UBC Theses and Dissertations

UBC Theses Logo

UBC Theses and Dissertations

The reduction of cupric salts in aqueous solution by molecular hydrogen Macgregor, Edwin Robert 1956

Your browser doesn't seem to have a PDF viewer, please download the PDF to view this item.

Item Metadata

Download

Media
831-UBC_1957_A7 M2 R3.pdf [ 4.41MB ]
Metadata
JSON: 831-1.0081192.json
JSON-LD: 831-1.0081192-ld.json
RDF/XML (Pretty): 831-1.0081192-rdf.xml
RDF/JSON: 831-1.0081192-rdf.json
Turtle: 831-1.0081192-turtle.txt
N-Triples: 831-1.0081192-rdf-ntriples.txt
Original Record: 831-1.0081192-source.json
Full Text
831-1.0081192-fulltext.txt
Citation
831-1.0081192.ris

Full Text

THE REDUCTION • OP CUPRIC SALTS IN AQUEOUS •. SOLUTION BY MOLECULAR HYDROGEN by EDWIN ROBERT MACGREGOR A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF APPLIED SCIENCE i n the Department of MINING AND METALLURGY We accept t h i s thesis as conforming to the required standard Members of the Department of Mining and Metallurgy THE UNIVERSITY OF' BRITISH COLUMBIA December, 195>6 i i ABSTRACT The t h e s i s d e s c r i b e s a k i n e t i c s t u d y of the r e a c t i o n s between c u p r i c s a l t s and hydrogen i n aqueous s o l u t i o n . The f i r s t p a r t i s concerned w i t h the homogeneous a c t i v a t i o n o f hydrogen i n s o l u t i o n by c u p r i c p e r c h l o r a t e - , as e v i d e n c e d by i t s c a t a l y t i c e f f e c t on the r e a c t i o n between d i c h r o m a t e and hydrogen. These s t u d i e s , w h i c h were conducted on s o l u -t i o n s of h i g h e r a c i d i t y t h a n p r e v i o u s l y i n v e s t i g a t e d , p r o v -i d e f u r t h e r s u p p o r t f o r the mechanism proposed f o r t h i s r e a c t i o n , i . e . , Cu + Ho — ^ CuH + H d. ^ : k - l C uH + + C u + + ——> 2Cu + + H + k 2 . ^ ++ 2Cu + s u b s t r a t e fas t * " ^ u + p r o d u c t s and the c o r r e s p o n d i n g r a t e e q u a t i o n : d t ( k - i / k 2 [H+] + |Cu + +] The second p a r t d e a l s w i t h the thermodynamics and k i n e t i c s of the p r e c i p i t a t i o n o f m e t a l l i c copper f r o m aqueous s o l u t i o n by Hg. The. r e d u c t i o n r e a c t i o n s were s t u d i e d i n b o t h c u p r i c p e r c h l o r a t e and c u p r i c s u l p h a t e s o l u t i o n s and the e f f e c t s o f a number o f v a r i a b l e s on the r e a c t i o n r a t e i n b o t h systems are compared. The r a t e , , w h i c h was i n i t i a l l y f a s t , d e c r e a s e d g r a d u a l l y u n t i l i i i an apparent steady-state c o n d i t i o n was approached. This behavior was common to both systems, although the r a t e s were much higher and the. r e s i d u a l copper conce n t r a t i o n lower i n the- sulphate .system. This apparent steady-state i s not p r e d i c t e d by thermodynamics and i s considered to r e s u l t from k i n e t i c factors„ The dependence of the r e a c t i o n r a t e on a number of v a r i a b l e s can be p r e d i c t e d from the mechanism C u + + + H 2 — ^ CuH + + H+ CuH+ + C u + + > 2Cu + + H + k 2 2 C u + Fist c*° + C u + + which corresponds to an o v e r a l l r e a c t i o n of C u + + + H 2 »- Cu° + 2H + Reasonably good q u a n t i t a t i v e agreement between the p r e d i c t e d ++ + and measured k i n e t i c dependence on Cu H and H 2 concen-t r a t i o n suggests that the- r a t e of r e d u c t i o n of c u p r i c s a l t s i s determined by the homogeneous a c t i v a t i o n of H 2 by C u + + or a c u p r i c complex. Although the i n i t i a l decrease i n r a t e can be p r e d i c t e d by t h i s mechanism, the apparent steady-state c o n d i t i o n approached cannot be f u l l y explained. The e q u i l i b r i u m constant f o r the r e a c t i o n 2Cu+ ^ =± C u + + + Cu° was a l s o determined •'experimentally at s e v e r a l temperatures and compared w i t h c a l c u l a t e d v a l u e s , In presenting t h i s thesis i n p a r t i a l f u l f i l m e n t of the requirements f o r an advanced degree at the University of B r i t i s h Columbia, I agree that the Library s h a l l make i t f r e e l y available f o r reference and study. I further agree that permission f o r extensive copying of t h i s thesis for scholarly purposes may be granted by the Head of my Department or by his representative. I t i s under-stood that copying or publication of th i s thesis f o r f i n a n c i a l gain s h a l l not be allowed without my written permission. Department of On*****^- '7ruZa£Cu«^ The University of B r i t i s h Columbia, Vancouver 8, Canada. Date QjLce^tJ^ 2./ , &  i v ACKNOWLEDGEMENT The author i s gr a t e f u l f o r the assistance,, advice and encouragement given by the members of the Department of Mining and Metallurgy, and es p e c i a l l y to Dr. J. Halpern, under whose d i r e c t i o n t h i s i n v e s t i g a t i o n was carried out. The author also wishes to thank the National Research Council and Defense Research Board of Canada f o r f i n a n c i a l assistance i n the form of Research Ass 1 s tantships held during the past year. V TABLE OP CONTENTS Page GENERAL INTRODUCTION 1 HYDROMETALLURGICAL HYDROGENATION REACTIONS 1 HETEROGENEOUS HYDROGENATION REACTIONS . . . . . . . 3 HOMOGENEOUS HYDROGENATION REACTIONS . . . . . . . . $ PURPOSE AND SCOPE- OP THIS INVESTIGATION 7 PART I STUDIES ON THE HOMOGENEOUS CATALYTIC ACTIVATION OP H 2 BY C u + + 1° INTRODUCTION 1° EXPERIMENTAL 12 MATERIALS . , . 12 APPARATUS 12 PROCEDURE x 3 RESULTS AND DISCUSSION IM-PART II THE REDUCTION OP CUPRIC SALTS IN AQUEOUS SOLUTION BY MOLECULAR HYDROGEN . . . . . 22 A. THERMODYNAMIC CONSIDERATIONS . 22 EXPERIMENTAL ' . 3k PROCEDURE 3k ANALYSIS 3k RESULTS AND DISCUSSION . . . . . 3^ Page B. THE PRECIPITATION OP METALLIC COPPER PROM AQUEOUS CUPRIC SALT SOLUTIONS BY HYDROGEN 38 INTRODUCTION 38 EXPERIMENTAL ^ 2 RESULTS AND DISCUSSION . . . . . . . . . . . . . . . . . 1. REDUCTION OF AQUEOUS CUPRIC PERCHLORATE SOLUTIONS BY H 2 .' kz E f f e c t of I n i t i a l C u + + Concentration k3 E f f e c t of H 2 P a r t i a l Pressure k& E f f e c t of I n i t i a l H + Concentration . . . . . . . i+6 Ef f e c t of Temperature: . i+9 E f f e c t of Me t a l l i c Copper Surface f>l 2. REDUCTION OF AQUEOUS CUPRIC SULPHATE SOLUTIONS BY H 2. . . . . . . . . . . . . . . . . 63 E f f e c t of I n i t i a l OuSO^ Concentration . . . . . 69 E f f e c t of I n i t i a l H2S0^ Concentration . . . . . 69 E f f e c t of H 2 P a r t i a l Pressure . . . . . . . . . 70 E f f e c t of Temperature- 70 E f f e c t of Adding Na 2S0^ . . . . . . . . . . . . . 70 E f f e c t of Metallic Copper Surface 7k CONCLUSIONS . . . . . . . . . . . 78 COMPARISON OF THE REDUCTION OP CUPRIC PERCHLORATE AND CUPRIC SULPHATE SOLUTIONS BY HYDROGEN . ". 78 MECHANISM OP THE REACTION 78 SUGGESTIONS FOR FURTHER WORK 79 BIBLIOGRAPHY . . 8 l v i i TABLES Page I. EFFECT OF PERCHLORIC ACID ON THE RATE OF REACTION BETWEEN H 2 AND DICHROMATE IN CUPRIC PERCHLORATE SOLUTIONS . . . . . . . . . . . . . . 17 I I . EFFECT OF CUPRIC PERCHLORATE CONCENTRATION ON THE RATE OF REACTION BETWEEN H 2 AND DICHROMATE IN AQUEOUS SOLUTION . . . . . 20 II I . SUMMARY OF THERMO CHEMICAL DATA . . . . . . . . . . . 2-£ IV. VARIATION WITH TEMPERATURE OF CALCULATED EQUILIBRIUM CONSTANTS FOR REACTIONS BETWEEN Cu + +, Cu + AND H 2 29 V. RESULTS OF AN EQUILIBRIUM EXPERIMENT FOR THE DETERMINATION OF = ( [cu +^/ £ C u+] 2 ) 36 VT. COMPARISON OF EXPERIMENTAL AND CALCULATED VALUES OF K 3 . . . . . . . . . . . . . 37 VII. RESULTS OF A TYPICAL REDUCTION EXPERIMENT CUPRIC PERCHLORATE SOLUTIONS . 1+4 VIII. INITIAL REDUCTION RATES FOR VARIOUS EXPERIMENTAL CONDITIONS . . IX. APPARENT EQUILIBRIUM CONSTANTS CALCULATED FROM EXPERIMENTAL DATA 62 X. REACTION RATES CALCULATED FOR FINAL STAGES OF EXPERIMENTS 61; XI. RESULTS OF A TYPICAL REDUCTION EXPERIMENT ON CUPRIC SULPHATE SOLUTIONS . . . . . . . . . 6£ v i i i LIST OP' FIGURES Pig. No. Page 1. Rate Curves f o r Cupric Perchlorate at Various Perchloric Acid Concentrations . . . _ ° 2. Dependence of the Rate on the Perchloric Acid Concentration . . . . . . . . . . . . . 3. Rate Curves f o r Cupric Perchlorate at Various Cupric Perchlorate Concentrations . . * . . ^ I4.. Dependence of the Rate on the Cupric Perchlor-ate Concentration 2.1 $. Vari a t i o n of A F ° with Temperature For Reduc-tion of Cupric Salts From Aqueous Solution by Hydrogen 28 6. Log K vs. V T For Four Reduction Reactions . . 30 7. Relation Between Equilibrium Concentrations of Cu**,, Cu + and H + i n Solution Under H 2 Pressure . . . . . . . . . . 32 8. T y p i c a l Rate Curve for the Reduction of Cupric Perchlorate Solutions With Hydrogen . . . . \\$ 9. Effect of I n i t i a l Cupric Perchlorate Concen-t r a t i o n on the Rate of Reaction . . . . . . k-7 10. E f f e c t of H 2 P a r t i a l Pressure on the Rate of Reaction lj-8 11. E f f e c t of I n i t i a l Perchloric Acid Concentration on the Rate- of Reaction $0 12„ E f f e c t of Temperature on the Rate of Reaction. $2 ix Pig. No. Page 13. E f f e c t of M e t a l l i c Copper Surface on the Rate of Reaction . „ . . . . . . . . . . . Dependence of the I n i t i a l Reaction Rate on Hg P a r t i a l Pressure 15. Dependence- on the I n i t i a l Reaction Rate on the I n i t i a l H + Concentration . . . . . . ^® 16. Dependence, of the I n i t i a l Reaction Rate on the I n i t i a l C u + + Concentration 6 0 • • o 17. Typical Rate Curve f o r the Reduction of Cupric Sulphate Solutions with H 2 . . . 66 18. Photomicrographs of Me t a l l i c Copper Product . . . . . . . . . . . . . . . . 68 19. E f f e c t of I n i t i a l CuSO^ Concentration on the Rate of Reaction 71 20. E f f e c t of I n i t i a l HgSO^ Concentration on the Rate of Reaction 72 21. E f f e c t of H 2 P a r t i a l Pressure on the Rate of Reaction . . . . . . . . . . . . . . 73 22. E f f e c t of Temperature on the Rate of Reaction . . . . . . . . . . . . . . . 23. E f f e c t of Sulphate on the Rate, of Reaction 76 2I4.. E f f e c t of Me t a l l i c Copper Surface on the Rate of Reaction . . . . . . . . . . . . 77 THE REDUCTION OP CUPRIC SALTS IN AQUEOUS SOLUTION BY MOLECULAR HYDROGEN GENERAL INTRODUCTION HYDROMETALLURGICAL HYDROGENATION REACTIONS In recent years there has been increasing i n t e r e s t i n the application of hydrometallurgical processes to the (1-ij.) {$) (6-9) recovery of metals such as Cu, Ni, Co , V and U from low-grade ores. This has resulted in the development of high pressure and high temperature techniques applicable to both leaching and p r e c i p i t a t i o n operations. Among the advantages of using high temperatures and pressures in hydrometallurgical reactions are: (1) favourable displacement of thermodynamic e q u i l i b r i a ( 2 ) increase of reaction rates to commercially a t t r a c t i v e values (3) use of v o l a t i l e or gaseous reagents such as hydrogen is permitted. The use of hydrogen as a reducing agent i n the p r e c i p i t a t i o n of metals and metallic oxides from aqueous (10) solution has been known f o r many years , but has only recently become commercially attractive' as a r e s u l t of improved hydrometallurgical techniques. Hydrogen can be employed to pr e c i p i t a t e metal l y i n g below or immediately above i t i n the electromotive s e r i e s , although the use of higher pressures and temperatures extends th i s range s l i g h t l y . The p r e c i p i t a t i o n of an insoluble compound of the metal i n a lower valence state -2-may also be advantageous. The p r e c i p i t a t i o n of metals or oxides from aqueous s o l u t i o n s by r e d u c t i o n w i t h hydrogen can be d i v i d e d i n t o three general classes which r e f l e c t the nature of the pro-cess by which hydrogen i s promoted i n t o a r e a c t i v e s t a t e : ( 1 ) r e a c t i o n s r e q u i r i n g an added heterogeneous hydro-genation c a t a l y s t (2) a u t o c a t a l y t i c r e a c t i o n s ( 3 ) homogeneous r e a c t i o n s The f i r s t c l a s s of r e a c t i o n s i n c l u d e s the p r e c i p i t a t i o n of ( 1 1 ) (jp) the oxides UOg and v^O^ • ; from aqueous carbonate leach s o l u t i o n s . In each case, the rate was found d i r e c t l y pro-p o r t i o n a l to the amount of m e t a l l i c n i c k e l powder present, i n d i c a t i n g that the r e d u c t i o n r e a c t i o n s are heterogeneously c a t a l y z e d . I t was also found that m e t a l l i c c o b a l t and platinum were e f f e c t i v e c a t a l y s t s f o r the r e a c t i o n s , and that no r e d u c t i o n took place i n the absence of these metals. Thus i t appears that the r e a c t i o n occurs at the c a t a l y s t surface where the reactants are chemisorbed, as i n other i n o r g a n i c heterogeneous c a t a l y t i c hydrogenation processes. The second c l a s s of hydrogenation r e a c t i o n s (13 ,14) includes the p r e c i p i t a t i o n of' n i c k e l and cobalt from aqueous a c i d i c or a l k a l i n e s o l u t i o n s pf t h e i r s a l t s by r e d u c t i o n w i t h hydrogen. The r e a c t i o n i s i n i t i a t e d by add-ing some f i n e metal powder or by. reducing some of the s a l t w i t h a reducing agent such as a f e r r o u s s a l t or hydrazine. Subsequent hydrogenation proceeds i n an a u t o c a t a l y t i c - 3 -manner i n view of the c a t a l y t i c properties of the metals formed. The third class of reactions includes the p r e c i p i -t i n (13) ( 1 6 ) tation of metals such as Cu , Hg and Ag from aqueous solutions by hydrogen reduction. These reactions do not require any added c a t a l y t i c surface to activate the hydrogen or to promote nucleation of the met a l l i c p r e c i p i -tate, but proceed homogeneously with reasonable rates at moderate, temperatures and hydrogen pressures. Extensive studies of the homogeneous c a t a l y t i c activation of hydrogen , (17,18) , (19) , ... (20) by cupric , mercuric and s i l v e r ions m aqueous solution have recently been undertaken i n .these laboratories. The investigation of the homogeneous reduc-tion of cupric s a l t s by hydrogen, which i s described i n this thesis, represents an extension of the above studies. HETEROGENEOUS HYDROGENATION REACTIONS In the absence of catalysts, most hydrogenation reactions proceed only very slowly under moderate conditions, This i s i l l u s t r a t e d by the above discussion concerning react-ions i n solutions and i s generally true also of uncatalyzed gas-phase hydrogenation reactions. This apparent inertness of the H£ molecule can be attributed to i t s high d i s s o c i a -tion energy (103 K c ai/mole) and to the closed-shell electronic configuration of i t s ground state. Thus i s repelled strongly on approaching most other species, and i t s homogeneous reactions are characterized by large a c t i v a t i o n energies. -k-Catalysts are defined as substances which lower the free energy of a c t i v a t i o n , usually by allowing the reaction to proceed by an alternative mechanism. The catalysts most commonly used in hydrogenation reactions are metals such as Ni, Co, P d a n d Pt and certain oxides such as Cu£0, ZnO and CrgOo. For example, the a c t i v a t i o n energy f o r the hydro-genation of acetylene i n the gas phase is l\Z Kca]_/mole but (22) only 6 to 7 Kca-|_/mole on the surface of metallic n i c k e l Various attempts have been made to correlate the c a t a l y t i c a c t i v i t y of these solids with such physical properties as semiconductivity, c r y s t a l structure and d-band character. The function of the heterogeneous catalyst i s gen-e r a l l y associated with the promotion of hydrogen into a reactive state, that i s , the hydrogen molecule forms a reactive complex with the catalyst surface. This activation of molecular hydrogen probably involves i t s adsorption and accompanying d i s s o c i a t i o n on the catalyst, and may be accompanied by the formation of covalent bonds between the two hydrogen atoms and two adjacent surface atoms. Hence a co r r e l a t i o n between the c a t a l y t i c a c t i v i t y and the l a t t i c e spacing and crystallographic orientation of the surface i s to be expected. This interpretation has received experi-( 2 2 23) mental support f o r several hydrogenation reactions. •'' Another in t e r p r e t a t i o n which has received much support relates the c a t a l y t i c a c t i v i t y of s o l i d catalysts ( 2 2 ) to t h e i r d-band character. Thus Beeck has shown that the rate of hydrogenation of ethylene over evaporated - 5 -porous metal films increased as the d-character of the metal (2k) bonds increases. Also, Couper and Eley found that the activation energy pf the parahydrogen conversion on p a l -ladium-gold alloys increases as the gold content increases. This corresponds to the f i l l i n g of the d-band with electrons, thus suggesting that vacant d-orbitals are ess e n t i a l f o r low-temperature c a t a l y s i s . This view is also supported by studies of the hydrogenation of styrene over Ni-Fe catalysts These re s u l t s imply the entry of electrons from the adsorbed reactant into u n f i l l e d e l e c t r o n i c bands of the catalyst. HOMOGENEOUS HYDROGENATION REACTIONS Homogeneous hydrogenation c a t a l y s i s i s pf consid-erable significance in that i t makes possible the study of the activation of molecular hydrogen i n r e l a t i v e l y simple systems, thus leading to more detailed i n t e r p r e t a t i o n of results than i s possible with heterogeneous c a t a l y s i s . In 1938, Ca l v i n v ; reported the f i r s t clear-cut example of a homogeneously catalyzed hydrogenation reaction i n solution. He showed that the reduction of cupric acetate by hydrogen i n quinoline solution was auto-c a t a l y t i c , the e f f e c t i v e catalyst being the cuprous salt produced by the reaction. The rate determining step, i n which hydrogen i s activated, involves the int e r a c t i o n of a hydrogen molecule with two cuprous acetate molecules. The f i r s t example of a homogeneous hydrogenati.cn reaction i n aqueous solution was observed by Dakers and (27) Halpern in 1953. These authors found that cupric -6-acetate reacts homogeneously with molecular hydrogen accord-ing to the equation: 2 Cu (0Ac) 2 + H 2 + H 20 > Cu 20 + IjflOAc (l) The reaction rate was found to be f i r s t order with respect to both H 2 and cupric acetate, and independent of s o l i d surfaces i n contact with the solution. The rate-determining step has been postulated as the reaction between a cupric acetate molecule and a hydrogen molecule to form an intermediate complex, as i l l u s t r a t e d by the following mechanism: k Cu Ac 9 + Ho — > Cu Ac„.Hp (2) c- <- slow 2 ^ Cu Ac 2-H 2 + Cu Ac 2 + H 20 f&s% C u- 20 + l^HAc^ • Peters, has further investigated t h i s system and has found that cupric acetate i s also capable of functioning as a homogeneous catalyst for reactions with H 2 of other sub-= ++++ -strates such as Cr 20y , Ce and 10-^  , which are thermo_ dynamically more reducible than CuAc£ i t s e l f . He also examined the k i n e t i c s of the homogeneous ac t i v a t i o n of hydrogen by a number of other cupric s a l t s . Similar studies have been conducted here by (19) Korinek v ' on the kin e t i c s of the homogeneous act i v a t i o n of hydrogen by mercurous and mercuric s a l t s i n aqueous solution. He observed that the k i n e t i c s are analogous to that observed f o r cupric s a l t s , that Is, f i r s t order in H 2 and f i r s t order i n H g + + or H g 2 + + . Concurrently, Webster has. shown that Ag + functions as a homogeneous catalyst f o r the reduction of CrgO^ - i n aqueous solution, the k i n e t i c s being f i r s t order In H2 and second order i n Ag + at low temperatures. However, at higher temperatures (above 8£°C) the k i n e t i c s appear to be analogous to those f o r C u + + and H g + + ( 2 9 ) . PURPOSE AND SCOPE OP THIS INVESTIGATION Peters! investigations ^0) Q n ^ Q a c t i v a t i o n of Hg by cupric salts were largely confined to solutions of low H + concentrations. He reported that the rate of Cu +- catalyzed hydrogenations of dichromate decreased s l i g h t l y with increasing H + concentrations but t h i s e f f e c t was not f u l l y elucidated. It seemed desirable to examine i n greater d e t a i l the k i n e t i c s of this reaction at high acid concentrations, f o r the following reasons: (1) the information thus gained might be useful i n elucidating the mechanism of the reaction and the nature of the activated intermediate which is formed in the rate determining step, (2) i t was proposed subsequently to study the d i r e c t reduction of cupric s a l t s to m e t a l l i c copper by hydrogen i n aqueous solutions. This reaction is accompanied by the formation of hydrogen ions. Hence a knowledge of the effects of varying the Cu and H concentrations on the rate of a c t i -vation of H 2 seemed desirable. Part I of t h i s thesis discusses the results of an i n v e s t i -gation along these l i n e s . - 6 V The investigation, which is described in Part II pf this thesis, was undertaken in an attempt to e s t a b l i s h ++ the k i n e t i c s and mechanism of the reduction of Cu to metallic cppper (Cu°) i n aqueeus splutipns by Hg. This reaction has been examined s u p e r f i c i a l l y by I p a t i e f f and cc-wprkers, and more'recently by Shaufelberger. Calcul-ations of the e q u i l i b r i a associated with the p r e c i p i t a t i o n ++ reactions indicated that v i r t u a l l y a l l Cu should reduce to m e t a l l i c copper at moderate temperatures and pressures, even at very high acid concentrations. Hence an i n v e s t i -gation which included the following considerations seemed j u s t i f i e d : (1) Previous studies were conducted on sulphate solutions. Since there i s l i t t l e information available concerning the d i s s o c i a t i o n constants f o r HgSO^ and CuSO^ solutions at high temper-atures, i t was considered desirable to study the reduction reactions i n a simpler system, i . e . , ++ perchlorate solutions where Cu i s believed to be e s s e n t i a l l y uncomplexed. The sulphate system was also examined for comparative purposes and because of i t s great p r a c t i c a l Interest. (2) The available thermodynamic data f o r the reaction were based on the extrapolation of room-temperature data and were thus subject to consid-erable uncertainty. Therefore, i t was of i n t -erest to check such data experimentally where possible. -9 -I n i t i a l experiments indicated that p r e c i p i t a t i o n o of copper occurred at reasonable rates above 130 C with moderate Hg pressures. Hence a series of experiments on cupric perchlorate solutions were undertaken to determine the effect of the following variables on the rate of copper reduction. The range over which each variable was exam-ined i s indicatedo (a) I n i t i a l C u + + concentration (0.0$ to 0.3M). (b) I n i t i a l HCIO^ concentration (0.01 to 0.2M) (c) Hg p a r t i a l pressure (10 to 2$ atm) (d) Temperature (l$0 to 17$°C) (e) E f f e c t of copper powder A similar series of experiments was conducted i n sulphate solutions and the two series were compared. Experiments were also conducted In both systems in an attempt to determine experimental equilibrium constants fo r the reaction. C u + + + Cu° ^ = ± 2 C u + , W These values were compared with constants calculated from available thermochemical data. The course of the reduction reactions was followed by determining at various time i n t e r v a l s the amount of Cu (II) remaining i n the solution a f t e r cooling, and estimating the solution composition under reaction conditions from a knowledge of the above equilibrium. PART I STUDIES ON THE HOMOGENEOUS CATALYTIC ACTIVATION OP H 2 BY C u + + INTRODUCTION The k i n e t i c s of the homogeneous c a t a l y t i c a c t i -vation of molecular hydrogen by cupric salts has been (30) extensively investigated by Peters . Of p a r t i c u l a r interest are his studies of the reduction of dichromate and other substrates i n aqueous cupric perchlorate solutions, i n which the C u + + ion i s believed to be. e s s e n t i a l l y uncpmplexed. The results of his research indicate that at constant temperature and hydrogen pressure,, the reduction of dichromate i s zero-order k i n e t i c a l l y , and proceeds according to the stoichiometric equation C r 2 0 ? = + 3 H 2 + 8 H + ^ — » 2 C r + + + + 7H20 (£) ++ = The concentration of Cu does not change while Cr 20y i s being reduced. The rate was found to be independent of the nature or concentration of the substrate. This sug-gests that the substrate species do not enter into the ra t e - c o n t r o l l i n g step. At low a c i d i t y , the k i n e t i c s were found to be of the form - d[H 2]/dt = k [CU++][H2] (6) suggesting that an H 2 molecule reacts homogeneously with a C u + + ion In the rate-determining step, leading to the formation of an "activated intermediate" which reduces -11-dichromate i n a subsequent fas t reaction. Peters found that the rate of this reaction decreased s l i g h t l y as the concentration of HCIO^ increased, the rate at 1.0M HCIO^ being about 30% of i t s value at low HC10.[j_ concentrations. To explain t h i s behavior i t i s proposed that the rate^determining step, i n which' H 2 i s converted to a reactive intermediate complex, releases an H + ioh and that the decrease i n rate i s due to the f a c t that the reversal of t h i s step> which regenerates H 2, competes with the sub-sequent reaction of the complex. The observed zero->order dependence of the rate on C r 2 0 7 ~ concentration, even at high H + concentrations, suggests that the complex does not react d i r e c t l y with the Gr 20y . The most l i k e l y a l ternative i s that i t reacts with another C u + + ion. These steps are incorporated i n the following mechanism: C u + + + H 2 . ^ > CuH + + H + (7) k - l CuH + + C u + + —7--* 2Cu + + H + (8) k2 6Cu + + + C r 2 0 7 = + li^ H + faat* 2 C r + + + + 6Cu + + + 7H 20 (9) Application of the steady-state treatment to the above sequence y i e l d s the following rate expression: dt - - r .n- • F — T - T v HQ) ( k _ l / k 2 ) [ H + ] + [ C u + + ] - 1 2 -This reduces to equation ( 6 ) at low H + concen-trations. Equation ( 1 0 ) implies a l i n e a r dependence of ( r a t e ) - 1 vs £H +] at constant [cu + +J and [H 2 ] • Peters' experimental re s u l t s are i n accord with t h i s . The above rate law predicts a s h i f t of the k i n e t i c dependence on the Cu** concentration from f i r s t toward + second order with increasing H concentration. The experi-ments described here were made with a view to t e s t i n g this prediction and to providing further evidence concerning the detailed mechanism of the reaction. EXPERIMENTAL MATERIALS Reagent grade G . P . Smith cupric perchlorate and Baker and Adamson perchloric acid were used in the prepar-ation of the solutions. Hydrogen and nitrogen gas were, supplied i n cylinders by Canadian Liquid A i r Company and were used without further p u r i f i c a t i o n . D i s t i l l e d water was used i n the preparation of a l l solutions. APPARATUS The reduction experiments were carried out i n a high-pressure stain l e s s s t e e l autoclave manufactured by Autoclave Engineers, Inc. The autoclave, which had a capacity of one gallon, was equipped with a motor-driven propeller^type s t i r r e r mounted on a v e r t i c a l shaft passing - 1 3 -through a pressure gland. A l l parts i n contact with the solution were made of titanium. The thermocouple w e l l , s t i r r e r and sampling tube were connected through the auto-clave l i d and extended into the solution,- which was contained by a titanium l i n e r . The external sampling l i n e passed through a cooling jacket so that f l a s h i n g of solution during sampling was eliminated. The hydrogen pressure above the solution was controlled by a standard gas regulator. The autoclave was heated externally by a ring-type gas burner and the temperature was controlled to within - 0 .3°C by a Leeds and Northrup Micromax c o n t r o l l i n g recorder. PROCEDURE Solutions (2„2 l i t e r s was used i n each experi-ment) were made up by dissolv i n g weighed amounts of reagents i n d i s t i l l e d water. The charge was placed i n the auto-clave j which was sealed, flushed with nitrogen gas and then heated tq temperature. When the desired temperature was reached, a number of samples were taken while the solution remained under steam or a small p a r t i a l pressure of n i t r o -gen. Hydrogen was then introduced and maintained at the desired p a r t i a l pressure throughout the experiment. Samples were taken p e r i o d i c a l l y during the experiment. The course of the reduction reaction was followed by analyzing the cooled solution samples f o r dichromate with a Beckman DU Spectrophotometer using the 350 m^ u absorption peak. The copper concentration was determined e l e c t r o l y t i c a l l y and pH measurements were made with a -111.-Beckman Model H-2 a.c. pH meter. After the completion of each experiment, a l l parts of the autoclave i n contact with the solution were soaked in d i l u t e n i t r i c acid and rinsed with d i s t i l l e d water. RESULTS AND. DISCUSSION A series of experiments were conducted i n order to e s t a b l i s h the dependence of the rate of reduction of dichromate on the H + concentration at 0.3M cupric perchlor-ate. T ypical rate plots are reproduced i n Pig. 1. Reaction rates were estimated from the slopes of the zero-order rate p l o t s , a f t e r making small corrections f o r contributions due to side reactions not involving hydrogen (presumably due to corrosion e f f e c t s ) . Table I summarizes the results of these experiments. Results of duplicate experiments are included to indicate the r e p r o d u c i b i l i t y . Equation (10) can be re-written as 1 1 k - l . g + (Rate)" = k l [Cu + +] [H 2] + k T 1 ^ * [ C u " ] * [H 2] (11) At constant [Cu + +]and [H2J , as shown i n Pig. 2, the experimental results give a reasonably good l i n e a r p l o t of ( R a t e ) - 1 vs [H"*] , as required by equation (11). At 110°C, the values of k x 'and ( k - l / k 2 ) , obtained from the The concentration of H 2 i n solution, |jH2] , was e s t i -mated from the measured p a r t i a l pressure, using the data of Pray and a s s o c i a t e s ( 3 1 ) . -15-i - 5 ' - i intercept and slope of t h i s plot,, are 9 . 4 x 1 0 ^ 1 . mole -1-sec 1 and 0 .25 , respectively. This agrees w e l l with values of k]_ = 9.$ x 1 0 " ^ 1 . mole" 1 sec" 1 and ( k-l/k2) = 0.26 . - ^  + + obtained by Peters from a s i m i l a r p l o t f o r O.lM.^nd Cu A related series of experiments was conducted to ++ e s t a b l i s h the dependence of the rate on the Cu concen-t r a t i o n at 1 . 0 M HCIO^. T y p i c a l rate plots are shown i n Pig. 3 and the res u l t s of these experiments are l i s t e d i n Table I I . Equation ( 1 1 ) also predicts a l i n e a r r e l a t i o n between [cu + +]/rate vs [pu + +J at constant temperature , should increase k - l k i k 2 N with [EPJ while the intercept remains constant. The r e s u l t s from Table II showing the dependence of the rate on [cu J at 1 . 0 M are plotted i n P i g . k w values of k]_ and (^ _l / k 2 ) estimated from the intercept and slope of the plot for 1 . 0 M HCIO^ are 9-4 x 1 0 — 1 . mole 1 s e c - 1 and 0 .25 , respectively. These values agree The well with those determined f o r the dependence of the rate prov: ,(32) on H* concentration, thus oviding convincing support f o r the proposed mechanism* " The plot f o r 0 . 0 1 M HCIO^ i s based on data of Peters and i s given for comparison. - 1 6 -o L 1 J t 0 20 k O 60 80 Time - Minutes Figure 1 T Reduction of Dichromate by H_ in the Presence of Cupric Perchlorate at d i f f e r e n t Perchloric Acid Concentrations. 0 . 3 M/L C u ( C 1 0 k ) 2 110°C 20 Atm* H 2 -17-TABLE I EFFECT OF PERCHLORIC ACID ON THE RATE OF REACTION BETWEEN HYDROGEN AND DICHROMATE IN SOLUTIONS CONTAINING CUPRIC PERCHLORATE 0.3 M/L Cu(C10^_)2 : 20 atm. Hg : 110°C Run No. f"M/L) d [cr207 =]/dt (M/L/min)xJ.o6 Rate = -d[ H2j/dt' :" (M/L/sec)xl07 Rate" 1(L.sec/F xlO-7 22 1.0 £.06 2.^3 0.396 30 1.0 £.09 2.$$ ( 0.393 31 1.0 $.0 2. £0 0.1+0 32 0.7$ $.$2 2.76 0.3^7 33 o.$o 2.98 0.336 3k 0.1$ 7.$2 3.76 0.266 3 . 0.01 9.20 L..60. 0.218 Calculated from the stoichiometric r e l a t i o n : - d CH£| /d t - 3 d I c^Oy - J/dt -18-O Rate vs H • (Rate)"'vs H 0 0.2 O.k 0.6 0.8 1.0 HCIO^ - Mole L i t e r s 1 Figure 2. Dependence of the Rate of Reaction of Hg on the H+ Concentration. 0.3 M/L Cu(C10^) 2 : 110°C : 20 Atm. Hg - 1 9 -220 Time - Minutes Figure 3. Reduction of Dichromate i n the Presence of Various Concentrations of Cupric Perchlorate. 1.0 M HCIO^ : 20 atm. H"2 : 110°C -20-TABLE II EFFECT OF CUPRIC PERCHLORATE CONCENTRATION ON THE REDUCTION OF DICHROMATE IN AQUEOUS SOLUTIONS BY HYDROGEN 1.0: M/L HCIO^ : 20 atm. H 2 : 110°C Run No. [Cu++j (M/L) _d [ c r 2 ° 7 ] / d t (M/L/min) x l O 6 Rate= - d [ H 2 ] / d t (M/L/sec) xl07 [ c u + + ] / r a t e ( L-sec )x l0-? fer3 0.10 • (i . 0.92 0.46 2.18 10 11 0.20 2.95 1.48 1.3.8 22 0.30 >.o5 2.53 1.19 3.3 19 0.I).0 7.20 3.60 1.04 2.5 24. 0.^0 8.60 • 4.30 1.16 2. 0 Average of three experiments. -21-3-0 Q OOIM/L HCIO+ O '>0 M/L HClO+ 0 ,. $ 10 [ C u + + ] _ 1 (M/L)" 1 Figure i | . Dependence of the Rate of Reaction of H 2 on the Concentration of C u + + 110°C : 20 Atm. H 2 PART II THE REDUCTION OP CUPRIC SALTS' IN AQUEOUS SOLUTION BY MOLECULAR HYDROGEN A, THERMODYNAMIC CONSIDERATIONS The degree to which the reaction aA + bB — c C + dD (12) proceeds under a given set of conditions can be related to the decrease i n free energy, - A P , accompanying the reaction, i . e . , A P = A P° + RT In ( ac)° ( a D ) d f (13) ( a A ) a ( a B ) b -where R = gas constant (I . 987 cal./deg/mole) T = absolute temperature and a denotes the a c t i v i t i e s (or fugacity i n the case of a gas)' of the species under reaction conditions. In the case where the system approaches i d e a l i t y , the a c t i v i t i e s and fu g a c i t i e s may be replaced by the concentration, and p a r t i a l pressures, respectively. The standard free energy change, A P ° , i s a constant f o r the reaction and i s related to the standard free energies of formation of each substance, i n Its standard state'"' by A F ° = c F c ° + d P D ° - a P A ° - b P B ° ( U J *':" The. standard state for solids usually refers to the' stable modification of the pure substance at 2£°C and 1 atm. pressure. For solutions, the standard state of each dissolved substance i s taken as unit a c t i v i t y expressed i n mo l e s / l i t e r . For gases, the standard state i s usually taken as unit fugacity expressed i n atmospheres. The requirement f o r a system i n equilibrium under a speci-f i e d set of conditions i s such that the free energy must be a minimum, that i s , AF = 0 (15) Hence, the standard free energy change f o r a reaction can be related to the equilibrium constant, K, by A P ° = - RT In K (16) c d where K = a£ x a D a x a b (17) at equilibrium. The following reactions may be written to describe the reduction of aqueous solutions of copper s a l t s with . hydrogen: C u + + ( a q ) + i H 2 ( g ) ^ Cu(" a q) + H + a q ) (18) Cu + +(aq) + H 2 ( g ) — Cu° ( s ) + 2 H + a q ) (19) 2C u + ( a q ) ^ C u + | a q ) + Cu? s ) (20) ° U + (aq) + * *2 ( S ) ^ C U ° ( S ) + (22r)_ Any two of these reactions i s s u f f i c i e n t to describe the equilibrium of the system. Neglecting a c t i v i t y c o e f f i c -ients , the following equilibrium constants may therefore L i t t l e i s known about a c t i v i t y c o e f f i c i e n t s i n these systems, p a r t i c u l a r l y i n the range of conditions of inte r e s t here ( i . e . high temperatures and f a i r l y con-centrated solutions). E q u i l i b r i a are therefore expressed in terms of concentrations and p a r t i a l pressures of reactantso This may introduce considerable error but i s not l i k e l y to affe c t the. q u a l i t a t i v e pattern of the re s u l t s . -2k-be written: K. K. 1 (22) 2 (23) [H+] ^ = M ( p H 2 ) t (2^) Where denotes the molar concentrations and P the p a r t i a l pressure in atmospheres. The free energies of formation, heats of form-ation and p a r t i a l molal entropies of copper and aqueous solutions of i t s ions i n t h e i r standard states are well known and are l i s t e d i n Table I I I . By using these data, i t i s possible to calculate K2,. K^, and f o r the reduction of C u + + ions i n solution by the H 2 at 25>°C. However, t h i s reaction attains measurable rates only at temperatures above 100°G. Since actual thermodynamic data are not available f o r elevated temperatures, extra-polation of room-temperature data becomes necessary. This i s usually done by c a l c u l a t i n g the standard free energy change (AF°) f o r a reaction at 2*~>QG and then extrapolating i t to the elevated temperature. Hence, the equilibrium constant can be evaluated at the reaction temperature and information regarding the effects of variables on the reduction e q u i l i b r i a can be obtained. - 2 ^ TABLE III SUMMARY OP THERMOCHEMICAL DATA"" (2£°C) Species AH f ( Kcal/mole) APg ( Kcal/mole) S° (cal/mole/°K) „ ++ ^ ( a q ) l£.lf l£.£ -23.6 C u ( a q ) 17> 1 12.0 9.1f C u ( s ) 0.0 0.0 8.0 H 2(g) 0.0 0.0 31.2 (aq) 0.0 0.0 0.0 A l l data are from: "Selected Values of Chemical Thermodynamic Properties," C i r c u l a r £00, National Bureau of Standards, Washington, D.C. (19£2) except those f o r Cu| a^), which are,taken from D.D.Wagman, J . Am. Chem. S o c , £1+63 (19£l). -26'r-The c a l c u l a t i o n of K^, Kg, K^, and K^ at elevated temperatures i s d i f f i c u l t because of uncertainties i n the eff e c t of temperature on the entropies of ions and undis-sociated molecules i n aqueous solution. There i s also the additional uncertainty i n determining the concentrations of each ionic species i n solution because the v a r i a t i o n of di s s o c i a t i o n constants f o r s a l t s and acids with temperature is not known. Shaufelberger"^" has overcome the former d i f f i c u l t y i n the case of equation (19) by making the approximation that the temperature changes i n entropies are assumed to cancel out,, since there are ions on both sides of the equation. Hence the entropy change f o r the reaction at temperature T, AS^, i s assumed equal to the reaction ,o entropy change at 25 C, AS^gO^- This approximation i s admittedly rather crude but w i l l be used here also, because, at least f o r reactions (18), (19)> and (21), no other method of evaluating A P ° i s available, Thus from the f a m i l i a r thermodynamic equation d C A P j . _ _ A S (26) c|T i t i s possible to derive the following r e l a t i o n between temperature and the standard free energy change of a reaction: A P T ° = A P 2 9 8 ° K ' A S 2 9 6 ° K <T " <2?> The data i n Table I I I was used i n the above equation to calculate the values of A P ° f o r reactions (18) to (21). -27-The v a r i a t i o n of A F ° with T f o r these reactions i s shown in Pig. 5\ Prom the calculated values of A F T ° the equilibrium constants K-j., K2,. K^, and were determined f o r selected temperatures between 25>°C and 200°C using equation (16). The results are tabulated i n Table IV and plotted i n Pig. 6. Upon examination of Figures 5 and 6, the following thermodynamic information i s evident: 1. The equations f o r any two of the four reactions shown are s u f f i c i e n t to describe the reduction pf copper i n solution by hydrogen. For example, equations (1.8) and (21) can be added to give equation (19), and the equilibrium constants and can be used to determine Kg and by K l E k = K 2 VK1 = K 3 2. Copper i s r e a d i l y formed from solutions of i t s salts by reduction with hydrogen at moderate pressures over the entire range of temperatures under cpnsideratipn. 3. The general e f f e c t pf increasing temperature i s tc cause an increase i n the tendency f o r reduction of copper solutions by H 2 to metallic copper to occur." .k. The apparent s t a b i l i t y of the Cu ion increases markedly with increasing temperature (reflected i n the (ificreasing value of K^) . r-28" T (°K) Figure £. Variation of A F ° with Temperature for the Various Reactions Involved in the Reduction of Aqueous Copper Salt Solutions by Hg. - 2 9 -TABLE IV ,: VARIATION WITH TEMPERATURE OP CALCULATED EQUILIBRIUM CONSTANTS FOR REACTIONS BETWEEN C u + + , Cu +, and H p T(°C) K l V K 3 : ^ 3 , 8 8 x l 0 2 . 2 . k 6 x l 0 — 1.63XJL06 6 . 3 k x l 0 8 $0 5 . k 8 x l 0 2 3 .28x10 1 0 l . k l x l O ^ 6 . 8 0 x l 0 7 1$ 5 . 8 0 x l 0 2 5 . 8 3 x l 0 9 1 . 7 3 x 1 0 ^ 1.OxlO 7 TOO 6 . 8 l x l 0 2 1 . 3 1 x l 0 9 2 . 8 2 x l 0 3 1.92x10 6 120 7 . 6 3 x l 0 2 k . 5 3 x l 0 6 7 - 7 3 x l 0 2 5.9kxlO^ IkO 8 . k 6 x l 0 2 1 . 7 k x l 0 8 2 . k k x l 0 2 2 . 0 6 x 1 0 ^ 1 5 0 8 . 8 7 x l 0 2 1 . 1 2 x l 0 8 l k 2 1,26x10^ 1 6 0 9 . 2 8 x l 0 2 7 . 3 0 x i o 7 8 k . 8 7 . 8 8 x 1 0 ^ 1 7 0 9 . 7 0 x l 0 2 k . ' 8 7 x l 0 7 5 1 . 9 5.03X10M-H$ 9 . 9 0 x l 0 2 k.OxlO 7 k 0 . 9 k . 0 5 x l O k I8O 1 0.IxlO 2 3-31X107 3 2 . k 3.28xlO k 200 1 0 . 9 x l 0 2 l , 6 0 x l 0 7 13.5 I . k 7 x l 0 k -30-12 0 J 1 ! _J I : I 2.0 2.3 2.6 2.9 3.2 3.£ 1000/T°K Figure 6. Plot of Log K vs l/T f o r the Four Reactions i n the Reduction of C u + + to Cu° by Hydrogen from Aqueous Solution -31-I t i s also of interest to calculate the equilibrium + ++ + concentrations of H , Cu and Cu under a specified hydro-gen pressure and reaction temperature. This may be done by using the equilibrium constants l i s t e d i n Table IV. Figure 7 i s an i l l u s t r a t i v e example showing the equilibrium concen-++ + , o 2 0 trations of Cu , Cu and H at 160 C under vatm. Hg pressure. The p l o t i s of the same form as that used by Shaufelberger 13 • and Roy to i l l u s t r a t e reduction e q u i l i b r i a of Ni, Co and Cu solutions. From Figure 7 i t i s evident that the reduction of copper from solution with H 2 w i l l go v i r t u a l l y to completion at a c i d i t i e s up to a pH of -3, or 1000 M. H +. While t h i s cannot be demonstrated experimentally, i t has been reported that copper can be completely p r e c i p i t a t e d from solutions containing up to 20$ H2S0[j.. The determination of experimental equilibrium ++ + + constants involving e q u i l i b r i a between Cu , Cu , H ions and dissolved Hg i s extremely d i f f i c u l t and was consid-ered beyond the scope of th i s investigation. However, the experimental determination of the equilibrium constant i s f e a s i b l e because reaction (26) i s independent of both pH and H 2 pressure. This equilibrium between C u + + and Cu + ions i n solution i n contact with m e t a l l i c copper has 33 received some attention i n the past. I p a t i e f f f i r s t attempted to calculate t h i s equilibrium constant from o reduction experiments on CuSO^ solutions by H 2 at 110 C and 100 atmospheres pressure, but h i s results are not ^32-Figure 7* Relation'Between Equilibrium Con-centrations, of CU++, CU+ and H + i n Aqueous solution 160°C : 20 atm. H~2 -33-3k r e l i a b l e . Heinerth has examined t h i s equilibrium i n both sulphate and perchlorate solutions i n the temperature range o 20 to 60 C. He has determined very r e l i a b l e data i n th i s temperature range and has shown that no change i n the quotient (Cu ) s/(Cu ) appears, which he has interpreted as indicating the a c t i v i t y c o e f f i c i e n t s f o r the cuprous and cupric ions tend to cancel. Kawassiadas has evidently studied the same equilibrium i n sulphate solutions over the temperature range 50 to 200 C but h i s data has been rejected 36 ' as unreliable . In view of the lack pf data at temper-atures above 60°C, a series of experiments were conducted in the temperature range l50 to 175°C to determine values of f o r reaction (20). These were based on the follow-ing p r i n c i p l e . • On heating cupric salt solutions i n contact with metallic copper, the reverse of reaction (20) occurs, i . e . , ++ o - + Cu + Cu * 2Cu u n t i l equilibrium i s attained. On sampling and cooling to room temperature, the reaction 2Cu + — > C u + + + Cu° proceeds u n t i l equilibrium i s again attained. Since the value of i s very large at 25°C ( > 1x10^), i t can be assumed that a l l copper remaining i n solution at room temperature i s i n the form of cupric ions. Thus the amount of metallic copper p r e c i p i t a t i n g pn cooling Is equal to one h a l f the amount of cuprous ion present at the higher temperature-, while the ampunt pf cupric ion at the higher temperature i s equal to the room-temperature value minus the amount of copper p r e c i p i t a t e . . EXPERIMENTAL PROCEDURE The equilibrium experiments were conducted i n the same apparatus as described previously ( i . e . , the high-pressure autoclave) except that a sintered glass f i l t e r was attached to the immersed end of the sampling tube. The; condenser was also removed so that solution passing through the sampling l i n e cooled as l i t t l e as possible. Two-liter charges were made up containing 0.1M Cu , 0.1N acid (HCIO^ or H2S0^) and approximately 30 gms;. of pure copper shot (Fisher S c i e n t i f i c Company) . The charge was sealed In the autoclave,, flushed with n i t r o -gen and heated to temperature. Samples of solution were collected i n b o t t l e s , which were immediately capped and then allowed to cool to room temperature. ANALYSIS: After cooling the solution volumes were measured and the precipitated copper was separated from the l i q u i d . by f i l t r a t i o n . The C u + + concentration of the f i l t r a t e was determined e l e c t r o l y t i c a l l y . The copper residue was redissolved and i t s weight determined i n a s i m i l a r manner. RESULTS AND DISCUSSION The results of an equilibrium experiment are shown in Table V, which l i s t s the determined values of [Cu , ( c u + J and K^. The deviations from the mean value are considered reasonable i n view of uncertainties i n the experimental techniques employed. Table VI l i s t s the mean values of the experimentally determined equilibrium constants f o r several temperatures, which are compared with values of Kg calculated from thermodynamic data. The equilibrium constants determined f o r cupric perchlorate solutions i n contact with metallic, copper agree f a i r l y well with the calculated values. Unfort-unately, experiments at temperatures above l60°C were not permitted because of the i n s t a b i l i t y of the C10|^ ~ ion, which + reduces to form C l . This combines with Cu to form i n s o l -uble CuCl. Equilibrium experiments on cupric sulphate solu-tions in contact with me t a l l i c copper gave mean values of Kg much greater than the calculated values. However, Kg was determined on the basis that no complexing of Cu and + Cu occurs. This assumption i s obvipusly u n j u s t i f i e d , but the values of complexity constants i n the temperature range under consideration are. unknown. Therefore, these constants cannot be used for any quantitative calculations, but do indicate that the s t a b i l i t y of the Cu + ion Increases with increasing temperature, as calculated from thermodynamic data. The f a c t that the experimental values of Kg are TABLE V RESULTS OP AN EQUILIBRIUM EXPERIMENT FOR THE DETERMINATION OF = ( |cu +*]/|Cu +] 2) ' l60°C : O.IOM Cu(C10^) 2 : O.IOM HCIO^ No. Time (min) Volume Solution ; (ml) Cu° ppt. (gm) Cu° (gm/1) KPP] .(gm/1) : locale j (M/L) P u c a l c ] (M/L) K 3 ' % Dev. From Mean 1 43 45=. 8 .0^1 1.13 6.81 .089 .035 71.4 - 14.7 2 7k 46-5 .048 1.03 6.84 .091 .032 87.6 + 1.1 3 : 118 40.7 .043 1.02 7-14 .096 .032 9 2 . £ + 10.. 5 4 176 48.2 .053 ; 1.10 7-32 .098 .035 81.3 - 2.9 236 42.3 .046 1.09 7.57 ; .10 .034 86.1 + 2.9 Mean Value K 3 = 83-7 -37-TABLE VI COMPARISON OF EXPERIMENTALLY DETERMINED VALUES OF K WITH CALCULATED VALUES " •1 , • . . . O.IM CB(C10^) 2 t O.IM HCIO^ T (°C) Av. (exp.) Calc. . l$o 113 142 160 83 .7 81+.8 O.IM CuSO^ • 0.1N H 2S0^ i$o 21$ 142 160 2$$ 81+. 8 170 20$ $1.9 175 20$ 1+0.9 Assuming no complexing of Cu and Cu . -38-consistently greater than the calculated ones, suggests that ++ + — Cu i s complexed to a greater extent than Cu , by SO^ . This i s not unexpected, since complexing generally increases with i o n i c charge. The accuracy of equilibrium measurements such as this on solutions above 100°C are limited by uncertain-t i e s i n the experimental techniques used. However, the results obtained f o r the perchlorate system suggest that the extrapolation of room temperature thermodynamic data to higher temperatures as done here i n not an unreason-able procedure. B. THE PRECIPITATION OF METALLIC COPPER FROM AQUEOUS CUPRIC SALT SOLUTIONS BY HYDROGEN INTRODUCTION The method of separating metals from aqueous solutions of t h e i r s a l t s by means of Hg at high temper-ature and pressures has been known f o r nearly 100 years. (37) Beketoff (1859) succeeded in p r e c i p i t a t i n g Ag and Hg from d i l u t e solutions of t h e i r s a l t s under moderate hydrogen pressures. However, attempts to produce me t a l l i c (38) (39) copper were unsuccessful u n t i l 1909 , when I p a t i e f f and co-workers began an extensive investigation of the displacement of metals and me t a l l i c oxides from solutions (10) by hydrogen at elevated temperatures and pressures It i s of interest to review b r i e f l y some of th i s early -39-work. I n i t i a l experiments, conducted on neutral solutions of Ag, Hg, Cu, Zn and Cd salts,, indicated that •H2 pressures of up to 600 atmospheres would not produce ' (39) any metal at 2£°C . However, by increasing the temper-ature, basic metal sal t s or metal oxides could be produced at moderate Hg pressures. For example, treatment of a neutral C u S O i solution under 25> atm. H 2 p a r t i a l pressure o , , (40,1*1) at 90 C yielded a basic s a l t , CuSO^- 2Cu(0H) 2 . When contact times were increased, separation of CugO took place, and after 7 days, the only product observed was metallic copper. On increasing the Hg p a r t i a l pressure to 200 atmospheres, pure Cu° was found to p r e c i p i t a t e r e a d i l y . I t was thus concluded that the formation of basic s a l t s and me t a l l i c oxides occurs by hydrolysis of the s a l t s , whereas the: p r e c i p i t a t i o n of the metals with Hg i s a separate reaction, Formation of the basic sa l t s and oxides of the metals Cu, Hi , Mn, Fe, Pb, As, Sb, etc. (41-47) were also reported between 1909 and 1931 Experiments conducted with Hg pressures bsetween 100 and 200 atmospheres and temperatures ranging from 100 to 35>0°C resulted i n p r e c i p i t a t i o n of pure Cu, N i , Co, (33,48) B i , Pb, Sb and As from solutions of t h e i r s a l t s + Of p a r t i c u l a r i n t e r e s t i s the study of the eff e c t of H ion on-the p r e c i p i t a t i o n of m e t a l l i c copper from i t s s a l t solution. The reaction of Hg with cupric formate solution under various conditions showed that the formation of Cu° - k o -rather than basic s a l t s or Cu 20 was favoured by the + (k9) presence of H . This was also true of experiments on the sulphate and acetate systems. Hence further evidence that basic s a l t s p r e c i p i t a t e by means of hydrolysis was advanced, since increasing H + concentration suppresses hydrolysis and results i n the formation of the pure metal. The fact that Cu20 p r e c i p i t a t e d i n low-acid solutions suggested that cuprous ions form i n cupric s a l t solutions during hydrogen reduction. A subsequent i n v e s t i -gation carried out to determine under what conditions + cuprous s a l t s form and what influence the presence of H ions i n solution has resulted i n the following conclusions regarding the p r e c i p i t a t i o n of metallic copper: 1. While cuprous ions are formed during* the reduc-t i o n of a cupric s a l t solution by H 2, under no conditions was i t possible to obtain a solution containing only Cu + ions. 2. The displacement of Cu° from solution by Hg never goes to completion, and the. residual ++ + solution contains both Cu and Cu ions. Formation of me t a l l i c copper w i l l not take place unless both io n i c species e x i s t i n the solution. 3. After the expiration of a certain time, the r a t i o of the concentrations of Cu and Cu remained approximately constant. These conclusions suggested that the entire system exists -41-i n a state of dynamic equilibrium involving Hg and both + ++ Cu and Cu ions, which i s l i k e l y governed by the follow-ing reactions: 2CUS0J+ * Hg CugSO^ + HgSOk (28) CuSOk + H 2 5 = i 2Cu° + HgSO^ (29) CuSO^ + C u 0 * ^ CugSO^ (30) I t was also observed that the rate of p r e c i p i -tation increased with Increasing cupric s a l t concentration, but decreased only s l i g h t l y with increasing a c i d i t y . These investigations, however> were l a r g e l y of a q u a l i t a t i v e nature and did not permit elucidation of the k i n e t i c s and mechanisms of the reactions involved. No further investigations of the reduction of metals from aqueous, solutions of t h e i r s a l t s by H 2 seem to have been conducted u n t i l about 1946, when the p o s s i b i l i t y of applying t h i s procedure i n d u s t r i a l l y to the recovery of metals from leach solutions was considered. Results of investigations by Chemical Construction Corporation on the production of Ni, Co and Cu have been reported by Shaufelberger. Although some calculations of reduction e q u i l i b r i a f o r these reactions have been attempted, most of the information reported i s of a q u a l i t a t i v e nature. In p a r t i c u l a r , the f a c t that Hg i s activated homogeneously by cupric s a l t s has only recently been recognized. The present investigation was undertaken with a view to elucidating the k i n e t i c s and mechanism of the reduction of aqueous cupric s a l t solution by molecular hydrogen. -1+2-EXPERIMENTAL The e x p e r i m e n t a l apparatus and p roced u re used i n t h i s p a r t o f the work i s i d e n t i c a l to t h a t employed i n s t u d i e s o f the r e d u c t i o n o f d i c h r o m a t e . However,, the course o f each experiment was f o l l o w e d by d e t e r m i n i n g e l e c t r o l y t i c a l l y the amount of copper r e m a i n i n g i n s o l u t i o n (on c o o l i n g ) a f t e r v a r i o u s time i n t e r v a l s . Temp-+ o e r a t u r e c o n t r o l was o n l y about - 1 C s i n c e these exper iments were, conducted at somewhat h i g h e r t e m p e r a t u r e s . RESULTS AND DISCUSSION 1. REDUCTION OP AQUEOUS CUPRIC PERCHLORATE SOLUTIONS BY MOLECULAR HYDROGEN The course of each experiment was f o l l o w e d by d e t e r m i n i n g the amount of copper r e m a i n i n g (as C u + + ) i n the coo led s o l u t i o n samples . U n l e s s o therwise s t a t e d , the k i n e t i c d a t a r e f e r to these measurements. However, as d e s c r i b e d p r e v i o u s l y , these v a l u e s must be c o r r e c t e d f o r the d i s p r o p o r t i o n a t i o n o f C u + on c o o l i n g , i n o r d e r to + ++ o b t a i n the t r u e c o n c e n t r a t i o n s o f Cu and Cu i n the s o l u t i o n s at the r e a c t i o n t e m p e r a t u r e . A l t h o u g h t h i s c o r r e c t i o n does not a f f e c t the e s t i m a t i o n o f the r a t e o f r e a c t i o n o f h y d r o g e n , i t does c o m p l i c a t e v e r y c o n s i d e r -a b l y the d e t a i l e d i n t e r p r e t a t i o n of the k i n e t i c r e s u l t s , the d e t e r m i n a t i o n of r e a c t i o n o r d e r s and the assessment of the e f f e c t s o f the d i f f e r e n t r e a c t i o n v a r i a b l e s . I n i t i a l experiments conducted on the reduction of cupric perchlorate solutions indicated that conveniently mea-surable rates were attained only at temperatures above l£ 0°C. Hence the reaction was studied between 150 and 1 7 5 ,C, 160 C being selected as the standard reaction temperature. The course of a t y p i c a l experiment i s represented i n Table VII and the corresponding rate curve i s plotted in Figure 8 , It w i l l be noted that the reduction rate f a l l s off with time from an i n i t i a l l y rapid value u n t i l an apparent equilibrium or steady-state i s reached a f t e r about 9 hours. This type of behaviour was observed i n a l l experiments and i t s significance w i l l be discussed more f u l l y l a t e r , pH measurements of cooled solution samples indicated that the H Ion concentration increased with time, r e f l e c t i n g the formation of perchloric acid as one of the reaction products. The only s o l i d reduction product detected was metallic copper. Thus the o v e r a l l reaction, following disproportionation of a l l the Cu + formed i n i t i a l l y , can be represented by C u + + + H 2 = Cu° + 2H + E f f e c t of I n i t i a l C u + + Concentration A series of experiments were conducted i n which the i n i t i a l cupric perchlorate concentration was varied from 0 . 0 £ to 0 . 3 0 M. The results are plotted i n Figure 9 . The i n i t i a l rate of reduction can be seen to increase ++ with increasing i n i t i a l Cu concentration. This i s consistent with the view that Cu p a r t i c i p a t e s i n the -kk-TABLE V I I RESULTS OP A TYPICAL REDUCTION EXPERIMENT I n i t i a l Cu ( C l O ^ ) g C o n c e n t r a t i o n Temperature : l 6 0°C : 2 0 atm. 0.1M Hg 'Pressure I n i t i a l HC10 C o n c e n t r a t i o n 0.1M R e a c t i o n Time (min.) Apparent Copper C o n c e n t r a t i o n (gm/l) pH 0 6 . 1 6 1 . 0 k 3 3 5 ^ 0 . 9 0 67 k . 9 2 1 0 0 k , 5 5 0 . 8 3 1 5 1 k . 2 0 1 9 1 k . 0 5 0 . 7 9 2 5 3 ' 3 - 7 9 3 1 2 3 . 6 0 0 . 7 3 3 7 0 3 - k 9 k 3 0 3 . k 5 0 . 7 0 k 9 0 3 . 3 9 5 5 5 3 - k 3 0 . 7 0 6 1 0 3 . 3 8 655^ 3 . k 5 0 . 7 0 6.0 + o, l + . O 2.0 1 -f=-1 0 0 60 120 21+0 Time -360 Minutes k$o 600 Figure 8. Typical Rate Curve for the Reduction of Cupric Perchlorate Solutions ~with I n i t i a l HCIO^, In i t i a l - CU ( C 1 0 L . ) p, 0 . 1 M / L . .IM/L . T , l 6 0 ° C , H 2 , - 20 atm. -k6-rate-determining step of the reaction. It w i l l also be noted that the apparent f i n a l equilibrium concentrations of Cu follow the same order as the i n i t i a l concentrations, although the differences between f i n a l concentrations are much smaller than between the I n i t i a l ones. This effect i s + attributable to the fact that, the amount of H ion produced in the reaction increases with increasing i n i t i a l C u + + concentration. E f f e c t of H.2 P a r t i a l Pressure Rate plots f o r a series of experiments i n which the Hg p a r t i a l pressure was varied from 10 to 2$ atmospheres are shown in Figure 10. It can be seen that there i s a s l i g h t decrease i n the apparent equilibrium concentration of copper and a s l i g h t increase i n the i n i t i a l rate with increasing hydrogen p a r t i a l pressure. However, th i s e f f e c t is small in comparison with the effects of the other var-iables studied. E f f e c t of I n i t i a l H + Concentration Since p e r c h l o r i c acid i s formed i n the reaction and since the perchlorate solutions were unbuffered, i t was of interest to determine the e f f e c t of varying the i n i t i a l H + concentration on the reaction rate, the apparent ++ equilibrium concentration of Cu and the f i n a l pH. There-fore, a series of experiments was conducted i n which the i n i t i a l acid concentration was varied from 0.01 to 0.20 M. perchloric acid. The r e s u l t s are shown in Figure 11. 3 o 18.0 -12.0 6 . 0 21+0 360 Time - Minutes Figure 9. E f f e c t of I n i t i a l Cupric Perchlorate Concentration 6n the Rate of Reaction I n i t i a l HCIO^: 0.1M/L , l60°C , 20 atm. H2 i i 60 , 120 2k0 360 L80 600 Time - Minutes Figure 10. E f f e c t of E? Partial"Pressure on the Rate of Reaction I n i t i a l Cu(C10jj_)2 0.1M/L ,. I n i t i a l HCIG^ 0.1M/L. l60°C. -1+9-It can be seen that increases i n the i n i t i a l HCIO^ concentration causes a marked decrease i n the i n i t i a l reaction rate and results in larger f i n a l concentrations of copper i n solution. Increasing acid concentration thus affects the reduction of copper adversely. Eff e c t of Temperature Experiments were conducted at temperatures between o 1^0 and 175 C i n an attempt to determine the temperature dependence of the reduction reactions. However, at temper-atures above 160°C, the perchlorate solution became unstable, possibly as the r e s u l t of reduction of perchlorate to chloride by Cu +. The chloride reacted with Cu + to form insoluble CuCl, which precipitated as a white powder. The p r e c i p i t a t e was Identified as a cuprous s a l t by d i s s o l v i n g i t i n ammoniacal solution, Acetylene was then bubbled through th i s solution and a red fl o c c u l e n t p r e c i p i t a t e of cuprous acetylide was obtained. Chloride could also be detected i n the solutions using s i l v e r n i t r a t e . P r e c i p i -tation of cuprous chloride was not detected at temperatures below l 6 0 ° C o Experiments at l£ 0°C and 160°C are shown i n Figure 1 2 . The e f f e c t of increasing temperature is to increase the I n i t i a l rate of reaction and to decrease s l i g h t l y the apparent equilibrium concentration of copper i n solution. Figure 1 1 . Effect of I n i t i a l HCTOK Concentration on the Rate of Reaction I n i t i a l C u ( C 1 0 ^ ) 2 0.1M, 160°C ,. 2 0 atm. H 2 -Si-E f f e c t of M e t a l l i c Copper Surface It was considered of p a r t i c u l a r importance to e s t a b l i s h the e f f e c t of m e t a l l i c copper surface on the reaction rate and on the apparent equilibrium attained. Hence a series of experiments was conducted i n which f i n e copper powder ranging In amount from zero to l i | gms/l. was added i n i t i a l l y to the solution. The rate p l o t s for this series are shown in Figure 13, I t was found that the addition of copper powder had no appreciable e f f e c t on the reaction rate or the f i n a l copper concentration. Thus i t can be concluded that the reaction i s not heterogeneously catalyzed to a s i g n i f i c a n t extent. It was considered of importance to e s t a b l i s h whether the e f f e c t s of the. variables discussed above on the reaction rate could be predicted by the mechanism ++ proposed f o r the- a c t i v a t i o n of H^ by Cu . As discussed previously, t h i s mechanism as applied to this system enta i l s the sequence of reaction steps: C u + + + H 2 '-T^* CuH + + H + k - l ++ + k2 + + Cu + CuH > 2Cu + H 2 C u + f i s t * C u + + + C u The corresponding rate law i s o R a t e * « .(k.jAa [ H + ] | + J0H++J Figure 1 2 . E f f e c t of Temperature on the Rate of Reaction I n i t i a l C u ( C 1 0 k ) 2 , . 0 . 1 M/L, I n i t i a l .HCTOh , 0 . 1 M/L. 20 atm. H 2 - 5 3 -For the purpose of testing t h i s rate-law, i n i t i a l reaction rates obtained from the d i f f e r e n t experiments were used. These were estimated from the i n i t i a l slopes of the experi-mental rate plots (-d [Hg] /dt - -d [pu++] /dt) and the precision of this procedure i s such that they can hardly be considered r e l i a b l e to better than + 10%. However, the rates thus determined re f e r to conditions where the con-++ + centrations of Cu , H and Hg are known with considerable p r e c i s i o n . In l a t e r stages of the reaction there i s some uncertainty, at least about the f i r s t of these. This uncertainty, together with the complexity of the integrated rate equation also makes i t d i f f i c u l t to interpret quanti-t a t i v e l y the v a r i a t i o n in rate during the course of each experiment„ The above rate-law predicts that the rate of reaction should be d i r e c t l y proportional to the hydrogen p a r t i a l pressure. The experimentally measured i n i t i a l o rates (see Table VIII) f o r the Hg series are plotted against Hg p a r t i a l pressure i n Figure l k . The r e s u l t i n g p l o t i s i n good agreement with this predicted r e l a t i o n and i t can be ++ concluded that the i n i t i a l rate of reduction of Cu i s f i r s t order with respect to Hg. The proposed mechanism also predicts a l i n e a r r e l a t i o n between (Rate ^) and the H + concentration. Data from the series of experiments on the acid dependence of the reaction rate, plotted as ( I n i t i a l Rate) 1 vs I n i t i a l [ H j i n Figure If?, are i n accord with t h i s . From 0 I I I J J I L L 1 L L_ 0 60 120 2k0 360 . k80 . 600 Time - Minutes Figure 13. E f f e c t of Metal l i c Copper Surface on the Rate of Reaction I n i t i a l C U ( C 1 0 K ) 2 , 0.1 M / L , I n i t i a l HCIO^, 0.1M/L. T, 160 C. 4 H 2, 20 atm. -CI-TABLE VIII INITIAL REDUCTION RATES FOR VARIOUS EXPERIMENTAL CONDITIONS T = 160°C Cupric Perchlorate Solutions Run No. I n i t i a l [Cu++] M/L I n i t i a l [H+] M/L P H 2 atm. I n i t i a l Rate x 1 0 6 M/L/S (Rate) _ 1xlO^ L.S/M [cu + +]/Rate xlO^ - Sec. 6 2 0 . 1 0 0 . 0 1 20 1 2 . 6 0 . 7 9 60 0 . 1 0 0 . 1 0 2 0 6 . 0 1 . 6 7 6 1 0 . 1 0 0 . 0 5 2 0 8 . 6 4 1.16 6 9 0 . 1 0 0 . 2 0 20 4 - 4 3 2.26 . 6 0 0 . 1 0 0 . 1 0 20 6 . 0 1 . 6 2 . 68 o . o 5 0 . 1 0 2 0 . 1 . 9 2 . 6 5 6 5 0 . 2 0 0 . 1 0 20 1 2 . 6 1 . 5 0 4 5 0.30 0 . 1 0 2 0 3 8 . 1 0 . 7 4 6 0 0 . 1 0 0 . 1 0 2 0 6 . 0 ' 7 0 0 . 1 0 0 . 1 0 15 7 . 2 3 7 1 0 . 1 0 0 . 1 0 10 4 . 0 67 0 . 1 0 0 . 1 0 2 5 1 0 . 0 Hydrogen P a r t i a l Pressure - Atm. Figure li+. Dependence of the Reaction Rate on the Hydrogen P a r t i a l Pressure. T: 160°C ; I n i t i a l I n i t i a l -57--3 the intercept and slope of t h i s p l o t , values of k]_ = 6.3x10 — 1 — 1 l.mole sec and (k-i/kg) = 0„97> respectively, were c a l -culated. Considering the poor precision of the procedures involved, the value of k]_ thus obtained, i s i n reasonably good agreement with the- value of [L..8X10—-5 1. mole sec obtained by extrapolating Peters' d a t a ^ " ^ f o r k^ to 160°C. It i s also of interest to compare the value of ( k - l / k 2 ) at 160°C (0,97) with that reported e a r l i e r at 110°C (0.26). Apparently the r a t i o of these two constants increases with increasing temperature, which implies that the reversal of the hydrogen activation step becomes increasingly favoured at higher temperatures. Thus the i n h i b i t i n g e f f e c t of acid on the reactions would appear to be greater at 160°C. On the basis of the proposed mechanism, i t can also be predicted that f o r constant H + concentrations and R"2 pressures there w i l l be a l i n e a r r e l a t i o n s h i p between |cu + +J/rate and Jcu4"^ 1 . The measured i n i t i a l rates of reaction f o r the C u + + series of experiments have been plotted i n thi s fashion i n Figure 16. Although there i s some scattering of points, the results are seen to conform i n general to the expected pattern. From the slope and intercept of the straight l i n e which was drawn through the points, values of k-j_ = 8.9x10 ^ 1. mole 1 sec 1 and ( k-l/k2) =1.7 have been calculated. These values d i f f e r appreciably from those estimated e a r l i e r from the H + dependence of the rate, but the agreement i s probably as -58-2 .5 r-i 1 O CD CO 2 . 0 —< "LA O r-i •l . S ' X -rH 1 w EH h-H 1 .0 <u NIT] o.5o — 0 i i 0 0 . 1 0 . 2 I n i t i a l [H + ] - M / L Figure l 5 . Dependence of the I n i t i a l Reaction Rate on the I n i t i a l H + Concentration. 0.1M C u + + : 160°C : 20 atm. H 2 -59-good as can be expected considering the uncertainties involved i n estimating i n i t i a l rates from experimental data. Mean values of k l and ( k - l / k 2 ) from both sets of experiments are 7.6x10"3 1. mole - 1 s e c - 1 and 1.3h> respectively. In a l l the reduction experiments,, the reaction rate gradually decreased with time u n t i l an apparent equilibrium was attained, although the thermodynamic considerations discussed e a r l i e r would suggest that the reactions should go ++ v i r t u a l l y to completion, leaving only trace amounts of Cu + and Cu ions In solution. Assuming the f i n a l stage of each experiment (i.e.,. a f t e r reduction had v i r t u a l l y stopped) to represent an equilibrium, apparent equilibrium constants ++ have been calculated from the f i n a l concentrations of Cu , Cu and H . F or the purposes of this calculation,' the equilibrium 2 C u + C u + + + Cu° ++ was assumed to apply and the true concentrations of Cu and Cu + at the reaction temperature were estimated from , the known equilibrium constant for this reaction, as des-cribed e a r l i e r . Values of these constants thus estimated are given i n Table IX^. and f o r purposes of comparison the values obtained by extrapolation of the available thermos-chemical data are also shown. It seems u n l i k e l y that the uncertainties involved i n the estimation of these constants, i.e.,. i n the extra-polation procedure, would t o t a l l y account f o r the very large discrepancies between the orders of magnitude of the - 6 1 -two sets of values. Hence; i t seems more, l i k e l y that the apparent equilibrium i s due to k i n e t i c f a c t o r s , i . e . , the slowing down of the reaction to the point where i t v i r -t u a l l y stops. Such a slowing down might be expected on the basis of the rate-law derived e a r l i e r -since the. concentration of Cu"^ decreases and that of H increases during the/ reaction. Using the mean values of k-^  and (k-l/k2) obtained previously the expected reaction rates f o r the f i n a l stages of several experiments have been calculated, and are l i s t e d i n Table X. It i s seen that the predicted rates while quite slow should s t i l l be measurable and thus the apparent l e v e l l i n g - o f f of the reactions i s not completely explained. However,, i t seems ++ l i k e l y that the reduction of Cu to m e t a l l i c copper by Hg follows the proposed mechanism pr a modification thereof since the dependence, of the experimentally determined rates on various reaction variables i s at least q u a l i t a t i v e l y i n accord with the behavior predicted by t h i s mechanism. One possible explanation of the f i n a l slowing doWn of the reaction is that the second step gives r i s e to the formation of an intermediate which may undergo a reverse + ++ reaction to regenerate CuH and Cu thus competing with the formation of Cu"1", i . e . , CuH + + C u + + ^ z z ± (intermediates) -J^.Cu* + H + TABLE IX APPARENT EQUILIBRIUM CONSTANTS CALCULATED PROM EXPERIMENTAL DATA Run No. Temp OC P H 2 atm. I n i t i a l [Cu++] M/L I n i t i a l [H+| M/L Apparent P i n a l [Cu++] M/L |cu++] M/L calc . M/L calc . P i n a l [M. M/L calc. pH K l x i o 3 k 2 2 xlO^ Calc exp. 45 160 20 0.280 0.1 0.12k 0.107 0.035 0.411 • 3S - 30.5 7.9 160 20 0.182 0 .1 0.073 0.059 0.029 0.314 .50 .69 35 .4 8 . 4 . 60 160 20 0.097 0 .1 o .o53 0.042 0.022 0.186 .73 0.69 21.9 4 . 1 68 160 20 0.05 0.1 0.039 0.030 0.019 0.121 •12 0.88 17.0 2.k 61 160 20 0.101 o .o5 0.042 0.032 0.019 0.169 • 77 .78 22.9 4 . 5 62 160 20 0.096 0.01 0.037 0.028 0.018 0.128 .SI .89 18.7 3 .0 69 160 20 0.097 0.20 0.063 0.05 0.025 0.269 .57 .53 29 .0 7.1: 67 160 25 0.104 0.10 0.054 0.043 0.022 .189 .71 .70 19.8 3.3 70 160 15 0.098 0.10 0.054 O.043 0.022 .188 .73 .79 25 .4 5 .5 71 160 10 0.09k 0.10 0.0606 0.047 0.027 .166 .78 .79 3 0 . 0 5 .8-75 150 20 0.093 0.10 0.054 0.045 0.018 .177 .TT .72 15.6 3.5 #91 160 20 .097 0.10 0.051 0.040 .022 .192 .72 • 79 23.5 4 . 6 #92 160 20 .097 0.10 0.052 0.041 .022 .174 .76 .79 20.7 3.7 #93 160 20 .107 - 0.10 0.046 0.036 .021 .221 .U .79 28.2 6.7 contained added amounts of copper powder. Values of Kn a n d K2 calculated from thermodynamic data *1(160) = 9 2 8 . * l ( 1 c t f ) ^ ( 1 6 d ) = 7 - 3 x l ° 7 tentf) = 887 = 1.12x10 - 6 3 -Su'ch a reaction may become increasingly important as the ++ + Cu concentration decreases and. the H concentration increases toward the l a t t e r stages of an experiment and would e f f e c t i v e l y contribute to the further slowing down of the reaction. .2. REDUCTION OF AQUEOUS CUPRIC SULPHATE .SOLUTIONS BY H 2 A comparison of the reduction of cupric perch-lorate and cupric sulphate solutions by Hg was of intere s t ++ •+ because complexing of Cu and H ions occurs i n the l a t t e r system. The c a t a l y t i c a c t i v i t y of the CuSO^ complex i s .++ (30) known to. be greater- than that of the simple Cu ion , so that reaction rates should be f a s t e r . Also, the s u l -phate system provides some buffering action through the formation of the r e l a t i v e l y stable HSO^- ion, which i s of significance since: H + Ions are formed during the p r e c i p i t a t i o n process. Hence, a series of experiments was conducted under conditions s i m i l a r to those employed in the reduction of perchlorate solutions by Hg. The resu l t s of this i n v e s t i g a t i o n are presented below. The course of a t y p i c a l reduction experiment Is given i n Table XI,. which shows the v a r i a t i o n of apparent ++ Cu concentration and pH with time. A corresponding rate curve i s plotted i n Figure 17, and a plot of a si m i l a r experiment using perchlorate solution i s shown f o r comparison. It can be seen that the i n i t i a l rate of reaction i s much f a s t e r than f o r the. corresponding reaction -61t-TABLE X REACTION RATES CALCULATED PROM CONCENTRATIONS OBSERVED IN PINAL STAGES OP EXPERIMENTS 160°C Run No. C u + + meas. (M/L) [ c u + + ] Calculated M/L . H2 Pressure (atm.) P i n a l [H+J calculated M/L Meas-ured* Rate (gm/ 1/hr) Calcul-ated-"-"* Rate (gm/1/ hr) 6 0 o . o 5 3 o. 0 4 2 2 0 0 . 1 8 6 0 0 . 2 1 62 0 . 0 3 7 . 0 . 0 2 8 2 0 0 . 1 2 8 0 . 0 6 0 . 1 3 6 5 0 . 0 7 3 0 . 0 5 9 2 0 0 . 3 1 4 • 0 . 1 4 0 . 2 4 67 o . o 5 4 0 . 0 4 3 2 5 0 . 1 8 9 0 . 0 3 0.26 68 0 . 0 3 9 0 . 0 3 0 2 0 0 0 1 2 1 0 0 . 1 6 6 9 0 . 0 6 3 0 , 0 5 1 2 0 0 . 2 6 9 0 . 1 1 0 . 2 2 Rate of reaction measured after, 9 hrs. under H2 pressure-. Rate; of reaction calculated using concentrations of .substances i n solution a f t e r 9 hrs. under HP pressure. - 6 5 -TABLE XI RESULTS OP A TYPICAL REDUCTION EXPERIMENT O.IM CuSO^ O.O5M H2SO^  1 6 0 GC 20 atm. H2 Sample Time (min) [Cu + +] app. g.p.l. pH 1 0 5 . 9 5 1 .38 2 30 2 . 04 3 62 1 .55 0 . 9 7 4 105 1 .31 5 i 5 o 1.12 0 . 9 1 6 220 0 . 9 9 7 286 0 . 9 8 0 . 9 0 8 346 0 . 9 0 9 431 0 . 8 4 0 . 8 9 10 482 0 . 8 9 11 544 0 . 8 5 0 . 8 9 Time - Min Figure 17. Typical Rate Curves f o r the Reduction of Cupric Salt Solutions with Hydrogen _ I n i t i a l Cupric S a l t , 0.1M. I n i t i a l Acid, 0 . 1 N . , T, 160 C. H 2, 20 atm. - 6 7 -++ In the perchlorate system. The apparent Cu concentration at which the rate, levels o f f i s also much lower. These two factors were generally true of a l l experiments with sulphate solutions. The only reaction products detected were H + ions,, whose concentration was measured by a pH meter, and m e t a l l i c copper. A photomicrograph of the l a t t e r product i s shown i n Figure l&a. The metallic copper pre c i p i t a t e shown In Figure l 8 a . i s t y p i c a l of that produced by reducing Cupric s u l -phate solutions with H 2. In most cases a l l the product was found as a fibrous mass loosely adhered to the i n t e r i o r of the autoclave: l i n e r . Microscopic examination shows that the m e t a l l i c copper' p r e c i p i t a t e s l a r g e l y as very ti n y needles,, although some small p l a t e - l i k e crystals are: also produced. The form of th i s p r e c i p i t a t e i s much • d i f f e r e n t from that observed by C o u r t n e y t o form on slow disproportionation of Cu + at 2£°C when ammonical solutions of Cuprous sulphate were a c i d i f i e d . The l a t t e r product appeared to be £ to 1 0 f l o e s of 0.5yU polyhedral c r y s t a l s . ^ ^ ) This p h y s i c a l difference i s probably due to the f a c t that the i n i t i a l H 2 reduction rate i n the s u l -phate system i s very rapid (approximately 80$ reduction takes place i n the f i r s t two hours),, r e s u l t i n g i n the formation of a large number of n u c l e i which can grow rapidly to form needles of m e t a l l i c copper. The p r e c i -p i t a t e of m e t a l l i c copper from the reduction of cupric perchlorate solutions by H 2 (Figure (lgtb)), however, was Figure l 8 a „ Photomicrograph of M e t a l l i c Copper Product on Reduction of Cupric Sulphate S o l u t i o n s w i t h H- x 3 0 0 dm v i Figure l 8 b Photomicrograph of M e t a l l i c Copper Product on Reduction of Cupric P e r c h l o r a t e S o l u t i o n s w i t h H 2 . x 8 0 - 6 9 -found to be comprised of metal globules,, adhering i n loosely-bound conglomerates of various siz e s . Some adher-ing of copper on the autoclave: l i n e r takes place, but this also appears as a network of copper globules. The marked difference between the physical form of the. metallic copper p r e c i p i t a t e reduced from sulphate and perchlorate- solutions by H can probably be attributed to the difference i n the rate of displacement of metal from solution i n each system. E f f e c t of I n i t i a l Cupric Sulphate Concentration A series of experiments were carried out in which the i n i t i a l concentration of cupric sulphate was varied from O.Of? to 0 . 3 0 M, The r e s u l t s , plotted i n Figure 19,- show that the reaction rate i s very rapid i n i t i a l l y , and increases with increasing cupric sulphate concentration. The concentrations at which the rates l e v e l off are much lower than f o r s i m i l a r experiments on cupric perchlorate solutions. Also, the difference between f i n a l concentrations i n sulphate- .solution are much smaller, presumably due to the buffering action of the- S 0 ^ = ions. Eff e c t of I n i t i a l HgS.O^ Concentration The e f f e c t of i n i t i a l HgSOj^ concentration, which was varied from 0.01 to 0.3M, on the, reaction rate i s shown i n Figure 2 0 . Increasing i n i t i a l acid concentration Is seen to cause, a marked decrease i n i n i t i a l rate: and -70-an e a r l i e r l e v e l l i n g off of the reaction,. I.e.,. a higher-++ f i n a l Cu concentration. Thus^. with increasing a c i d i t y , the behavior approaches that of the perchlorates:. This i s to be expected since the effectiveness of complexing and of buffering are both reduced' Ef f e c t of H"2 P a r t i a l Pressure The: e f f e c t of varying Hg p a r t i a l pressure on the reduction reactions i s shown i n Figure 21. Higher Hg pressures re s u l t i n s l i g h t l y higher i n i t i a l rates and ++ lower f i n a l Cu concentrations. However, from a p r a c t i -c a l standpoint, the effect of increasing Hg p a r t i a l pressure seems to be small compared with the e f f e c t of varying the i n i t i a l acid or i n i t i a l cupric sulphate" con-centrations. Ef f e c t of Temperature-The r e s u l t s of reduction experiments carried out at four temperatures between 150 and 175>°C are. reproduced in Figure 22. Increasing temperature produces increased i n i t i a l rates and lower f i n a l copper concentrations. The amount of copper reduced i s increased by about %% f o r a ,o 25 C r i s e i n temperature. Ef f e c t of Adding N^SO^ The e f f e c t of S0^~ concentration was of inte r e s t f o r two reasons. I t i s known that the c a t a l y t i c a c t i v i t y of the CuSO^ complex i s about 7 times that of the: simple Cu ion. Hence, by increasing the concentration of SO]^  , Time - Minutes Figure 19. E f f e c t of I n i t i a l CuSOj, Concentration on the Rate of Reaction I n i t i a l R^SO^, 0.05M. T, 160°C. H 2, 20 atm. I ->] IV) I Time - Minutes Figure 20. E f f e c t of I n i t i a l H2S0k- Concentration on the Rate of Reaction I n i t i a l CuSO, , 0.1 M/L, T, I 6 O 0 C . H 2, 20 atm. 600 Time - Minutes Figure 2 1 . E f f e c t of Hp P a r t i a l Pressure on the Rate of Reaction I n i t i a l CuSO^, 0 . 1 M/L,- I n i t i a l H 2 S 0 ^ , 0 . 0 £ M/L. T, 1 6 0°C. -7k-an increase. In rate of reduction would be expected due to further' complexing of the Cu ion. Also, increasing S0|^  concentration should increase the buffering action Of the solution and this should also cause an Increase i n the. reaction rate. Therefore,- a series of experiments was conducted i n which various amounts of NagSO^ ranging from zero to 0.2M. were added to the so l u t i o n j the r e s u l t i n g rate plots are shown In Figure 23. The: most s i g n i f i c a n t trend observed is that the f i n a l concentration of copper i n solution decreases markedly with increasing sulphate concentration. Thus complete reduction of copper i n about 6 hours i s attained by the addition of 0.2 M/L of NagSO^ to the solution. It can also be seen that the i n i t i a l rates increase with increasing sulphate concen-tration,, as predicted. E f f e c t of Metallic Copper Surface Various amounts of f i n e copper powder were added to cupric sulphate solutions i n order to determine the effect of copper product P h the reaction rate. The r e s u l t s , plotted i n Figure 21)., show that the addition of copper metal has no s i g n i f i c a n t effect on the reduction process, which supports' the view that the reaction rate i s homogeneously^determined. 0 I I I I j I J J j J L_ 0 60 120 2k0 360 k80 600 Time - Minutes Figure 22. E f f e c t of Temperature on the Rate of Reaction I n i t i a l CuSOj, : 0.1 M/L,. I n i t i a l H 2S0^ : 0.0$ M/L, H 2 : 20 Atm. 3 i o I 0 60 120 21+0 360 Time - Minutes h.80 600 Figure 23. E f f e c t of Na o S0l, Additions on the Rate of Reaction I n i t i a l CuSog : 0.1 M / L , - I n i t i a l HgSOj, : 0.0£ M / L , . T:l60 C ^ Ho : 20 atm. COPPER POWDER O - None • - 3-5 OM/L A - 7.0 GM/L O - lk.O GM/L QQ M-o o 0 60 120 2k0 360 Time - Minutes k80 600 Figure 2k. Effec t of Metallic Copper Surface on the Rate of Reaction I n i t i a l CUSOK : 0.1 M/L, I n i t i a l H 2 S 0 h : 0.05 M/L, T : 160°C, ^ H 2 : 20 Atm. - 7 8 -CONCLUSIONS COMPARISON OF THE REDUCTION OF CUPRIC PERCHLORATE AND CUPRIC SULPHATE SOLUTIONS BY Hg The; general shapes of the rate, plots f o r reduction of cupric perchlorate and cupric sulphate; were s i m i l a r . The rate, which was i n i t i a l l y fast,, gradually decreased 0 with time u n t i l an apparent equilibrium or steady state condition was approached. Reaction rates were much f a s t e r i n the sulphate system, and the f i n a l copper concentrations lower, presumably due to the greater r e a c t i v i t y pf the CuSO^ complex and to the buffering action of the s u l -phate. In both systems, the rate of reaction was apparently homogeneously-determined as evidenced by the fact that addition of copper powder was without appreciable e f f e c t . Increasing i n i t i a l cupric s a l t concentration, temperature and H 2 pressure r e s u l t s i n higher i n i t i a l reaction rates, the l a t t e r e f f e c t being the least pronounced. Increasing i n i t i a l acid concentration markedly decreases the rate and increases the f i n a l concentrations of copper i n solution. This variable has the most pronounced e f f e c t on the reaction rates i n both systems. Additions of Na2S0j^ to cupric sulphate solutions permitted more complete, reduction to copper metal, presumably as a re s u l t of increased buffer-ing. MECHANISM OF THE REACTION The dependence; of the reaction rate on each of the -79-variables considered i s i n q u a l i t a t i v e agreement with that predicted by the mechanism: C u + + + H 2 CuH + + H + k - l CuH + + C u + + ,k2"> 2Cu+ + H + r e s u l t i n g in an o v e r a l l reaction of: ++ o + Cu + H 2 > Cu + 2H Reasonably good quantitative agreement between the pre-++ + dieted k i n e t i c dependence on i n i t i a l Cu and H concen-trations and Hg p a r t i a l pressure- and that found from me as-sured i n i t i a l reduction rates has also been demonstrated. Hence i t .seems probable that the rate of reduction of C u + + to me t a l l i c copper i s determined by the homogeneous activation of molecular hydrogen by C u + + or a cupric com-plex. While the observed'decrease i n rate as the reaction proceeds can be accounted f o r by the above mechanism, the f i n a l attainment of an apparent equilibrium or steady state,, remains to be f u l l y explained. SUGGESTIONS FOR FURTHER WORK In order to elucidate the ley e l ling'-off of the reaction rate to an apparent steady-state condition, the following work could be carried out. F i r s t , i t would be useful to determine more accurately the. temperature co-e f f i c i e n t of ( k-l/ k2') , which would e n t a i l carrying out a series of experiments at various temperatures. Also, -80-exchange studies using DgO-enriched water as the solvent could be useful i n providing further support f o r the f i r s t step i n the reaction sequence,, as well as a measure of (k-l/kg) r Experiments along this l i n e are presently being conducted by Webster on the homogeneously-catalyzed reduction of Gr p0„ i n the. presence of Ag +. -81-BIBLIOGRAPHY 1. P.A. Forward,. C.S. Samis and V. Kudryk, Trans. C.I.M., 5_1, 350 (1948). 2. F.A. Forward, Trans. C.I.M., £ 6 , 363 (1953). 3. F.A. Forward and V.N. Mackiw,, Trans. A.I.M.E. . 203. 457 (1955). 4. F.A. Forward^ Mining Congress Journal,- 40* 49 (1954). 5. J.A, Butler, Eng. and Min. Jour., l52 , 56 (195l) . 6. F.A. Forward and J. Halpern, Trans. C.I.M., 56, 344 (1953). 7. E. Peters and.J. Halpern,. Trans. C.I.M., , 5 6 , 350 (1953). 8. F.A. Forward and J. Halpern,. Trans. A.I.M.E., 200. 1408 (1954)• 9. F.A. Forward and J . Halpern, Trans. A.I.M.E., 203. 463 (1955). 10. V.N. I p a t i e f f , J . Chem. Educ ., 3.0, 110 (1953). 11. F.A. Forward and J. Halpern, Trans. C.I.M., 5_6, 354 (1953). 12. R.N. O'Brien, F.A. Forward and J. Halpern,- Trans. C-.-I.M.,. £ 6 , . 359 (1953). 13. F.A. ShaufeTberger and T..K. Roy, Trans. The Inst, of Min. and Met., 6}±, 375 (1955). 14. F.A. .Shaufelberger,, Trans. A.I.M.E., 206,. 695 (1956). 15. G.J. Korinek, "The Kinetics of the Reduction of Mercuric S a l t s by Molecular Hydrogen i n Aqueous Solution,"- PhD. Thesis, University of B r i t i s h Columbia (1956). 16. H. Webster,. Unpublished Data. 17. J. Halpern and E. Peters,. J . Chem. Phys., 23, 605 (1955). 18. J. Halpern and E. Peters, J . Phys. Chem., 59, 793, (1955). - 8 2 -19. G.J. Koririek and J. Halperny J . Phys. Chera., 6 0 , 285 (1956). ~*~ 2 0 . A.H. Webster and J . Halpern,. J. Phys. Chem., 60 , 280 (1956). 21 . R.N. Pease, Jour. Am. Chem. Soc., j ^ . , 1877 (1932). 22. 0. Beeckj. Diss. Faraday S o c , 8 , 122 (1950) . 23 . 0. Beeck and A.W. R i t c h i e , Ibid.. 8 , l 5 9 (1950). 2i+. A. Couper and D.D. Eley... Ibid., '8-,. 17'2 (1950). 25. D.A. Dowden and P..W. Reynolds,. I b i d . 8 , , X8J4. (1950) . 26. M. Calvin,: Trans. Faraday Soc.,. 2k*• H 8 l (1938). 27. R.G. Dakers and J . Halpern, Can. Jour. Chem., ^ 2 , 969 (1954)• 28. E. Peters and J . Halpern, Ibid., 21» 356 (1955). 29. A.H. Webster and J . Halpern, Trans. Far, Soc., In Pre,s s. 30. E. Peters, "The Homogeneous C a t a l y t i c A c t i v a t i o n of Molecular Hydrogen By Cupric Salts i n Aqueous Solution," PhD. Thesis, University of B r i t i s h Columbia- (1956) . 31 . E.A. Moelwyn-Hughes, "The Kinetics of Reactions i n Solutions," 2nd e d i t i o n , Oxford University Press, London (1947) pp. 68-77. 32. J . Halpern, E.R. Macgregor and E. Peters, J . Phys. Chem., 60, 1455 (1956). 33- V. I p a t i e f f and V. I p a t i e f f , J r . , Ber., 62,, 386 (1929). 34. Heinerth,, Z. Elektrochem,21, 61 (1931). 35. C.T. Kawassiadis, Praktika Akad., Athenon,. 10,, 391 (1935). 36. Z.Z. Hugus, Jour. Am. Chem. S o c , • jQ, £4^9 (195l) • 37- M.N. Beketoff, Compte Rendu, 1^ 8,, 44-2 (1859). 38. G. Tammann and W.. Nerst, Z. Phys. Chem., 9 , 1 (1892). 39. V". I p a t i e f f and W. Werchowsky, Ber., 1+2, 2078 (1909). -83-kO. V. I p a t i e f f and W. Werchowsky, Ber. ,. kjj., 1755 (1911). k l . V. I p a t i e f f , Ber., ^9, l k l 3 (1926). k2. V. I p a t i e f f , Ber., kjk_, 3452 (1911). . 43. V. I p a t i e f f and N. Kondyrew, Ber., 5 .^, 1421 (1926). kk. V. I p a t i e f f and A. Kisselew, Ber., 1I4.I8 (1926) . k5. V. I p a t i e f f and N. Nikolajew, Ber. ,.J^9_, Lk23 (1926) . k6. V. I p a t i e f f and B. Muromzew, Ber.,, 60, 1980 (1927). 47- V. I p a t i e f f and A. Starynkewitsch, Ber.,- £6 , - 1663 (1923). ' 48. V.V. I p a t i e f f , J r . , Ber., 6k_, 2725 (1931). i | 9 . V. I p a t i e f f and V. I p a t i e f f , J r . , Ber., 60_, 1982 (1927) . 50. Welby G. Courtney, J. Phys. Chem... 60, l 4 6 l (1956). 

Cite

Citation Scheme:

        

Citations by CSL (citeproc-js)

Usage Statistics

Share

Embed

Customize your widget with the following options, then copy and paste the code below into the HTML of your page to embed this item in your website.
                        
                            <div id="ubcOpenCollectionsWidgetDisplay">
                            <script id="ubcOpenCollectionsWidget"
                            src="{[{embed.src}]}"
                            data-item="{[{embed.item}]}"
                            data-collection="{[{embed.collection}]}"
                            data-metadata="{[{embed.showMetadata}]}"
                            data-width="{[{embed.width}]}"
                            async >
                            </script>
                            </div>
                        
                    
IIIF logo Our image viewer uses the IIIF 2.0 standard. To load this item in other compatible viewers, use this url:
https://iiif.library.ubc.ca/presentation/dsp.831.1-0081192/manifest

Comment

Related Items