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The kinetics of the reaction between cupric acetate and molecular hydrogen in aqueous solution Dakers, Ronald Gill 1953

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THE KINETICS OF THE REACTION BETWEEN CUPRIC ACETATE AND MOLECULAR HYDROGEN IN AQUEOUS SOLUTION by RONALD GILL DAKERS A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF SCIENCE in the Department of Mining and Metallurgy at the University of British Columbia We accept this thesis as conforming to the standard required from candidates for the degree of 8Master of Science' i n Metal Chemistry Members of the Department of Mining and Metallurgy THE UNIVERSITY OF BRITISH COLUMBIA SEPTEMBER^ 1953. ABSTRACT A kinetic study was made of the reaction between cupric acetate and hydrogen in aqueous solution. The reaction was followed by determining the concentration of cupric acetate remaining in solution and was found to proceed in accordance with the equation 2CuAc2 + H2 + H20 Cu20 + AHAc The reaction was found to be second order kinetically, the rate being proportional to the concentrations of cupric acetate and molecular hydrogen. It was established that the reaction is homogen-eous, both the stainless steel vessel and the cuprous oxide product being without effect. The rate increased slightly when sodium or ammonium acetate in concentrations of 0.25 moles per liter was added to the solution but remained unchanged with further addition of these salts. The rate was also independent of the concentration of acetic acid and the pH of the solution over a wide range. The activation energy was found to be 23,400 cal. per mole and the frequency factor -12 -1 -1 1.0 x 10 1, mole sec. corresponding to an entropy of activation -1 =1 of - 5.9 cal. mole deg„ . The following mechanism has been proposed to account for these kinetic results, the first step in the reaction sequence being rate determining slow 1. CuAc2 + H2 -* CuAc2 . H2 fast 2. CuAc2 . H2 + CuAc2 -» ' 2CuAc + 2HAc fast 3. 2CuAc + H20 ~* Cu20 + 2HAc This appears to be the first established instance of a homogeneous hydrogenation reaction in aqueous solution. ACKNOWLEDGEMENT The author is indebted to the National Research Council of Canada for the financial assistance given, without which this work could not have been carried on0 The author is grateful for the assistance, advice and encouragement given by the members of the Department of Metallurgy and especially to Dr0 J 0 HALPERNj, under whose direction this invest-igation was carried out0 TABLE OF CONTENTS Page Introduction 1 Hydrogenation Reactions 1 Properties of Molecular Hydrogen 2 Reduction of Metal Salts in Aqueous Solutions 4 Homogenous Hydrogen Reactions in Solution 6 Object and Scope of the Present Investigation 7 Experimental 10 Materials 10 Equipment 10 Procedures 11 Results and Discussion 13 Chemistry of the Reaction 13 Kinetics of the Reaction 14 Reproducibility of the Rate Measurements 19 Effect of Stirring Velocity and Solution Volume 21 Effect of initial Cupric Acetate Concentration 21 Effect of Surface 26 Effect of Cuprous Oxide Product 26 Effect of Acetate Salts 29 Effect of Acetic Acid and pH 2 9 Effect of Hydrogen Partial Pressure 35 Effect of Temperature 43 Conclusions Summary of Kinetic Results Mechanism of the Reaction Form of the Reacting Cupric Acetate General Conclusions Appendix A Spectrophotometric Method for Determining Copper in Aqueous Solution Appendix B Summary of Kinetic Results Appendix C X-ray Analysis of Cuprous Oxide Product Bibliography LIST OF FIGURES Figure No. Page 1 0 Rate Curve showing the Reproducibility of the Reaction 1 7 2 . F i r s t Order Plot showing the Reproducibility of the Reaction 18 3 o Rate Curves at Various I n i t i a l Concentrations of Cupric Acetate 2 2 4. Fi r s t Order Plots at Various I n i t i a l Concentrations of Cupric Acetate 2 3 5 . Rate Curve showing the Effect of Surface 2 7 6. First Order showing the Effect of Surface 28 7 . Rate Curve showing the Effect of Cu 2 0 Product 3 0 8. F i r s t Order Plot showing the Effect of Cu a 0 Product '"' 3 1 9 o Rate Curves at Various Concentrations of Sodium Acetate 3 2 1 0 c First Order Plots at Various Concentrations of Sodium Acetate 3 3 1 1 . Rate Curves showing the Effect of Acetic Acid and pH 3 6 1 2 . F i r s t Order Plots showing the Effect of Acetic Acid and pH 3 7 1 3 o Rate Curves at Various Partial Pressures of Hydrogen 3 8 (continued next page) LIST OF FIGURES (continued) Figure No,, Page 1 4 o First Order Plots at Various Partial Pressures of Hydrogen 3 9 1 5 o Effect of Hydrogen Partial Pressure on the Rate of Reaction 40 1 6 o Rate Curves at Various Temperatures 4 4 1 7 » First Order Plots at Various Temperatures 4 5 1 8 o Arrhenius Plot 4 7 1 9 o Absorption Spectrum of Copper in Ammonia Solution 6 0 2 0 , Dependence of the Optical Density of Copper on the Ammonia Concentration in Aqueous Solution 6 0 2 1 0 Calibration Graph for Copper in Ammonia Solution 6 1 22„ X-ray Diffraction Photographs of Cu 20 Product 6 6 LIST OF TABLES Table No. Page I Results of a Typical Kinetic Experiment 15 I I Reproducibility of Rate Measurements 20 I I I Effect of I n i t i a l Copper Acetate Concentration on the Rate of Reaction 24 IV Effect of Volume, Surface and of Cu20 Product on the Rate of Reaction 25 V Effect of pH, Acetic Acid, and Acetate Salts on the Rate of Reaction 34 VI Effect of Hydrogen P a r t i a l Pressure on the Rate of Reaction 41 VII Effect of Temperature on the Rate of Reaction 46 VIII X-ray D i f f r a c t i o n Measurements on the Cu 20 Product 65 1. THE KINETICS OF THE REACTION BETWEEN CUPRIC ACETATE AND MOLECULAR HTDROGEN IN AQUEOUS SOLUTION INTRODUCTION Hydrogenation Reactions. Because of the many important applications of hydrogenation reactions and because of interest in the fundamental problems of kin-etics and catalysis connected with them, the study of these reactions has received a good deal of attention over the past half century. Most of the applications of hydrogenation reactions lie in the field of organic chemistry and include such well known processes as the reduction of oils to edible fats, the reduction of fatty acids to higher alcohols, the conversion of unsaturated petroleum fractions to gasoline and the formation of methanol from carbon monoxide. There are also some important applications in the f ield of inorganic chem-istry, a notable example being the high temperature synthesis of ammonia by reaction between nitrogen and hydrogen. Application of hydrogenation reactions in the field of metallurgy include the production of metals by the reduction of oxides. More recently interest has been focused on the reactions of hydrogen with metallic salts in aqueous solutions. Although i t has been known for many years (l) that molecular hydrogen, under suitable conditions 2 . of temperature and pressure will react with the salts of many metals, such as Gu, Ni, Co, Pb and Bi in aqueous solution, reducing them to the metallic state or to compounds of lower valence, such reactions were not considered to be of practical interest, except perhaps as a method of preparation of metallic catalysts (2). Recently i t has been shown that i t is possible to use this method to produce pure metals, such as Ni and Co, on a commercial scale and to recover metals such as uranium and vanadium as oxides of lower valence from leach solutions ( 3 ) . Processes of this nature have been subjected in recent years to intensive investigation, chiefly by workers in the Metallur-gical Laboratories of Sherritt Gordon Mines Ltd,, Chemical Construction Corporation and the University of British Columbia. This research, although largely empirical in nature, has brought to light many im-portant facts and applications. The work described in this thesis was undertaken primarily with a view to obtaining information of a more fundamental nature, relating to the kinetics and mechanism of such hydrogenation reactions in aqueous solution. Properties of Molecular Hydrogen. The configuration and properties of the hydrogen molecule are well established. The molecule consists of two hydrogen atoms, held together by forces which may be thought of as a covalent bond, arising from electrostatic and exchange interactions of the two elec-trons in the molecule with each other and with the protons which constitute the nuclei of the hydrogen atoms. The bond dissociation energy of the molecule (i.e. the energy necessary to separate the molecule into two hydrogen atoms), derived either from quantum mechan-ical calculations (4) or from experimental date (5) is 103 kilocal-ories per mole. In reactions involving reduction by molecular hydrogen the products are either hydrides in which the hydrogen atoms become bonded to the atoms or molecules being reduced, or as is more general-ly the case in aqueous solutions, ions. In either case, i t is evident that the bond in the molecule must be broken and the two hydrogen atoms separated in the course of the reaction. Because of the high bond dissociation energy involved, such reactions would norm-ally be expected to have activation energies of the order of 50 kilo-calories per mole or higher , and hence to require temperatures of the order of 500°C. It has been found possible to reduce this activation energy, in the case of many hydrogenation reactions, by employing suitable catalysts. The best known catalysts for this purpose are finely divided heavy metals, notably Pt, Ni and Co and certain oxides such as ZnO and Cr 20 3. Reference has also been made to the catalytic properties of the crystals of certain complex organic molecules, notably phthalocyanine (7). Attempts have been made to relate this catalytic activity to properties such as the semiconductivity and d band character of the solids involved, but many features are s t i l l incompletely understood. In general, i t is believed that the catalytic effects in _ , , , , , , • •• — For example,the activation energy of the reaction Hp + CO —• H2C0 has a value of 61 kilocalori.es per mole. See Poianyi 'Reference 6) 4. hydrogenation reactions reflect the activation of molecular hydrogen by the catalyst. The mechanism probably involves the adsorption of mqlecular hydrogen on the surface of the catalyst and i t s accompanying dissociation. Where the catalyst i s a metal i t i s l i k e l y that in this process the hydrogen atoms become chemically bonded to the sur-face metal atoms. Thus the spacing of the metal atoms must be such that the adsorption of the molecule i s f a c i l i t a t e d and the two atoms readily form bonds with adjacent metal atoms. This i s supported by the observation that the (100) and (110) plane surfaces of a nickel crystal, in which the separation pf adjacent atoms i s 3.52 A., are catalytically more active than the other surfaces (8). Other metal catalysts are also known to have interatomic spacings of this order. Wherever these geometrical requirements are f u l f i l l e d and where the metal atoms form f a i r l y strong bonds with the hydrogen atoms i t i s l i k e l y that pronounced catalytic activity w i l l be observed. It has been suggested (6) that the activation energy for the dissociation of hydrogen in such systems i s substantially lowered by resonance favour-ed by the high density of energy states (i.e. high degree of degen-eracy) i n the metal. This would account for the fact that activation energies as low as 5-10 kilo calories per mole are often observed in metal catalyzed hydrogenation reactions. Reduction of Metal Salts in Aqueous Solutions. As implied above most known hydrogenation catalysts are solids and where these are employed the hydrogenation reactions are heterogeneous in nature,occurring on the catalyst surface. This ap-plies to reactions i n which the reacting species are dispersed in solution as well as to gas phase hydrogenation reactions. For example studies of the reduction of nickel salts in aqueous solution, suggest that the metallic nickel product provides the catalytic surface on which the nickel ions are reduced. In the absence of such a surface, reduction by molecular hydrogen is slow and i t is generally necessary to nucleate the reaction by adding some finely divided nickel or by adding another reagent which will reduce some of the dissolved nickel salt to the metallic form. In view of the well known catalytic activity of metallic nickel i t is not surprising that the reduction of its salts should proceed in this manner. It might also be anticipated that the reduct-ion of salts of other metals, which do not themselves show any cata-lytic activity, should require the presence of an added catalyst. In conformity with this i t has been found (3) that the reduction of uranyl and vanadate salts in aqueous solution by hydrogen occurs only in the presence of a catalyst such as finely divided metallic nickel. On the other hand i t has been found (1) that the salts of certain other metals, such as Cu, Pb and Bi, which in metallic form do not generally display catalytic activity in hydrogenation reactions, can also be reduced by molecular hydrogen in aqueous solution, without addition of any other known catalytic agent. The nature of the catalysis in these reactions, and the mechanism by which the molecul-ar hydrogen becomes activated was considered to be of great theo-retical interest. In particular, i t was felt that some consideration should be given to the possibility that such reactions might be homo-geneous in nature and involve the homogeneous activation of molecular hydrogen by some molecule or ion in the solution. Unfortunately none b. of these reactions had been examined kinetically or otherwise studied in sufficient detail to indicate whether this might be the case. Homogeneous Hydrogenation Reactions in.Solution. Although most theories of hydrogenation catalysis refer only to solid catalysts, a number of workers, notably Polanyi and Calvin, have considered the question as to whether i t might be pos-sible to effect such catalysis homogeneously. Such a process would be of considerable significance since i t would make i t possible to study the activation of molecular hydrogen in a relative simple system, leading to more interpretable results than can be obtained with heter-ogeneous reactions involving solid catalysts. It would also demon-strate effects which are not accounted for by present theories of hydrogenation catalysis. It is also apparent that the application of such a process might be of some practical interest. Although the possibility of homogeneous activation of molecular hydrogen has been suggested in connection with several reactions, only one or two cases appear to have been definitely estab-lished. These involve reactions in organic solvents. As early as 1938, Calvin (9) showed that cuprous acetate, dissolved in quinoline was capable of activating molecular hydrogen and catalyzing such reactions as the reduction of cupric acetate and quinone. He sug-gested that the reactions involved were homogeneous. This has recent-ly been confirmed by Y/eller and Mills (10) who examined the same system and obtained similiar results. The reaction appears to involve the formation of a complex between hydrogen and a dimer of cuprous acetate, as the rate determining step. The hydrogen becomes activated ' 7. i n t h i s step, so that i t s subsequent reaction with the dissolved substrate ( i . e . cupric acetate or quinone) i s rapid. The cuprous acetate appears to act as a true, catalyst and i s regenerated. The reaction appears to be r e s t r i c t e d to a very narrow range of solvents, a l l of which are nitrogen containing compounds, s i m i l i a r to quinoline . Thus there are indications that the solvent molecules are also i n -volved i n the complex that i s formed between hydrogen and cuprous acetate. The only other instance of a reaction apparently involving homogeneous hydrogenation Is that between hydrogen and Co 2(C0)g lead-ing to the formation of HCo(C0)4 (11), which takes place i n organic solvents such as ether. This process has not been as extensively studied as the cuprous acetate reaction and the evidence f o r i t s homogeneous character i s less conclusive. There i s very l i t t l e i n -formation r e l a t i n g to i t s k i n e t i c s or mechanism or to the species responsible f o r a c t i v a t i n g the hydrogen. Object and Scope of.the Present Investigation. The investigation.which i s described i n t h i s t h e s i s , was o r i g i n a l l y undertaken with the object of examining the k i n e t i c s of some reactions between molecular hydrogen and metal s a l t s i n aqueous solution, p a r t i c u l a r l y with a view to establishing whether any of these reactions might be homogeneous i n character. The particular reaction, which was selected for study i n the f i r s t instance, was the reduction of cupric acetate by hydrogen to form cuprous oxide, which was f i r s t observed by I p a t i e f f ( l ) . Among the factors governing A — • — — Cuprous acetate cannot be dissolved i n aqueous solution without decomposition, disproportionation or hydrolysis (to Cu 20). 8. this choice, which in the light of subsequent work appeared to be justified, were the following. 1. As indicated above, the reduction of cupric acetate had already been investigated by Calvin and others and shown to occur homogeneously in certain organic solvents, although the catalytic agent responsible for activating the hydrogen in those systems was cuprous acetate and not cupric acetate itself. It was recognized that the reaction in aqueous solution is fundamentally different from that in organic solvents in that insoluble cuprous oxide rather than soluble cuprous acetate is formed as the product. 2. The possibility that a metallic product might be re-sponsible for catalyzing the reaction, as is the case with the reduct-ion of salts of nickel and other metals, is eliminated since no metal-lic copper but only cuprous oxide is formed. However, the possibility that the cuprous oxide might have some catalytic activity also had to be considered. . * 3. Solutions of cupric acetate do not hydrolyze as readily as those of other copper salts, such as cupric sulphate, when heated to the high temperatures (100-150°C.) necessary to effect reaction with hydrogen. The complicating features which might be introduced by such hydrolysis and the resulting precipitation of basic salts, could thus be more readily prevented. 4. By adding convenient amounts of acetic acid and acetate salts, the solutions could be effectively buffered and the possible complications arising from the formation of acid during the reaction 9. thus eliminated. 5. The system was a very convenient one from the exper-imental point of view. The reagents could be obtained in pure form and the solutions conveniently prepared and handled. The reaction could be very readily followed by determing the amount of copper remaining i n solution. The preliminary results obtained on this reaction were considered encouraging and a comprehensive investigation of i t s kinetics was undertaken. The reaction was carried out in an autoclave in which the solution of cupric acetate was maintained under control-led temperatures and pressures of hydrogen. Rates of reaction were determined by periodic sampling and analysis of the solutions. The effect on the rate of the following variables was determined. (a) cupric acetate concentration (b) concentrations of acetic acid and acetate salts . ' ( c ) pH (d) stainless steel surface (e) cuprous oxide product (f) pa r t i a l pressure of hydrogen (g) temperature The results of this investigation are presented and dis-cussed below. 10. , EXPERIMENTAL A. Materials. (i) Reagent grade cupric acetate, ammonium acetate, sodium acetate, acetic acid and other chemicals used in preparing the solutions and in the analytical determinations were supplied by-Nichols Chemical Company. (ii) Hydrogen and Nitrogen gas, were supplied in cylinders by the Canadian Liquid Air Company and were used without further puri-fication. ( i i i ) Distilled water was used in the preparation of a l l the solutions. B. Equipment. (i) Autoclave. The precipitation experiments were conducted in an auto-clave manufactured by Autoclave Engineers Inc., and designed for work-ing pressures up to 2000 pounds per square inch. A l l parts in contact with the solution were made of 316 stainless steel. The vessel had a capacity of one gallon, corresponding to a diameter of five inches and a height of fourteen inches, inside dimensions. It was equipped with a set of three stirrers, each of 2 1 /2 inches diameter, mounted on an externally driven shaft, whose speed of rotation could be varied from 400 to 900 R.P.M. The thermocouple well, stirrer, cooling coil and sampling tube were connected through the l i d of the autoclave and extended below the surface of the solution. The sampling tube was 11. fitted with a fine porosity stainless steel f i l t e r and the samples withdrawn for analysis were thus free of cuprous oxide product. The pressure of hydrogen above the solution was controlled by a standard gas regulator, allowance being made for the partial pressure of the solution. The autoclave was heated externally by a ring type gas burner and the temperature of the solution recorded and maintained to within *1°C. by a Leeds and Northrop Micromax controlling recorder. (ii) Spectrophotometer. A Beckman D0U0 Spectrophotometer was used to determine the copper concentration of the solutions. ( i i i ) pH Meter. pH measurements were made with a Beckman Model H-2 A.C. pH Meter. C. Procedures. (i) Analytical. Copper in the solutions was determined spectrophotometric-ally after addition of ammonia. Details of the procedure and the calibration curve are given in Appendix A. (ii) Kinetic Experiments. The solution was prepared by dissolving weighed amounts of cupric acetate and other salts in 3 liters of distilled water and adjusting the pH to the desired value by addition of glacial acetic acid. This charge was then placed in the autoclave which was sealed, flushed with nitrogen and heated to temperature. At temperature, the autoclave was thoroughly flushed with hydrogen which was then intro-12. -duced and maintained at a desired pressure throughout the experimentT The reaction was followed by withdrawing 25 ml. samples of the solution from the autoclave at measured time intervals, generally o f t h e o r d e r of 5 minutes, and determining the copper concentration and pH. The experiments were generally continued until analysis indicated that-about 98 percent of the copper had reacted and been precipitated, After cooling, the charge was removed from"the autoclave, and the pro-duct separated from the solution by filtration. It was washed with distilled water, dried in a vacuum dessicator and analyzed for copper. At the completion of each experiment the autoclave was cleaned through-ly with dilute nitric acid and rinsed with water to remove any adher-ing cuprous oxide. 13. RESULTS AND DISCUSSION Chemistry of the Reaction. The solutions of cupric acetate used for the reaction studies, generally contained acetic acid and either sodium or ammonium acetate, added as buffering agents. The i n i t i a l pH of the solutions was generally in the range 4.5 to 5.0, and could be adjusted by vary-ing the amount of acetic acid. If the i n i t i a l pH was allowed to rise above a value of 5.0, some hydrolysis of the cupric acetate, resulting in the formation of cupric oxides and basic cupric salts, was observed when the solution was heated to the reaction temperature. When hydrogen was introduced into the autoclave and ab-sorbed by the solution, the immediate formation of a cuprous oxide precipitate was detected. As the reaction proceeded the amount of this precipitate increased and the concentration of copper remaining in solution correspondingly f e l l off. At the same time the pH of the solution also decreased, reflecting the formation of acetic acid as one of the reaction products. The reaction proceeded in every case until the amount of copper remaining in solution was less than 0.2 . g./l. This generally reflected 98-99 percent reaction of the copper originally present. The only reduction product of cupric acetate that could be detected in any of the experiments was cuprous oxide. This was ident-ified both by chemical analysis and by X-ray diffraction measurements. No metallic copper or cupric salt impurities could be detected in the precipitate. An observation of particular interest was that the ap- . pearance of the cuprous oxide was different, depending on whether the solution contained.sodium or ammonium acetate. In the former case the product was bright red while in the latter i t ranged from dark purple to black in colour. Both products were crystalline solids and no difference between them either in composition or in crystal struct-ure could be detected. The difference in colour was presumably due to differences in the grain size. Reference is made to similar varia-tion in the colour of cuprous oxide, in the literature (13). Anal-yses of some of the products are listed in Appendix B, and are seen to conform closely to theoretical composition of C U 2 O Cu). Some typical X-ray diffraction data are given in Appendix G. The products were a l l found to have X-ray patterns identical with that recorded for Cu20 (face centered cubic structure; lattice parameter = 4.26 A . Ref.14). These observations point conclusively to the reaction re-presented by the following equation; 2 Cu Ac2 + H2 + H 20— Cu20 + 4 HAc * (l) As indicated above, the reaction goes essentially to completion, with Cu20 and HAc being the only detectable products. Kinetics of the Reaction. The course of a typical experiment is depicted in Table I where values of the copper concentration and pH at various time inter-— _ _ The symbols CuAc2 and HAc will be used in this thesis to denote cupric acetate and acetic acid, respectively. 15. TABLE I. Results of a Typical Kinetic Experiment. Initial Cu - 15 g./l. (0.236 m./l.) NaAc - 0.25 m./l. HAc - 0.44 m./l. H2; - 13.6 atm. Temp. - 125°C". Reaction Time min. Optical Density 1/20 Dilution Cu Concentration g./l. Log M pH 0 0.615 14.89 1.173 4.49 5 ' 0.495 11.98 1.079 4.41 10 0.377 9.13 0.960 4.36 15 0.273 6.73 0.828 4.30 20 0.204 4.94 0.694 4.20 25 0.148 3.58 0.554 4.19 . 30 0.108 2.62 0„418 4.14 35 0.076 1.84 0.265 4.10 40 0.053 1.28 0.107 4.10 4f 0.040 0.97 - 0.013 4.10 50 v 0.028 0.68 - 0.167 4.10 55 . 0.021 0.51 - 0.292 4.10 60 0.016 0.39 - 0.409 4.10 65 0.010 0.24 - 0.620 4.10 70 0.007 0.17 - 0.769 4.09 16. -vals after the hydrogen is introduced into the reaction vessel, are listed. A corresponding typical rate curve showing how the cupric acetate concentration decreases with time is shown in Figure 1. It can be seen that the rate falls off as the reaction proceeds. A plot of the log of the concentration of cupric acetate against time is linear as shown in Figure 2 . This was the case for a l l the experiments which were conducted. The linear relation implies that at constant temperature and partial pressure of hydrogen, the rate of reaction is first order with respect to the concentration of cupric acetate. This may be expressed mathematically in the form of the well known f i r s t order rate equation. - d CcuAc?] = k,(~CuAc27 . . . . . . . ( 2 ) dt On integration, and evaluation of the integration constant, this be-comes log fcuAcpJo = k.t . . . . . . . ( 3 ) rCuAc2J ?.103 where [CuAcpJo and £"CuAc2J are the concentrations of GuAc2 at zero - time and time t, respectively, and k, is the f i r s t order rate constant of the reaction defined by equation ( 2 ) . Equation ( 3 ) is seen to correspond to a linear relation between log )_CuAc2J and t, such as that shown in Figure 2 . The rate constant is related to the slope of this plot as follows: kt = - 2 . 3 0 3 x slope (k) 16. 20 30 40 T I M E - M I N U T E S FIGURE: 1. Typical rate curves for the reaction between Cupric Acetate and hydrogen in aqueous solution, showing reproducibility of measurements. NH^ .Ac, 0,25 m/j , HAc, 0.18 m/1. Initial pH , 4.5. H2, 13.6 atm. T., 125 18 TIME - MINUTES FIGURE 2. Typical first order plots for the reaction at constant H2 partial pressure. 19. This relation was used to evaluate the rate constant for each exper-iment from the slope of the corresponding first order plot. The slight i n i t i a l increase in the rate constant, apparent in Figure 2, is believed to reflect the fact that the zero.on the time scale refers to the time when hydrogen is first introduced into the reaction vessel. A few minutes will be required for the solution i t -self to become saturated with hydrogen, and during this period such an increase in the rate is to be expected. All the experiments showed a similar effect. However, after about five minutes, when the solu-tion had presumably become saturated with hydrogen, no deviation from the linearity of the rate plots could be detected. Reproducibility of the Rate Measurements. The reproducibility of the rates, determined in this way, is illustrated in Figures 1 and 2 and in Table II, where the results are listed for two series, each one consisting of three experiments conducted under identical conditions. In the fi r s t series the max-imum deviation from the mean rate was 7.1 percent and the mean deviation - 4.6 percent. In the second series the corresponding deviations were 0.7 percent and - 0.6 percent. In general i t is felt that the accuracy of individual rate determinations is about - 5 percent or better. This accuracy is believed to be limited in the system used, by the precision with which the temperature could be regulated. 20. TABLE II. Reproducibility of Rate Measurements. Initial Cu Concentration % 15 g./l. (0.236 m./l.) Temperature : 125°C H2 Pressure s 13.6 atm. Experiment . NaAc . NH^ Ac HAc Rate Constant No. m./l. m./l. m„/l„ k i rrlnT1 Deviation fron "ean -41 0.25 0.44 0.0643 +' 4.1 54 0.25 - 0.44 0.0573 - 7.1 36 0.25 0.44 0.0633 + 2.6 Mean Value 0.25 0.44 0.0617 ± 4.6 8 0.25 0.18 0.0617 - 0.7 14 - 0.25 0.18 0.0619 - 0.4 rr ! 0.25 0.13 0.0626 + 0.7 Mean Value - 0 .25 0.18 0.0622 ± 0.6 2 1 . Effect of Stirring Velocity and Solution Volume. It was established that under the experimental conditions employed, the rate of reaction was independent of the stirring velocity over a considerable range and also remained unchanged when a set of baffles was mounted on the inner walls of the reaction vessel. This confirms that the dissolution of hydrogen is fast in relation to the rate of reaction, and that the solution therefore remains effect-ively saturated with hydrogen throughout the reaction. Under these conditions there is no change in the concentration of hydrogen in the solution during each experiment. It was also established (see Table IV) that the reaction -rate was independent of the volume of solution. Thi3 adds support to the conclusions noted above. Effect of Initial Cupric Acetate Concentration. Two series of experiments were made in each of which the i n i t i a l concentration of cupric acetate was varied from 0.079 to 0.315 moles per lit r e (5 to 20 g./l» of Cu). The results for one of these series are plotted in Figures 3 and 4 and a l l the results are summarized in Table III. The first order rate constants are seen to be independent of the i n i t i a l cupric acetate concentration, as well as of the cupric acetate concentration throughout each exper-iment. This confirmsthat at constant temperature and partial pres-sure of hydrogen the rate of reaction is first order with respect to cupric acetate. It should be noted that most of the experiments were j 20 0 I 1 J I I I 0 10 20 30 40 50 TIME - MINUTES FIGURE 3 . Rate curves for the reaction at different; i n i t i a l cupric acetate concentrations, NaAc, 0,25 m/1. HAc, 0 , 4 4 m/1. Initial pH, 4«5, H2, 13,6 atm. t , 125°C FIGURE, 4. Effect of i n i t i a l cupric acetate concentration on rate. First order plots. TABLE III. Effect of Initial Copper' Acetate Concentration on Rate of Raaetioriq H2 - 13.6 atm. Temp. - 125°C. Initial NaAc NH^ Ac Initial Initial Rate Cu Cone. HAc pH Constant. g./l. m./l. m./l. m./l. m./l. k8s, min."^ 5 0.079 0.25 = 0.31 4.6 0.0643 10 0.157 0.25 - 0.38 ' 4.5 0.0668 15 0.236 0.25 - 0.44 4.5 0.0643 20 0.315 0.25 - 0.40 4.5 0.0654 10 0.157 - . 0.50 0.18 5.1 0.C65S 15 0.236 - 0.50 0.18 5.1 0.0617 20 0.315 - 0 .50 0.18 5.0 0.0645 25. TABLE IV. Effect of Volume,, Surface and of  G'uaO Product on Reaction Rate. Initial Cu Concentration - 15 g./l. (O.236 m./l.) H2 Pressure - 13.6 atm. Temperature - 125°C. Volume of Solution litres Stainless Steel Powder1 Added g./l. Cu20 Product Added g./l. Rate Constant k,- min."1 2.0 0.0643 3.0 - - 0.0643 3.0 NONE - 0.0617 3.0 20 g./l. - .0.0617 3.0 — NONE 0.0643 3.0 - 15 g./l. 0.0604 Estimated surface area = 100 cm /gm. 26. i continued until at least 95% of the cupric acetate originally present had reacted and that the first order relation persisted over this whole region. Effect of Surface. It was considered of particular importance to establish that the surface of the autoclave which was constructed of 316 stain-less steel had no catalytic effect on the reaction. An experiment was therefore conducted in which 20 grams per litre of similar stainless steel powder was added to the solution. The surface area of this powder was estimated to be about 100 cm per gram. The total area of the added powder was at least five times the surface area of the autoclave in contact with the solution. It was found that the addition of this powder had no ef-fect whatsoever on the rate. The results which show this are given in Figures 5 and 6 and in Table IV. This is considered as conclusive evidence that the reaction is not heterogeneous and is not catalyzed by any metallic surface in contact with the solution, but rather that i t occurs homogeneously in the solution itself. Effect of the Cuprous Oxide Product. It was also considered of importance to establish that the cuprous oxide product which was formed had no effect, either of a catalytic or retarding nature, on the reaction. To this end an ex-periment was conducted in which some 15 grams per litre of cuprous oxide product obtained in earlier experiments, was added to the sol-27. TIME - MINUTES _5_. Rate curves showing effect of surface. NaAc, 0 .25 m/] HAc, 0 . 4 4 m/1. Initial p H , 4 . 5 , H 2 > 13.6 atm. t, 125»C. 28 FIGURE 6 , Effect of surface on the rate of reaction. First order plots. 2 9 . -ution. Again no effect whatsoever on the rate could be detected. This is shown in Figures 7 and 8 and in Table IV. This result also points to the completely homogeneous character of the reaction. Effect of Acetate Salts. A series of experiments was conducted in which different amounts of sodium acetate or ammonium acetate, ranging from zero to 0.75 m./l., were added to the solution. Some comparative rate plots are given in Figures 9 and 10 and the results are summarized in Table V. The following effects may be noted. 1. Sodium acetate and ammonium acetate had identical ef-fects on the reaction rate, when present in similiar concentrations. The reaction does not appear to be influenced by the nature of the cation of the acetate salt. 2;. While the rate was independent of sodium or ammonium acetate concentrations over a fairly wide range (at least from 0 .25 to 0.75 m./l.), i t was slightly lower when neither, of these salts was present. This suggests that while the acetate ion itself does not participate in the rate determining step of the reaction, the copper which reacts is in the form of an acetate containing complex. The evidence for the existence of such a complex and its nature will be discussed in a subsequent section. Effect cf Acetic Acid and pH. ^ Since acetic acid is one of the products of the reaction, FIGURE 7. Effect of Cu20 product on the rate. NaAc, 0.25 m/-T • HAc, 0.44 m/3. Initial pH, 4.5, K2, 13.6 atm. T,125°C. 31. FIGURE 8. Effect of Cu20 product on the rate. First order plots'. FIGURE 9 Rate curves showing effect of sodium acetate concentration Initial pH, 4 . 5 , H 2 , 13.6 .atm., % 125°C. 33. FIGURE 10. Effect of sodium acetate concentration on the rate. First order plots. 3 4 . TABLE V. Effect of pH, Acetic Acid.and Acetate Salts on the Rate of Reaction. Initial Cu Concentration - 15 g . / I . (0.236 m./l.) Hg Pressure - 13.6 atm. Temperature - 125°C. NH^ Ac m./1. NaAc m./l. HAc m./l. Rate Constant k| - min.-l Initial Final 0 0 0.24 4.4 3.3 0.0451 0,25 tt 0 . 0 9 5.1 4.3 0.0613 0.50 tt 0.18 5.1 4.6 0.0617 0,75 « 0.26 5.1 4.7 0.0617 0 0 0.24 4.4 3.3 0.0451 tt 0.25 0.44 4.5 4.1 0.0643 it 0.425 O.64 4.5 4.2 0.0629 0 0.25 • 0 .29 4.8 4.3 0.0631 tt It 0.44 4.5 4.1 0.0643 n It 1,17 4.0 3.8 0.0691 w tt 2.35 3.7 3.6 0.0696 35. the effect of its concentration and the resulting pH of the solution, on the rate was considered of interest. A series of experiments was therefore conducted in which the i n i t i a l concentration of acetic acid was varied from 0.29 to 2.35 m./l. The corresponding i n i t i a l pH values of the solution, ranged from 4.8 to 3.7 and the final values from 4.3 to 3.6. Within experimental error this variation in the acetic acid concentration and in the pH appeared to have no appreci-able effect on the reaction. This is shown by the results i n Figures 11 and 32 and in Table V. Since the variations examined were much larger than occur-red during any individual experiment, i t may be concluded that the acetic acid produced in the course of the reaction, and the consequent change in the pH of the solution are also without effect. At the same time i t would appear that neither acetic acid molecules nor hydrogen ions participate in the rate determining step of the reaction. Effect of Hydrogen Partial Pressure. The dependence of the rate on the par t i a l pressure of hy-drogen was established i n a series of experiments in which the part-i a l pressure was varied from 6.8 to 3 4 . 0 atmospheres (100 to .500 p s i ) . The rate curves are shown in Figures 13 and 14 and the results sum-marised in Figure 15 and Table VI. The rate of the reaction, represented by the f i r s t order rate constant, is seen to be directly proportional to the par t i a l pressure of hydrogen. Since the. sol u b i l i t y of hydrogen in asosi-aqueous solutions, in the range of pressure investigated, follows Rate curves showing effect of acetic acid concentration and pH . .NaAc, 0 . 2 5 H 2 , 1 3 . 6 atm. T. f125°C. 37. FIGURE 12. Effect of acetic acid concentration and pH on the rate. First order plots. / 38, 16 Ha PARTIAL PRESSURE M a 1 o 3 o o • V ^> A 6.8 Atm. 13.6 Atm. 20.4 Atm. 27.2 Atm. 34.0 Atm, 12 24 36 4S 60 TIME MINUTES F I G U R E 13. Rate curves showing effect o f H2 partial pressure. NaAc, 0.25 m/7„ HAc, 0.44 m/1. Initial pH, 4.S„. T„ 125°C. FIGURE 14o- Effect of H2 partial pressure on the rate of reaction. First order plots. 40. 0 10 20 30 H2 PARTIAL PRESSURE - ATM. .FIGURE 15. Plot of rate as a function of H2 partial pressure. 41. TABLE VI. Effect of H;> Partial Pressure on Reaction Rate. Initial Cu Concentration - 15 g./l. (0.236 m./l.) NaAc 0.25 m./l. HAc 0.44 m./l. Temperature L25°C H2 Partial Rate Constant Rate " Pressure atm. (k.) min."*-1- ? 2 k2, Pressure min."1 atm.-l 6.8 0.0334 0.00491 13.6 0.0643 0.00473 20.4 0.0990 0.00485 2 7 . 2 0.132 0.00485 34.0 0.163 0.00479 Mean value of kj,-0.00483 42. Henry's Law (15), this signifies that the rate is also proportional to the concentration of dissolved hydrogen in the solution. The first order rate constant k,, defined in equation ( 2 ) can thus be expressed in terms of the concentration of hydrogen^£H2J, in the solution, as follows. " * 2 fej . . . . . . . ( 5 ) where k 2 is a second order rate constant, independent of the concentrations of both CuAc2 and H 2. Combining equations (3) and (5), the rate of the reaction becomes, - d fcuAc J = k 2 [ C U A C 2 J [ H 2 J ( 6 ) dt This represents a second order reaction, where the rate is first order with respect to the concentrations of both CuAc2 and H 2. As shown earlier, when the concentration of one of the reactants (i.e. H2) is kept constant during any given experiment the reaction follows first order kinetics. It is also possible to express the rate of the reaction in terms of the partial pressure of hydrogen instead of its concentration. The modified rate equation becomes, - d [cuAcJ = k 2 f3uAc27 P„ . , ..(7) dt L- 2 where ?^ is the partial pressure of hydrogen above the solution, k2; is a modified rate constant defined by, 43. k 2 = k 2 a .(8) and a is Henry's constant denoting the solubility of hydrogen in the aqueous solution and defined by £H2J| = c ^ H2 ..•«..»(9) The value of a, for solutions of the composition in these experiments, has not been determined but the corresponding value denoting the solubility of H2 in pure water is known to be 7.14 x 10"^ moles atm."^ I."""*" (15). It is of interest that this value has been found to be independent of temperature in the range 60° to 100°G,, where some of the present experiments were conducted. Such determinations as have been made of the effect of acetic acid and other salts on the solubility of hydrogen in water, would suggest that the value of a, for solutions of the composition used in these experiments (i.e. approximately 0.5 M. HAc and 0.25 M. NaAc or NH/^ Ac), would be about 10 percent lower than for pure water or approximately 6.4 x 10*"^  moles atm."1 1. \ in the temperature region 60° to 100°C. It is believed that the error involved in as-suming this value for purposes of the present calculations is negligib-ly small. Effect of Temperature. The rate of the reaction was determined at six different temperatures, ranging from 75° to 135°C, over which its value varied by a factor of some 140. The rate curves, for these experiments are shown in Figures 16 and 17 and the results summarized in Table VII. Log k, is plotted against l/T in Figure 18 and i t is seen that an FIGURE 16. Rate curves at various temperatures. NaAc, 0.25 m/1. ..' HAc, 0 . 4 4 m/1. Initial pH, 4.5 H2, 13.6 atm. FIGURE 17. First order plots showing rate of reaction at different temperatures. TABLE VII, Effect of Temperature on the ReactionRate^ Initial Cu Concentration - 15 g./l. (0,236 m./l.) NaAc - 0.25 m./l. HAc - 0.44 m./l. H2 Partial Pressure - 13.6 atm. Temperature °C. 1 Rate Constant k,, min." 1 I Log kB j T°k 75 0.00287 0.00106 ! - 2.975 95 0.00272 0.00606 - 2.217 105 0.00264 0.0161 - 1.793 1 115 0.00258 0.0313 - 1.505 125 0.00251 0.0643 - 1.192 135 0.00245 0.149 - 0.823 FIGURE. 18., Arrhenius plot showing rate of reaction as a function of temperature. 48. excellent Arrhenius plot, which remains l i n e a r over the entire region, i s obtained. The value of the activation energy, calculated from the slope of t h i s plot using the r e l a t i o n — E = 2.303 x R x slope . . . . . . . ( 1 0 ) was found to be 23,400 calories per mole. This value i s believed to be accurate to within * 500 calories per mole. No attempt was made to correct t h i s value f o r the i n f l u -ence of temperature on the s o l u b i l i t y of hydrogen ( i . e . on a ). Such a correction i s necessary since the p a r t i a l pressure of hydrogen and not i t s concentration i n the solution, was maintained at a constant value, throughout the series of experiments at different temperatures. However, as Indicated e a r l i e r there i s evidence that the s o l u b i l i t y of hydrogen remains pearly constant over at least a considerable portion of the temperature range involved, and hence that the cor-rection f o r t h i s effect would be n e g l i g i b l y small and probably within the experimental error of the determined ac t i v a t i o n energy. 49. CONCLUSIONS Summary of the Kinetic Results, It has been shown that the reaction between cupric acetate and hydrogen i n aqueous solution, i s a second order process, f i r s t order with respect to the concentration of each of the reactants. The rate can thus be expressed by the relations Rate a - d CcuAcajf = k 2 rCuAc27/H2"] »....•.(6) dt J or = k 2 £CuAc] 2 P H .......(7) j The values of k 2 or k 2 were found to be independent of the concentra-tions of acetic acid, acetate salts and of the pH of the solution over a wide range. They were also found to be independent of any surface effects connected with the walls of the reaction vessel or the pres-ence of the cuprous oxide powder. The dependence of k;x on the temperature i s given by the ex= pressions k a » A exp - 23400 j .......(11) L HT. J Combining equations 5, 9 and 11, the following relation i s obtained. k i a PH«i " A e x p f ~ 23400 1 .......(12) L R T J t At 75°C and 13.6 atmospheres H 2 partial pressure, the value of k was -1 found to be 0.00106 min. . Using this value, and assuming the value , -4 -1 -1 of a to be 6.4 x 10 moles atm. 1. , from considerations discussed previously, equation (12) permits A, the frequency factor of the 5 0 o reaction to be evaluated. The value was found to be. 1 2 - 1 - 1 A = 1 . 0 x 1 0 ... 1 . mole sec. . . . . . . . ( 1 3 ) A can further be related to the entropy of activation of the reaction, Z\S , by means of the Eyring r e l a t i o n ( 1 6 ) . kz m e kT e e ^ ^ • » . . . . . ( 1 4 ) h whence A. » f e kT e . . . . . . . (15 ) h -.16 - i where k i s Boltzmann's Constant (1.38 x 1 0 erg deg. ) - 2 7 and h i s Planck's Constant ( 6 . 6 x 10 erg sec.). A The value o f A S calculated from equation ( 1 5 ) iss A - 1 i As. r = - 5.9 c a l . mole deg.~ x the standard state being chosen as 1 mole per l i t r e . I t might be noted at t h i s point that t h i s value of Z^S^ and the corresponding value of A, derived e a r l i e r , are very close to the values expected f o r a bimolecular reaction i n solution. Mechanism of the Reaction. A l l the k i n e t i c r e s u l t s reported and discussed above sug-gest that the rate determining step of the process investigated i s a bimolecular reaction between cupric acetate and molecular hydrogen s which occurs homogeneously i n the solution. The form of the cupric e i n these solutions, w i l l be discussed subsequently s but f o r the present i t w i l l be as iumed to exist and react as the molecule 51. CuAc2, The product of this bimolecular reaction is a complex between CuAc2 and H2 in which the hydrogen has apparently become activated so that its dissociation and further reaction with cupric acetate to give the observed products is rapid. The sequence of steps involved in the overall reaction may thus be represented as follows. k 2 1. CuAc2 + H 2 -* CuAc2 . H2 (16) " slow 2. CuAc2 , H 2 + CuAc2 2CuAc + 2HAc .#..,..(17) 3. 2GuAc + H20 fa-^. Cu20 + 2HAc . . . . . . . ( I B ) Overall Reaction: 2CuAc2 + H2 + H20 -> Cu20 + l&Ac (1) The f i r s t step apparently is the slowest one, determing the rate of the overall process. The second and third steps involv-ing the reduction of the cupric acetate to cuprous acetate and the hydrolysis of cuprous acetate to cuprous oxide apparently follow rapidly and go essentially to completion. That the f i r s t step in the above sequence is rate determ-ining and that i t is a homogeneous reaction is supported by the f o l -lowing evidence. 1. Kinetically the reaction is fi r s t order with respect to the concentrations of both cupric acetate and molecular hydrogen. This suggests the rate determining step is bimolecular involving one 52. molecule of each of these reactants. 2. The fact that the rate is essentially independent of the concentrations of acetic acid and acetate salts, the only other components of the solution, suggests that cupric acetate and hydrogen are the only reactants involved. The role of water cannot be deduced from the kinetic results, since i t is always present in large and essentially constant excess. It is possible that the solvent does enter in some way into the activation process. 3 . The homogeneous character of the rate controlling step is confirmed by the fact that the rate is independent of the surface areas of stainless steel and cuprous oxide* the only two solids in contact with the solution. 4 . The values of the frequency factor and the entropy of activation, determined from the kinetic results also point to a homo-geneous bimolecular rate determining process. 5. There is no indication that cuprous acetate, which might be formed as an intermediate product in the reaction, has any effect. This is of particular importance since i n some experiments by other workers referred to earlier, i t was found that cuprous acetate in certain organic solvents, has the property of being able to activate molecular hydrogen and act as a homogeneous catalyst for hydrogenation reactions. The suggestion that cuprous acetate might be having a similar effect in the present aqueous system appears to be excluded on the following grounds. 53. (a) » As shown above the kinetic results quantitatively confirm that only cupric acetate and molecular hydrogen are involved in the rate controlling step. (b) „ The concentration of cuprous acetate in the aqueous solutions used, appears to be negligible^ its hydrolysis to cuprous oxide apparently being essentially complete. (c) „ If such small traces of cuprous acetate as might be present in the solution, had any effect on the reaction, such as being responsible to any significant extent for the activation of molecular hydrogen, then the apparent constancy of the first order rate con-stants which was observed, would demand that the concentration of cuprous acetate in the solution remained constant over very wide variations in the pH and in the concentrations of acetic acid and acetate salts. In view of the known sensitivity of cuprous acetate to hydrolysis, this is considered to be most unlikely. (d) . There appears to be l i t t l e relation between the' •present: system arid those organic systems in which cuprous'acetate was found to be catalytically active. The kinetic features of the .^act-ions in the two cases are fundamentally different, and in addition the catalytic activity of cuprous acetate appears to be confined to certain nitrogen containing organic solvents. ^;iSu There would thus appear to be very convincing evidence to support the rate determing step which has been proposed for'this reaction. The kinetic results of course provide no information about the subsequent sequence of steps in the reaction, since these are 54. presumably fast and have no influence on the rate 0 However, the nature of the rate controlling step and other known features of the chemistry of the reactions would appear to indicate a sequence of reaction steps such as has been postulated above in equations 16 to 18. Form of the Reacting Cupric Acetate. In the previous discussion i t was assumed that the cupric acetate was present in the solutions largely as the undissociated CuAc2 molecules and that i t reacted in this form. The evidence for this is fairly strong although largely of an indirect nature. 10 It is known from electrometric and spectrophotometric studies (17)„that the cupric ion forms complexes with acetate ions + • in aqueous solutions. The existence of the complexes CuAc and CuAcg has been definitely established and reference has also been made to the possible existence of the higher acetate complexes CuAc3 and CuAc^ . . Pedersen investigated the equilibria! ++ + Cu + Ac~^f CuAc ( 1 9 ) CuAc+ * Ac" . 5 * . CuAcg ' .......(20) and obtained the following approximate values for the equilibrium constants at room temperatures [puAc+2__ K K, - 56 .......(21) F 7 M 1 55. [cuAc?3 B K2 = 6.7 (22) According to these values, i n a solution of cupric acetate containing about 0.5 moles per l i t r e of sodium or ammonium acetate, the copper would be present predominantly i n the form of the undissociated CuAc 2 molecule, with some CuAc" present as w e l l . This would be the case f o r most of the solutions used i n these experiments. Unfortunately these values apply to room temperature, and no quantitative inform-ation i s available i n the temperature range, 75° to 135°G. i n which the k i n e t i c studies were conducted. However, i t seems reasonable to assume that the s i t u a t i o n w i l l be at least q u a l i t a t i v e l y s i m i l a r under these conditions since the temperature co e f f i c i e n t s of ioni z a t i o n e q u i l i b r i a of t h i s type are generally small. 2. Some features of the k i n e t i c results also support t h i s suggestion and point to the conclusion that the cupric acetate i s present predominantly as the undissociated CuAc2 molecule and that i t reacts i n t h i s form. The fact that the rate of reaction i s e s s e n t i a l l y independ-ent of the acetate ion concentration, over most of the range invest-igated, suggests that i n t h i s range the copper i s present predominant-l y i n the form of one of these possible species i e . Cu + +, CuAc~, or CuAc 2 and that t h i s species i s the one p a r t i c i p i t a t i n g i n the reaction. I f two or more species were present i n comparable concentrations,, then i t would be expected that the equilibrium between them would be shi f t e d s i g n i f i c a n t l y by a change i n the acetate ion concentration, and unless both species happened to be equally reactive, a change i n 56. the rate would result. Similarly i f one of the forms was present in predominant concentration, but the other form, present in.smaller amount was the reactive one, its concentration and consequently the rate of reaction would be sensitive to changes in the acetate ion concentration. 3. The magnitude of the frequency factor of the reaction also suggests that most of the cupric acetate is in the form of the reactive species. If only a small fraction of the copper were in this . form, then the observed frequency factor would become abnormally large for a bimolecular reaction. 4. The fact that a lower rate of reaction was obtained when the acetate ion concentration was very low (i.e. in the absence of added sodium or ammonium acetate) suggests that i t is the undis-sociated form of cupric acetate which is reactive (the CuAc2 molecule). If the reactive species were Cu or CuAc then an increase of rate at low acetate concentrations would be expected. Most of the evidence thus appears to favour the conclusion that the copper reacts as the undissociated CuAc2 molecule, and that i t is present predominantly in this form in the solution. Another conclusion which would also appear to be consistent with a l l the results, is that CuAc2 and CuAc+ are both present in comparable quantities, but that they are identically reactive. In such a case the observed rate would not be affected by changes in their relative concentrations, and would thus also be independent of the acetate ion concentration over a wide range. However, this is considered less likely than the interpretation suggested earlier. 57 It should be noted that there is no indication, either, from the kinetic results or from other studies on cupric acetate solutions, of the existence of a dimer molecule or ion involving two copper atoms. All the evidence appears to rule out this possibility. This is of interest since the only two other instances in which i t appears likely that homogeneous activation of molecular hydrogen oc-curs, both involve dimeric molecules i.e. Cu2Ac2 and Co2(C0)g, each of which contains two metal atoms. It has been suggested (10) that such a dimeric configuration is essential, i f a molecule is to pos-sess the property of being able to activate hydrogen, but the results of the present study would appear to be in contradiction to this suggestion. General Conclusions. It has been shown that cupric acetate reacts homogeneously with molecular hydrogen in aqueous solution. The significance of this result is enhanced by the fact that only one or two other cases of homogeneous hydrogenation reactions are known and this would appear to be the f i r s t instance of such a reaction having been shown to occur in aqueous solution. Although the kinetic results which were obtained in the course of this investigation, suggest the general pattern of the mechanism of the reaction and the sequence of steps involved, the detailed configuration of the activated complex and the manner in which the hydrogen molecule interacts with cupric acetate to become activated, is s t i l l not clearly understood. It is hoped that simi liar 58. investigations on related, systems together with theoretical examin-ation of the problem, will throw further light on these questions i and on the mechanism of hydrogenation catalysis in general. \ In conclusion, reference might be made to some of the possible applications arising from this study. 1. The reaction investigated appears to afford a con-venient method of preparing cuprous oxide, and may have some advantages in this connection over other known methods. 2, The fact that i t has been demonstrated that cupric acetate is capable of activating hydrogen, suggests that i t might • serve as a catalyst for the homogeneous hydrogenation of other mater-ials in aqueous solution. Such a catalyst might have important advantages over heterogeneous catalysts, in permitting higher rates and better control to be achieved. In addition i t might find applic-ation where most solid catalysts are unsuitable because of the presence of poisons in the solution. 59. APPENDIX A. Spectre-photometric Method for Determine Copper in Aqueous Solution. Because of the number of samples taken during the course of each experiment, a relatively simple and rapid method of determining copper was desirable. The method chosen was the spectrophotometric determination of the cuprammine ion formed upon addition of ammonia to cupric salt solutions. This method was particularly suitable in the concentration range of interest ( 0 . 1 to 20 grams per lit e r of Cu). The following procedure was used: To a 5 ml. aliquot of the solution, generally containing from 0 .5 to 100 milligrams of Cu, was added 6 ml. of concentrated NH4.OH, and the total volume adjusted to 100 ml. with distilled water. The optical density of the solution was measured with a Beckman D.U. O Spectrophotometer using a wavelength of 6100 A., which corresponded to the maximum light absorption of the cuprammine complex (see Figure 1 9 ) . The concentration of copper was determined from the calibration graph shown in Figure 2 1 , which was prepared using standard copper samples. The calibration graph shows that the solution obeys Beer's Law within the range of concentrations considered, the optical dens-ity remaining directly proportional to the copper concentration. Figure 20 shows that the optical density is insensitive to the concentration of ammonia over a wide range. Only minor errors are thus introduced by slight variations in this concentration. Where no interfering ions are present, as was the case with the solutions used in these experiments, the method is considered to be accurate to within about * 2 percent. O.br o.; 10 MG. Cu PER 100 ML OF SOLUTION '4000 WAVE LENGTH - A. FIGURE 19. Absorption spectrum of copper, i n ammonia solution. 0.6 0.4 0.2 10 Mg. Cu - O 5 Mg, Cu - O 1 Mg. Cu -O-^ o-_1_ L. -o-JL 2 • 8 10 CONCENTRATED.NH^ OH -:ML. PER 100..ML. OF SOLUTION FIGURE 20. .Effect of ammonia concentration on optical density of cupric salts. 4 61. Mg. of Cu p e r 100 ml. of solution FIGUHS21. Calibration curve for determination of copper after addition of NH 3. e X = 6100 A. 62 APPENDIX B o Summary of Kinetic Results. Exp Initial Cu NHA.AC NaAc HAc PF Hp T Pro-duct No. g./l. m./l. m./l. m./l. m./l. i n i t i a l final atm. °c. min"^ % Cu 1 10 0.157 0 . 5 - 0.18 5 .0 4.5 13.6 125 0.0672 8 8 . 0 2 10 0.157 0.5 - 0.18 5,1 4.6 13.6 125 0,0658 8 8 . 1 8 15 0.236 0.5 - 0,18 5.1 4,6 13.6 125 0.0617 8 8 . 5 9 20 0.315 0.5 - 0.18 5,0 4.5 13.6 125 0.0645 8 8 . 2 10 15 0.236 0.5 0.18 5.1 4,7 13.6 125 0.0612 88.7 l l 4 15 0.236 0.5 0.18 5 ,0 4.6 13,6 125 0.0617 -12 .;5~ 0.079 0.5 0.18 5.1 4.9 13.6 125 0.0631 88.5 13 15 0.236 1.0 0.18 5.1 solut Ion hy drolj /zed i 14 ' 18 15 0.236 0.25 - 0.09 5 . 1 4,3 13.6 125 0.0613 88.4 15 0.236 0.75 - 0,29 5.1 4.7 13.6 125 0.0617 -20 5 0.079 - 1.0 0.49 5,1 4.9 13.6 125 0.0614 88.6 23 15 0.236 0.25 - 1.17 4,1 3 . 8 .13.6 125 0.0691 -27 15 0.236 1- 0.25 1.17 4.0 3 . 8 13.6 125 0.0691 28 15 0.236 0.25 2.35 3,7 (3,6 13.6 125 0.0702 29 15 0.236 - 0.25 2.35 3,7 3.6 13.6 125 O.O696 31' 10 0.151 0.5 - 1,75 5.1 4.7 13.6 125 0.0686 ± Solution contained 20 g./l. 316 stainless steel powder. Baffel mounted on inside walls of the reaction vessel. Theoretical composition of Cu20 is 88.8% Cu and 11.2% 0 . APPENDIX B (Continued) Exp Initial Cu NH^ Ac NaAc HAc PH Hp T k. Pro-duct No. K . / l . m./l. m./l. m./l. m./l. i n i t i a l final atm. °C. irin" 1' % Cu 34 5 0.079 - 0.25 0.31 4.6 4.5 13.6 125 0.0693 36 15 0.236 - 0.25 0.18 4.8 4 . 4 13.6 125 0.0630 37 20 0.315 - 0.25 0.40 4.5 4 . 1 13.6 125 0.0654 41 15 0.236 - 0.25 0 . 4 4 4.5 4 . 1 13.6 125 0.0643 42 15 0.236 0.25 0 . 4 4 4.5 4 . 1 20.4 125 0.0990 43 15 0.236 - 0.25 0 . 4 4 4.5 4 . 1 27.2 125 0.132 44 15 0.236 - 0.25 0 . 4 4 4.5 4 . 1 34.0 125 0.163 45 15 0.236 - 0.25 0 .44 4.5 4 . 1 6.8 125 0.0334 47 15 0.236 - 0.25 0 . 4 4 4.5 4 . 1 13.6 135 0.149 48 15 0.236 - 0.25 0 . 4 4 4. 5 4 . 1 13.6 115 0.0313 49 15 0,236 - 0.25 0 . 4 4 4. 5 4 . 1 13.6 105 0.0161 50 15 0.236 - 0.25 0 . 4 4 4.5 4 . 1 13.6 95 0.00606 53^ 15 O.236 - 0.25 0.29 4.5 4 .1 13.6 125 0.0643 5 5 10 0.157 - 0.25 0.38 4.5 4 . 1 13.6 125 0.0668 56 20 0.315 - 0.25 0 .44 4.5 4 . 1 13.6 125 0.0632 5 9 * 15 0.236 - 0.25 0 . 4 4 4 . 4 4 . 0 13.6 125 0.0603 6 3 1 15 0.236 - 0.25 0 . 4 4 4 . 4 4 . 2 13.6 75 0.00106 64 15 0.236 - - O.24 4 . 4 3 .3 13.6 125 0.0451 ,65 15 O.236 - 0.38 0.53 4.5 4 . 2 13.6 126 0.0736 69 15 0.236 - 0 . 5 O.64 4.5 4.2 13.6 125 0.0629 Volume of solution 2 liters. ft Solution contains 15 g./l. Cu20. 1 Reaction not taken to completion. 64. APPENDIX C. X-ray Analysis of the Cuprous Oxide Product. The X-ray diffraction measurements on the cuprous oxide products were made with a Norelco X-ray unit manufactured by North American Phillips Co. Inc., using a copper source and nickel filter. A powder technique was employed and the photographic patterns were obtained with a 11 cm. Straumanis Camera. Typical powder patterns obtained for two representative products and for a standard sample V S i . i . ' . of cuprous oxide are shown on the next page. A detailed analysis of one of these patterns is given below. The recorded value of the o lattice constant of cuprous oxide is 4<>26 A. This agrees with the values found. 6 5 . TABLE VIII. X-ray Diffraction Measurements on the CupO Product. e d Plane Lattice Parameter degrees I 0 A 18.325 2.44766 (111) 4.240 21.275 2.12103 (200) 4.242 30.850 1.50002 (220) 4.243 36.900 1.28128 (311) 4.248 38.825 1.22733 (400) 4.250 52.050 0.97488 (331) 4.250 53.950 0.95077 (420) 4.252: 54.100 0.95135 (420) 4.255 62.300 0.86820 (422) 4.253 62.575 0.86829 (422) 4.253 69.900 0.81853 (333) 4.253 70.325 0.81837 (333) 4.252 Lattice parameter not corrected for absorption. Figure 22. X-RAY DIFFRACTION PATTERNS Standard C112O Product from Solution Containing Ammonium Acetate Product from Solution Containing Sodium Acetate 67. BIBLIOGRAPHY l o V.N. Ipatief and W. Werchowsky, Ber. Z£, 2078 ( 1909) . 2. K. Marx and H. Behncke, (to I.G. Farbenindustrie), German Patent No. 5,216,828, March 21, 1932. 3. F0A„ Forward and J. Halpern, Canadian Mining and Metallurgical Bulletin, IN PRESS. 4. H.M. James and A.S. Coolidge, J 0 Chem. Phys,, 1, 825 (1933). 5. A.G. Gaydon, 'Dissociation Energies and Spectra of Diatomic Molecules', Chapman Hall, Ltd., London, 1947, p, 78. 6. M. Polanyi, Scientific Journal of the Royal College of Science, 2, 21 (1937). 7. M, Calvin, Cockbain and M„ Polanyi, Trans. Faraday Soc, 3.2, 1436 (1936). M. Calvin, D.D. Eley and M. Polanyi, ibid., 32 , 1443 (1936). 8. 0. Beeck, A. Wheeler and A.E. Smith, Phys. Rev., j>5_, 601 (1939). See also S. Gladstone, K.J. Laidler and H. Eyring, 'Theory of Rate Processes', McGraw-Hill 3ook Co., New York, N.Y., 1941, p. 346. 9. M. Calvin, Trans. Faraday Soc., 3J*» 1181 (1938); J. Am. Chem. Soc, 61, 2230 (1939). 10. S. Weller and G.A. Mills, J. Am. Chem. Soc, 7£, 769 (1953). 68. 11. I. Wender, M. Orchin and H.H. Storch, J. Am. Chem. Soc., 72, 4842 (1950). 12. A.I. Vogel, 'Quantitative Inorganic Analysis", Longmans, Green and Co., Toronto, 1951, p. 518. 13. J.W. Mellor, 'A Comprehensive Treatise on Inorganic and Theoretical Chemistry", Vol,III, Longmans, Green and Co., London, 1928, p. 120, 14. 'Handbook of Chemistry and Physics", Chemical Rubber Publishing Co., Cleveland, 1944, p. 1949, 15. A. Seidell, 'Solubilities of Inorganic and Metal Organic Compounds', Vol. I, D. Van Nostrand Co. Inc., New York, N.Y., 1940, p. 553, 16. S, Gladstone, K.J. Laidler and H, Eyring, 'Theory of Rate Pro-cesses', McGraw-Hill Book Co., New York, N.Y., 1941, P. 198. 17. K.J. Pedersen, Kgl. Danske Videnskabernes Selskab, Mat. - Y.S;,, Medd. XXII, No. 12, 1945. 


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