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Kinetics of the dissolution of copper in aqueous aliphatic amines, and catalytic reduction of nickel.. 1959

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KINETICS OF THE DISSOLUTION OF COPPER IN AQUEOUS ALIPHATIC AMINES And CATALYTIC REDUCTION OF NICKEL ION BY MOLECULAR HYDROGEN by SISIR COOMAR SIRCAR A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF SCIENCE i n the Department of MINING AND METALLURGY We accept t h i s t h e s i s as conforming t o the standard required from candidates f o r the degree of MASTER OF SCIENCE. Members of the Department of Mining and Metallurgy. THE UNIVERSITY OF BRITISH COLUMBIA May 1959 ABSTRACT Ki n e t i c s of the D i s s o l u t i o n o f Copper i n Aqueous A l i p h a t i c Amines. An i n v e s t i g a t i o n was made of the d i s s o l u t i o n of copper metal i n aqueous solutions of methyl, e t h y l and n-butyl amine and aranonia, as w e l l as i n solutions of t h e i r aminium and ammonium ions, under oxygen pressure. Study of the k i n e t i c s of the d i s s o l u t i o n was c a r r i e d out over a wide range of concentrations. I t was observed that there are two regions of oxygen pressure dependency of the r a t e . The range of oxygen pressure, where rate i s independent of oxygen pressure, was inv e s t i g a t e d i n d e t a i l . The rate process has been established to obey the r e l a t i o n kjJAV 1 L J . f o r a l l the systems studied. A mechanism f o r the d i s s o l u t i o n reaction has been proposed. ABSTRACT C a t a l y t i c Reduction of N i c k e l Ion by Molecular Hydrogen. The rate of p r e c i p i t a t i o n of n i c k e l from s a l t s o l u t i o n by molecular hydrogen i n the presence of ca t a l y s t has been i n v e s t i - gated over a wide range of concentrations. The v a r i a b l e s studied are concentration of n i c k e l , hydrogen pressure, amount of c a t a l y s t , concentration of- hydrogen ion and temperature. The rate of reduction i s found t o obey an equation of the form: dt The a c t i v a t i o n energy i s found to be 25*2 kcal/mol. The equilibrium constant k f o r the reaction N i + + + H 2 Ni° + 2 H + i s found t o be 4.65 x 10"6 atm" 1 at 140°C. I n p r e s e n t i n g t h i s t h e s i s i n p a r t i a l f u l f i l m e n t o f t h e r e q u i r e m e n t s f o r an a d v a n c e d d e g r e e a t t h e U n i v e r s i t y o f B r i t i s h C o l u m b i a , I a g r e e t h a t t h e L i b r a r y s h a l l make i t f r e e l y a v a i l a b l e f o r r e f e r e n c e and s t u d y , I f u r t h e r a g r e e t h a t p e r m i s s i o n f o r e x t e n s i v e c o p y i n g o f t h i s t h e s i s f o r s c h o l a r l y purposes, may be. g r a n t e d by t h e Head o f my D e p artment o r by h i s r e p r e s e n t a t i v e s . I t i s u n d e r s t o o d t h a t c o p y i n g o r p u b l i c a t i o n o f t h i s t h e s i s f o r f i n a n c i a l g a i n s h a l l n o t be a l l o w e d w i t h o u t my w r i t t e n p e r m i s s i o n . D epartment o f M l n l n g a n d Notallurgy The U n i v e r s i t y o f B r i t i s h C o l u m b i a , V a n c o u v e r S, C a n ada. TABLE OF CONTENTS Ki n e t i c s of the D i s s o l u t i o n o f Copper i n Aqueous A l i p h a t i c Amines. Page INTRODUCTION 1 Scope of Present Work 3 EXPERIMENTAL 6 Preparation of Copper Sample 6 Apparatus 6 Temperature Control 8 Chemical Reagents 6 A n a l y t i c a l Procedure 9 Experimental Procedure 11 Preliminary Results • 11 RESULTS AND DISCUSSION 16 Free Amines 16 E f f e c t of Acid 2k Ammonia 31 Reaction Mechanism 31 REFERENCES 39 APPENDIX I : 67 Data Pertaining to the D i s s o l u t i o n of Copper by Aqueous A l i p h a t i c Amines 67 C a t a l y t i c Reduction of N i c k e l Ion by Molecular Hydrogen INTRODUCTION kO Scope of the Work k2 EXPERIMENTAL 43 Catalyst ^ 3 Reagents . • k3 Table of Contents (cont»d.) Page Analysis . . . 45 Experimental Procedure . . . . . . . . . . 45 Preliminary Experiments • . 47 RESULTS AMD DISCUSSION . 52 (a) Dependence of the rate on nickel concentration . . . . . 52 (b) Effect of the amount of catalyst 52 (c) Effect of hydrogen pressure 57 (d) Dependence of the rate on hydrogen ion concentration • . 57 (e) Temperature dependence . • 57 Use of Equilibrium N i c k e l Concentration for Calculation of the Thermodynamic Equilibrium Constant 61 REFERENCES 66 APPENDIX II 73 Data Pertaining to the Reduction of Nickel Ion by Molecular Hydrogen 73 •FIGURES Kinetics of the Dissolution of Copper in Aqueous Allphatio Amines, No, Page 1, Schematic diagram of the autoclave , * . • . , 7 2, Calibration curve for copper analysis • 10 3, Typical curve for the dissolution of copper . 12 4, Pressure dependence plot for ethyl amine 14 5, Pressure dependence plot for n-butyl amine 15 6, Typical family of rate curves for methyl amine t 17 7, Plot of rate vs concentration for methyl amine 18 8, Plot of rate vs concentration for ethyl amine 19 9, Plot of rate vs concentration for n-butyl amine , , , , , , , 20 10. Plot of QAJ /R VS [A] for methyl amine 21 11. Plot of [A] 2/R VS CA] for ethyl amine 22 12, . Plot of [A] 2/R VS [A] for n-butyl amine , 23 13, Plot of rate vs concentration of methyl amine at constant pH , 25 14,. Relation between the rate and methyl aminiura ion concentration, 27 3.5, Relation between the rate and ethyl aminlum ion concentration , 28 3.6, Relation between the rate and n-butyl aminium ion concentration . . . . , , . , , , , * » 29 17, Plot of £A] 2/R VS [[A] for methyl amine at constant pH , . , . 30 18• Plot of rate vs concentration of ammonia 32 19. Plot of [A] 2/R VS [A] for ammonia , , , , . 33 20, Relation between the rate and ammonium ion concentration , , , 34 FIGURES Catalytic Reduction of Nickel Ion by Molecular Hydrogen, No. Page 1. Schematic diagram of arrangements for holding catalyst and filtering under pressure . . . . . . . . . . 44 2. Calibration curve for the determination of nickel . . . . . . . 46 3. Typical plot showing two types of rate curves 50 4. Typical family of rate curves . , 51 5. Effect of i n i t i a l concentration on the i n i t i a l rate . . . . . . 53 6. Plot of log (initial rate) vs log (initial concentration) . . . 54 7. Effect of amount of catalyst on the i n i t i a l rate . . 55 8. Effect of hydrogen pressure on the i n i t i a l rate . . . . . . . . 56 9. Effect of i n i t i a l hydrogen ion concentration on the i n i t i a l rate. 58 10, Plot of log (initial rate) vs reciprocal of temperature . . . . 59 11, A plot showing approach to equilibrium concentration from "dissolution and reduction of nickel . . . . 60 12, Plot of equilibrium nickel concentration vs reciprocal of hydrogen pressure . . . . . 62 13, Plot of equilibrium nickel concentration vs square of final hydrogen ion concentration 63 TABLES Ki n e t i c s of the D i s s o l u t i o n of Copper i n Aqueous A l i p h a t i c Amines. No. Page 1, Acid d i s s o c i a t i o n constants f o r aminium ion . . . . . . . . . 4 2, Values of the constants k^ and k 2 i n Equation 1 f o r the amines studied. . . . . . . . . . . . . * 24 3, Values obtained f o r k 3 , 26 4, Values of constants f o r ammonia 31 5, Selected values of constants 36 ACKNOWLEDGEMENTS The author i s grateful to the National Research Council f o r the financial aid which enabled this project to be carried out. The author is grateful to the members of the Department of Mining and Metallurgy for their assistance, and i s especially grateful to Dr. D.R. Wiles, under whose guidance and supervision this work has been carried out. The author takes pleasure in expressing gratitude to Dr. W.C. Lin for suggesting the part of work on the reduction of nickel, and taking interest in the problem. Suggestions made by Mr. I.H. Warren are gratefully acknowledged. KINETICS OF THE DISSOLUTION OF COPPER IN AQUEOUS ALIPHATIC AMINES, INTRODUCTION The corrosion of metals has attracted the att e n t i o n of s c i e n t i f i c i n v e s t i g a t o r s f o r over a hundred years. The main e f f o r t and i n t e r e s t i n t h i s problem, however, have been dire c t e d towards prevention of corrosion. In the past few decades, the att e n t i o n of in v e s t i g a t o r s has been focussed^ more on the nature and causes of corrosion processes. Up to the present time, however, d e t a i l e d understanding of the k i n e t i c process involved i n corrosion 2 3 ix 5 has been gained i n only a very few instances, Among the systems which have been studied with a view to determining the k i n e t i c aspects of the corrosion process, the d i s s o l u t i o n of copper has been the most thoroughly investigated, Yamasaki^ was one of the f i r s t t o study the problem of d i s s o l u t i o n of copper i n ammonia so l u t i o n i n the presence of p u r i f i e d a i r . He observed the d i s s o l u t i o n r a t e to be independent of the concentration of ammonia. He concluded ̂ moreover, that the d i s s o l u t i o n process i s a u t o c a t a l y t i c , the cupric ammonia complex a c t i n g as c a t a l y s t . From studies of the d i s s o l u t i o n of copper i n ammonium compounds, Z a r e t s k i i and Akimov'' concluded that three important steps are involved i n the d i s s o l u t i o n process. The steps are ( i ) electrochemical d i s s o l u t i o n of copper with formation of cuprous ammonia,(ii) oxidation of t h i s species to the cupric ammonia complex and ( i i i ) electrochemical reduction of the cupric ammonia complex at the metal surface. - 2 - 8 More recent work of Lane and McDonald on dissolution of copper in ammonia solution led them to believe that no autocatalytic effect i s present and that the dissolution rate is i n fact dependent on the concentration of ammonia. A l l of these investigators have observed that the rate of dissolution is strongly dependent on the s t i r r i n g rate. A detailed study of dissolution of copper in ammonia and in o mixtures of ammonia and ammonium ion has been made by Halpern, 7 He found two regions in the dependence of rate of dissolution of copper on the oxygen pressure. In the range of low oxygen pressure, he observed that the rate of dissolution of copper is dependent on the oxygen pressure and that the rate of dissolution is independent of oxygen pressure at higher pressure. The activation energies found by him for the oxygen pressure dependent rate and oxygen pressure independent rate are 1.33 and 5.54 kcal/mol respectively. From the low activation energy for rates at low oxygen pressure, he concluded that the rate controlling step is a physical process, such as diffusion of oxygen from the bulk of solution to the metal surface. The higher activation energy for the oxygen pressure independent rate has been taken as indicating that some chemical step is involved in the rate controlling process. From his work in the region of chemically-controlled rate, he has proposed the following mechanism: 1. Adsorption of dissolved oxygen on the surface. fast Cu + 1/2 0 2 * Cn - 0 2. Reaction of ammonia or ammonium ion with, the copper-oxygen complex - 3 - on the surface, followed by desorption o f the complex slow Cu - 0 + NH 3 -> ^NH 3 C u + ^ or r Cu - 0 + NH^+ -+ i,Cu++ H f a s t + KOH -* Product f a s t + 0H + -» Product Studies on the k i n e t i c s of the d i s s o l u t i o n of copper by c h e l a t i n g agents were c a r r i e d out by M i l a n t s , ^ i n the range of chemically-controlled r a t e , to elucidate the k i n e t i c process f u r t h e r . His studies i n d i c a t e d that the s t a b i l i t y constants of chelate complexes and d i s s o l u t i o n rate are r e l a t e d to one another. Due to the small number of systems studied, the exact corre- l a t i o n of complexity constants and d i s s o l u t i o n rates has not been established. The mechanism of the k i n e t i c process was found to be analogous to t h a t of ammonia. Scope of Present Work In the l i g h t of the r e s u l t s obtained f o r d i s s o l u t i o n of copper by c h e l a t i n g agents, i t was considered worthwhile t o study the k i n e t i c s of the attack of the copper surface by non-chelating agents, such as the a l i p h a t i c amines, to elucidate the mechanism. It was a l s o a purpose of t h i s work to look f o r possible differences i n the k i n e t i c processes involved i n case of ammonia and a l i p h a t i c amines. The rate of a heterogeneous r e a c t i o n may be c o n t r o l l e d by one or more of s e v e r a l p h y s i c a l or chemical processes (adsorption of gaseous reactant i n s o l u t i o n , transport of d i s s o l v e d reactants from bulk of s o l u t i o n to the surface, r e a c t i o n on the surface, and desorption of reaction product). The present work was planned p r i m a r i l y to f i n d out the chemical processes involved i n the d i s s o l u t i o n of copper by a l i p h a t i c amines. The complexing agents used i n t h i s work were methyl, e t h y l , and n-butyl amines, and ammonia. Present knowledge of the nature of the complexes formed i n aqueous so l u t i o n between a l k y l amines and copper ion i s not w e l l advanced. Ammonia i s known to form several d i f f e r e n t amine complexes with copper,^" and 12 i n d i c a t i o n s f o r complex formation with a l k y l amines have been noted. The constants f o r the d i s s o c i a t i o n RNH 3 + xi RNH2 + H+ K = M faNHj] where R = H, or an a l k y l group, are given i n Table I. i Table I 'Acid d i s s o c i a t i o n constants f o r aminium ions. K (M) (13) (14) (14) (14) Ammonia Methyl amine E t h y l amine n-Butyl amine 5.68 x 10"° 2.30 x I O - 9 1.80 x 1 0 - 9 2.L6 x 10" 9 The order of s t a b i l i t y of copper complexes may be c l o s e l y r e l a t e d t o the s t a b i l i t y of the RNH 3 + species. Since the main aim of the present work was the study of the chemical aspects of the r e a c t i o n , the va r i a b l e s studied were concentration of reactants (amine and aminium i o n concentration) and hydrogen ion concen- t r a t i o n . Some observations were made of the e f f e c t of v a r i a t i o n of oxygen pressure i n order to ensure that the experiments were not ca r r i e d out i n the range of oxygen pressure where d i f f u s i o n of oxygen i s r a t e c o n t r o l l i n g . - 6 - EXPERIMENTAL Preparation of the Copper Sample. In t h i s work were used the same copper specimens as were used by M i l a n t s . ^ These specimens were prepared from conductivity copper, by melting and casting i n a spectroscopic graphite c r u c i b l e and grain r e f i n e d by cold working and annealing. F i n a l l y they were mounted i n bak e l i t e leaving one exposed face whose surface area was about 2 cm. . The p u r i t y ( i n t h i s case 99.99$) and the. grain s i z e were found by Milants not to influence the r a t e . The average grain s i z e of the specimen used i n t h i s work was 0.12 mm. (A.S.T.M. nonferrous grain size standard). Before each experiment, the specimen was polished with 4/0 emery paper, and was l i g h t l y etched with a sodium persulphate ammonia mixture. The geometric area of the exposed surface was measured by means of a microscope with a c a l i b r a t e d eye piece ( c a l i b r a t e d to 0.01 mm.). The area was computed from the mean of several measurements. In p r a c t i c e , the specimens were always etched just before the experiment, and when not i n use, were stored i n a desiccator. Apparatus. The experiments were c a r r i e d out i n an autoclave, shown i n - Figure 1, designed f or working at pressures up to S atmospheres. The auto- clave was made from 316 s t a i n l e s s s t e e l . A titanium l i n e r was used t o contain the s o l u t i o n to avoid corrosion of the s t e e l by the reagents used. The bakelite-mounted copper specimen was held i n place by means of a s t a i n l e s s s t e e l rod, t h i s being screwed i n t o the l i d of the autoclave. H A = shaft B = thermoregulator w e l l C = cooling c o i l D = sampling tube E = copper sample F = impeller G = pump H = r e l a y I = heater J = thermostat F i g . 1, Schematic diagram of the autoclave and temperature c o n t r o l system. - 8 - Ag i t a t i o n was provided by a double turbine-type impeller which swept a c y l i n d r i c a l volume 8 cm. i n diameter. It was driven at 750 RPM f o r a l l the experiments. It was f o u n d ^ that under the experimental conditions used, the s t i r r i n g r a t e had no e f f e c t on the rate of d i s s o l u t i o n of copper. Oxygen, obtained from Canadian L i q u i d A i r Company, was supplied to the autoclave from a c y l i n d e r through a pressure-regulating gauge. An a d d i t i o n a l pressure gauge was mounted on the oxygen supply l i n e t o check the pressure. Temperature Control. The temperature of the r e a c t i o n s o l u t i o n was held at 25-0.1°C by f o r c i n g controlled-temperature water through a s t a i n l e s s s t e e l c o i l f i t t e d i nto the autoclave. The temperature of the c i r c u l a t i n g water was c o n t r o l l e d by a mercury-contact thermoregulator. The water c i r c u l a t i o n was maintained with the help of a c e n t r i f u g a l pump. Chemical Reagents. The methyl, e t h y l and n-butyl amines as w e l l as the carbamate were Eastman Kodak white l a b e l grade. Ammonia and sodium hydroxides were chemically pure q u a l i t y supplied by Nichols Chemical Company. Anhydrous reagent grade sodium perchlorate was supplied by the G. Frederick Smith Chemical Company. The experimental solutions were prepared by d i l u t i n g solutions of known strength t o the required concentration. A l l the solutions were checked by t i t r a t i o n against standard sulphuric a c i d t o the methyl-red end point. In 9 order to avoid changes i n the s a l t e f f e c t at low concentrations, 0.1 or 0.2 mol/1. of sodium perchlorate was added to autoclave s o l u t i o n . The volume of - 9 - the s o l u t i o n was shown^ to have no e f f e c t on the rate of d i s s o l u t i o n . A n a l y t i c a l Procedure. Determination of copper was done c o l o r i m e t r i c a l l y by means of the yellow chelate compound formed by sodium d i e t h y l thiocarbamate. This method has been described by s e v e r a l a u t h ors^>^»-^ It has been reported by these authors that the s t a b i l i t y of the c o l l o i d a l copper carbamate complex i s a f f e c t e d by d a y l i g h t . To avoid t h i s e f f e c t , gum arable was added to s t a b i l i s e 17 the colour. The carbamate-gum arabic s o l u t i o n was prepared i n the f o l l o w i n g way: 0*2 gms. of carbamate was dissolved i n 200 ml of d i s t i l l e d water and t h i s s o l u t i o n was subsequently added to 500 ml of 1% gum arabic s o l u t i o n . The resultant s o l u t i o n was d i l u t e d to 1000 ml and 1 ml of toluene was added. I t was f i n a l l y f i l t e r e d i n the dark and stored i n a c o l d , dark place. A l l the a n a l y t i c a l work was done i n an a r t i f i c i a l l y l i g h t e d room. The s t a b i l i t y of 17 the chelate complex i s reported t o be independent of pH only w i t h i n the range of 7.5 to 9.2. For t h i s reason, a l l analyses were c a r r i e d out at pH 9»0. Samples f o r analysis were prepared i n the following way: To an a l i q u o t of the sample taken from the autoclave, 5 ml of 20% ammonium c i t r a t e was added to complex a c c i d e n t a l traces of i r o n and the pH was adjusted to 9.0 by a d d i t i o n of 1:1 ammonia or 1:1 p e r c h l o r i c a c i d . Then 10 ml of carba- mate s o l u t i o n was added, and the mixture was d i l u t e d i n a volumetric f l a s k to 50 ml. It was mixed c a r e f u l l y and i t s o p t i c a l density was determined with a Beckman model-Dk2 r a t i o - r e c o r d i n g spectrophotometer, at a wave length of 437 m i l l i m i c r o n s . The solutions were found to obey Beer's law over a s u f f i c i e n t range of concentration f o r convenient a n a l y s i s . The c a l i b r a t i o n curve i s given i n Figure 2. The r e s u l t s from t h i s method were found to be reproducible to w i t h i n :*0.5%. I t was a l s o observed that the presence of the amines used - 10 - Copper concentration (mg/l). Calibration ?urve for analysis of copper. - 11 - i n t h i s work does not have any e f f e c t on the o p t i c a l density. Experimental Procedure. The procedure used f o r measuring the rate of d i s s o l u t i o n of copper i n d i f f e r e n t s o l u t i o n s was as follows: The autoclave was f i r s t charged with the desired s o l u t i o n and, with the copper sample i n place, the s o l u t i o n was allowed t o stand under a nitrogen atmosphere u n t i l the required temperature was reached. Then the autoclave was flushed several times very q u i c k l y with oxygen and the oxygen pressure r a i s e d t o the f i n a l desired value. This l a s t was taken as zero time. To f o l l o w the course of the r e a c t i o n , measured volumes of s o l u t i o n were removed from the autoclave by means of a sampling tube, at various time i n t e r v a l s . These samples were analysed f o r copper. Using these r e s u l t s , the known surface area of the copper and the volume of the s o l u t i o n at the time cf sampling, the t o t a l amount of copper disso l v e d per unit area f o r each sampling time was computed. The experiments u s u a l l y l a s t e d about 75 minutes. In a t y p i c a l experiment, the amount of copper dissolved was u s u a l l y between 0,5 and 2.0 x 10~^ mol/1. This i s too small a quantity to cause any s i g n i f i c a n t decrease i n the concentration of the complexing agents. Preliminary Results. In a s e r i e s of preliminary experiments, i t was hoped to determine ( i ) the nature of the rate p l o t s , and ( i i ) the nature of the dependence of the rate on the oxygen pressure. A t y p i c a l rate p l o t i s shown i n Figure 3. The l i n e a r form of 9 10 t h i s plot i s i n agreement with the findings of Halpern and Milants and indicates that the r e a c t i o n i s of zero order i n d i s s o l v e d copper. I t was - 12 - •» ! 1 1 • , r Time (minutes) Conditions: 0.801 (M) methyl amine; sti r r i n g rate 750 RPMj Temperature 25°Cj NaClO^ 0.1 M; oxygen pressure 7.9 atm. Fig. 3, Typical curve for the dissolution of copper. - 13 - found that the rates, as determined from the slope of the r a t e p l o t s , were reproducible to b e t t e r than *2%, q if) E a r l i e r work 7' showed that f o r the r a t e of d i s s o l u t i o n , there exists two separate regions of pressure dependence. In the low pressure region, the r a t e i s proportional to the pressure of oxygen, while at high pressure, the r a t e i s independent of oxygen pressure. Results of experiments performed i n t h i s work and shown i n Figures 4 and 5 show conformity with those of 9 previous observations. This type of behaviour has bean shown to be the r e s u l t of a t r a n s i t i o n from d i f f u s i o n c o n t r o l to c o n t r o l by the reaction on the surface. Since i n the present work, the main aim was to study the chemical attack on the copper surface, a l l experiments were c a r r i e d out at s u f f i c i e n t l y high pressures of oxygen to eliminate d i f f u s i o n c o n t r o l . - 14 - F i g . 4« Pressure dependence curve.for e t h y l amine. \ - 15 - Pressure of oxygen (atm.) Conditions; Temp. 25°C; s t i r r i n g rate 750 RPM; NaClO* O.IM. F i g . 5» Pressure dependence curve f o r b u t y l amine. - 16 - RESULTS AND DISCUSSION Free Amines The values of the rates measured f o r methyl, e t h y l and n-butyl amines are given i n Table IA, IIA and IIIA i n the Appendix. A t y p i c a l f a m i l y of rate curves i s given f o r methyl amine i n Figure 6. The r e l a t i o n between the rate and the amine concentration i s given i n Figures 7, 8, and 9, f o r methyl, e t h y l and n-butyl amines r e s p e c t i v e l y . The s i m i l a r i t y among these curves i s qu i t e evident. The form of these curves indicates a r a t e dependence approximately second-order at low amine concentra- t i o n and approaching f i r s t order at higher concentration. This behaviour suggests that some adsorption pre-equilibrium may be important i n c o n t r o l l i n g the r e a c t i o n rate. I t was found t h a t the data can be best f i t t e d by an expression of the form d [ cu + + ] - R = k l ^ . - ( k 2 L A 3 ) (1) dt k$'J+ 1 where R stands f o r rate o f d i s s o l u t i o n , tAjstands f o r concentration of amines and k]_ and k 2 are constants, the s i g n i f i c a n c e of which w i l l be discussed l a t e r , Equation ( l ) can also be expressed i n the form [£= W + j _ ( 2 ) R k 2 k ^ 2 from which i t i s possible to evaluate the constants and k 2. I t i s seen i n Figures 10, 11, and 1 2 that Equation (2) f i t s the data quite w e l l . The values of k 2 f o r the three amines were evaluated from the slopes of l i n e a r portion of the curves 7, 8 and 9, and from the slopes of the l i n e a r p l o t s of Equation ( 2 ) . The constant k i was evaluated i n each case from the intercept - 17 - 28 Cone, of methyl amine g © o to (0 •H X) ,o 24 20 16 1 2 h • 0.195 M 7.9 atm. o 0.396 M i t 1 1 o 0.568 M 1 1 i t A 0.600 M r t 1 1 0.801 M I t i t • 1.117 M t I t t 0 1.492 M l i i t o 1.492 M t t 1 1 X 1.490 M 5.5 '» 30 40 50 Time (minutes) Conditions: Temperature 25°Cj s t i r r i n g rate 750 RPMj NaCIO* 0.1 Mi 7b- F i g . 6. T y p i c a l f a m i l y of rate curves f o r methyl amine - IS - Concentration of methyl amine (M) Conditions: Temperature 25°C; s t i r r i n g rate 750 RPM; NaClOi,. 0.1 M; oxygen pressure 7.9 atm. F i e . 7. Plot o f rate vs concentration f o r methyl amine - 19 - 0 0.50 1.00 Concentration pf e t h y l amine (M) Conditions: Temperature 25°C; s t i r r i n g rate 750 RPM; NaClO^ 0.1M; oxygen pressure 7.9 atm. F i g . 18 , Plot of rate vs concentration of e t h y l amine ' •••• "" • 1 1 i y • 9Y ^ XT / • 0 0.50 1.00 Concentration of n-butyl amine (M) Conditions: Temperature 25°C; S t i r r i n g rate 750 RPM. NaCIO*,. O.IM; oxygen pressure 6.9 atm. 1.30 F i g . 9. Plot of rate vs concentration of n-butyl amine - 21 - 0 0 . 5 0 1 . 0 0 Concentration of methyl amine (M) Conditions; Temperature 2 5°C; s t i r r i n g rate 7 5 0 RPMj NaClO^ O.lMj oxygen pressure. 7 . 9 atm. 2 F i g . 1 0 . Plot of jAj /R.vs [_AJ f o r methyl amine. - 22 - Fig. 11. P l o t o f |jQ2/R vs 00 f o r e t h y l amine. 0.10 • 0.05c. 0.50 1.00 1.50 Concentrate of n-butyl amine (M) Conditions; Temperature 25°Cj s t i r r i n g rate 750 RPM. NaClO^ 0,10Mj oxygen pressure 6.9 atm. F i g . 12. Plot of [Al 2/R VS [A] f o r n-butylamine - 24 - of the s t r a i g h t l i n e i n Figures 10, 11, and 12. The values of k± and k 2 are given i n Table I I . Table II Values of the constants k]_ and k 2 i n Equation (1), f o r the amines studied. Complexing agent k^M" 1) k 2 (Eig.ciir 2h- 1M- 3-") k^mg.cm-^h-lM-1) from in t e r c e p t s of from F i g s . 7, from Figures F i g s . 10, 11 and 12. 8, and 9. 10, I I and 12. Methyl amine 1.86 18.20 20.8 E t h y l amine 2.27 19.20 20.0 n-Butyl amine 1.85 20.00 23.0 It i s evident from a comparison of the two sets of values f o r k 2 - \ that curves i n Figures 7, 8, and 9 had not yet become quite l i n e a r , so that the values from Figures 10, 11, and 12 are l i k e l y to be the more c o r r e c t . I t i s apparent from the values given i n Table I I f o r the constants that the three amines are very nearly i d e n t i c a l i n t h e i r behaviour. I t i s doubtful whether any r e a l s i g n i f i c a n c e can be attached to the small d i f f e r - ences observed. i E f f e c t of Acid. The e f f e c t of a c i d was studied i n two waysj f o r methyl amine only, at constant pH, constant r a t i o of aminium to amine and f o r a l l three amines at constant free amine concentration. The r e s u l t s of experiments performed at constant pH are shown i n Figure 13 (the rate values are given i n the Appendix, Table IB). This indicates that the v a r i a t i o n of the hydrogen ion concentration has no e f f e c t on the q u a l i t a t i v e nature of the rate dependence. - 25 - l : i r Concentration of methyl amine (M) Conditions; Temperature 25°C; s t i r r i n g rate 750 RPM; Oxygen pressure 7.9 atm. F i g . 13o Plot of rate; vs concentration of methyl amine at constant pH e - 26 - The r e s u l t s of experiments at constant amine concentration are shown i n Figures 14, 15 and 16 f o r methyl, e t h y l and n-butyl amines respect- i v e l y (see Tables IC, IIB, and IIIB i n the Appendix f o r rate v a l u e s ) . The l i n e a r r e l a t i o n s h i p between rate and aminium ion (AH +) concentration suggests that Equation ( l ) be modified s l i g h t l y : R = k l V ^ kltA>l k 2 + k 3 a j (3) where a = ̂ -MSiD. This r e l a t i o n can also be expressed as C j f = CAT + l R k 2 + k 3 a k ^ k 2 + k 3a) (4) A l i n e a r p l o t i s shown f o r methyl amine i n Figure 17. The values of k 3 obtained f o r the three amines are given i n Table III. Table III Values obtained f o r k 3 Complexing Agent Methyl amine E t h y l amine n-Butyl amine Values of k^(mg.cmg2}r1M'"1') From Figures From Calc. from . 13. 14. 15. Equation (4) Equation (3) 140, 153 93 103 160 x 146 93 94 ft These values of k 3 are mean values. Values of a l l concentrations studied are given i n Tables ID, IIB, and IIIB i n the Appendix. The values f o r k 3 determined by d i f f e r e n t methods are i n good agreement with each other. The value of k 3 f o r methyl amine i s s t r i k i n g l y d i f f e r e n t from the values f o r the other amines. - 27 - Concentration of methyl aminium ion (M) Conditions; Temperature 25°Cj s t i r r i n g rate 7 5 0 RPM; Oxygen pressure 7 . 9 atm. F i g . H Relation between the rate and methyl aminium ion concentration. - 28 Fig. 15 Relation between the rate and ethyl aminium ion concentration. - 29 - Conditions: Temperature 25°C; s t i r r i n g rate 750 RPM; Oxygen pressure 6.9 atm. F i g . 16. Relation between rate and n-bntyl aminium ion concentration. - 30 - 0.04 Pi 0.03 0.02 A l AH+3 = 0.10 & and are c a l c . values . O Conditions; 0.25 0.50 Concentration of methyl amine (M) Temperature 25°C; s t i r r i n g rate 750 RPM; Oxygen pressure 7.9 atm. . 17. Plot of [AjP/R VS [Xj for methyl amine at constant pH. - 31 - Values of the rates obtained at constant high methyl amine concen- t r a t i o n and at d i f f e r e n t amounts of aminium species are given i n Appendix.I. These rate values observed i n higher range of concentration do not appear to fol l o w Equation (3). The reason f o r t h i s i s n o t yet c l e a r . Ammonia. It was found that the ammonia system showed a pattern of a k i n e t i c behaviour very s i m i l a r t o that described f o r the a l i p h a t i c amines. The experimental data are presented i n Figures 16, 19, and 2 0 and Tables IVA, IVB, and IVC i n the Appendix. The values of the constants f o r ammonia are given i n Table IV. Table IV. Values of constants f o r ammonia. klClT1) k 2 (mg. c m - 2 ^ 1 ^ 1 ) k 3 (mg, c n f 2 ^ 3 ^ 1 ) From Figure IB From Figure 19 From Figure 2 0 From Equation (3), Mean value x 4 . 6 5 x assumed 4 . 6 5 6 9 . 5 7 7 . 0 x assumed 7 7 . 0 2 4 4 5 , 2 4 2 0 2 4 5 0 From Ref. 9. 8 4 From a l l the experimental r e s u l t s , i t i s found that the nature of the k i n e t i c process f o r d i s s o l u t i o n of copper i s the same f o r ammonia and a l i p h a t i c amines. However, the values of the constants f o r ammonia are . quite d i f f e r e n t from those of the amines. Reaction Mechanism. The form of Equation (3) suggests that the t o t a l r e a c t i o n may involve an adsorption e q u i l i b r i u m followed by t w o p a r a l l e l independent f i r s t - 32 - 0T2I 0T50 Concentration of ammonia (M) Conditions: Temperature 25°C; s t i r r i n g rate 750 RPM; NaClO^ 0.10; oxygen pressure 7.9 atm. Fig. IB. Plot of rate vs concentration for ammonia. - 33 - 0.010 cv 0.005 . 0 0.25 0.50 Concentration of ammonia (M) Conditions: Temperature 25°C; s t i r r i n g rate 750 RFM; MaClO^ 0.10; oxygen pressure 7.9 atm. F i g . 19. Plot of [ T J /R v s \ A 3 f o r ammonia. - 34 - 30. - Cone, of ammonia. • 0.202M A 0.303 M O 0.404 M 0 0 .005 Concentration of ammonium ion (M) Conditions; Temp. 25°C; s t i r r i n g rate 750 RPM; N a C 1 0 4 0.10-j Oppress. 7 . 9 atm. F i g . 20. R e l a t i o n between the rate and ammonium Ion concentration. order steps. These steps may be formulated as follows (S stands f o r copper atom on the s u r f a c e ) . f a s t r S A 1 Step I. S + A SA; k x = L ; . . slow f a s t „ ' Step I I . (a) SA + A ^ SA 2 — > Product slow ^ f a s t (b) SA + AH+ -—v SA 2 + K+ > Product k 3 In addition to these steps, the oxygen adsorption equilibrium must be considered. This i s assumed to be fast and e s s e n t i a l l y complete. There i s no way of deciding unambiguously whether oxygen adsorption step precedes or follows the amine adsorption step. However, i t i s thought l i k e l y . t h i s step precedes the amine adsorption step under the conditions used i n the present study. The occurrence of a p a r a l l e l e q uilibrium step i n v o l v i n g adsorption of AH + species i s not indicated by the observationsthat the d i s s o l u t i o n of copper does not proceed i n the presence of aminium or ammonium ion alone and t h a t the plots of rate vs^AH^are l i n e a r (Figures 14, 15, 16 and 20). From consideration of the above mentioned steps, the r a t e law w i l l be given by the expression R = k 2 [SA^fjO + k 3(SAl(AH+l (5) Considering the surface balance S D = S + SA, and the equilibrium Step 1, the f i n a l rate expression can be formulated as, R = K J A ~ ^ (k£k\+ k3JAH+)3 (°) k $ j l L where R i s expressed i n rate per unit area. This i s i d e n t i c a l with - 36 - Equation (3) derived from the k i n e t i c data. Selected values f o r the constants k f, k 2 and k 3 f o r methyl, e t h y l , n-butyl amine and ammonia are given i n Table V, Table V. Selected values c f constants. Complexing agents k^M" 1) ^ ( m g . c n r 2 ^ 1 ^ 1 ) k 3 (mg. cm-2h""lM-l) k 3 / k 2 Ammonia 4.65 77 2450 31.80 Methyl amine 1.86 20.8 146 7.00 E t h y l amine 2.27 20.0 93 4.65 n-butyl amine 1.85 23.0 94 4.10 I t i s evident from the values of the constants that the length of the carbon chain exerts very l i t t l e influence on the r a t e . I t i s , however, quite apparent that the e f f e c t due to addition of a proton to form the AH + i o n i s very pronounced i n a l l the cases studied. S i g n i f i c a n c e of k]_: According to the proposed scheme f o r the re a c t i o n , k^ i s the equil i b r i u m constant f o r the adsorption process. I t i s apparent from the k^ values that ammonia i s much more strongly adsorbed than the other amines. The adsorption c h a r a c t e r i s t i c s of methyl, e t h y l and n-butyl amines appear to be e s s e n t i a l l y the same. Sig n i f i c a n c e of k 2: The s i g n i f i c a n c e of the values of k 2 i s as an i n d i c a t i o n of attack of surface adsorbed species by those i n the s o l u t i o n . The values of k 2 f o r the amines are almost the same f o r the three amines and widely d i f f e r e n t from - 37 - that of ammonia. I t i s d i f f i c u l t to draw any d e f i n i t e conclusion about the dif f e r e n c e between the k 2 values of amines and ammonia. It may be perhaps l a r g e l y due to a s t e r i c f a c t o r ; the ammonia has a smaller chance of unsuccess- f u l c o l l i s i o n with the re a c t i o n s i t e than have the amines. S i g n i f i c a n c e of k 3: The value of k 3 i n d i c a t e s the attack of surface adsorbed species by AH + or NHV1" species from s o l u t i o n . The di f f e r e n c e of values i n k 3 between the amines and ammonia i s perhaps also due t o a s t e r i c e f f e c t . The diff e r e n c e i n surface attack by A and AH + species are evident from the differ e n c e i n value of k 2 and k 3. Perhaps f o r the l a t t e r species, the e l e c t r o n exchange i s much f a c i l i t a t e d due t o presence of charge on the AH + species. In the l i g h t of the foregoing d i s c u s s i o n , i t appears l i k e l y that the dif f e r e n c e i n order of rea c t i o n observed i n the case of che l a t i n g a g e n t s ^ and that i n the case of amines and ammonia may a r i s e from the d i f f e r e n c e i n the number of f u n c t i o n a l groups present i n these two types of compound. For the chelating agents, consequently, i t appears to be a unimolecular reaction,' the adsorption step being the only observable step. The reaction can be represented by the fo l l o w i n g sequence of steps: f a s t Step I S + S x slow f a s t Step I I S — x — — S i > x ~* Product — x — - stands f o r bidentate c h e l a t i n g agent. These steps are e s s e n t i a l l y s i m i l a r t o those proposed f o r amines except that the slow step does not involve a second molecule. -38 In view of the present r e s u l t s , i t appears that f u r t h e r work ought to be done to extend the same work to other complexing systems. In t h i s sense i t would be i n t e r e s t i n g t o investigate the foll o w i n g : (1) Rate measurements with mixtures of amine and aminium ion (^A"\ + ^AH"^ = constant) over a wide range of pH. (2) Rate measurements with mixtures of primary amines and secondary amines. (3) Investigation of the rate process with 3- and 4-carbon chain diamines. - 39 - REFERENCES 1. Evans, U.R., Introduction to Corrosion, Edward Arnold and Co, (1951). 2. King, C.V. and Hillner, E,, J . Elec. Chem. Soc. 103. 261 (1956), 3. Gatos, H.C., J. Elec. Chem. Soc. 103. 286 (1956). 4. Cretella, M.C. and Gatos, W.C., J. Elec. Chem. Soc. 105. 487 (1958). 5. Smith, Tennyson and H i l l , R. George, J. Elec. Chem. Soc. 105. 117 (1958). 6. Tamasaki, E., Science reprint; Tohuku Imperial University, 9_> 169 (1920). 7. Zaretskii, E. and Akimov, G., J. Appl. Chem. USSR 11, 1161 (1938). 8. Lane, R.W. and McDonald, H.J., J. Amer. Chem. Soc. 68, 1699(1946), 9. Halpern, J., J. Elec. Chem. Soc. 100. 421 (1953). 10. Milants, H.Y., M. Sc. Thesis, Department of Mining and Metallurgy, University of British Columbia (1958). 11. Bjerrum, J . , Metal Amine Formation in Aqueous Solution. Copenhagen (1941), 12. Sidgewick, N.V., The Chemical Elements and Their Compounds. Vol. I f Oxford University Press (1950). 13. Conway, B.E.., Electrochemical data. Elsevier Publishing Co. (1952). 14. Kolthoff, I.M. and Stenger, V.A., Volumetric Analysis Vol.1, Interscience (1942). 15. Hoar, T.P., Analyst. 62, 657 (1937). 16. High, T.H., Analyst 72, , 60 (1947). 17. Sandell, E.B,, Colorimetric Determination of Traces of Metals, Interscience (1950). V CATALYTIC REDUCTION OF NICKEL ION BY MOLECULAR HYDROGEN INTRODUCTION The p o s s i b i l i t y of reduction of metal ions from s o l u t i o n by molecular hydrogen has been recognize'd f o r a long time. Among e a r l y 1 2 3 in v e s t i g a t o r s , Ipatiew ' ' has perhaps made the greatest c o n t r i b u t i o n . In recent years, more d e t a i l e d studies of these reduction processes have k 5 been made by Shaufelberger, Mackiw, L i n and Kunda, and Knacke, Fawlek and Sussmuth.^ An excellent review of hydrometallurgical processes at high 7 pressure i s by Forward and Halpern. E a r l y attempts of Ipatiew to displace metals from solutions at high pressures and temperatures met with varying degrees of success. His work was u s u a l l y done at high temperature (300°C) and pressure (7000 p s i ) and f o r times as long as se v e r a l days. He succeeded i n p r e c i p i t a t i n g by hydrogen reduction, e i t h e r as the metal or as a basic s a l t , the fol l o w i n g metals; platinum, i r i d i u m , copper, n i c k e l , cobalt, lead, t i n , arsenic, antimony and bismuth. Shaufelberger^ - has investigated the p o s s i b i l i t y of reduction of copper and n i c k e l by molecular hydrogen. From examination of the thermo- dynamic data, he concluded that such reduction i s possible at high tempera- ture (•vr200°C) and at s u i t a b l e hydrogen i on concentration. The e f f e c t j o f hydrogen ion concentration and complex formation were investigated by him to a r r i v e at a suitable system f o r the complete removal of the metal i o n from s o l u t i o n . - u - The k i n e t i c s of the reduction of n i c k e l i n ammoniacal s o l u t i o n was studied by Mackiw, L i n and Kunda.5 The conditions under which t h i s work was done were s i m i l a r to those adopted by S h e r r i t t Gordon Mines Ltd., f o r recovery of n i c k e l . Knacke, Pawlek and Sussmuth 0 investigated the k i n e t i c s of the homo- geneous reduction of s i l v e r and copper and heterogeneous c a t a l y t i c reduction of n i c k e l , from t h e i r sulphate s o l u t i o n s . Tho r e s u l t s f o r p r e c i p i t a t i o n of 8 9 s i l v e r and copper are not in agreement with those found by l a t e r workers. ' The rate of reduction of copper and s i l v e r was found to be dependent on the pressure of hydrogen while according to Pawlek and co-workers, i t depends on the square root of hydrogen pressure. For the reduction of n i c k e l , the value of the a c t i v a t i o n energy was found 0 t o be 4.12 kcal/mol, which suggests that i n h i s work, the reduction rate was d i f f u s i o n c o n t r o l l e d . The rate equation proposed by these workers • v _ 2 1/2 -4.12 - J k - 5.59 x 10 2 C N 1 + + P H 2 e does not suggest that hydrogen ion has e f f e c t on the rate of reduction. 8 Halpem and Peters have investigated the homogeneous a c t i v a t i o n of hydrogen by cupric i o n , and the reduction of copper from the s o l u t i o n . The product can be copper or cuprous oxide depending on the pH of the s o l u t i o n and the temperature of the reacti o n . I t has also been found from a study of the copper sulphate system,^ that the r e a c t i o n process i s r e v e r s i b l e . This r e v e r s i b i l i t y i s strongly hydrogen-ion-concentration dependent. Reduction of the metals, copper, s i l v e r and mercury has been found to be homogeneously catalysed by the metal ions and by t h e i r complex ions. But f o r reduction of n i c k e l or cobalt, i t has always been necessary to add a heterogeneous c a t a l y s t . Among c a t a l y s t s used are n i c k e l powder, ferrous sulphate, c o l l o i d a l graphite and hydroquinone,^ In the work reported here, n i c k e l powder i s used as a c a t a l y s t . Scope of the Work, In the present work an attempt was made to elucidate the k i n e t i c s of reduction of n i c k e l from s a l t solutions by molecular hydrogen i n the presence of n i c k e l powder c a t a l y s t . In the l i g h t of previous work, i t was thought worthwhile to study the reduction of n i c k e l from a buffered s o l u t i o n and a few unbuffered s o l u t i o n s . It i s also an aim of t h i s work to evaluate the equilibrium constant f o r the equilibrium N i + + + H 2f=^ Ni° + 2 H +, at high temperature. EXPERIMENTAL Apparatus. The reaction studies were conducted in a one-gallon stainless-steel 12 autoclave manufactured by Autoclave Engineers inc. Additional fittings included an arrangement to drop the catalyst from a sealed glass capsule and an arrangement in the sampling tube to collect the sample under pressure. These are shown in Figure 1. Since for the system studied, the stainless steel autoclave may act as a catalyst, a glass liner with a titanium baffle was used to contain the solution. The impeller, thermocouple well and sampling tubes were also made of titanium. The autoclave was heated by a gas ring burner, and the temperature was controlled to within il . 5°C using a Leeds and Northrup micromax controller. A l l pH measurements were made at room temperature, using a Beckman model GS pH meter. Catalyst. The catalyst used in these experiments was standard *A» carbonyl nickel powder, prepared by the Mond Nickel Company, Ltd. These are spherical particles, and of 100% -350 mesh. The surface area was found* to be 1.85 nrVgm. Reagents. The nickel acetate, acetic acid, and anhydrous sodium acetate used in these experiments were a l l r -rent grade, obtained from the Allied Chemical and Dye Company and the Niclyls Chemical Company. Other incidental chemicals were also reagent grade. Hydrogen gas supplied by the Canadian Liquid Air Company was used without further purification. A l l the solutions for •k The author is grateful to Mr. Peter Hennes for surface area determination. - hk A = glass capsule B = 0•ring arrangement C = sampling tube D = autoclave w a l l E = gas i n l e t s t o push the capsule F = wire gauge G = glass wool H = asbestos I = t o t a l f i l t e r i n g device. F i g . 1. Schematic diagram f o r holding c a t a l y s t and f i l t e r i n g under pressure. - 45 - experiments were prepared by diluting solutions of known strength. Measured volumes of sodium acetate, acetic acid and nickel acetate were transferred to a 2000 ml volumetric flask and diluted up to volume to get a solution of the desired concentration. Distilled water was used for preparation of a l l the solutions. Analysis. Determination of pickel was done colorimetrically using dimethyl glyoxime. ' A Beckman Model DK -2 ratio-recording spectrophotometer was used for measurement of optical density. Measurements were made at a wave- length of 450 millimicrons. The analysis was performed in the following way: To an aliquot of the sample were added 2 ml of saturated bromine water 4.5 ml of concentrated ammonia, 7 ml of 1% solution of dimethyl glyoxime in alcohol, and 5 ml of alcohol, and the whole was diluted to 25 ml. Measurements were made within 10 minutes of mixing. Beer's law is obeyed, as seen from Figure 2, over a sufficient range of concentration to permit analysis following appropriate dilution. Experimental Procedure, The experiments were done in the following way; the required amount of catalyst was weighed and placed in a glass capsule and the capsule was fi l l e d with distilled water and sealed. Then i t was placed in the autoclave as shown in Figure 1, The autoclave was charged with the required solution. After the autoclave was sealed, i t was flushed several times with hydrogen and was heated under hydrogen at a pressure of about 200 psi. It was flushed again at about 80°C, and finally heated up to the required temperature under 200 psi of hydrogen pressure. Samples were removed after attainment of the final temperature to ascertain the zero time concentration. The capsule of - U6 - - 4 7 - catalyst was injected by hydrogen pressure into the autoclave where i t was smashed by the impeller. The hydrogen pressure was then adjusted to the required value. The time when the catalyst was dropped into solution was taken as zero time. Samples were collected through the sampling device (through which samples are fi l t e r e d out from solution under hydrogen pressure), at suitable time intervals for nickel analysis by the method already described. Preliminary Experiments. Preliminary experiments were conducted to examine (i) the nature of the reduction process in unbuffered solutions (solutions of nickel sulphate and nickel perchlorate), ( i i ) the variation of the pH of the solution result- ing from the reduction of nickel ion, ( i i i ) the variation of the reduction rate using different types of catalysts and (iv) the nature of the rate curve. (i) Reduction from nickel sulphate and perchlorate solutions: Experiments using nickel perchlorate and sulphate solutions were found to give only a very small amount of reduction (10% of a 0.1 M solution at 160°C). It was observed that during this process, the pH of the solution f e l l to about 2.4 and the reaction stopped. For perchlorate solutions s t i l l less reduction was found. For this reason, further work described in this thesis was done with buffered solutions. ( i i ) Variation of pH: A l l the pH measurements were done at room temperature. It was found that the i n i t i a l and f i n a l pH of the buffered system after reduction differed only slightly (for example, from pH 4.25 to 4.15). This difference increases as the amount of nickel reduced increases. In spite of the small change i t was.felt desirable to use only the i n i t i a l rate in studying the mechanism since only i n the i n i t i a l part can the conditions for reaction be known with - 48 - any certainty. ( i i i ) Different types of catalyst: Three types of nickel powder were tried as catalyst. (a) Dense nickel powder: These were dense granular nickel powders obtained from Sherritt Gordon Mines Ltd. The particles used for this test were 20-30 microns in diameter. These were annealed in a hydrogen atmosphere before use. It was found that this catalyst was not active enough to make the reduction of nickel ion from solution possible. Instead, dissolution of nickel powder continued for a few hours. It was found that starting with two different i n i t i a l concentrations of nickel, the dissolution approaches a type of pseudoequilibrium. This powder was not used because of i t s lack of catalytic activity. (b) 'Nucleating powder': This powder, obtained from Sherritt Gordon Mines Ltd., was less than three microns i n diameter and of irregular shape. This catalyst was found to be very active, but reproducibility of rate measure- ments was poor. (c) Carbonyl nickel powder: (This catalyst has already been described). It was found to be quite active and as well gave good reproduci- b i l i t y . The difference in the reproducibility of rate measurements for catalyst (b) and (c) i s perhaps due to the nature of the catalyst surface and it s shape. The carbonyl nickel catalyst was used for a l l the experiments described hereafter. (iv) Nature of the reduction curves: The rate curves i n which nickel concentration i s plotted against time, were found to show an induction period of variable length. These curves were of two types, (a)'at high s t i r r i n g rate (75° RPM), the nickel concentration in solution was found to increase considerably and then to - 49 - decrease, (b) at lower s t i r r i n g r a t e (400 RPM), the n i c k e l concentration remained constant f o r a few minutes and then decreased. In a l l cases, however, neither the measured reduction rate, nor the f i n a l equilibrium concentration was dependent on the s t i r r i n g r a t e . The two types of rate curve are shown i n Figure 3. I t i s also made clear i n Figure 3 how the rates were evaluated from the data. The reason f o r there being two types of rate curves was not ascertained. However, i t may be conjectured that the cause may l i e i n the presence of an oxide layer on the ca t a l y s t surface, which i n one case i s being dissolved before being reduced and i n the other case i s reduced i n s i t u . I t seems inconsistent, however, that pre-treatment of the cat a l y s t (the powder was treated with warm a c e t i c a c i d , washed with a l c o h o l , d r i e d at 60°C) d i d not appear to a f f e c t the duration or the nature of the induction period. This may s i g n i f y that e i t h e r the oxide f i l m r e s i s t s t h i s treatment or forms a f t e r t h i s treatment at a very r a p i d r a t e . It was found that although the duration of the induction period was v a r i a b l e , measured values of the i n i t i a l rates were consistent. Table I i n the Appendix gives values showing s t i r r i n g e f f e c t , r e p r o d u c i b i l i t y and e f f e c t of chemical treatment. Figure 4 gives a t y p i c a l f a m i l y of rate curves ( a l l at 400 RPM). Independent evaluation of rates from these curves gave values which, though d i f f e r e n t by as much as 20%, were consistent among themselves. Time (minutes) Conditions: Temperature 140°C; Catalyst 10 gms/1400 ml; Hydrogen 13.55 atm; Acetic acid 0.2 M; Total acetate 0.06 M; Initial, nickel 0.005 M. Fig. 3. Typical plot showing two types of rate curves. - 51 - Time (minutes) Conditions: Temp. 140°C; Catalyst 10 gms/1400 ml; Hydrogen 13.55 atm; Acetic a c i d 0.2 (M); T o t a l acetate 0.06 (M); S t i r r i n g rate 400 RPM; i n i t i a l pH 4.25. F i g . 4. T y p i c a l f a m i l y of rate curves RESULTS AND DISCUSSION The variables studied i n an attempt to elucidate the mechanism of reduction of n i c k e l from s o l u t i o n by molecular v-\irogen i n the presence of a n i c k e l c a t a l y s t were (a) i n i t i a l n i c k e l concentration, (b) amount of c a t a l y s t , (c) pressure of hydrogen, (d) i n i t i a l hydrogen ion concentration and (e) reaction temperature. The experimental data are given i n Tables I to VTII i n Appendix I I . (a) Dependence of the rate on n i c k e l concentration. The i n i t i a l rate f o r any concentration was obtained from the rate p l o t s as shown i n Figure 4. These runs were made at 140°C, 13.55 atm. hydrogen pressure, 10.00 gms of c a t a l y s t . i n 1400 ml s o l u t i o n and. s o l u t i o n pH of 4.25 and varying i n i t i a l concentration of n i c k e l . The v a r i a t i o n i n i n i t i a l rate with i n i t i a l n i c k e l concentration i s shown i n Figure 5. This p l o t shows a l i n e a r dependence of the rate on the n i c k e l concentration. The plo t passes through the value of the equilibrium n i c k e l concentration f o r the conditions used. The f i r s t order dependence on n i c k e l concentration i s confirmed by a pl o t of log (cone, of n i c k e l ) vs log ( i n i t i a l rate) shown i n Figure 6. The measured slope of t h i s l i n e , 0.93 i s s a t i s f a c t o r i l y close to 1. (b) E f f e c t of the amount of c a t a l y s t . A number of runs were done varying the amount of c a t a l y s t , at a constant i n i t i a l n i c k e l concentration o f 0.005 (M) with other conditions being the same as given above. F i r s t order dependence of the rate on the c a t a l y s t area i s shown by the l i n e a r i t y of the p l o t given i n Figure 7. The f a c t that the st r a i g h t l i n e passes through the o r i g i n i s consistent with the observation that no r e a c t i o n occurs i n the absence of a c a t a l y s t . - 53 I n i t i a l n i c k e l concentration (M) Conditions; Temperature 140°C; Catalyst.10 gms/1400 ml; Hydrogen 13.55 atm; Acetic acid 0.2 (M); T o t a l acetate 0.06 (M); I n i t i a l pH 4.25. F i g . 5. E f f e c t of i n i t i a l concentration on the i n i t i a l r a t e . -4 .31 - — i — J 1 1 1 — - 2 . 9 - 2 . 7 - 2 . 5 - 2 . 3 - 2 . 1 - 1 . 9 log C Ni++ Condition's; Temperature 140 PC; Catalyst 10 gms/1400 ml; Hydrogen 13 .55 atm; Acetic a c i d 0 . 2 (M); T o t a l acetate 0 .06 (M); I n i t i a l pH 4 . 2 5 . F i g . 6. Plot of l o g ( i n i t i a l rate) vs l o g ( i n i t i a l c o n e ) . 5 10 15 Amount of catalyst (gms/1400 ml) Conditions; Temperature 140°C; Hydrogen 13.55 atm; Acetic acid 0.2 (M); Total acetate 0.06 (M); Initial nickel concentration 0.005 (M); Initial pH 4.25. Effect of amount of catalyst on the i n i t i a l rate. - 56- (c) E f f e c t of hydrogen pressure. Results of experiments conducted at d i f f e r e n t hydrogen pressures, showed that the rate i s d i r e c t l y proportional to the square root of the hydrogen pressure. The data are presented i n Figure 8. At s u f f i c i e n t l y high pressures, the rate i s seen to drop from t h i s dependency, which may i n d i c a t e saturation of ca t a l y s t surface. This type of dependence suggests that the rate-determining step may be one i n v o l v i n g atomic hydrogen. (d) Dependence of the rate on hydrogen ion concentration. Experiments designed to t e s t the dependence of the rate on hydrogen i o n concentration were performed under the same conditions as described above, but with d i f f e r e n t i n i t i a l hydrogen ion concentrations. The r e s u l t s shown i n Figure 9 show that the rate decreases i n an approximately l i n e a r fashion with increasing concentration of hydrogen i o n . (e) Temperature dependence. Temperature v a r i a t i o n was studied to determine the a c t i v a t i o n energy and thereby perhaps to be able to make some d e c i s i o n as to the nature of the rate l i m i t i n g step. The measurements were.made over the temperature range 130 to l60°C. The a c t i v a t i o n energy determined from Arrhenius' p l o t shown i n Figure 10 appears t o be 25*2 kcal/M. This value f o r the a c t i v a t i o n energy strongly supports the view that the r a t e process measured i s not one which i s c o n t r o l l e d by d i f f u s i o n . From the above experimental r e s u l t s , the rate equation can be expressed as i M l i l - - k ^ M ^ t N i ^ . k a C s ] (V) 4 t - 57 - Conditions? Temperature 140°C; Catalyst 10 gms/L400 ml; I n i t i a l n i c k e l concentration 0.005 M; Acetic a c i d 0,2 (M); T o t a l acetate 0.06 M; I n i t i a l pH 4.25. F i g . 8. E f f e c t of hydrogen pressure on the i n i t i a l r a t e . Fig, 9« Effect of i n i t i a l hydrogen ion concentration on the i n i t i a l rate. - 59 - l/T x 10 3 Gondii ions: Catalyst 15 gms/1400 ml; Hydrogen 13,55 atm; Initial concentration of nickel 0,005 M; Acetic acid 0.2 (M); Total acetate 0.06 (M); Initial pH 4,25. Fig. 10. Plot of log (initial rate) vs reciprocal of temperature. I n i t i a l pH 4.25. F i g . 11. A plot showing approach to equilibrium from dissolution of nickel and reduction of nickel. - 61 - where ^Ni + +~] and [H +]are: concentrations and Qsl is the catalyst surface area. Although the mechanism of the reduction reaction can not be deduced with any assurance from the above presented information, i t i s now clear that the , N i + + activated complex is of the form surface £ . The reverse reaction - N H nickel dissolution appears to go through an activated complex of the type / H+ H i Pawlek; but i n this work, importance of hydrogen ion is clearly shown to be surface / , The form of the rate equation is similar to that found by X N involved i n the kinetics of rate of reduction. Use of the equilibrium nickel concentrations for calculation of the thermo- dynamic equilibrium constant; The experiments from which the nature of the dependence of the rate on the pressure of hydrogen and on the hydrogen ion concentration has been derived, were usually prolonged for a sufficient time to obtain the equilibrium nickel concentration. A plot showing the approach to equilibrium nickel concentration from reduction of nickel as well as dissolution of nickel i s shown i n Figure 11. From the values of the equilibrium nickel concentration and corresponding hydrogen pressure and hydrogen ion concentrations, an attempt was made to evaluate the ; equilibrium constant for the reaction • N i + + + H 2 ̂  Ni° + 2 H + M 2 where (̂ Ni"1""̂  i s -the equilibrium nickel concentration £H +~\ i s the hydrogen ion concentration ^PH 2~\ i s the' pressure of hydrogen - 62 - 1/P„ (atm - 1) x 10 2 Conditions; I n i t i a l pH 4.25; temperature 140°C; i n i t i a l nickel concentration 0.005. Fig. 12. Plot of equilibrium nickel concentration vs reciprocal of hydrogen pressure. 0 k 6 8 10 12 Concentration of (H+) (M) x 10 8 Conditions; Temp. 140CC; catalyst 15/gms/1400 ml; Hydrogen 13.55 atm.; i n i t i a l concentration of nickel 0.005 (M). Fig. 13. Plot of equilibrium nickel concentration vs final hydrogen ion concentration. - 64 - A plot of the equilibrium concentration of nickel in solution against reciprocal of hydrogen pressure (at constant \^H+]) is given in Figure 12, It is seen to be a straight line passing through the origin in agreement with expectation,. The equilibrium nickel concentration was found to be a linear function of the square of hydrogen ion concentration as is seen in Figure 13. The straight line does not pass through the origin, since the values used for the hydrogen ion concentrations are those measured at room temperature. The difference will be a constant factor related to the intercept.., Hydrogen ion concentration from the plot was found to be 1.6 x 10"^ a t 140°C (the room -5 temperature value being 7 x 10 ). The valuejof constant k calculated with the help of corrected hydrogen ion concentration and slopes of plots 12 and 13, are 4.7 x 10"^ and 4.6 x 10"^ respectively. The value of constant k was calculated from thermodynamic data available for room temperature (extra- polated to 140°C, using relation A F = A F T - (T 2 - T n ) A S ) 0 The value T-z x l obtained from this calculation is 7.6 x 10"^0 The agreement between experi- mental and calculated values is satisfactory, since validity of extending thermodynamic data to higher temperature, and of the assumption of the activity of nickel to be equal to that of total nickel present are very uncertain. The agreement indicates, however, that the assumptions used are not grossly incorrect. £ Refs. 15 and 16. FNi++ = -11.53 kcaly^ HNi++ = -15.3 kcal/^y S N i + + ~ -38.1 cal/degree SH2(gas) = +31»2 cal/degree % i ( s o l i d ) ^ +7.1 cal/degree - 65 - I t i s apparent t h a t be fore f u r t h e r work i s done on t h i s problem, t h e n a t u r e o f t h e c a t a l y s t s u r f a c e must be s t u d i e d i n an attempt t o understand and e l i m i n a t e t h e v a r i o u s u n c e r t a i n t i e s and anomalies w h i c h p e r s i s t e d throughout t h e present work . A f t e r e l i m i n a t i o n o f t h e t r o u b l e s a r i s i n g out of u n c e r t a i n t y i n the nature o f catalyst s u r f a c e , f u r t h e r work may be e n l i g h t e n i n g i n o b t a i n i n g a more q u a n t i t a t i v e approach t o t h e r e a c t i o n mechanism. REFERENCES 1. Ipatiew, V. and Verkhovskii, V., Ber. dtsch. Chem. Ges 4Jt, 1755 (1911). 2. Ipatiew, V., J . Soc. Phys. Chem (Russ.) 43_, 1746 (1911). 3. Ipatiew, V. and Ipatiew, V.V. ( J r . ) , Ber. ctsch. Chem, Ges. 62B. 386 (1929) 4. Shaufelberger, F.A., J . Met. 8, 695 (1956). 5. Mackiw, V.N., L i n , W.C. and Kunda, W,, J . Met. 9, 786 (1957). 6. Knacke, 0, Pawlek, F. and Sussmuth, E,, Z . f . Erzbergbau u, Metallhuttenws, 2, 566 (1956). 7. Forward, F.A. and Halpem, J , , Trans, Inst, Min, and Met., 66, Part 5, 191 (1956-57). 8. Peters, E, and Halpem, J . , Can, J . Chem, 3Jt» 554 (1956). 9. Webster, A.H, and Halpem, J , , J , Amer, Chem, Soc, 60, 280 (1956). 10. Macgregor, E.R. and Halpem, J . , Trans. A.I.M.E. 212. 244 (1958). 11. Kaneko, Thomas M. and Wadsworth, M.E., J . Phys. Chem, 60, 457 (1956), 12. Peters, E,, Ph.D. Thesis, U n i v e r s i t y of B r i t i s h Columbia, 1956, 13. Mitchell," A.M. and Mellon, M.G., Ind. Eng. Chem, Annal, Ed, 17., 380 (1945) 14. Sandell, E.B,, Colorimetric Determination of Traces of Metals. Interscience, 2nd Ed. (1950), '15, Latimer, W.M., Oxidation P o t e n t i a l , 2nd Ed., Prentice H a l l (1952). 16, Kubaschewski, 0. and Evans, E.L., M e t a l l u r g i c a l Thermochemistry. Butterworth Springer (1951), - 67 - APPENDIX I Experimental data pertaining to dissolution of copper Table I A Methyl amine Run No. Oxygen pressure (Atm. ). Ionic Strength (M) Amine £AJ concentration (M) Rate (R) (Mg.cm~2h":L) [A) 2 R 1" 7.9 0.10 0.568 5.80 0.05560 2 11 1.117 15.60 0.07998 3 11 0.801 10.20 PJD6290 4 • t » » 0.600 6.60 0.05450 5 t • 0.396 3.40 O.O4613 6 11 0.195 1.20 0.03169 7 11 1.492 22.60 0.09850 7R 11 » » 11 22.70 0.09800 23 5.5 11 22.90 0.09700 Table IB Methyl amine Oxygen pressure - 7.9 atm. . . Run I o n i c . , Free amine [A] Aminium C&H+J Rate (R) [j^j No. s t $ | n g t h Conc'n. (M) Conc'n. (M) (Mg.cm-V1) R 16 0.10 0.192 0 , 0 4 0 2.10 0,0176 17 0.20 0.384 0 . 0 8 0 7.80 0.1880 18 " 0.576 0 . 1 2 0 16.70 0.1980 19 " 0.480 0.100 12.60 0.183 25 0.1 0.600 0 . 0 6 0 11.60 0.0310 26 »t 0.400 0.040 6.20 0.0258 27 •» 0.700 0.070 15.20 0.0322 29 , f 0.500 0.050 8,60 0,0291 - 68 - Appendix I (cont'd.) Oxygen pressure - 7.9 atm. Table IC Methyl amine Run No. Ionic strength Free amine CKJ Concentration (M) Aminium Ion [\F +J' Concentration (M) Rate (Mg.cmSh-1) 25 28 30 20 21 17 0.10 11 11 0i20 t • 11 0.600 _ f t m t f 0^384 11 t ? 0.060 0.100 0.030 0.120 0,040 0.080 11.60 15.20 9.20 10.30 5.70 7.80 Table ID Methyl amine Oxygen pressure - 7,9 atm. . Run Ionic Free amine, Aminium ion QtHf) kaOng.cm^h-^-M-1) No. strength Concentration (M) Concentration (M) where (M) ki=1.86,k2=20.8 25 0.100 0.600 0.060 .140 28 ••»•» 0.030 160 30 I I I I f t 137 26 f t 0,400 0.040 156 27 11 0.700 0.070 124 29 1 f 0.500 0.050 136 20 0.200 0.384 0.120 140 21 11 11 0.040 . " ' 142 17 11 11 0,080 135 16 0.100 0.192 0.040 100(excluded from mean) 18 0.200 0.576 0.120 169 19 11 0.480 0.100 . 167 - 69 - Appendix I (cont'd.) Table IE Methyl amine. Ionic strength - 0.20(M) Oxygen pressure = 7 . 9 atm. • , Run Free amine Aminium ion Rate No. Concentration (M) Concentration (M) (mg.cm~2h~l) 8 0.800 0.195 50.0 9 »» 0.050 18,0 10 *» 0.125 52.0 11 •» 0.025 15.0 12 • " 0.040 17.0 13 " 0,100 55.6 14 " 0.075 48.8 15 " 0.05 18.0 Table IIA E t h y l amine Ionic strength = 0.10(M) Run Oxygen Amine £k) •Rate (R) C a 1 2 No. pressure (atm.) Concentration (M) (mg.cm h~l) . H 1 7.9 0.203 1.30 6-0316 2 11 0.400 3.60 0.0445 3 11 0.757 9.00 0.0636 4 t» 0.916 12.60 0.0666 5 1.71 11 6.90 6 2.74 11 11.00 : O.R. 7.90 0.535 6.40 0.0447 11 11 1.040 16.60 0.0650 - 70 - Appendix I (cont'd.) Oxygen pressure = 7.9 atm. Table IIB E t h y l amine Run Ionic Free amine No. Strength Conc'n. (M) (M) Aminium ion Conc'n. (M) Rate mg. cm' ~ 2 h ~ l where k 3 (mg.cm^h - 1^ 1) kx=2.27, k2=20.0 8 0.350 0.210 0.100 4.50 97.0 9 0.200 0.210 0.050 2.80 90.0 10 0.200 0.210 0.150 5.80 92. Ionic strength = 0.10 (M) Table IIIA n-Butyl amine Run Oxygen Amine [£) Rate (R) [JJ2 No, pressure Concentration (M) (mg.cm~2h~l) ~ (atm.) . 1 6.9 0.326 2. 80 .0381 2 » » 0.958 14.30 O.064I 3 » » 15.00 0.0620 4 • « 0.630 7.60 0.0523 5 « » 0.187 1.10 0.0318 6 » « 1.641 32.30 0.0844 7R • » 1.293 20.70 0.0809 8 1.71 11 10.20 9 2.74 »t 15.80 - 71 - Appendix I (cont'd.) Table IIIB n-Butyl amine Ionic strength = 0 .20 (M) Oxygen pressure = 6 . 9 atm. Run Free amine Aminium ion Rate (R) k 3(mg.cm- 1h" 1M- 1) No. Concentration (M) Concentration (M) (mg.cm'^h"1) where lq.-l.85; k2=23.0 12 0.202 0.050 2.60 98.0 15 0.202 0.025 1.80 80.0 17 0.398 0.025 5.00 105. Table IVA Ammonia Oxygen pressure = 7 .9 atm, Ionic strength = 0.10(M) Run Ammonia [£\ Rate (R) C ' A3 2 No. Concentration (M) (mg.cnr^h -!) R 1 0.65 5.45 0.00500 2 0.655 38.00 0.01130 3 0.310 14.10 0.00680 4 0.086 2.10 0.00352 Table IVB Ammonia Oxygen pressure = 7 .9 atm. Ionic strength = 0 .10 (M) Run No. . Free ammonia 00 Concentration (M) Ammonium ion CNH^^ Concentration (M) Rate (R) (mg.cm^h"-'-) R 5 0.404 0.0040 26 .40 0.00615 6 11 0.0080 32.80 0.00495 7 0.303 0.0030 17.80 0.00515 8 0.0060 22 .20 0.00414 9 0.202 0.0020 10.00 0.00407 10 o 11 0.0040 12 .60 0.00322 Appendix I (cont'd.) Table IVC. Ammonia Run Free ammonia Ammonium ion k 3 (mg.cm - 2^ 1^ 1) No. Concentration (M) Concentration (M) where •kj- 4.65, k 2 = 77 5 0.404 0 .0040 2350 6 0 .0080 2410 7 0.303 0.0030 2340 8 _ 11 0.0060 2420 9 0.202 0.0020 2600 10 11 0.0040 2540 APPENDIX I I Experimental data pertaining to c a t a l y t i c reduction of n i c k e l ion by hyrirogen. Table I Test on r e p r o d u c i b i l i t y , s t i r r i n g e f f e c t , and treatment of c a t a l y s t . Run S t i r r i n g rate I n i t i a l c o n c ' n . No. RPM of n i c k e l (M) R 400 0 . 0 0 5 R 2 " R 3 7 4 0 5 5 0 R 6 4 0 0 » » R^ »• " Conditions; Concentration of a c e t i c acid Concentration of t o t a l acetate Temperature Catalyst PH 2 - • . I n i t i a l pH I n i t i a l rate Remarks (M.min-1) x 10h 1.95 1.85 1.80 1.80 1.90 0.3 gms.FeAc3 added 1.95 Catalyst treated with a c e t i c a c i d . - 0.2M - 0.06M - 140° ±1.5 - 10 gms/1400 ml. - 13.55 atm. . - 4.25 -7k- Appendix II (cont'd.) Table I I Values of i n i t i a l r a t e and i n i t i a l concentration. Run No. I n i t i a l concentration I n i t i a l rate of n i c k e l 00 (M.min-1) x 10^ R, 0.005 1.95 R 2 '» 1.85 R 5 0.003 1.20 R 7 0.0015 0.65 R 8 0.010 3.90 Conditions; Concentration of a c e t i c a c i d Concentration of t o t a l acetate Pressure of hydrogen Temperature I n i t i a l pH Catalyst S t i r r i n g rate - 0.2M - 0.06M - 13.55 atm. ," - u o ° c - 4.25 - 10 gms/1400 ml. - 400 RPM. Table I I I I n i t i a l rate and amount of c a t a l y s t Run No. Amount of cat a l y s t I n i t i a l rate i n gms. f o r 1400 ml. (M.min - 1) x 10^ of s o l u t i o n RR]_ RR2 RR3 R]_ and R 2 6 20 15 10 1.00 4.00 2.70 1.95, 1.85 Conditions: Concentration of acetic a c i d - 0.2M Concentration o f t o t a l acetate - 0.06M Temperature - 140°C Pressure o f hydrogen - 13 .55 atm. I n i t i a l concentration of n i c k e l - 0.005M I n i t i a l pH - 4.25 S t i r r i n g rate - 400 RPM - 75 - Appendix II (cont'd.) Table IV I n i t i a l rate and hydrogen pressure Run No. PH ? i n Catm») . I n i t i a l rate ( I L a i n " 1 ) x 104 P, 27.2 5.22 2.20 P 2 6.70 2.59 1.30 P 3 3.30 1.82 0.90 R]_ and R 2 13.55 3.68 1.95, 1.85 Conditions: Concentration cf a c e t i c acid Concentration of t o t a l acetate Temperature Catalyst S t i r r i n g rate I n i t i a l pH 0 .2H 0.06M 140° C' lOgms/1400 ml. 400 RPM 4.25 Table V I n i t i a l rate and i n i t i a l hydrogen i n concentration. Run No Conctn. of ac e t i c a c i d (M) I n i t i a l conc'n. of hydrogen ion (M) I n i t i a l rate (M.min-1) x 10 4 H 2 H 3 RR3 0.80 0.40 0.60 0.20 3.16 x ICrh 1.60 x 10-^ 2.40 x 1 0 - 4 5.6 x 10-5 0.65 1.60 1.10 2.70 Conditions: Catalyst Temnerature % I n i t i a l concentration of n i c k e l Total concentration of acetate 15 gms/1400 ml. 140°C 13.55 atm, 0.005M 0.06M. - 7 6 - Appendix II (cont'd.) Table VI Initial rate and temperature Run .No. Temperature °C. 1 x 103 T 00 (M.min-1) log ( i n i t i a l rate) Ti T 3 RR3 Hi 1 6 0 1 5 1 1 3 0 1 4 0 1 4 0 2.31 2.36 2.48 2.42 2.42 1.1 x 10"? 5.1 x 10"* 1.4 x 10"f 2.5 x 10"? 2.7 x 10"* -2.96 -3.29 - 3 . 8 5 -3.60 - 3 . 5 7 (rate converted to 15 gm/HC0 ml) Conditions: Concentration of acetic acid Concentration of total acetate — Catalyst Initial pH Initial concentration of nickel - Pressure of hydrogen 0.2M 0 . 0 6 M 15gms/1400 ml. 4.25 0.005M 13.55 atm. ' Table VII Final hydrogen ion concentration and corresponding equilibrium nickel ion concentration. . Run No, pH CH + 2(final) Equilibrium nickel Initial Final Concentration (M) x 103 - 77 - Appendix II (cont'd.) :) Table VIII Hydrogen pressure and corresponding equilibrium nickel concentration. Run No. Pressure of H2 1 x 10 2 Equilibrium concentration. (atm.). P H of nickel (M) X 103 P, 27.2 3.67 0.30 P 2 6.70 14 .90 0.85 P 3 3.30 30 . 3 0 1.55 RR2 13.55 7.40 0.35 PUBLICATIONS: S i s i r Coomar SIRCAR 1. Studies on the behaviour of bi-univalent salts in aqueous solution, Part VII, Copper Acetate, (with S. Aditya and B. Prasad), J. Indian Chem. S o c , 50, 655, 1955- 2. Studies on the behaviour of bi-univalent salts in aqueous solution, Part X, Zinc Perchlorate, (with B. Prasad), J. Indian Chem. Soc, 31, 483, 195^. 3. Determination of solubility product of copper oxide, (with B. Prasad),.J. Indian Chem. Soc,.1955- h. Studies on the behaviour of uni-bivalent salts in aqueous solution, .Part II, Na2S04, (NEL4) 2S0 4, (with B. Prasad), J. Indian Chem. S o c , (in press). 5> Studies on the behaviour of uni-bivalent salts in aqueous solution, .Part III, Na 2C 20 4, (with B. Prasad), J. Indian Chem. S o c , (in press). 6. Studies on the behaviour of bi-bivalent salts in aqueous solution, .Part I I , ZnSC-4, (with B. Prasad), J.. Indian Chem. S o c , (communicated). 7- Studies on the behaviour of bi-bivalent salts in aqueous solution, Part IV, MnS04, BeS04, CuS04,.(with B. Prasad), J. Indian Chem. Soc,(communicated).

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