation of cobalt borohydride. Cobalt (II) strongly polarizes the borohydride, resulting in cobalt hydride and borane. The produced cobalt hydride is unstable and decomposes into cobalt metal and hydrogen. Catalysis by water may be involved. Borane (BH3) can undergo two reactions: hydrolysis and decomposition into boron and hydrogen with the mixing metal cobalt to form the pseudo-boride due to the catalysis of highly dispersed metals. Dependent on the relative rates of above two competing reactions related to B H 3 , the composition of the pseudo-boride should be variable and lie between C02B and Co. The reaction could be affected by temperature and solvents (water or ether). At room temperature, the compositions of the pseudo-cobalt borides approach that of C02B. The reaction mechanism is expressed as: 2Co2+ + ABH4 -> 2Co(BH4)2 (2-53) 2Co(BH4)2 -> 2CoH2 + 4BH3 (2-54) 2CoH2 -> 2Co + 2H2 (2-55) BH3 CMalyzed > B + \.5H2 (2-56) 3BH3 + 9H20 -> 3B(OH)3 + 9H2 (2-57) 2 C o 2 + + 4BH4~ + 9H20 = 2CoB + 3B(OH)3 + \25H2 2. Glavee et al. [14, 15] reported that the reduction of CoCi2(aq) with sodium borohydride yields ultrafme C02B as the primary product i f the two reagents are rapidly mixed and the product is handled under argon. The procedure must be followed rigorously. Concentrations, B H 4 - M ion ratio, pH, method of mixing, and rate of mixing can determine the composition particle size and reactivity of the product. A broad main peak of Co 2 B appeared on 36 the X-ray diffraction pattern of the fresh precipitate suggesting amorphous C 0 2 B . After heat treatment at 500 °C under argon, the precipitate was identified as C 0 2 B by X-ray diffraction. The overall reaction can be expressed by equation (2-37). The following reaction mechanism was proposed: Co(H20)62+ o [(H20)5CoOH]+ + H+ (2-58) [(H20)5CoOH]+ + H+ + BH~ -> H2 + [(H20)5 Co - OH - BH3 ] + (2-59) [(H20)5Co-OH-BH3]+ + Co(H20)2+ -> 3+ H2 + [(H20)5Co-OH-BH2-OH-Co(H20)5] 3BH4- + [(H20)5Co-OH-BH3-OH-Co(H20)5r -> (H20)5Co-OH-BH2-OH-Co(H20)5 + 3BH3 + \.5H2 (2-60) (2-61) 3 5 ^ + 9H20 m p i d ) 3B(OH)3 + 9H2 (2-62) [(H20)5-Co-OH-BH2-OH-Co(H20)5] -> Co2B + \2H20 (2-63) 2Co(H20)62+ + ABHA~ + 9H20 Co2B +\2.5H2 + 3B(OH)3 The equilibrium between C o ( H 2 0 ) 6 and [ C o ( H 2 0 ) 5 O H ] 2 + is a well established example o f the reason why aqueous solutions of transition metal ions are acidic. In the second step, H + and BH4" may react with the [(H20)5CoOH]+ species. This hydrated cobalt-borane adduct would then react with another [Co ( H 2 0 ) 6 ] 2 + species. The resulting species ([(H20)5Co-OH-BH2-OH-C o ( H 2 0 ) 5 ] 3 + ) is triply charged. Reduction could occur by reaction with 3 more equivalents of BH4" with concomitant formation of more H2 (Equation 2-61). Finally, in the sixth step, the adduct could yield C02B which precipitates. The sum of all these steps yields the overall equation (2-37) with the experimentally observed stoichiometry based on Co and BH4", and the 37 formation of C02B. However, the 12.5 moles of H2 predicted is a lightly lower than the values observed. Borohydride hydrolysis was invoked to explain this. In fact the hydrogen adsorbed on the precipitate affected the collection of hydrogen gas. The above reaction mechanism is really hard to realize because the intermediate products [ ( H 2 0 ) 5 C o - O H - B H 2 - O H - C o ( H 2 0 ) 5 ] 3 + and ( H 2 0 ) 5 - C o - O H - B H 2 - O H - C o ( H 2 0 ) 5 are almost impossible to form and the reaction step 2-61 would require more than one step. When the reduction was carried out in a nonaqueous medium (diglyme), the primary products were cobalt metal, H2 and B 2 H 6 [150]. In the presence of the catalytic cobalt particles, some o f B2H6 was catalytically decomposed into boron and H2 and the precipitate contained less boron than that formed in aqueous media. This suggests that water also plays an important role in the formation of C02B. 3. Gomez-Lahoz et al [41, 42] reported that the reduction of cobalt ions with borohydride released H + , the molar ratio o f C o to B in the precipitate was 2, and borohydride underwent hydrolysis. Therefore the following reactions were proposed: 2Co1+ + BH- -> Co2B +l/2H2 + 3H+ (2-64) 3BH4~ + 9H2Q +3H+ -» 3H3B03 + 10H2 (2-65) 4BH4~ + 2CoI+ + 9H20 -» 3H3B03 + Co2B + \25H2 The sodium hydroxide in stabilized borohydride solution was found to be the main factor in efficiency improvement, due to the inhibition of the hydrolysis of borohydride. 4. Dragieva et al [151, 152] reported that the amount o f boron incorporated in the powders depended on the concentration of sodium borohydride in the reducing solution. It was impossible to establish the phase composition of the amorphous powders using electron and X -ray diffraction. The crystallization of powders was carried out at 650 K for one hour under 38 vacuum (10~6 Torr). The cobalt powders contained about 25 at. % B or had an approximate composition C03B. Therefore the following reaction was proposed for the reduction of metal ions: The addition of complex-forming agents usually decreases the amount of boron incorporated in the precipitate. The role of the counter ions of the initial cobalt salt was investigated and the boron contents of powder reduced from C0SO4 (natural p H 6.7), C 0 C I 2 (natural p H 5.4) and C o ( N 0 3 ) 2 (natural p H 6.7) were 6.07, 5.57, 5.25 wt.% respectively. Wi th increasing p H of the initial salt solutions, the boron content of the powders increased considerably. The content of hydrogen in the precipitate decreased with increasing initial p H (5.8-7.8) from 0.538 to 0.268 wt. %. 5. Saida et al [13] observed no distinct X-ray diffraction peak corresponding to a crystalline phase. A halo peak appeared in the vicinity of 45 degrees for all the samples prepared in the molar ratios (2-10) of K B H 4 to C o 2 + ( N i 2 + and Fe 2 + ) , indicating that a mostly single amorphous phase was formed for the Co-B or N i - B alloys, and the structure of the two alloy systems was independent of the molar ratios of K B H 4 to C o 2 + . The temperature did not have an appreciable influence on the amorphous structure which implied that there was no significant change in the reaction mechanism with temperature. Considering the charge and mass balance o f the reaction and high B content of the precipitates, the following principal reactions were thought to be dominant: (2-66) 4 C o 2 + + BHf + WH~ -> 4Co + B02 + 6H20 (2-67) 4 C o 2 + + 2BH4~ + 60H~ -» 2Co2B + 6H20 + H2 (2-68) 39 The actual generation of hydrogen seems to support the appropriateness of the reaction (2-68). Reaction (2-67) indicates that the existence of B O 2 " in the residual aqueous solution is expected. The hydrogen adsorbed on the precipitate was not measured and so it is not reasonable to use the amount of evolved hydrogen gas to support the reaction mechanism. 6. Jianyi et al [12] studied the reductions of C 0 C I 2 , N i C l 2 and FeSC>4 in aqueous solution with potassium borohydride solution (pH 12), and suggested three independent reactions: BHA~ +2H20 = B02 + 4H2 (2-69) BH~ + 2M2+ + 2H20 = 2M + BO{ + 4H+ + 2H2 (2-70) BH~ + H20 = B + OH~ +2.5H2 (2-71) The relative extent of these reactions depends on the metal ions and the reaction conditions, and therefore there is no single reaction equation that can be used to express the overall reactions. Taking reaction (2-69) as a by-product reaction, the ratio of reaction (2-70) to reaction (2-71) is usually 1-1.5. The corresponding reduction reactions would be: 2BH~ +2M2++ 2H20 = M2B + HB02 + 2H+ + 45H2 (2-72) 5BH- + 6M2+ + 6H20 = 2M3B + 1HB02 + 7H+ +11H2 (2-73) The overall reaction sequence can maintain the p H of the reaction mixtures at 3-4.5 by themselves because reactions (2-70) and (2-71) produce H + and O H " ions, respectively which together form water. The reactions promote each other. The mole ratio of BH4" to metal ions for equivalent reactions was found to be 1.8 to 1.9. 40 Both the addition rate and concentration of B H 4 " solution affected the boron content in Fe-B and N i - B but did not affect the boron content in C o - B . This effect is due to the different orders in BFLf of reactions (2-70) and (2-71) for the preparation of Fe-B and N i - B . Except for the thin layer of oxides of metals and boron on the surfaces (formed by passivation), the particles consisted of amorphous structures with average particle sizes from 40 to 60 nm. 7. Awadalla et al [47] studied the reaction of cobalt with sodium borohydride and found that the product of cobalt reduction was amorphous cobalt boride. The stoichiometry of cobalt reduction was expressed by Equation (2-37) and some ions such as Z n and C l " had a negative effect on the reduction of cobalt which is described by the data listed in Table 2-6. From Table 2-6, it can be seen that all impurities have negative effects on cobalt reduction. The complexing agents form complexes with cobalt and make cobalt difficult to be ^ I j o j j r\ j reduced and hence decrease the cobalt reduction efficiency. Ca , N a , M g , N H 4 and Z n possibly form the corresponding borohydride and interfere with the combination of cobalt ions and borohydride ions. Zinc ions have a strong negative effect of cobalt reduction and 100 ppm zinc ions decreased the reduction efficiency from 97 % to 32 %. The zinc ions have a strong negativ% effect on the iron reduction, no effect on copper reduction [47] and a positive effect on gold (I) reduction [25]. The presence of 10 mg/1 zinc ion in thiourea solution containing 10 mg/1 A u decreased sodium borohydride consumption by 83 %. The possible cause is that zinc ions formed zinc borohydride which decomposes into zinc hydride and hydrogen. Zinc hydride hydrolyzes rapidly and so cobalt reduction efficiency is decreased, i.e., zinc ions competed with cobalt ions for borohydride ions. The higher than the stoichiometric requirement of sodium borohydride was not beneficial because the excess of the reductant was consumed by hydrolysis. The reduction reaction was very fast and was completed in less than 3 minutes. 41 Table 2-6 Effect of additives on the reduction of cobalt [47]. Initial Co concentration^ 00 P P M , Initial p H =6, the mol ratio of NaBtLj to Co 2 + was 2, p H was not controlled. Sodium borohydride was added as a solid powder. N o . Materials added Amount added Reduction Efficiency, % 1 N a O H 0.0424 g/1 96 2 CaO 0.05g/l 72 3 M g ( O H ) 2 0.05 g/1 75 4 N H 4 O H 0.15 g/1 91 5 N a 2 C 0 3 0.0848 g/1 82 6 C l " 100 ppm 87 200 ppm 87 300ppm 86 7 Z n 2 + 100 ppm 32 8 Sodium acetate 0.1391 g/1 70 9 Sodium citrate 0.4990 g/1 0.0 10 Sodium tartrate 0.1952 g/1 66.7 11 E D T A 0.3158 g/1 0.0 12 N o impurities added - 97 * Nos. 1-5 were added to adjust p H 2.3 to p H 6.0 ** For N o . 7, 89% of zinc was precipitated in the form o f zinc hydroxide and no zinc metal was precipitated. *** For Nos. 7-11, the initial p H was not given ( p H is the range of 5.0-7.5) 8. Polyakov et al [114, 115] reported that the stoichiometry of cobalt reduction can be expressed by Equation (2-34) and at a consumption of 110 % of the stoichiometric requirement, cobalt could be reduced from 20 ppm to 0.06 ppm [114] or 0.005 ppm [115] from zinc sulfate solution (120 g/L Zn) in the presence of trienthanolamine at p H 3. The reduction efficiency increased with increasing reaction time and decreased with increasing temperature and decreasing p H . 42 2.4.3.3 K I N E T I C S O F C O B A L T R E D U C T I O N W I T H B O R O H Y D R I D E There are few reports on the kinetics of cobalt reduction with borohydride because the reduction o f cobalt ions from aqueous solution in the absence o f some ligands is too fast to follow. In addition the reaction is very complicated. Thus far it appears that only Gomez-Lahoz [41] reported on the kinetics of cobalt reduction with borohydride. He found an autocatalytic effect on cobalt reduction and gave the following kinetic expression: - - ^ — ^ = kl[BH4][Co2+] + k2[BH4-][Co2+][Co2B] (2-74) at B y combining equations (2-33) and (2-74) and using a numerical method, Gomez-Lahoz obtained k i and k 2 . The reaction rate was also affected by p H . However, the p H changed greatly and so did the form of the cobalt ions (e.g. C o ( H 2 0 ) 2 + , [ C o ( H 2 0 ) 5 O H ] + , Co(OH) 2(aq).). The reaction rate was affected becoming lower with increasing p H . 2.5 S U M M A R Y Since sodium borohydride has unique redox features ( a low reducing equivalent weight of 4.75 g/ mole e", a high reducing power and utility in different media such as water and organic solvents, and acidic, neutral and alkaline conditions), it has advantages over other reducing agents. In addition since it is the cheapest of all the metal borohydrides, it has been widely used for various purposes. N o w sodium borohydride reduction of metal ions is the basis of several commercial processes such as the preparation of selective catalysts and magnetic materials, the recovery of precious metals, the removal and recovery of heavy metals from waste water, and electroless plating. Wi th respect to cobalt, sodium borohydride is widely used to prepare cobalt-boron catalysts and magnetic materials and to remove cobalt from waste water. Cobalt removal with 43 sodium borohydride has been selected as the best available technology for the treatment of cobalt-bearing waste water. The efficient removal of cobalt from zinc electrolyte continues to be a problem in the purification of the electrolyte and so zinc production plants have been looking for a better method to replace the present processes. If sodium borohydride could be used to remove cobalt from zinc sulphate solution, the purification process would have many advantages over the present technology, for example, high efficiency, the low reaction temperature, the short reaction time. Poyakov et al reported [114, 115] that sodium borohydride can reduce cobalt in zinc sulphate solution to 0.01 ppm in the presence of triethanolamine. However, Cominco's staff [116] has reported that cobalt can not be reduced from zinc sulphate solution with sodium borohydride. Awadalla [47] also found that zinc ions have a strong negative effect on cobalt reduction. Hence there are conflicting claims. The reaction of cobalt reduction with borohydride is very sensitive to the reaction conditions such as p H , concentration, solution composition, the method of mixing reactants and the treatment of the precipitate. Therefore various authors have obtained different reaction stoichiometrics and have proposed different reactions. Because of the extreme rapidity of the reaction and other complications, relatively little systematic work has been done in this area. Most authors reported that the stoichiometry of cobalt reduction with sodium borohydride is dependent on the reaction conditions and that there is no definite stoichiometry. Usually the reduction of one mole of cobalt needs about two moles of sodium borohydride and only Polyakov [114. 115] reported that 1.1 moles of sodium borohydride can reduce 4 moles o f cobalt ions. A more thorough study of cobalt reduction by borohydride is needed. Such a study could lead to a better understanding of the stoichiometry, and i f possible, to increasing the reduction efficiency. Finally, the possibilities of using sodium borohydride for cobalt removal should be studied. 44 3. E X P E R I M E N T A L 3.1 E X P E R I M E N T A L O B J E C T I V E S A N D M E T H O D S The main objective of this research was to investigate the stoichiometry o f the reduction of cobalt ions with sodium borohydride and the effects of the addition of hydroxide, p H , temperature and the rate of sodium borohydride addition on the stoichiometry. Another key objective was to study the feasibility of removing cobalt from zinc sulphate electrolyte with sodium borohydride. Sodium borohydride solutions with different concentrations of sodium hydroxide were used as reducing agents in order to study the effect o f the mol ratio of N a B H 4 to N a O H on cobalt reduction efficiency. A solution of sodium borohydride was added in order to maintain the addition at a uniform and continuous rate. To improve the reproducibility, a batch reactor was used. The p H was controlled manually by adding dilute sulphuric acid or sodium hydroxide solutions because the constituents of buffer solutions might affect the cobalt reduction with borohydride. Cobalt sulphate solution was used for all experiments because the ultimate intended application was cobalt removal from zinc sulphate solution with borohydride. 3.2 E Q U I P M E N T A two-liter Pyrex glass reactor was immersed in a water bath where the temperature was controlled to within ± 0.2 °C. The water in the bath was heated with heaters connected to an Omega C N 900A temperature controller, and cooled with cold water. A magnetic stirrer bar was used for agitation. A n acrylic l id atop the reactor was used to minimize air in the reactor and evaporative loss. The solution was sparged with highly pure nitrogen to eliminate oxygen and agitated with a plastic-coated impeller connected to a stirrer motor (Caframo Model R Z R 1 ) . The agitation (rotation) rate was checked with a phototachometer (Cole-Parmer 08210). A Corning p H probe and a Bach-Simpson Ltd. P H M 8 2 standard p H meter were used to measure p H which was recorded by computer. The p H was controlled manually by adding 0.05 M sulphuric acid or sodium hydroxide. A Fisher platinum indicating electrode for redox titration and a saturated calomel reference electrode and Bach-Simpson Ltd. P H M 8 2 Standard p H meter were also used to measure the solution redox potential which was recorded by a computer. The reducing solution 45 was pumped into the reactor (Cole-Parmer pump model 7519-20A equipped with a digital variable-speed console drive). Figure 3-1 is a schematic diagram of the experimental setup. The syringes and Cameo II 25 mm filters (DDN-02T25-50, 0.22 [i m pore size) were used to take the solution samples. The solutions were analyzed for cobalt by a colorimetric method using Nitroso-R-salt with a Bausch & Lomb Spectronic 20 spectrophotometer, and for nickel with a Perkin-Elmer 306 Atomic Absorption Spectrophotometer. The precipitates were obtained by filtering the reaction mixtures through a Mil l ipore GS 0.22 jxm membrane filter held on a Mil l ipore Funnel/Support assembly. The precipitates were washed with deoxygenated water under a nitrogen atmosphere and then put under acetone. The precipitates were dried in a desiccator under vacuum at 0.001 mm H g with a type JR-150 rotary vacuum pump produced by Japan Electron Optics, Tokyo Laboratory Co. . The precipitates were analyzed by the International Plasma Lab Ltd. for Co and B . The heat-treatment of the precipitate was carried out in a tube furnace at 10"6 Torr. The following equipment was used to obtain high vacuum: Balzers Type D U O 56 rotary pump as fore pump, Edwards high vacuum diffusion pump, and a Balzers K V 313 vacuummeter was used for measuring the vacuum. Fig . 3-2 is a schematic diagram o f the precipitate heat-treatment experimental setup. 3.3 M A T E R I A L S (1) Fisher reagent grade sodium borohydride and sodium hydroxide were used to prepare the required sodium borohydride solutions. (2) Fisher reagent grade zinc sulfate, cobalt sulfate, nickel sulfate, copper sulfate, lead chloride, cadmium sulfate, antimony tartrate, sodium arsenite, and deionized water were used to prepare the various solutions. (3) Reagent grade triethanolamine and polyacrylamide were used as surface active agents. Polyacrylamide was used to prepare 1% polyacrylamide water solution. (4) Reagent grade anhydrous zinc chloride, sodium borohydride and anhydrous tetrahydrofuran were used to prepare zinc borohydride. (5) Highly pure nitrogen was used as an inert gas to remove dissolved oxygen. 46 To temperature g controller Hole for sampling and adding solution To temperature controller ^ -Acrylic lid Glass reactor O Stirring plate Reducing solution Figure 3-1 Schematic diagram of the experimental setup 4 7 High vacuum gage connected to vacuum meter Thermocouple connected to furnace temperature controlk z 1- Sealed mechanism for thermocouple protection tube 2- O-ring sealed quartz tube connector 3- O-ring sealed quartz tube connector for specimen entry port 4- Quartz viewing window Figure 3-2 Schematic diagram of sample heat treatment apparatus 48 3.4 E X P E R I M E N T A L P R O C E D U R E 3.4.1 P R O C E D U R E F O R T H E R E D U C T I O N O F C O B A L T (1) A one-liter solution was prepared according to the required compositions. (2) The solution was added to the reactor. (3) The solution was sparged with nitrogen gas for over 30 minutes to eliminate dissolved oxygen and heated until the temperature was within the required range. (4) The surfactant solution was added to the reactor about 3 minutes before the sodium borohydride was added. (5) 0.05 M H 2 S 0 4 or N a O H solution was used to adjust p H to the required value. (6) The sodium borohydride solution was continuously and uniformly added to the reactor with a pump within the required time. If p H had to be controlled, sulphuric acid and sodium hydroxide solution were added manually to the reactor. For the kinetics of cobalt reduction, the sodium borohydride was added to the reactor with a syringe within one second. (7) The solution samples were taken using syringes and syringe disk filters. The precipitates were taken by filtering the solution through a Mil l ipore membrane filter under vacuum and were washed with deoxygenated water under a nitrogen atmosphere and then submerged under acetone. 3.4.2 H E A T T R E A T M E N T O F T H E P R E C I P I T A T E S (1) The precipitates were dried in a vacuum desiccator in which the pressure was maintained at 0.001 mm H g for one day. (2) The dried precipitates were put into a boat which was pushed into the tube furnace. The furnace travelled along the quartz tube and the movable cooling system was positioned where the samples were so that the samples would not be heated. (3) When the vacuum reached 10"6 Torr, the furnace was turned on. (4) Due to the rise of the temperature in the tube, the gas was released and the pressure rose slightly. When the vacuum reached 10"6 Torr again at the required temperature, the heating part of the furnace was moved into position over the sample, and the sample was heated. 49 (5) After two hours, the furnace was moved so that the sample was no longer in the zone of the furnace. The movable cooling system was also put in place to cool the sample and so the sample temperature was lowered rapidly. (6) Because the sample is readily pyrophoric, the furnace was cooled for two days and argon was passed through the tube holding the sample. Finally the cover was removed and the sample was withdrawn. 3.4.3 X - R A Y D I F F R A C T I O N OF T H E P R E C I P I T A T E S (1) The heat treated precipitates were ground to fine powder with a mortar and pestle. (2) The samples were put onto a glass plate and dispersed uniformly with acetone. (3) The sample was put into the diffraction chamber. The radiation wave length was 1.504 A (CuKJ. 3.4.4 S C A N N I N G E L E C T R O N M I C R O S C O P I C A N A L Y S I S OF T H E P R E C I P I T A T E S (1) A fresh powder sample was directly put into acetone. Heat-treated samples were first ground with a mortar and pestle and then put in into acetone. (2) The sample was placed onto a conductive plate and put into the chamber of the scanning electron microscope for analysis. (3) The voltage of the electron beam was 20 k V . 3.4.5 T R A N S M I S S I O N E L E C T R O N M I C R O S C O P I C A N A L Y S I S O F T H E P R E C I P I T A T E S (1) Fresh powder samples were put directly into acetone. Heat-treated samples were first ground with a mortar and pestle and then put into acetone. (2) The beaker holding acetone solution was put in a ultrasonic bath to separate particles. (3) The sample was put onto a thin, electron transported, carbon support film. (4) The fi lm was put into the chamber of the T E M and examined at 200 kv. 50 4. RESULTS AND DISCUSSION 4.1 S T O I C H I O M E T R Y O F C O B A L T R E D U C T I O N W I T H S O D I U M B O R O H Y D R I D E The cobalt reduction with borohydride was carried out by adding the reducing solution continuously and at a uniform rate to the reactor. Therefore the amount of sodium borohydride added was proportional to the addition time. A 0.25 M N a B H 4 + 0.05 M N a O H solution was used to reduce cobalt. The cobalt reduction fraction, p H and potential of the reaction solution vs. time are shown in Figure 4-1. The total addition time was 10 minutes. Figure 4-2 shows similar data for a 5-minute addition time. The reaction is very fast and almost immediate and so the addition time can represent the reaction time. From Figures 4-1 and 4-2, it can be seen that as the sodium borohydride was added to the mixture, the recorded p H values of the solution in the first several seconds were higher than the initial p H . This is simply due to the basic sodium borohydride solution increasing the concentration of OH". The p H then went down because of the reaction of cobalt with borohydride which released hydrogen ions. After some time the p H stabilized. This could be accounted for by a balance between the release of hydrogen ions and borohydride hydrolysis which consumes hydrogen ions. After most of the cobalt was reduced, the p H went up rapidly because the main reaction became the hydrolysis of borohydride. The faster the addition rate, the greater the p H dropped (from 7 to 6.7 for a 10-minute addition time and from 7 to 6.2 for 5 minutes). The faster addition of sodium borohydride released more hydrogen ions per unit time. Several seconds after the beginning of the addition, the potential went down and then stabilized. A s the p H went up, the potential went down reflecting the dependence of potential on the solution p H . The measured potential at p H 6.4 was about -0.63 V S C E . This potential is much higher than the redox potential of H 3 B 0 3 / B H 4 " (-1.05 V S C E at [ B H 4 ] = [H 3 B0 3 ] ) and approximates the redox potential of H + / H 2 (-0.624 V SCE) . The concentration of borohydride is very low and so the measured solution potential is mainly dependent on the redox potential of H + / H 2 couple. FigureS 4-3 and 4-4 are the same type of diagram for the cobalt reduction at pHjn; tjai=5.5 and 8.0. Considering that the reaction released hydrogen ions, reducing solutions with different sodium hydroxide concentrations (0, 0.05, 0.125 and 0.25 M ) were used to reduce cobalt ions. For each case, different initial p H conditions were also applied (pH 5.5, 7.0, 8.0). From Figure 4-51 1, and Figures 4-3 to 4-13, the p H of the solution increased with increasing concentration of sodium hydroxide in the reducing solution. Also with increasing concentration o f sodium borohydride, the reduction efficiency increased greatly from 2.1-2.4 moles o f sodium borohydride needed to reduce 1 mole of cobalt ion to 1.0 mole (see Table 4-1). Table 4-1 Moles of sodium borohydride needed to reduce one mole of cobalt ions for different basic solutions and different initial p H values. Initial p H 0.25 M N a B H 4 0.25 M N a B H 4 + 0.25 M N a B H 4 + 0.25 M N a B H 4 + 0.05 M N a O H 0.125 M N a O H 0.25 M N a O H 5.5 2.4 1.95 1.6 1.1 7.0 2.2 1.85 1.5 1 8.0 2.1 1.8 1.5 - 1 N a O H in the reducing solution neutralized the hydrogen ions released by the cobalt reduction reaction. This suppressed the borohydride hydrolysis and increased the reduction efficiency. The initial p H also has some effect on the cobalt reduction efficiency. A low initial p H leads to lower cobalt reduction efficiency probably because of the hydrolysis of more borohydride. However, in the p H range of 5.5-8, the initial p H has less effect on the cobalt reduction efficiency than the hydroxide concentration of the reducing solution. The higher the concentration of sodium hydroxide, the less the effect of initial p H . For 0.25 M N a B H 4 solution (no added NaOH) , 2.4, 2.2, and 2.1 moles of sodium borohydride were needed to reduce 1 mole o f cobalt ions respectively at initial pH's of 5.5, 7.0, 8.0 respectively. However, for 0.25 M NaOH+0.25 M N a B H 4 solution, the reduction efficiencies are almost the same. Apparently because the rate of cobalt reduction is much higher than that for the hydrolysis at pH's greater than 7, p H has a little effect on the reduction efficiency for the same composition of reducing solution. The N a O H in the reducing solution played the main role in increasing the reduction efficiency and decreasing borohydride hydrolysis. The hydrogen ions released by the cobalt reduction act to promote the hydrolysis. The hydrogen ions released were consumed by N a O H (neutralization) or N a B H 4 (hydrolysis). The more basic the reducing solution, the less the hydrolysis of N a B H 4 . Gozm-Lahoz [41] reported that sodium hydroxide in the sodium borohydride solution can improve the reduction efficiency. 52 A t an initial p H of 5.5 (see Figures 4-3, 4-5, 4-8 and 4-11), as sodium borohydride solution was added to the reaction solution, the p H first rose rapidly, which suggests that sodium borohydride hydrolysis was dominant, then stabilized and finally rose rapidly. A t an initial p H of 7.0 (see Figure 4-1, 4-6, 4-9 and 4-12), the p H went up a little to start, then dropped, stabilized and finally rose rapidly for the reducing solutions with 0.0, 0.05, 0.125 M N a O H . But for the reducing solution with 0.25 M N a O H , p H first went up a little, stabilized, and then finally went up rapidly. A t an initial p H of 8 (see Figures 4-4, F ig . 4-7, F ig . 4-10 and Fig.4-13), the p H first went up a little, decreased, then stabilized and finally went up rapidly. However the reducing solution with 0.25 M N a O H , the p H went down by less than 0.20, which might be caused by the buffering effect o f the boric acid product. It is possible that the reaction o f one borohydride and one cobalt ion produced one hydrogen ion. A n attempt was made to use more basic reducing solutions (with 0.5 M and 0.75 M NaOH) to reduce cobalt ions, but it was found that the p H increased rapidly and some blue cobalt hydroxide was precipitated. Apparently i f the reaction of cobalt reduction released 2 or more hydrogen ions and a reducing solution with a mole ratio o f N a O H to NaBH4 =2 was added to the reaction solution, the p H would not go up so fast and cobalt hydroxide would not precipitate. 53 (a) Mol ratio of N a B H 4 to C o + 2 9.0-0.0 0.4 0.8 1.2 1.6 2.0 8.5 H !.0H X 7.5 H CL 7.0 H 6.5 H 6.0 ->—r - potential - p H H 0.2 Time, min. —r-10 H -0.8 (b) Figure 4-1 (a) Cobalt reduction fraction, (b) pH and potential vs. time, [Co 2 +]i ni ti ai = 30 mg/L. Note that time is directly proportional to the amount of added borohydride. 54 Mol ratio of NaBH 4 to C o 2 + 0.0 0.4 0.8 1.2 1.6 2.0 - 1 ' 1 ' 1 ' 1 ' 1 ' 1 -~\ 1 1 1 1 1 1 1 1 • r 0 1 2 3 4 5 Time, min. (a) Mol ratio of NaBH 4 to C o + 2 0.0 0.4 0.8 1.2 1.6 2.0 0 1 2 3 4 5 Time, min. (b) Figure 4-2 (a) Cobalt reduction fraction, (b) pH and potential vs. time, [ C o 2 + ] j n j t j a i = 30 mg/L. Note that time is directly proportional to the amount of added borohydride. 55 Mol ratio of NaBH 4 to C o 2 + Time, min. (a) o.o 0.4 Mol ratio of NaBH 4 to C o 2 + 0.8 1.2 1.6 6 H 35 °C, 500 rpm stirring rate, pH y 4 6 Time, min. (b) 20 0.2 0.0 H -0.2 LU O CO -0.4 -j? CD O CL H -0.6 -0.8 10 Figure 4-3 (a) Cobalt reduction fraction, (b) pH and potential vs. time, [Co2 +]ini tiai = 30 mg/L. Note that time is directly proportional to the amount of added NaBH 4. Figure 4-4 (a) Cobalt reduction fraction, (b) pH and potential vs. time, [Co ]initial = 30 mg/L. Note that time is directly proportional to the amount of added borohydride. Figure 4- 5 (a) Cobalt reduction fraction, (b) p H and potential vs. time, [Co ] initial = 30 m Note that time is directly proportional to the amount of added borohydride. Figure 4- 6 (a) Cobalt reduction fraction, (b) p H and potential vs. time, [Co Hinitial = 30 mg/L. Note that time is directly proportional to the amount of added borohydride. 59 Mol ratio of N a B H 4 to C o 2 + 0.0 0.4 0 .8 1.2 1.6 2 .0 2 .4 —1 ' 1 1 1 1 1 1 1 1 1 1 T -1 • 1 > 1 > 1 1 1 1 1 1 r 0 2 4 6 8 1 0 1 2 Time, min. Figure 4- 7 (a) Cobalt reduction fraction, (b) p H and potential vs. time, [Co ]initial = 30 mg/L. Note that time is directly proportional to the amount of added borohydride. Figure 4-8 (a) Cobalt reduction fraction, (b) p H and potential vs. time, [Co "Jinitial = 30 mg/L. Note that time is directly proportional to the amount of added borohydride. 61 Figure 4-9 (a) Cobalt reduction fraction, (b) p H and potential vs. time, [Co ]initial = 30 mg/L. Note that time is directly proportional to the amount of added borohydride. 62 Figure 4- 10 (a) Cobalt reduction fraction, (b) p H and potential vs. time, [Co ]]initial = 30 mg/L. Note that time is directly proportional to the amount of added borohydride. 63 Mol ratio of NaBH 4 to Co 1 1 I ' I 1 I 1 I 1 T" 0 2 4 6 8 1 0 Time , min. Figure 4-11 (a) Cobalt reduction fraction, (b) p H and potential vs. time, [Co ]initial = 30 mg/L. Note that time is directly proportional to the amount of added borohydride. 64 Figure 4- 12 (a) Cobalt reduction fraction, (b) p H and potential vs. time, [Co J initial - 30 mg/L. Note that time is directly proportional to the amount of added borohydride. 65 Figure 4-13 (a) Cobalt reduction fraction, (b) p H and potential vs. time, [Co ] initial = 30 mg/L. Note that time is directly proportional to the amount of added borohydride. 66 The industrial zinc electrolyte p H is usually controlled below 4.5. Therefore one attempt was tried to reduce cobalt ions with sodium borohydride at p H 4 in the absence of zinc ions in solution. From Figure 4-14, after the mol ratio of N a B H 4 to Co exceeded 4, the reduction fraction was not increased. This means that all the borohydride after this point was consumed in the hydrolysis. The rate of Co (II) reduction became equal to the rate of redissolution and cobalt played the role of catalyzing the hydrolysis. The redox potentials E ( H + / H 2 ) =- 0.244 V S H E at p H 4 and E(Co 2 + /Co)=-0.277 V S H E for ac0(ii) = 1 and -0.377 V S H E for ac0(ii) = 0.0005, indicate that the H + should first be reduced and then cobalt ions. However, cobalt ions react with borohydride ions much faster than do hydrogen ions and this makes it possible for cobalt ions to be reduced by borohydride. The stoichiometry of cobalt reduction is dependent on the reaction conditions, p H and basicity of the reducing solutions. Figure 4-14 The reduction fraction vs. time and the mol ratio of N a B H 4 to Co . The reducing solution: 0.25 M N a B H 4 + 0.25 M N a O H . p H was manually controlled by adding dilute sulphuric acid solution. 67 4.2 T H E E F F E C T OF p H O N C O B A L T R E D U C T I O N E F F I C I E N C Y W I T H B O R O H Y D R I D E p H has a great effect on cobalt reduction efficiency below p H 5. From Figure 4-15, the cobalt reduction efficiency is seen to increase with increasing p H . A t p H 2, cobalt reduction efficiency is almost zero and at p H 6, the reduction efficiency is 96 %. A t a lower p H , from equation (2-27), the borohydride hydrolysis (reaction (2-26)) becomes faster and in accordance with the Eh-pH diagram for the cobalt-water system (Figure 2-6), cobalt is oxidized by H + and redissoves. Co + 2H+ -+Co1+ +H2 (4-1) Therefore the reduction efficiency is low under these conditions. PH Figure 4-15 Effect of p H on cobalt reduction efficiency. The reducing solution was 0.25 M NaBH4 + 0.25 M N a O H . p H was manually controlled by adding dilute sulphuric acid solution at p H 2 ± 0.05, 3 ± 0.07, 4 ± 0.1, 5 ± 0.15,and 6 ± 0.22. [ C o 2 + ] i n i t i a i = 30 mg/L, M o l ratio of N a B H 4 to C o 2 + = 2. 68 4.3 T H E E F F E C T O F A D D I T I O N R A T E O N C O B A L T R E D U C T I O N E F F I C I E N C Y A t lower p H the cobalt metal product is readily redissolved, and so slow addition leads to the redissolution o f more cobalt and a low reduction efficiency. Figure 4-16 shows the effect of the addition rate of sodium borohydride solution on the cobalt reduction efficiency. Fast addition of sodium borohydride increases the rate of the hydrolysis of sodium borohydride, However, the reaction with cobalt (II) is much faster than that of hydrogen ions and most of the sodium borohydride is used to reduce cobalt ions. Overall, therefore the shorter addition time increases the reduction efficiency. Figure 4-16 The effect of addition rate of borohydride on cobalt reduction efficiency at 35 °C, the reducing solution: 0.25 M N a B H 4 + 0.25 M N a O H , 500 rpm, the p H was manually controlled by adding dilute sulphuric acid solution. 69 4.4 T H E E F F E C T O F T E M P E R A T U R E O N C O B A L T R E D U C T I O N E F F I C I E N C Y Increasing temperature has a negative effect on cobalt reduction. The effect o f temperature is shown in Figure 4-17. From 25 to 45 °C, the extent of cobalt reduction changes little with increasing temperature. Above 45 °C, the cobalt reduction efficiency decreases greatly with increasing temperature. A t higher temperatures, the hydrolysis of borohydride and the redissolution become more active and the reduction efficiency becomes poorer. It can be seen that the shorter addition time increases the cobalt reduction efficiency because it decreases the amount o f the redissoved cobalt. Figure 4-17 The effect of temperature on cobalt reduction efficiency. The reducing solution was 0.25 M + 0.25 M N a O H , mol ratio of N a B H 4 to C o 2 + = 2, and [ C o 2 + ] i n i t i a i = 30 mg/L. 70 4.5 T H E R E D I S S O L U T I O N OF R E D U C E D C O B A L T A t p H 4 the redox potential of Co (U)/Co is below the redox potential of H + / H 2 (see F ig . 2-6), so the reduced cobalt w i l l be reoxidized by H + ions. A s shown in Figure 4-18, the cobalt concentration decreased to 5 mg/L about 20 minutes after sodium borohydride was added to the solution. Further addition of sodium borohydride hardly decreased the cobalt solution concentration. A t this point the cobalt reduction rate was equal to the rate of redissolution and sodium borohydride was consumed by hydrolysis. After the addition of sodium borohydride ceased, the cobalt concentration increased with time, due to the cobalt redissolution. Because the cobalt particles were small, the reduced cobalt was quickly redissolved, almost completely within 30 minutes. After the addition of sodium borohydride, the solution potential dropped to about-0.49 V (SCE) and stabilized. This potential is a little lower than the potential of H + / H 2 (-0.48 V (SCE) at p H 4, higher than the standard potential of Co (U)/Co (-0.517 V (SCE)) and much higher than the potential of H 3 B 0 3 / B H 4 ~ a t p H 4 (-0.927 V (SCE) assuming [H 3 B0 3 ]=[BH 4 "] ) . Since the solution potential was measured with a platinum electrode which catalyzed borohydride hydrolysis [123], H 2 , B H 4 " and probably the reduced cobalt particles adsorbed on the electrode and the measured potential is the mixed potential of H + / H 2 , H 3 B 0 3 / B H 4 " and C o 2 + / C o . The concentration of B H 4 " was very low and so the measured potential was determined mainly by the H + / H 2 redox couple. It was found that the measured potential did not reproduce well and was also determined by the electrode treatment method and electrode materials. 71 Time, min. Figure 4-18 The redissolution of reduced cobalt. The reducing solution was 0.25 M NaBFL; + 0.25 M N a O H . After the addition of sodium borohydride ceased, the p H was controlled at 4 ± 0.05 by adding dilute sulphuric acid. 72 4.6 I D E N T I F I C A T I O N O F T H E P R E C I P I T A T E 4.6.1 M O R P H O L O G Y A N D P R O P E R T I E S OF P R E C I P I T A T E S Figures. 4-19 and 4-20 are scanning electron micrographs revealing the shape and morphology of the fresh precipitates. The very small particles coalesced together and formed larger granules with sizes in the range of 10 to 20[im. After the precipitates were treated ultrasonically in acetone, the small particles became separated to some extent. The particles were then immediately collected on a thin evaporated carbon fi lm for transmission electron microscopy. The particle sizes of the resultant materials were in the range of 20-100 nm. After heat treatment at 500 °C and 10"6 torr for two hours, the particles became larger (about 200-500 nm). During the heat-treatment, the precipitates recrystallized and small particles combined to form larger particles as shown in F ig . 4-22. The precipitates exhibited magnetic properties both in the as-precipitated and heat-treated conditions. Cobalt boride is generally chemically inert and does not dissolve in sulphuric acid, but only in hot concentrated nitric acid. These precipitates showed high chemical reactivity and dissolved readily in dilute sulphuric acid. Even after the heat-treatment, the precipitate still dissolved in 0.1 M sulphuric acid solution. This chemical reactivity may be caused by the very small size which produces a large area-volume ratio and hence a high surface free energy-volume ratio. A sample of precipitate was immersed in boiling water. After 80 minutes, the atom ratio of Co to B in the residues was 2.5. During this treatment, some boron was dissolved. The probable reaction may be written as: Co2B + 3H20^> 2Co + H3B03 + l.5H2 (4-1) Figure 4-19 Surface morphology of the precipitate prepared under the conditions: pH 5.5-7.5, 35 0 C, 500 rpm stirring rate, mol ratio of borohydride to cobalt ion =2, 10- minute addition time and initial [Co (II)] = 30 mg/L, 200 x. Figure 4-20 Surface morphology of the precipitate prepared under the conditions: pH 5.5-7.5 and 35 °C, 500 rpm stirring rate, mol ratio of borohydride to cobalt ion = 2, 10- minute addition time, initial [Co (II)] = 30 mg/L, 4,000 x. Figiure 4-21 Transmission electron micrograph of the fresh precipitate prepared under the conditions: p H 5.5 -7.5, 500 rpm stirring rate, mol ratio of sodium borohydride to cobalt ion = 2, 10- minute addition time, initial[Co (U)] = 30 mg/L, 50,000 x. Figiure 4-22 Transmission electron micrograph of the heat-treated precipitate prepared under the conditions: p H 5.5 -7.5, 500 rpm stirring rate, mol ratio o f sodium borohydride to cobalt ion = 2, 10- minute addition time, initial [Co (II) ]= 30 mg/L, 100,000 x. 75 4.6.2 I D E N T I F I C A T I O N B Y X - R A Y D I F F R A C T I O N The X-ray diffraction pattern of the fresh precipitate (see Figure 4-23) showed a broad peak at 45 ±2 °. The strongest peak in C 0 2 B appears at 45.7 0 (d 2 n = 1.983 A ) and the strongest peak in Co appears at 44.3 °. (dm = 2.0467 A ) . There was no peak corresponding to B around 21.8 °. The precipitate was probably C o 2 B , Co or a mixture of C o 2 B and Co. The chemical analysis of the fresh precipitate showed that the atom ratio of Co to B was 2.1 close to C o 2 B . The broad peak in the X-ray diffraction pattern indicates that the crystalline phase had very low periodicity, typical of very fine crystallite size and approaching an amorphous material. After the precipitate was heat-treated for 2 hours at 500 °C in a high vacuum of 10"6 Torr, the X-ray diffraction pattern of the precipitate showed well defined peaks due to C o 2 B and also the main peak of Co. Furthermore, the preparation of cobalt boride from cobalt and boron powders requires a temperature of over 1000 °C and 500 °C is not enough to make cobalt boride. Therefore the precipitate likely consisted of amorphous C o 2 B and Co. Under the heat treatment, the diffusion allowed the periodicity to increase giving a longer crystalline region (grain) of C o 2 B and excess cobalt and particles became larger. After fresh precipitates were placed in boiling water for 80 minutes, some boron was dissolved and the atom ratio of Co to B was 2.5, and the broad amorphous peak in the X-ray diffraction pattern (Figure 4-25) became larger. After a two-hour heat-treatment at 500 °C, a stronger main peak of cobalt appeared (Figure. 4-26), which is consistent with the composition of the precipitate. After heat treatment at 250 °C for two hours, the precipitate still showed a broad X-ray diffraction peak, consistent with its being amorphous (Figure 4-27). 76 2 - T h e t a - S c a l e UHIUERSITY OF BRITISH COLUMBIA 0 2 - O c t - 1 9 9 5 15:59 1—'———•—i—•—•—•——i—1——•—1—r 5 10 15 20 25 D S O 0 O . D M T A J M L - 2 . R A U J M L - 2 ' C T : 0 . 8 S 30 35 3S :0 .Q20d9 , U L : I ' 1 1 ' I 40 45 1 , 5 4 0 6 H O ) 50 55 Figure 4-23 X-ray diffraction pattern of the precipitate prepared under the conditions: p H 5.5-7.5 and 35°C, 500 rpm stirring rate, mol ratio of N a B F L to C o 2 + = 2,and 10-minute addition time. 2 - T h e t a - S c a l e 19 -Dec-1995 12:13 Background subtraction cc S u 10 15 20 25 30 35 40 45 C : N D 5 0 0 0 \ D A T A \ J M L - 1 . R A W J M L - 1 ( C T : 0 . 8 s , S S : 0 . 0 2 0 d 9 , U L : 1 .5406Ao) Figure 4-24 X-ray diffraction pattern o f the precipitate prepared under the conditions: p H 5.5-7.5 and 35 °C, 500 rpm stirring rate, mol ratio o f N a B F L to Co 2 + =2,10-minute addition time and 500 °C heat-treatment for 2 hours. 77 2 - T h e t a - S c a l e 2 0 - O c t - 1 9 9 5 12:25 6 5 7 0 1 . 5 4 0 3 I H O J Figure 4-25 X-ray diffraction pattern o f the precipitate prepared under conditions: p H 5.5-7.5, 35 °C, 500 rpm stirring rate, mol ratio o f NaBELj to C o 2 + = 2, 10-minute addition time, and heating boil ing water for 80 minutes. 2 - T h e t a - S c a l e 2 0 - O c t - 1 9 9 S 09:41 25 30 35 40 45 C : \ D 5 0 0 0 \ D A T A \ J M L - 4 . R f l U JML-4 ( C T : 50 55 1 . 0 s , S S : 0 .020d9 , 65 70 1.5409AO) Figure 4-26 X-ray diffraction pattern o f the precipitate prepared under the conditions: p H 5.5 -7.5, 35 °C, 500 rpm stirring rate, mol ratio o f NaBFL; to C o 2 + = 2, 10-minute addition time, boi l ing water for 80 min., and 500 °C heat-treatment for 2 hours. 78 2-Theta - Scale 20-Oct-1995 12:34 IS IS M c 3 O o Uncorrected - no background subtraction IS IS IS 5 10 15 20 D 5 0 0 0 D H T A J I 1 L - 3 .RHII J M L - 3 'CT : 25 30 35 40 4 O . a i , S 3 : 0 . 0 2 0 d s , U L : 1 , 5 4 0 9 H O ) 5 50 55 Figure 4-27 X-ray diffraction pattern of the precipitate prepared under the conditions: p H 5.5-7.5, 35 °C, 500 rpm stirring rate, mol ratio of NaBFL; to C o 2 + = 2, 10-minute addition time, and 250 °C heat treatment for 2 hours. 4.6.3 T E M I D E N T I F I C A T I O N OF T H E P R E C I P I T A T E The transmission electron microscopy pattern o f a cluster of fresh precipitates is shown in F ig . 4-28. It does not exhibit a sufficiently well-defined ring pattern that allows crystal identification. However, the X-ray energy spectrum showed a strong cobalt peak ( k - a ) . After the precipitate was heat-treated for two hours at 500 °C, a single crystal transmission electron diffraction pattern was obtained from individual particles (see F ig . 4-29). A n evaporated gold film (polycrystalline) was used to obtain an electron diffraction ring pattern from which the camera constant was calculated and this constant was used to index the unknown single crystal 79 electron diffraction pattern of the precipitate. The gold ring electron diffraction pattern is shown in F ig . 4-30, and the details of the calculation are listed in table 4-2. Table 4-2 Calculation of the camera constant (Au calibration) for T E M h 2 + k 2 +12 h k 1 d(h,k,l) A radius of rings for A u pattern (rn>k,i) (mm) camera constant rh,k.ixd(h,k,i) (mm A ) 3 1 1 1 2.355 12.75 30.026 4 2 0 0 2.039 14.6 29.769 8 2 2 0 1.442 20.7 29.849 11 3 1 1 1.230 24.5 30.135 12 2 2 2 1.1774 25.5 30.024 Average 29.96 The camera constant is given by the product r*d, where r is the ring diameter and d is the interplanar spacing of the diffraction ring. The average value of the constant is 29.96 mm A . The d-spacings of the reflection on the unknown pattern obtained from the camera constant and the distance (r) o f the diffraction spot from the center are listed in Table 4-3 and compared with the known d-spacings for C 0 2 B , and the indices (h, k, 1) were obtained. According to the indices, the angles o f the planes were calculated using electron diffraction pattern software [153]. A protractor was used to measure the angles between the rel-vectors of the diffraction pattern The values of the angles are listed in Table 4-4. It can be concluded that the single crystal electron diffraction pattern is consistent with that of C 0 2 B and that the precipitate was cobalt boride. The zone [u, v, w] of Figure 4-29 is identified as [110]. u: v: w = ( k i l 2 - k 2 l i ) : (lih2 - l 2hi): (h,k 2 - h2k,)=(0x2 - 1x2): (2x1-2x0): (0xl-lx0)=l:(-l):0. So the plane (h, k, 1) belongs to the zone [110] and hu + kv + lw = 0 hiu + k i v + l i w = 0x1 + 0x(- l ) + 2x0 = 0 h 2 u + k 2 v + l 2 w = l x l + lx(-l) +2x0 = 0 h 3 u + k 3 v + l 3 w = 2x1 + 2x(-l) + 2x0 = 0 h4U + k4V + l 4 w = 3x1 +3x(-l) + 3x0 = 0 80 h 5 u + k 5 v + l 5 w = l x l + lx(-l) + 0x0 = 0 Table 4-3 The calculation of the d-spacings of the precipitate radius (mm) the calculated d-spacing (A) the theoretical d-spacing (A) the indices (h,k,l) n = 14.2 di=29.96/14.2=2.110 2.113 (002) r 2 = 16.5 d2=29.96/16.5=1.816 1.815 (112) r 3 = 22.1 d3=29.96/22.1=1.356 1.357 (222) r 4 = 29 d4=29.56/29.0= 1.033 1.032 (332) r 5 = 8.3 d5=29.96/8.3=3.625 3.55 (110) Table 4-4 Calculated theoretical angles and measured angles Calculated angle (using software) Measured angle c<(002)/(ii2) = 30.753° 31° a(oo2)/(222) = 49.959° 50° O((002)/(332) = 60.742° 61° O£(002)/(110) = 90° 90° Of(112)/(222) = 19.206° 20° a ( i i2)/ (332) = 29.989° 30° ot(i i2)/(i io) = 59.247° 60° C<(222)/(332) = 10.783° 10° O£(222)/(110) = 40.041° 40° Oi(332)/(ii0) = 29.258° 30° Figure 4-28, Transmission electron micrograph of the fresh precipitate prepared under the conditions: p H 5.5 -7.5, 500 rpm, the molar ratio o f sodium to cobalt ion =2, the addition time=10 minutes, [Co]=30 mg/1, 200 kv. Figure 4-29 Transmission electron micrograph of the precipitate prepared under the conditions: pH 5.5 -7.5, 500 rpm, mol ratio of sodium borohydride to cobalt ion = 2, 10-minute addition time, [Co (II)] = 30 mg/L and 500 °C heat-treatment for 2 hours, 200 kv. Figure 4-30 Transmission electron pattern of gold film, 200 kv 84 4.6.4 T H E E F F E C T OF T E M P E R A T U R E O N T H E C O M P O S I T I O N OF T H E P R E C I P I T A T E The reaction temperature has some effect on the composition of the precipitate. From Fig . 4-31, the atom ratio of Co to B increases with increasing temperature. A t higher temperatures, the hydrolysis of borohydride (or borane) becomes more active and dominates over the cobalt-catalyzed formation of elemental boron. After the precipitate was boiled in water for 80 minutes, the atom ratio of cobalt to boron of the precipitate increased from 2.05 to 2.55, which means that some boron was oxidized by water and went into solution. A t higher p H , the differences between the redox potential of C o 2 + / C o and H 2 O / H 2 or H3BO3/BH4" become more positive and so cobalt is easier to reduce. This also favors a higher atom ratio of Co to B . A t p H 4 some cobalt redissolved with the result that the atom ratio of Co to B in the precipitate became lower. Figure 4-31 The effect of temperature on the composition of the precipitate produced with 500 rpm stirring rate, with a reducing solution of 0.25 M NaBH4 + 0.25 M N a O H , initial [Co] = 30 mg/L. The p H was controlled by adding dilute sulphuric acid solution. 85 4.6.5 THE EFFECT OF ADDITION RATE ON THE COMPOSITION OF THE PRECIPITATE From Fig. 4-32, the addition rate (time) has little effect of the composition of the precipitate. 2.8-2.6 H CO 2.4-o o ^ 2.2 H ' o o « 2.0-1 E o < 1.8-i 35 °C, 500 rpm stirring rate Mol ratio of N a B H 4 to Co+ 2 — • — pH4±0.15 I—O—pH 5.5-7.5 1.64 —f— 10 Total addition time, min. Figure 4-32 Th% effect of the addition rate on the composition of precipitate produced with 500 rpm stirring rate, with a reducing solution of 0.25 M NaBH 4 + 0.25 M NaOH, initial [Co] = 30 mg/L. The pH was controlled by adding dilute sulphuric acid solution. 86 4.7 P O S S I B L E R E A C T I O N S D U R I N G C O B A L T R E D U C T I O N W I T H B H 4 " X-ray diffraction patterns show that the precipitate was a mixture of cobalt boride and cobalt. The atom ratio of cobalt to boron increases from 2.05 to 3.2 with increasing temperature from 25 to 75 °C, but is quite insensitive to the addition rate. From the p H vs. time results obtained by using sodium borohydride solution with different concentrations o f sodium hydroxide, the reduction of cobalt with borohydride may release one hydrogen ion. One mole of borohydride can reduce at most one mole of cobalt ions without the precipitation of cobalt hydroxide. Therefore, according to the inner-sphere mechanism proposed by Khain [132] and similar to Equation 2-42 to 2-47, the main reactions might be the following reactions: The formation of intermediate compound, Col+ (aq) + BH4~ -> [CoBH4]+ + (aq) (4-2) The formation of a bridging bond, [CoBH4]+ Coz+ -(H~ -BH3) (4-3) The transfer of a hydrogen atom and rupture of the bridging bond, Col+ - (H~ -BH3)^> CoH+ + BH3 (4-4) The electron transfer and bond rupture, CoH+ -+Co + H + (4-5) Summing up the above reactions, we get, C V + (aq) + BH4 -+Co + H+ + BH3 + (aq) (4-6) 87 The catalytic decomposition of some borane by cobalt, BH3—^>B + 3I2H2 (4-7) Most of the produced cobalt condensed together with boron to form amorphous C 0 2 B . 2Co + B^Co2B (4-8) The hydrolysis of the other borane, BH3 + 3H20 -> H3B03 + 3H2 (4-9) The hydrolysis of borohydride, BHA ~ +3H2 0 + H+ -> H3B03 + 4H2 (4-10) The atom ratio of cobalt to boron depends principally on reactions (4-7) and (4-9). A t a high temperature, the hydrolysis of borane (BH3) would become more active and less boron is formed. Therefore the atom ratio of cobalt to boron is high. The overall reaction can be written as: 2Co2+ + 2BH~ + 3(1 + 0.5x)H2O -> (4-11) (1 - 0.5x)Co2B + xCo + 2H+ + (1 + 0.5x)H3BO3 + (4.5 + 0.75x)H2 where x is the range of 0-2. If the ratio of reaction (4-7) to reaction (4-9) is 1, that is, x = 0, the overall reaction may be written as: 2Co2+ + 2BHA~ + 3H20 -> Co2B + 2H+ + H3B03 + 4.5H2 (4-12) 88 If the hydrogen ions are not neutralized, combining reactions (4-10) and (4-12), we can get the following overall reaction: 2Co2+ + 4BH4~ + 9H20-> 2Co2B + 3H3B03 + \2.5H2 (4-13) If the hydrogen ions are neutralized by sodium hydroxide, reaction (4-10) can be suppressed. The rate of hydrolysis is much slower than that of reaction (4-6) at lower temperatures (below 35 °C) and high p H (above 7). If sodium borohydride is added slowly to the cobalt solution and the hydrogen ions are neutralized by sodium hydroxide in the reducing solution, reaction (4-10) can be suppressed to the maximum extent and sodium borohydride reduction efficiency for cobalt ions can be optimized. The reaction can then be written as: 2Co2+ + (4 - y)BH~ + yOH~ + [3(3 + 0.5x) - 4y]H20 -> (1 - 0.5x)Co2B + xCo + (3 + 0.5* - y)H3BO} + (12.5 + 0.75x - 4y)H2 where y is in the range of 0-2. I f y is 2, we can get the following reaction. 2Co2+ + 2BH~ + 20H~ + (1 + \5x)H20 -> (1 - 0.5x)Co2B + xCo + (1 + 05x)H3BO3 + (4.5 + 0.15x)H2 (4-13) (4-14) If y = 0, i.e. [OH"]/[BH4-] is 0 in sodium borohydride solution, 2 moles of borohydride are needed for reducing 1 mole of cobalt ions; i f y = 0.4, i.e. [OETj/fBtLf] = 0.2, 1.8 moles of borohydride for 1 mole o f cobalt ions, i f y = 1, i.e. [OHVfBELV] = 0.5, 1.5 moles of borohydride for 1 mole of cobalt ions, and i f y = 2, i.e. [OH"]/[BH 4~] = 1, 1 mole o f borohydride is needed for 1 mole o f cobalt ions. These results are well consistent with the data at p H 7, and 8, listed in Table 4-1. A t lower p H e.g. 5, because of the hydrolysis of borohydride, the amount of borohydride needed for the reduction of cobalt ions is higher than the above calculated values. 89 4.8 K I N E T I C S T U D I E S O F C O B A L T R E D U C T I O N W I T H B O R O H Y D R I D E 4.8.1 F E A T U R E S O F K I N E T I C S O N C O B A L T R E D U C T I O N The reduction of cobalt ions with borohydride is almost instantaneous. In fact as soon as sodium borohydride solution was added to the solution at normal temperature (from 25 -35 °C), a black precipitate appeared, the amount of which did not increase with time. This means that the reduction was almost immediate and was completed within a few seconds. When the temperature was lowered to 5 °C, it was observed that after 5 seconds the black precipitate appeared and then the solution became black immediately. Because the reaction was too fast to follow, it was difficult to measure the cobalt concentration in the solution as a function of time. From p H vs. time data, we can see how the reaction progressed. Figure 4-33 shows the measured p H and potential vs. time at 15 °C. The reaction can be divided into three stages. In the first stage, alkaline sodium borohydride solution was added to the reactor and raised the p H of the solution and only a small proportion of cobalt ions was reduced. The p H should have gone up faster and to higher values, but it did not. The two causes of this result are that the reduction of cobalt released hydrogen ions, and cobalt sulphate had some buffering effect. The buffering effect of cobalt sulphate was the main cause. Figure 4-34 shows the p H of * various solutions to which different additions were made over time. When 0.25 M N a B H 4 + 0.05 M N a O H or 0.05 M N a O H was added to pure water, the p H went up very quickly and after 4 seconds, the measured p H exceeded 9. The p H for 0.05 M N a O H went up a little faster and reached a little higher value than that for 0.25 M +0.05 M N a O H . Sodium ions may affect p H measurement because the measured p H of 0.05 M , 0.125 M and 0.25 M N a O H were 12.44, 12.91, 13.14 respectively which were a little higher than the p H of 0.25 N a B H 4 + 0.05 M N a O H , 0.25 N a B H 4 + 0.125 M N a O H , and 0.25 M N a B H 4 + 0.25 M N a O H solutions, 12.37, 12.88, 13.13 respectively at 23 °C. Another cause may be the buffering effect of the produced boric acid. The p H vs. time plots for 0.05 M N a O H and 0.25 M N a B H 4 + 0.05 M N a O H added to 30 mg/L C o 2 + solutions are almost the same within the first two seconds, which means that cobalt sulphate acts as a buffer and keeps the p H from rising rapidly. After about two seconds, the p H for 0.05 M N a O H continued to increased and the p H for 0.25 M N a B H 4 + 0.05 M N a O H 90 decreased because the cobalt reduction released hydrogen ions. However, after 10 seconds the pH was higher than for the former because of sodium borohydride hydrolysis and the consumption of cobalt ions. The final pH's for 30 mg/L Co 2 +and 0.25 M N a B H 4 + 0.05 M NaOH are much lower than that for pure water and 0.25 M N a B H 4 + 0.05 M NaOH although cobalt ions were reduced. This is because the boric acid produced buffered the solution. In the second stage, cobalt(H) was reduced, and the reduction was autocatalyzed by reduced cobalt metal. The hydrolysis of borohydride was also catalyzed by reduced cobalt. After this stage, the measured concentration of cobalt ions was less than 1 mg/L. By the third stage, most of the cobalt ions were reduced and the main reaction became the hydrolysis of BH4" which caused the pH to rise rapidly. T 1 1 1 1 1 1 1 1 1 1 r 0 10 20 30 40 50 Time, seconds Figure 4-33 pH and potential of the solution vs. time. The reducing solution was 0.25 M N a B H 4 + 0.05 M NaOH , the mol ratio of N a B H 4 to C o 2 + was 2, [Co 2 + ] i n i t i a i = 30 mg/L. 91 i ' 1 1 1 ' 1 1 1 1 r 0 10 20 30 40 50 Time, seconds Figure 4-34 p H vs. time for solutions with different compositions and addition of different additives 15 °C, and 2500 rpm stirring rate. 92 4.8.2 E F F E C T O F T E M P E R A T U R E O N R E A C T I O N R A T E The p H vs. time data at different temperatures are shown in Figure 4-35. Wi th increasing temperature the reaction rate increased greatly. The reduction of most of the cobalt needs 24 seconds at 5 °C, 5 seconds at 15 °C, and about two seconds at 25 °C and less than two seconds at 35 °C. Above 15 °C the reaction rate was dependent on mass transfer and the p H measurement was affected by the time required for the p H probe to reach equilibrium and by the time required to add the reducing solution. The p H meter response time was about 1 second and so the real peak in p H was not detected. Figure 4-35 p H vs. time at different temperature, the reducing solution (4 ml) contained 0.205 M N a B H 4 + 0.05 M N a O H and the solution was added within 1 second, pHinitiai = 5.5 and [Co 2 + ] i n i t i a i = 30 mg/1. 93 4.8.3 E F F E C T OF p H O N R E A C T I O N R A T E A t higher p H , cobaltous hydrocomplexes such as C o ( O H ) + and Co(OH )2 (aq) form and the concentration of free Co (aq) 2 + decreases. This slows the reaction rate of cobalt reduction with borohydride. The time needed for the reduction of most of the cobalt is 5 seconds for an initial p H of 5.5 (see Figure. 4-33), 8 seconds for an initial p H of 6 (Figure. 4-36), 9 seconds for an initial p H of 7 (see Figure. 4-37), 14 seconds for an initial p H of 7.5 (see Figure. 4-38) and 27 seconds for an initial p H of 7.8 ( see Figure. 4-39) respectively at 15 °C. A t 35 °C the time required is less than 2 seconds for an initial p H of 5.5 (see Figure. 4-35), 2 seconds for an initial p H of 7.0 (see Figure. 4-40), 3 seconds for an initial p H of 7.8 (see Figure. 4-41). A t lower temperatures, p H has a large effect on the rate and at higher temperatures, it has less effect on. i 1 1 1 I 1 i 1 i 1 r i 1 1 1 1 < 1 ' 1 < r 0 10 20 30 40 50 Time, seconds Figure 4-38 p H vs. time at initial p H 7.5, 4ml of reducing solution and [Co ] in i t ia l = 30 mg/L 8.5 8.0 Q. 7.5 7.0 15 OC, 2500 rpm stirring rate, pH 10 1 ' 1 1 1 -20 30 40 50 Time, seconds Figure 4-39 pH vs. time at initial pH 7.8, 15 °C, 4 mL of reducing solution and [Co 2 + ] i nj tj a i 30 mg/L. 8.5 8.0 X Q. 7.5-7.0 35 °C, 2500 rpm stirring rate, pHjnitia|=7.0 10 20 30 40 - 1 -50 Time, seconds Figure 4-40 pH vs. time at initial pH 7.0, 35 C, 4 mL of reducing solution and [Co ]initial 30 mg/L. 96 i — i , 1 , 1 , 1 1 1 1 1 0 10 20 30 40 50 Time, seconds Figure 4-41 p H vs. time at initial p H 7.8, 35 °C, 4 m L of reducing solution and [ C o 2 + ] i n j t i a i = 30 mg/L. 4.8.4 A U T O C A T A L Y T I C E F F E C T OF C O B A L T O N C O B A L T R E D U C T I O N The reduced cobalt has an autocatalytic effect on cobalt ion reduction. A n initial 0.66 ml of sodium borohydride solution was added to a solution containing 35 mg/L C o 2 + it was found that cobalt reduction released hydrogen ions and caused a drop in p H . A t the 25th second, 4 ml of sodium borohydride solution was added to the solution containing 30 mg/1 C o 2 + . From the p H vs. time (Fig. 4-42), only 2 seconds were needed to reduce most o f cobalt ions. In another experiment 5 mg/L of cobalt ions was first reduced by adding 0.6 m L of sodium borohydride solution and the p H was adjusted to 6. Then 4 m L of sodium borohydride solution was added to the solution containing 30 mg/1 cobalt ions. From the p H vs. time graph (see F ig . 4-43), it can be seen that the time for the reduction of most of the cobalt ions was 4 seconds compared with the 8 97 seconds for the solution without pre-reduced cobalt present. Apparently, the reduced cobalt has a catalytic effect on cobalt reduction. i CL 5 H 15 °C, 2500 rpm stirring rate, pH i n i t i a l=6.0 10 20 30 Time, seconds 40 50 Figure 4-42 pH vs. time data for the autocatalytic effect on cobalt reduction 8H C L 15 °C, 2500 rpm stirring rate, pH i n i t i a l = — • — without pre-reduced cobalt — • — with 5 mg/l pre-reduced cobalt 10 15 Time, seconds 20 25 Figure 4-43 pH vs. time data demonstrating the autocatalytic effect on cobalt reduction. 98 4.9 C O B A L T P R E C I P I T A T I O N F R O M Z I N C S U L P H A T E S O L U T I O N 4.9.1 I N H I B I T O R Y E F F E C T OF Z I N C IONS O N C O B A L T A N D N I C K E L R E D U C T I O N The strong negative effect of zinc ions on cobalt reduction is surprising. The results are shown in Figure 4-44. When the zinc sulfate concentration exceeds several 10 (a,mol/L zinc, cobalt ions cannot be reduced by borohydride. The inhibitory effect is dependent on the addition rate (time) of sodium borohydride solution. The concentration of zinc sulphate at which no cobalt reduction occurs is 30 |j,mol/L for 10-minute addition time, 40 umol /L for 5 minutes, 70 jxmol/L for 2 minutes and 120 u.mol/L for 1 minute. Zinc sulphate also has a strong negative effect on nickel reduction. From Figure 4-44, the concentration of zinc sulphate for no cobalt reduction is 150 (xmol/L for 10-minute addition time, 200 (a,mol/L for 5 minutes, 250 u,mol/L for 2 minutes and 300 umol/L for 1 minute. It was also found that zinc sulphate has a strong effect on cadmium Figure 4-44 The inhibitory effect of zinc ions on cobalt reduction with the mol ratio o f NaBH4 to C o 2 + = 2 and 0.25 M N a B H 4 + 0.25 M N a O H reducing solution, p H was controlled by adding dilute sulphuric acid solution. 99 and lead reduction. A t [ZnS0 4 ] > =0.05 or [ZnCl 2 ] > 0.2 M , lead (30 mg/L initially) was not reduced and no precipitate formed at p H 4. However copper ions were reduced at [ZnS0 4 ] > 2 M . This means that zinc sulphate has a strong negative effect on reduction of electropositive metals and less effect on reduction electronegative metals. In order to verify that zinc ions suppress cobalt reduction, zinc chloride, magnesium sulphate and sodium sulphate were added to the reaction mixture. Zinc chloride has the same inhibitory effect on cobalt reduction, but magnesium sulphate and sodium sulphate have less effect on cobalt reduction. From Figures 4-46 to 4-49, it can be seen that only when their concentrations exceed 0.1 M was cobalt reduction inhibited. Under these conditions, the cations M g 2 + or N a + may combine with borohydride ions to form associated sodium borohydride salts. Figure 4-45 The inhibition effect of zinc ions on nickel reduction with the mol ratio of N a B H 4 to N i 2 + = 2 and 0.25 M N a B H 4 + 0.25 M N a O H reducing solution, p H was controlled by adding dilute sulphuric acid solution. 100 and so decrease the concentration of free borohydride ions. The SO4 " may be complexed by Co to form C0SO4 (aq) and thereby decrease the reactivity and potential of cobalt. Figure 4-49 shows the p H vs. time data for 30 mg/L C o 2 + («0.0005 M ) + 0.0001 M Z n 2 + after the addition of sodium borohydride solution. Compared to the p H vs. time plot for C o 2 + alone, the p H for the C o 2 + + Z n 2 + system rose more slowly due to the buffering effect of both zinc sulphate and cobalt sulphate, but it did not exhibit the characteristic initial rise and fall observed when C o 2 + is reduced. The reaction was suppressed by zinc ions. M u c h less cobalt powder was precipitated compared to C o 2 + in the absence of zinc ions and some cobalt was deposited on the glass reactor wall . From the above results, it is clear that zinc ions have a negative effect on cobalt reduction with borohydride. Figure 4-46 Effect of the additives ( Z n S 0 4 , Z n C l 2 , M g S 0 4 , N a 2 S 0 4 ) on cobalt reduction. The reducing solution was 0.25 M N a B t L + 0.25 M N a O H , the p H was manually controlled by adding dilute sulphuric acid solution. 101 Additive concentration, mole/L Figure 4-47 Effect of the additives ( Z n S 0 4 , Z n C l 2 , M g S 0 4 , N a 2 S 0 4 ) on cobalt reduction. The reducing solution was 0.25 M N a B H 4 + 0.25 M N a O H , the p H was manually controlled by adding dilute sulphuric acid solution. 102 rq 1 1 i i i i 111 1 1—i i i i 111 1 1—i i i i 111 1 1—i i i i 111' 1E-5 1E-4 1E-3 0.01 0.1 Additive concentration, mole/L Figure 4-48 Effect of the additives (ZnS04, ZnCl 2, MgS0 4 , Na 2S0 4) on cobalt reduction. The reducing solution was 0.25 M NaBH 4 + 0.25 M NaOH, the pH was manually controlled by adding dilute sulphuric acid solution. 103 •TJ 1 1—i i i 1111 1 1—i i i 1111 i 1—i i i 1111 1 1—i i i 1111 r 1E-5 . 1E-4 1E-3 0.01 0.1 Additive concentration, mole/L Figure 4-49 Effect of the additives (ZnS04, ZnCl 2, MgS0 4 , Na 2S0 4) on cobalt reduction. The reducing solution was 0.25 M NaBH 4 + 0.25 M NaOH, the pH was manually controlled by adding dilute sulphuric acid solution. 104 T 1 1 < 1 ' 1 1 1 1 r 0 10 20 30 40 50 Time, seconds Figure 4-50 p H vs. time data for 30 mg/L C o 2 + and 30 mg/L C o 2 + 0.0001 M Znz+ after the addition of 0.25 M N a B H 4 + 0.05 M solution. 105 4.9.2 E F F E C T O F O T H E R A D D I T I V E S O N C O B A L T R E D U C T I O N 1. Triethanolamine and polyacrylamide(0-200 mg/L) were added to the solution (0-120 g/L Z n 2 + , 30 mg/1 C o 2 + ) at temperatures (15-75 °C). Polyakov et al [114, 115] reported that the cobalt concentration was lowered to 0.06 mg/L in the presence of these two additives. But the additives had no effect on the reduction of cobalt and on the protection of cobalt from redissolution. 2. Zinc dust was added to the solution to help reduce cobalt ions, but there was not any positive effect on the reduction of cobalt in the zinc sulphate solution. 3. Antimony tartrate, sodium arsenite, and copper sulphate are used to activate the reduction of cobalt with zinc dust in zinc electrolytes in industry. Therefore these reagents were added to test the solutions to help reduce the cobalt ion. Unfortunately these additives did not show any positive effect on cobalt reduction. When sodium borohydride was added to the solution containing arsenite and antimony tartrate as well as zinc and cobalt ions, almost no precipitate appeared. The possible cause is that antimony tartrate and arsenite reacted with borohydride to produce stibine, SbH3,and arsine, A s H 3 . It was found that copper was first reduced and cobalt remained in solution. 4. Hydrazine sulphate which is a strong reducing agent and is sometimes used as a corrosion inhibitor, was added to the zinc sulphate solutions. However, it not show any positive effect on the reduction of cobalt with borohydride. 5. The precipitate prepared by cobalt reduction with sodium borohydride from cobalt sulphate solution was added to the zinc and cobalt sulphate solutions to promote the nucleation o f cobalt. However, the cobalt reduction with borohydride was still not improved. This means that the negative effect of zinc ions on the cobalt reduction is not the nucleation inhibition by zinc ions. As02- + BHA + H20-> AsH3 + H3B03 (4-14) SbO+ + BH4~ + 2H20 -> SbH3 + H3B03 + 2H2 (4-15) 106 4.9.3 P O S S I B L E I N H I B I T O R Y M E C H A N I S M OF Z I N C IONS Zinc ions might have some catalytic effect on the hydrolysis of sodium borohydride. When sodium borohydride was added to the solution, the local solution became cloudy and gas evolution occurred possibly due to hydrogen gas formation. The potential of the solution increased with increasing concentration of zinc ions. After sodium borohydride was added to the solution, zinc ions competed with cobalt ions for borohydride ions to form zinc borohydride. The rate of the formation of zinc borohydride is probably much faster than that of the cobalt reduction. According to the literature [135], the following reactions may occur: Zn1+ +2BH4~ -+ Zn(BH4)2 (4-17) Zn{BH4)2 ZnH2 + 2BH3 (4-18) ZnH2 + 2H+ -+Zn2+ + 2H2 (4-19) BH3 + 3H20^> H3B03 + 3H2 (4-20) According to this scheme, the zinc borohydride decomposes into zinc hydride and borane both of which hydrolyze rapidly. The hydrolysis of zinc hydride produces zinc ions which further react with borohydride ions and thus catalyze the hydrolysis of borohydride. Wi th increasing p H the hydrolysis of zinc ions increases and the rate of the formation of zinc borohydride decreases and even reaches zero. A t a high p H the zinc hydride is relatively more stable and so the catalytic effect on the hydrolysis of borohydride is low. A s the sodium borohydride solution (0.25 M N a B H 4 + 0.25 M NaOH) was gradually added to the solution containing 30 mg/L C o 2 + a n d 10 mg/L Z n 2 + , the p H increased and when it exceeded 7, the black precipitate began to appear. From Figures 4-43 and 4-44, the inhibitory effect of zinc ions decreased with increasing addition rate of sodium borohydride. For example, the concentration of zinc ions required to completely stoO cobalt reduction was 3 x 10"5 M for a 10-minute addition time, 4 x 10"5 M for 107 5 minutes, 7 x 10"5 M for 2 minutes and 1.2 x 10"4 M for 1 minute. Apparently i f a sufficient excess of borohydride is present in solution, some cobalt reduction can occur. 108 5. C O N C L U S I O N S The stoichiometry o f cobalt reduction with sodium borohydride is dependent on reaction conditions such as temperature, p H , and composition of the reducing solution. Cobalt reduction with sodium borohydride causes the p H o f the solution to drop indicating the release of hydrogen ions. The extent of the p H drop depends on the relative rates o f cobalt reduction and the hydrolysis of sodium borohydride. The progress of the reaction can be divided into three stages. When sodium borohydride is added to the solution, the p H first decreases, which indicates that the reduction of cobalt is the dominant reaction rather than the hydrolysis o f sodium borohydride. Next the p H reaches some value and stabilizes briefly. A t this stage there is a balance between the cobalt reduction and the hydrolysis of sodium borohydride and the rate of the production of hydrogen ions is equal to the rate of its consumption. Finally the p H increases rapidly, meaning that most of the cobalt is reduced and the main reaction becomes the hydrolysis of borohydride. Because the reaction of cobalt reduction releases hydrogen ions, which catalyze the hydrolysis of sodium borohydride, sodium hydroxide was added to the solution o f sodium borohydride to neutralize the hydrogen ions produced by the cobalt reduction. It was found that with increasing content of sodium hydroxide in the sodium borohydride solution, the cobalt reduction efficiency increased. The reduction o f 1 mole o f cobalt ions required 2.4 moles of sodium borohydride for the sodium borohydride solution without sodium hydroxide, 1.8-1.95 moles for the solution with 0.05 M sodium borohydride solution, 1.5-1.60 moles for the solution with 0.125 M sodium hydrooxide and 1.0-1.1 moles for the solution with 0.25 M sodium hydroxide. When more basic sodium borohydride solutions (0.50 and 0.75 M sodium hydroxide) were tried it was found that the reduction efficiency increased, but some cobalt was precipitated in the form of cobalt hydroxide. So the best reduction efficiency without the precipitation of cobalt hydroxide is about one mole of sodium borohydride to reduce one mole of cobalt ions. If the reaction solution p H is controlled at 4, the maximum reduction efficiency is about 81 % for alO-minute addition time and further addition of sodium borohydride added did not increase the extent o f reduction. This means that the rate o f redissolution o f cobalt metal was equal to the rate of the reduction. 109 The reduction efficiency of cobalt increases with increasing p H . A t p H 2 almost no cobalt was reduced. A t p H 6, 96 % of cobalt was reduced i f two moles of sodium borohydride were used to reduce one mole of cobalt ions. If the p H was controlled at a low value, the reduction efficiency increased with increasing rate of cobalt addition. For example, at p H 4, 2 moles of sodium borohydride can reduce 0.614. mole of cobalt for a 10-minute addition time, 0.633 mole for 5 minutes, 0.743 for 2 minutes and 0.771 for 1 minute. The shorter addition time can reduce the amount of cobalt redissloved but increases the hydrolysis of sodium borohydride. Overall it has a positive effect. The reduction efficiency decreased with increasing temperature above 35 °C. For example, at p H 4 the reduction efficiency decreased from 61.4% at 35 °C to 7 % at 75 °C for a 10-minute addition time. X-ray diffraction was used to identify the precipitates, but only a broad peak was obtained near the main peak for C 0 2 B . After the precipitate was heat-treated for two hours, the X-ray diffraction pattern showed the peaks of C 0 2 B and a weaker main peak of cobalt, suggesting that the precipitate consisted of C 0 2 B and Co. Transmission electron microscopy was also used, but failed to identify the fresh precipitate. After two hours of heat-treatment at 500 °C, the single crystal transmission pattern obtained was consistent with that for C 0 2 B . The atom ratio of cobalt to boron in the precipitate increased with increasing temperature which is consistent with the tendency of borohydride hydrolysis to increase with increasing temperature. The reduction of cobalt ions is completed within several minutes. The reaction process can be divided into three stages. During the first stage as sodium borohydride solution is added to the solution, the p H rises quickly and little cobalt is reduced. Because of the buffering effect of cobalt sulphate the p H did not go very high. In the second stage, cobalt is reduced and the reaction is autocatalyzed by the reduced cobalt. The p H decreases and the main reaction is cobalt reduction. In the third stage after the reduction of most of the cobalt, the hydrolysis o f borohydride, which also catalyzed by the reduced cobalt, becomes the main reaction and the p H rises rapidly. The initial p H o f the solution has some effect on the reaction rate of cobalt reduction. A high p H decreased the rate of cobalt reduction. The reaction time of cobalt reduction 110 decreased with increasing temperature from 24 seconds at 5 C to less than 2 seconds at 35 C. Cobalt reduction was strongly catalyzed by reduced cobalt. Attempts were made to use sodium borohydride to remove cobalt from zinc sulphate solution, but it was found that zinc ions had a dramatic inhibitory effect on cobalt reduction. Several tens of p.mol/1 of zinc ions lead to no cobalt reduction. The main cause of inhibition is likely that zinc ions compete for borohydride ions with the cobalt ions. Zinc borohydride forms and hydrolyzes rapidly producing zinc ions which react again with borohydride. Zinc ions may catalyze the hydrolysis of borohydride. Zinc ions also had an inhibitory effect on nickel and lead reduction. I l l 6. R E C O M M E N D A T I O N S Cobalt reduction with borohydride is very complicated with its stoichiometry depending on the reaction conditions and the procedures. This research has shown that the reaction stoichiometry and the precipitate composition are mainly dependent on pH and temperature. So far different reaction mechanisms have been proposed in the literature to account for the results such as the effects of pH, the solution composition, the rate of the addition of borohydride, the amount of the hydrogen gas evolved and the mol ratio of borohydride to cobalt ion. 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Kerby, "Manganese Powder Purification of Zinc Electrolyte- Cobalt and Nicke l Removal from Zinc Electrolyte by Manganese Powder Cemenatation", Cominco Ltd. , Technical Research Report, Zinc Plant Section, Report No . 1, 1976. 113. R. C . Kerby, "Manganese Powder Purification of Zinc Electrolyte- Cobalt and Nicke l Removal from Zinc Electrolyte by Manganese Powder Cementation", Cominco Ltd. , Technical Research Report, Zinc Plant Section, Report No . 2, 1976. 114. M . L . Polyakov, O. M . Samsonova and E . M . Polyakova, " Removal o f Nicke l , Cobalt, and Other Metals from Zinc Sulphate Solutions by Reduction with Sodium Borohydride", Izv. Vyssh. Uchebn. Zaved., Tsvetn. Metall . (Russian), No. 6, 1988, pp. 47-51. 115. M . L . Polyakov, E . M . Polyakova and O. M . Polyakova, " Removal of Metals from Zinc Sulphate Solution by Reduction with Sodium Borohydride in the Presence o f Triethanolamine", U S S R Patent 1,212,951, June, 1986. 116. R. Lew, 1995, private communication between R. Lew, Cominco Ltd. , Trial B . C . and D . B . Dreisinger, University of British Columbia. 117. N . N . Greenwood and A . Earnshaw, Chapter 6, "Boron", in Chemistry of the Elements, Pergamon Press, New York, 1984, pp. 155-240. 118. F . A . Cotton and G . Wilkinson, Part Two, "Boron' , in Advanced Inorganic Chemistry, Third Edition, Interscience Publishers, New York, 1972, pp. 223-259. 119. N . N . Greenwood, Chapter 11 ("Boron"), V o l . 1, in Comprehensive Inorganic Chemistry, edited by J. C . Bailar, H . J. Emeleus, S. R. Nyholm, and A . F. Troman-Dickenson, Pergamon Press, N e w York , 1973, pp. 665-991. 125 120. N . P Nies and G . W . Campbell, "Inorganic Boron-Oxygen Chemistry", Boron, Metallo-Boron Compounds and Boranes, edited by Roy M . Adams, Intersicience Publishers, New York , 1964, pp. 53-288. 121. N . P. Nies and G . W . Campbell, "Inorganic Boron-Oxygen Chemistry", Boron, Metallo-Boron Compounds and Boranes, edited by Roy M . Adams, Intersicience Publishers, N e w York, 1964, pp. 53-196. 122. R. M . Adams and A . R. Siedle, "The Hydroboron Ions (Ionic Boron Hydrides)", Boron, Metallo-Boron Compounds and Boranes, edited by Roy M . Adams, Intersicience Publishers, N e w York, 1964, pp. 373-506. 123. B . D . James and M . G . H . Wallbride, " Metal Tetrahydroborate", in Progress in Inorganic Chemistry", V o l . 11, edited by Stephen J. Lippard, 1970, Interscience Publishers, N e w York , pp. 99-232. 124. P. K . Liao and K . E . Spear, " B - C o (Boron-Cobalt) phase diagram", Binary A l l o y Phase Diagrams, edited by T. B . Massalski, 1990, American Society for Metals, W i l l i a m W . Scott Publisher, V o l . 1, pp. 344-347. 125. E . Winberg and W . Henle, "The Properties of Ether-soluble Zinc Borohydride", Z . Naturforschg. 7b, 1952, pp. 579-580. 126. W . H . Stockmayer, D . W . Rice and C. C. Stephenson, " Thermodynamic Properties of Sodium Borohydride and Aqueous Borohydride Ion", Journal of the American Chemical Society, Vol .77, 1955, pp. 1980-1983. 127. D . D . Wagman, W . H . Evans, V . B . Parker, R. H . Schumn, S. M . Bailey, I. Hakow, K . L . Churney and R. L . Nuttail, "Selected Values of Chemical Thermodynamic Properties", C R C Handbook of Chemistry and Physics, 70th Edition, 1989-1990, C R C Press, pp. D50-D93. 128. A . G . Turnbull and M . W . Wadsley, The CSIRO Thermochemistry System ,Version V , Divis ion of Mineral Chemistry, P. O. Box 124, Port Melbourne, Victoria 3207, Australia. 126 129. R. E . Mesmer and W . L . Jolly, "The Hydrolysis of Aqueous Hydroborate", Inorganic Chemistry, V o l . 1, 1962, pp.608-609. 130. M . M . Kreevoy and J. C. Hutchins, "H2BH3 as an Intermediate in Tetrahydroborate Hydrolysis", Journal of the American Chemical Society, Vol .94, 1972, pp. 6371-6376. 131. H . C. Brown and C. A . Brown, "New Highly Active Metal Catalysts for the Hydrolysis of Borohydride", Journal of the American Chemical Society, V o l . 84, 1962, pp.1493-1494. 132. V . S. Khan, "Some Characteristic Features of Redox Reactions Involving the B H 4 ~ Ions in an Aqueous Medium", Russian Journal of Inorganic Chemistry, V o l . 28, 1983, pp. 1410-1413. 133. N . N . Mal'tseva, Z . K . Sterlyadkina, I. G . Erusalimchik, and V . I. Mikheeva, " Reaction of Sodium Tetrahydroborate with Salts of Noble Metals", Russian Journal o f Inorganic Chemistry, V o l . 24, N o . 3, 1979, pp. 459-461. 134. V . S. Khain, A . A . Vo lkov and V . F . Martynova, "Reduction of Silver (I) with B H 4 " Ions", Izvestiya Akademii Nauk SSSR, Neorganicheskie Materialy, V o l . 27, 1991, pp. 1527-1532. 135. K . N . Mochalov and N . V . Tremasov, "Nature of the Products and Mechanism of the Reaction of B H 4 " with Heavy Metals", Trudy Kazan. K h i m . - Tekhn. Inst., 1967, V o l . 36, pp. 48-55. 136. K . N . Mochalov and N . V . Volkov, "The Role of the Nature of Metals in Reaction with Tetrahydridoborate", Trudy Kazan, K h i m . Technol. Inst. V o l . 40, 1969, pp. 181-185. 137. N . N . Mal'tseva, Z . K . Sterlyadkina and V . I. Miheeva, "Reaction of Sodium Borohydride with Nicke l Chloride in Aqueous Solution", Russian Journal of Inorganic Chemistry, V o l . 11, No.4, 1966, pp. 392-396. 138. Z . K . Sterlyadkinam, N . N . Mal'tesva, O. N . Kryukova and V . I. Mikheeva, "Reaction between Sodium Borohydride and Nicke l Chloride on Heating", Russian Journal of Inorganic Chemistry, V o l . 11, 1966, pp. 531-533. 127 139. W . L . Jolly, "The Preparation of the Volatile Hydrides of Groups TV-A and V - A by Means of Aqueous Hydroborate", Journal of the American Chemical Society , V o l . 83, 1961, pp. 335-337. 140. L . Berka, T. Briggs, M . Mi l la rd and W . Jolly, "The Preparation of Stibine and the Measurement of Its Vapor Pressure", Journal of Inorganic Nuclear Chemistry, V o l . 14, 1960, pp. 191-194. 141. S. R. Gunn, W . L . Jolly and L . G . Green, "The Heats of Decomposition of Arsine and Stibine", Journal of Physical Chemistry, V o l . 64, pp. 1334-1335. 142. V . E . Mironov, Y u . A . Makashev, I. Y a . Maverina, and M . M . Kryzhanovski, "Outer-sphere and Inner-sphere Complexes of Cobalt (II), Nickel (U),and Copper (II)", Russian Journal of Inorgainc Chemistry, V o l . 15, 1970, pp. 668-670. 143. J. A . Bolzan and A . J. Arvia , "Hydrolytic Equilibrium of Metallic Ions-I The Hydrolysis of Co(II) Ion in N a C 1 0 4 Solution", Electrochimica Acta, V o l . 7, 1962, pp. 584-599. 144. G . V . Makovskaya and V . B . Spivakoskii, "Precipitation p H of Basic Salts and Hydrooxide of Copper, Nicke l , Cobalt, and Iron (UI) from Chloride and Nitrate Solutions", Russian Journal of Inorganic Chemstry, V o l . 19, 1974, pp. 585-589. 9-1-145. J. Shankar and B . C. De Sousa, "Short Communications, Hydrolysis o f Co (aq) and Ni 2 + (aq) Ions", Australian Journal of Chemistry, V o l . 16, 1963, pp. 1119-1122. 146. D . D . Perrin, "Ionisation Constants of Inorganic Acids and Bases in Aqueous Solution", 2nd Edition, I U P A C , Chemical Data Series, No . 29, Pergamon Press, New York , 1982, p. 28. 147. G . R. Hedwig and H . K . J. Powell , " A Reinvestigation of the Enthalpy Changes for the Interaction o f the Sulfate Ion with Some Transition-Metal Ions in Aqueous Solution", Journal the Chemical Society, Dalton Transactions, 1973, pp. 798-801. 128 148. H. K. J. Powell, "Entropy Titrations: A Reassessment of Data for the Reaction of the Sulfate Ion with Protons and with Bivalent Metal Ions", Journal the Chemical Society, Dalton Transactions, 1973, 1947-1951. 149. R. M. Izatt, D. Eatough, J. J. Christensen and C. H. Batholomew, "Calorimetrically Determined Log K, AH° and AS 0 Values for the Interaction of Sulfate Ion with Several Bi-and Ter-valent Metal Ions", ", Journal the Chemical Society, Dalton Transactions, 1969, pp. 47-53. 150. G. N. Glavee, K. J. Klabunde, C. M. Sorensen and G. C. Hadjipanayis, "Sodium Borohydride Reduction of Cobalt Ions in Nonaqueous Media. Formation of Ultrafine Particles (Nanoscale) of Cobalt Metal", Inorganic Chemistry, Vol. 32, 1993, pp. 474-477. 151.1. Dragieva, M. SL. Slvacheva and D. T. Buchkov, "Structural Characteristics of Fine Boron-Containing Metal Particles", Journal of the Less Common Metals, Vol. 117, 1986, pp. 311-317. 152.1. Dragieva, "Crystallization Processes in Boron-and Hydrogen-Containing Ultrafine Powders", Jouranl of the Less Common Metals, Vol. 158, 1990, pp. 295-299. 153. P. J. Goodhew, Electron Diffraction, Engineering Materials Software Series, Ml.2, 1987. 129 A P P E N D I X 1 J o i n t C o m m i t t e e P o w d e r D i f f r a c t i o n ( J C P D S ) C a r d of a R a n d o m l y O r i e n t e d C o b a l t ( C u b i c ) 15-806 Co JCPDS-ICDD Copyright (c) 1991 Quali ty: * Cobalt Rad: CoKal Laabda: 1.78897 F i l t e r : Fe Cutoff: Int: Dif f ractoteter I/Icor: Ref: Na t l . Bur. Stand. (U.S.) Honogr. 25, 25 10 (1965) d-sp: ! d A Int. h k 1 ! 2.0467 100 1 1 1 ! 1.7723 40 2 0 0 ! 1.2532 25 2 2 0 ! 1.0688 30 3 1 1 •! 1.0233 12 2 2 2 Sys: Cubic a: 3.5447 A: Ref: Ib id. Dx: 8.789 Da: ea: Ref: nwB: S .G . : Fo3ffl (225) c: A: C: Z: 4 SS/F0M: F5=224C,004,5) ey: Sign: 2V: Color: Dark Pattern at 2! iray C. CAS no.: 7440-4B-4. Saaple was prepared at NBS by heating cobalt oxalate in H2 for 10 sinutes at 300 C. Spectroscopic analys is : 0.1 to 1.02 each of Ni and Sbj and 0.01 to 0.12 each of Al and Fe Ed. , p. 272. Cu type. Tungsten used as internal standard 58.93. VoluaetCD]: 44.54. Herck Index, 8th PSC: cF4. Hwt: Strong l ines : 2.05/X 1.77/4 1.07/3 1.25/3 1.02/1 0.00/1 0.00/1 0.00/1 130 A P P E N D I X 2 J o i n t C o m m i t t e e P o w d e r D i f f r a c t i o n ( J C P D S ) C a r d of a R a n d o m l y O r i e n t e d C o 2 B ( t e t r a g o n a l ) 25-241 JCPDS-ICOO Copyright <c) 1991 Quali ty: i Co B 2 Cobalt Boride d-sp: Rad: Cu Laibda: 1.54056 F i l t e r : Ni Cutoff: Int.- Diffractoiiieter I/Icor: ( 1 9 7 3 ) a V 1 " 9 a ' P h l l i P S R e s e a r c h L a b " ' E i n d h o v e n f Netherlands, .Private Conunicat ion, Sys: Tetragonal S . G . : I4/BCB (140) a: 5.015 b: c: 4.220 A: B: C: Z: 4 Ref: Havinga et a l . , J . Less-Common Met., 27 169 (1972) Dx: 8.053 Da: SS/FOM: F22=17t.045,29) C: .8415 up: ea: Ref: nwB: ey: Sign: 2V: A12Cu type. PSC: tI12. To replace 3-878. Hut: 128.68. VoluietCDl: 106.13. d A ! Int. ! h k 1 3.55 6 1 l 0 2.51 : 18 2 0 0 2.113 30 0 0 2 1.983 100 2 1 1 1.815 7 1 l 2 1.775 3 2 2 0 1.616 13 2 0 2 1.588 15 3 1 0 1.357 1 2 2 2 1.268 8 3 1 2 1.254 2 4 0 0 1.192 20 2 1 3 1.183 7 3 3 0 1.169 18 4 1 1 1.079 3 4 0 2 1.055 3 0 0 4 1.032 11 3 0 i i. 0.9727 •3 2 0 4 0.9202 12 4 1 3 0.9095 9 5 1 0.9068 1 2 2 4 0.8915 1 c J 1 Strong l ines : 1.98/X 2.11/3 1.19/2 2.51/2 1.17/2 1.59/2 1.62/1 0.92/1 131 A P P E N D I X 3 J o i n t C o m m i t t e e P o w d e r D i f f r a c t i o n ( J C P D S ) C a r d of a R a n d o m l y O r i e n t e d B ( r h o m b o h e d r a l ) 12-377 JCFDB-ICDD Copyright (c) 1991 Quali ty: ! 8 d A Int. h k 1 ! 4.25 45 0 0 0 I ! Boron 4.07 100 1 0 i : 3.544 20 0 1 2 : 2.545 70 1 0 4 : 2.479 10 1 1 0 ! ! Rad: CuKa ! ^ : „ ; - . : 1.5418 F i l t e r : Ni d-sp: ! Cutoff: Int: Diffractoaeter I/ Icor: 2.180 4 0 1 5 ': ! Ref: Decker, Kasper, Acta Crys ta l logr . , 12 503 (1959) 2.132 4 1 1 3 ! 2.113 55 0 2 i : 2.029 <2 0 ! Sys: Rhoabohedrsl (Hex) S . S . : R-3a (166) 2 L. 1 1.666 4 1 0 7 : ! a: 4.908 b: c: 12.567 A: C: 2.5605 ! A: B: C: Z: 36 ep: 1.634 8 2 0 5 ! ! Ref: Ib id. 1.603 8 i. 1 i : 1.482 8 0 1 s : i Dx: 2.465 Da: 2.460 SS/FOH: F30=3(.267,41) 1.438 15 2 1 4 ! 1.424 n 3 0 0 j 0 ! ea: nwB: ey: Sign: 2V: j Ref: 1.403 4 0 0 9 : 1.376 8 0 n 7 ! 1.359 10 1 i CAS no.: 7440-42-8. Single-crysta l data taken. Referred to as alpha-rhoabo. L. • i 1.346 15 3 0 3 : ! Boron, low teaperature fora. hkl indices based on t r i p l y p r i a i t i ve hexagonal ! c e l l : a=4.90B, c=12.567. Rhoabohedral c e l l : a=5.057, alpha=58.06. PSC: hR12. j Plus 16 ref lect ions to 0.778. Hwt: 10.81. VoluaetCDl: 262.16. 1.268 12 2 0 8 | 1.230 2 2 2 o : 1.199 <2 2 1 7 : 1.178 4b 2 2 3 ! 1.161 <2 3 1 1 1 L i 1.125 <2 1 2 8 ! d A ! Int. h k 1 ! d A ! Int . ! h k 1 ! d A ! Int . ! h k 1 1.105 ' "i t i . 1 4 ! 0.996 ! 4 ! 3 0 9 ! 0.913 ! <2 ! 4 0 7 1.081 ! <2 ft 10 ! 0.987 ! <2 ! 1 w 7 ! 0.909 ! 2 ! 2 3 5 1.060 ! / I t \L 4 0 1 ! 0.944 ! 2 : 3 i 8 ! 0.906 1 '"i 1 I L I 4 1 3 1.048 t <"i ( L 0 4 ") i. ! 0.932 '. 15 ! 3 0 i. 4 ', 0.880 ! 2 ! 0 2 13 1.008 I / --i t \ L 4 0 4 ! 0.922 ! <2 ! 2 rJ i. 9 1 1 1 ( I Strong l i nes : 4.07/K 2.55/7 2.11/6 4.25/5 3.54/2 1.44/2 1.35/2 0.93/2 132 A P P E N D I X 4 P r o c e d u r e f o r C o l o r i m e t r i c C o b a l t A n a l y s i s The colorimetric cobalt analysis is based on the complexation of cobalt with Nitroso-R-salt ( C i o H 4 ( O H ) ( S 0 3 N a ) 2 N O ) . Cobalt and Nitroso-R salt form a soluble coloured cobalt complex, whose intensity was measured colorimetrically with a Bausch and Lomb Spectronic 20. spectrophotometer. B y comparing the measured intensity with those of known standards, the cobalt concentration was determined. C O L O R I M E T R I C M E T H O D : 1. Pipet 2 ml of sample into a clean, dry 125-ml flask. 2. A d d 20 ml of distilled water. 3. A d d 5 ml o f Nitroso-R-salt solution. 4. B o i l 1.5 minutes. 5. A d d 1ml potassium bromate solution to oxidize and destroy uncomplexed nitroso-R salt. 6. B o i l 1 minute. 7. A d d 1 m l of concentrated nitric acid to destroy other organic complexes other cobalt. 8. Cool solution. 9. Transfer solution to test tube. 10. Measure intensity of the colour with colorimeter at 520 nm wavelength. 11. Compare intensity with intensity of standards. P R E P A R A T I O N OF S T A N D A R D S : Standards and Nitroso-R-salt solution were made up every day, because the colour of cobalt Nitroso-R-salt complex deteriorates with time. Nitroso-R-salt: 0.5%. Dissolve 0.5 g salt in 100 ml distilled water. 133 Potassium Bromate solution: 3 wt.% Dissolve 3.0 g salt in 100 ml distilled water. Blank Disti l led water. Standards 5, 15, 30 ppm cobalt as cobalt sulphate in zinc sulphate solution. The standards were prepared in the same manner as the sample. The relationship between the intensity o f the standards and the cobalt concentration was linear and consistent. 

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