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The homogeneous catalytic activation of molecular hydrogen by cupric salts in aqueous solution Peters, Ernest 1956

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THE HOMOGENEOUS CATALYTIC ACTIVATION OP MOLECULAR HYDROGEN BY CUPRIC SALTS IN AQUEOUS SOLUTION by  Ernest Peters - - 0O0  -  A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHY in METAL CHEMISTRY - - - oOo - - We accept this thesis as conforming to tbe standard required from candidates for tbe degree of Doctor of Philosophy  Members of the Department of Mining and Metallurgy  T  H  E  U  N  I  V  E  R  S  I  T  Y  O  F  B  R  I  April, 1956  T  I  S  H  C  O  L  U  M  B  I  A  liititiersittr al ^Mttslr Clnhtmbia Faculty of Graduate Studies  P R O G R A M M E OF T H E  Jflnt&l GDral Jzxmxmmtian lax tire Jleitxee al Jtottxa* al  ^ilazayixxf  of  ERNEST PETERS B. A.Sc. (University of British Columbia) 1949 M.A.Sc. (University of British Columbia) 1951 T H U R S D A Y , M A Y 3, 1956, at 3:30 p.m. IN T H E P H Y S I C A L M E T A L L U R G Y B U I L D I N G R O O M 204 COMMITTEE IN CHARGE DEAN H . F . ANGUS,  Chairman  J . HALPERN  C . A. MCDOWELL  F . A . FORWARD  M . KIRSCH  C . S. SAMIS  D . S. SCOTT  L . G. R. CROUCH  S. E . READ  External Examiner: PROFESSOR M . CALVIN  University of California, Berkeley, Cal.  T H E H O M O G E N E O U S C A T A L Y T I C A C T I V A T I O N OF M O L E C U L A R H Y D R O G E N B Y CUPRIC SALTS I N A Q U E O U S SOLUTION ABSTRACT  Hydrogen, which is relatively inert at ordinary temperatures, was found to be activated homogeneously in aqueous solution by dissolved cupric salts, as shown by their catalytic effect on the reactions between H and reducible substrates such as C r 0 , I 0 , and C e . From kinetic studies of the Cu(II)- catalyzed hydrogenation of C r 0 , it was shown that the catalytic activity of C u i s greatly influenced by complex-forming reagents. The catalytic activities of cupric complexes were found to decrease in the order: butryate, propionate > acetate > sulphate > chloride > H 0 (i.e. the uncomplexed C u ion) > glycine, ethylenediamine. 2  =  2  —  7  ++  +  +  3  = =  2  7  ++  2  + +  In all systems that were studied, the reaction was found to be second order kinetically, the rate being proportional to the concentrations of Cu(II) a n d H . 2  The occurrence of an H ion as a product of the initial step of the reaction was postulated to account for the perchloric acid dependence of the rate. The promoting effect of various negative ions, which follows the order of their basicities, was thus explained by assigning to them the role of stabilizing the H ion. The following mechanism was postulated to account for the observed kinetics: ki Cu(II) + H.. CuH -f H + +  +  +  k—i  CuH  +  -f- substrates _ J l _ _ C u ( II) - f products 2  >  CuH , the activated intermediate suggested by this mechanism, also seems on energetic grounds to be the most plausible of the possible intermediates having reasonable classical structures. +  It is suggested that the catalytic activity of C u is related to its electron affinity. A n attempt is made to extend this interpretation to other homogeneous and heterogeneous hydrogenation catalysts. ++  PUBLICATIONS Studies in the Carbonate Leaching of Uranium Ores. II. Kinetics of the Dissolution of Pitchblende: E . Peters and J. Halpern. Transactions of the Canadian Institute of Mining and Metallurgy, 56: 350-4. 1953. Homogeneous Reaction between Molecular Hydrogen and Mercuric Acetate: J. Halpern, G. J. Korinek and E . Peters. Research, 7: s61-2. 1954. A Thermodynamic Study of the Fuming of Zinc from Lead Blast Furnace Slag: R. C. Bell, G. H . Turner and E. Peters. Journal of Metals, 7:472-8. 1955. Kinetics of the Cupric Acetate Catalyzed Hydrogenation of Dichromate in Aqueous Solution:. E . Peters and J. Halpern. Canadian Journal of Chemistry, 33: 356-64. 1955. Mechanism of the Catalytic Activation of Hydrogen by Metal Ions: J. Halpern and E . Peters. Journal of Chemical Physics, 23: 605-6. 1955. Homogeneous Catalytic Activation of Molecular Hydrogen by Cupric Perchlorate: E . Peters and J. Halpern. Journal of Physical Chemistry, 59: 793-6. 1955. Effect of Complexing on the Homogeneous Catalytic Activation of Hydrogen by Cupric Salts: E . Peters and J. Halpern. Canadian Journal of Chemistry, 34: 554-62. 1956. Nature of the Activated Intermediate in the Homogeneous Catalytic Activation of Hydrogen by Cupric Salts: J. Halpern, E . R. Macgregor and E . Peters. Submitted to Journal of Physical Chemistry.  G R A D U A T E STUDIES Field of Study: Metal Chemistry Metallurgical Thermodynamics and Kinetics Theory of Metal Reactions Structure of Metal Compounds Metal Surface Chemistry Theory of Alloys...  C. S. Samis J . Halpern _ J . Halpern H . G. V . Evans J. G. Parr  Other Studies: Advanced Physical Metallurgy Topics in Physical Chemistry.... Chemical Kinetics Colloid Chemistry Industrial Kinetics and Catalysis Chemical Physics Quantum Mechanics: . Magnetic Properties of Metals  _  W. M . Armstrong ..J. G. Hooley and B. A . Dunell W. A . Bryce M . Kirsch D. S. Scott A . J . Dekker 0. Theimer . H . P. Myers  ii ABSTRACT  Hydrogen, which i s relatively inert at ordinary temperatures, was found to be, activated homogerieously i n aqueatis s o l u t i o n try dieeolvedcupric salts, as shown by their catalytic effect on the reactions between and reducible substrates such as  2*2^7 » ^3  »  an  '^  ^  e  "  From kinetic  studies of the Cu(ll)-catalyzed hydrogenation of Cr^O^", i t was shown that 1 1  the catalytic a c t i v i t y of Cu reagents.  i s greatly influenced by complex-forming  The catalytic a c t i v i t i e s of cupric complexes were found to  decrease i n the order: chloride ^  HgO  butyrate, propionate  ( i . e . the uncomplexed C u  + +  ^> acetate ion)  ^>  ^>  sulphate  ^>  glycine , ethylene-  diamine . In a l l systems that were studied, the reaction was.found to be second order k i n e t i c a l l y , the rate being proportional to the of Cu(ll) and  concentrations  Hg.  The occurrence of an H  +  ion as a product of the i n i t i a l step of  the reaction was postulated to account for the perchloric acid dependence of the rate.  The promoting effect of various negative ions, which follows,;the  order of their b a s i c i t i e s , was thus explained by assigning to them the role of s t a b i l i z i n g the H  +  ion.  The following mechanism was postulated to  account for the observed kinetics: k Cu(ll) + Hg —^>  CuH + H +  XT  kg CuH + substrates > Cu(ll) + products CuH , the activated intermediate suggested by this mechanism, also +  +  seems on energetic grounds to be the most plausible of the possible i n t e r -  "  mediates having reasonable classical structures. It i s suggested that the catalytic a c t i v i t y of Cu i t s electron a f f i n i t y .  i s related to  An attempt i s made to extend this interpretation to  other homogeneous and heterogeneous hydrogenation catalysts.  iv ACKNOWLEDGEMENT  The author i s grateful for the assistance, advice, and encouragement given by members of the Department of Mining.and Metallurgy, and especially to Dr. J . Halpern for h i s inspiring direction of the research reported i n this thesis, and for his constructive criticism of the manuscript during i t s preparation. The author also wishes to thank the National Research Council of Canada for financial support of the research and.for a Studentship held by the author i n 1954/55•  Thanks are also due to the Consolidated Mining and  Smelting Company f o r the Cominco Fellowship, held i n 1953/54.  TABLE OF CONTENTS Page INTRODUCTION  «  o  ©  o  o  o  o  o  o  o  o  o  o  o  o  »  o  o  o  o  ©  c  .  o  o  ©  «  HYDROMETALLURGICAL APPLICATIONS OP HYDROGENATION REACTIONS  I  «  ..  1  HYDROGENATION CATALYSIS . . . . . . . . . . . . . . . . . . . .  2  HOMOGENEOUS ACTIVATION OF HYDROGEN IN SOLUTION  4  . . .  HOMOGENEOUS HYDROGENATION REACTIONS IN AQUEOUS SOLUTIONS EXPERIMENTAL  .  .  .  .  .  .  o  .  .  o  o  .  .  .  .  .  .  .  .  .  .  .  .  .  . . . .  .  6 9  .  RESULTS AND DISCUSSION . . . . . . . . . . . . . . . .  11  A. HOMOGENEOUS CATALYTIC ACTIVATION OF MOLECULAR HYDROGEN BY CUPRIC PERCHLORATE . . . . . . . . . . . .  11  B. HOMOGENEOUS CATALYTIC ACTIVATION OF MOLECULAR HYDROGEN BY CUPRIC ACETATE . . . . . . . . . . . . . . .  25  C. EFFECT OP COMPLEXING ON THE HOMOGENEOUS ACTIVATION OP MOLECULAR HYDROGEN BY CUPRIC SALTS . . . . . . . . Cupric Carboxylate Complexes  37  .  37  Cupric Sulphate . . . . . . . . . . . . . .  42  Cupric Chloride Complexes . . . . » . . . . ,  46  Cupric Bthylenediamine Complexes  48  Cupric Glycinate Complexes  . . . . .  . . . . . . . . .  Summary of Effects of Complexing  .  . . . . . . .  51 55  D. HOMOGENEOUS ACTIVATION OF HYDROGEN BY SALTS OP OTHER METALS E. MECHANISM OF HYDROGEN: ACTIVATION  56 5  8  CONCLUSIONS  . . . . . . . . . . . . . . . . . . . .  63  REFERENCES  . . . . . . . . . . . . . . . . . . . . . . . . . . .  65  APPENDIX A  . . . . . . . . . . . . . . . . . .  69  APPEND TX B  o  o  o  o  o  o  o  o  o  o  o  o  o  o  o  o  o  o  o  »  o  »  »  «  »  #  #  71  TABLES Page I II III IV  RATES OP REACTION BETWEEN HYDROGEN AND DICHROMATE IN CUPRIC PERCHLORATE SOLUTIONS . . . . . . . . . . . . . . . . . . . .  22  RATES OP REACTION BETWEEN HYDROGEN AND DICHROMATE IN CUPRIC ACETATE SOLUTIONS . . . . . . . . . . . . . . . . . . . . . .  32  RATES OP REACTION BETWEEN HYDROGEN AND DICHROMATE IN SOLUTIONS OF VARIOUS CUPRIC CARBOXYLATE SALTS . . . . . . . .  38  RELATIVE CATALYTIC ACTIVITIES OF VARIOUS CUPRIC COMPLEXES  57  . .  vii LIST OF FIGURES Pig, No.  Page  1.  Typical Rate Plots for Cupric Perchlorate . . . . . . . . .  12  2.  Comparison of Rate Plots for Various Substrates  14  3.  Dependence of Rate on Cupric Perchlorate Concentration  4.  Rate Curves for Cupric Perchlorate at Different Hydrogen Partial Pressures . . . . . . . . . . . . . . . .  5.  . .  15  16  Dependence of the Cupric Perchlorate Rate on the Hydrogen Partial Pressures . . . . . . . . . . . . . .  17  6.  Solubility of Hydrogen i n Water as a Function of Temperature  19  7.  Log (k/T) vs. l/T for the Cupric Perchlorate-Catalyzed *  20  8.  Dependence of the Rate on the Perchloric Acid Concentration  24  9.  Typical Rate Curves for Cupric Acetate  26  10.  Rate Curves for Different Cupric Acetate Concentrations . .  27  11. 12.  Dependence of Rate on Cupric Acetate Concentration . . . . Dependence of Cupric Acetate Rate on the Hydrogen Partial Pressure . . . . . . . . . . . . . . . « . . . . • • • • •  28 29  13.  Log (k/T) vs. l/T for the Cupric Acetate-Catalyzed Reaction  31  14.  Dependence of the Cupric Acetate Rate on pH .  33  15.  Concentrations of Cupric Acetate Complexes as a-Function of pH  33  16.  Effect of Various Salts on the Cupric Acetate Rate  35  17.  Rate Plots showing Consecutive Reduction of Dichromate and  R@£LCt»10Il  o  o  Cupric Acetate  o  c  o  o  o  o  o  o  o  o  o  o  o  o  o  o  e  o  v  t  c  *  ....  . . . . . . . . . . . . . . . . . . . . . .  36  18.  Dependence of Rate on Cupric Propionate Concentration . . .  40  19.  Log (k/T) vs. l/T f o r Cupric Propionate and Cupric Sulphate  41  20.  Typical Rate Plots for Various Cupric Complexes .  43  21.  Dependence of Rate on Sulphate Concentration  44  22.  Dependence of the Cupric Sulphate Rate on Cu(ll) Concentration  45  viii Fig. No.  Page  23.  Dependence of the Rate on Chloride Concentration  24.  Dependence of the Cupric Chloride Rate on  25*  Effect of pH on the Cupric Ethylenediamine Rate  26.  Dependence of the Cupric Ethylenediamine Rate on the Concentration . . . . . . . . . . . . . . . .  27.  Effect of pH on the Cupric Glycinate Rate  28.  Dependence of the Cupric Glycinate Rate on Glycine Concentration . . . . . . . . . . . . . . . .  54  29.  Schematic Potential Energy Diagram . . . . .  60  30.  Absorption Spectra of Sodium Dichromate and Chromic Acetate  SolU/fc l O I l S  o  o  e  o  o  o  o  o  o  o  »  o  e  o  o  o  Cu(ll)  47 Concentration  50  Cu(ll)  . .  52 53  .  o  49  o  *  A  *  »  «  *  e  TO  THE HOMOGENEOUS CATALYTIC ACTIVATION OF MOLECULAR HYDROGEN BY CUPRIC SALTS IN AQUEOUS SOLUTION  INTRODUCTION  HYDROMETALLURGICAL APPLICATIONS OF HYDROGENATION REACTIONS The use of hydrogen as a metallurgical reducing agent i n the preparation of metals from their compounds has been appreciated for many years, but commercial interest i n this application i s of comparatively recent origin. This interest has resulted largely from the increased use of hydrometallurgical "leaching" processes i n the treatment of certain low grade ores of copper, nickel, cobalt, uranium, and vanadium, which y i e l d solutions from which the metals or their lower oxides can be conveniently displaced by hydrogen.  From thermo-  dynamic considerations of the metal precipitation reactions, i t i s apparent that the application of this procedure i s essentially limited to those metals which l i e below or immediately above hydrogen i n the electromotive series.  Some  measure of favourable displacement of the reduction equilibrium can usually be achieved by increasing the partial pressure of hydrogen or the temperature. In addition the formation of an insoluble compound (such as an oxide) of the metal i n a lower valence state, may sometimes be u t i l i z e d to advantage. From a kinetic standpoint, molecular hydrogen i s r e l a t i v e l y inert, and usually requires the presence of a catalyst to react at temperatures much below 600°C.  The most active catalysts are the transition metals including  Ni, Pt, and Pd, and certain metallic oxides such as Gr^)y their mixtures ( l ) .  ZnO, CuO, and  Early attempts by Ipatieff and h i s co-workers (2) to d i s -  place metals from solution at temperatures up to 300°C and with hydrogen pressures up to 600 atm. therefore met with varying degrees of success,  -2depending on the catalytic a c t i v i t y of the dissolved metal salts, the precipitated products, or the walls of the container.  In the course of these i n -  vestigations, Ipatieff succeeded i n precipitating Cu, B i , Pb, Co, Ni, and Fe by hydrogen, either i n metallic form or as oxides of lower valence, from aqueous solutions of their salts.  In recent years a number of processes have  been developed which u t i l i z e hydrogenation reactions of this type to y i e l d metallic Ni and Co as well as the oxides DOg and VgO^ from alkaline leach solutions (3°4i>5)° involved (6).  In each case a heterogeneous hydrogenation reaction i s  In the precipitation of the oxides an extraneous c a t a l y s t  9  such  as metallic Ni, must be present to catalyze the reduction of the dissolved salt (4,5)o HYDROGENATION CATALYSIS Both from the standpoint of applications such as these, and as a matter of fundamental interest, great importance attaches to the understanding of the mechanism of hydrogenation catalysis.  In general, the detailed mech-  anism by which hydrogenation catalysts function i s not f u l l y understood, i n spite of their widespread application and the extensive study to which they have been subjected.  The low reactivity of molecular hydrogen has been a t t r i -  buted to i t s high dissociation energy (103 kilocalories (7,8)) coupled with i t s closed shell ground state electronic configuration.  Thus i n the homogeneous  gas phase reactions of hydrogen with other electronically saturated molecules, a substantial portion of the dissociation energy must usually be expended i n the activation process, as evidenced by the suggested value (9) of 61 k i l o calories for the activation energy of the H^ - CO reaction, by the high ignition temperature  (580°C.) of hydrogen-oxygen mixtures, and by the immeasurably low  rate of the H^ - CO^ reaction at temperatures below 600°C. (10). The heterogeneous activation of molecular hydrogen, which i s believed to enter into the catalysis of hydrogenation reactions, i s most commonly ex-  -3plained i n terms of a mechanism that involves homolytic s p l i t t i n g of the molecule on the catalyst surface with simultaneous formation of covalent bonds between the H atoms and two adjacent surface atoms ( l l ) .  Application of the  absolute reaction rate theory to this model has shown that the activation energy of the adsorption process should depend on the distance between adjacent surface sites ( l l , 1 2 j .  This interpretation implies a correlation  between catalytic a c t i v i t y and the l a t t i c e parameter of the catalyst.  Evidence  for such a correlation has been observed experimentally for several hydrogenation reactions  (13).  An alternative approach to the interpretation of hydrogenation catalysis relates the a c t i v i t y of solid catalysts to their electronic Thus Couper and Sley (14)  properties.  have proposed that d.-band vacancies ( i . e . unpaired  _d electrons) are essential for the c a t a l y t i c a c t i v i t y of metals.  This i s help-  f u l i n explaining the high catalytic a c t i v i t i e s usually observed for the transition metals, and the decrease i n catalytic a c t i v i t y which occurs when these metals are alloyed with  elements such as Cu, rig, Pb, B i , etc., whose  valence electrons enter into and f i l l the _d-band of the transition metal  (l5)»  Thus the rate of parahydrogen conversion on Pd-Au alloys decreases markedly when the Au content reaches 60$, f i l l e d (14).  the concentration at which the d-band i s just  Similar observations have been made for the hydrogenation of  styrene and for the decomposition of formic acid and methanol on various alloys (16). These results have been interpreted to imply the formation of covalent bonds between the metal catalyst and the adsorbed reactants,  or,  alternately, the entry of electrons from the l a t t e r into the u n f i l l e d electronic levels or bands of the catalyst.  The l a t t e r view has received particular  support from kinetic studies by Schwab (17) on the dehydrogenation of formic acid on s i l v e r alloys, where the activation energy was  observed to increase  -4linearly with the square of the electron concentration of the Hume-Rothery a phase.  This has been interpreted to imply that catalytic activation  involves  the entry of two electrons from the formic acid molecule into the conduction band of the metal.  Among the intermetallic Hume-Rothery phases, the y-phase  (in which the conduction band i s nearly f u l l ) showed a maximum activation energy. Similar considerations have been applied to the interpretation of the catalytic a c t i v i t y of metal oxides, the most active of which are defect semi-conductors such as ZnGrgO^, which can readily accept electrons  (18).  Dowden (19,20) has attempted to provide a theoretical basis for the electronic interpretation of the catalytic a c t i v i t y of solids, and has shown how  the catalytic a c t i v i t y may be related to the electronic work function,  the energy density of electron levels at the Fermi surface [g ( E ) _ _ ] , e  the gradient of the l a t t e r [d g(E) / dEJg _ _Q = G.  g  0  P  and  The energy density i s a  maximum when a d_-band i s p a r t i a l l y f i l l e d , and the gradient, G, depends on the extent of f i l l i n g of the whole B r i l l o u i n zone.  HOMOGENEOUS ACTIVATION OF HYDROGEN IN SOLUTION By virtue of their much greater simplicity, the study of homogeneous hydrogenation catalysts should provide results which are more amenable to detailed interpretation and capable of discriminating more c r i t i c a l l y between alternative theories, than i n the case of solid catalysts. The f i r s t clear-cut example of the homogeneous activation of hydrogen in solution, observed by Calvin i n 1938, and cupric salicylaldehyde, salts (21, 22).  was the hydrogenation of cupric acetate  dissolved i n quinoline, to the corresponding cuprous  The autocatalytic nature of the reactions suggested that the  cuprous salt being produced was the effective catalyst, a view which was  con-  firmed when i t was demonstrated that other substrates such as quinone could also be hydrogenated homogeneously i n the presence of dissolved cuprous acetate.  The kinetics of this reaction were found to be f i r s t order i n dissolved hydrogen and between f i r s t and second order i n the concentration  of the cuprous s a l t .  The l a t t e r result was originally interpreted as indicating that a dimer of cuprous acetate was the effective catalyst responsible for activating the hydrogen. . Subsequent detailed studies of this system (23)24,25,26,27) substantiated the essential results and conclusions of the original work, demonstrating conclusively that hydrogen i s activated homogeneously.  However, i n  the light of recent work (27) i t appears l i k e l y that the rate determining step, i n which hydrogen i s activated, involves the interaction of a hydrogen molecule with two cuprous acetate molecules (rather than with a dimer).  An  intermediate, hydrogen-carrying complex i s thus formed which reacts rapidly with the reducible substrate, or which leads to parahydrogen conversion or hydrogen-deuterium exchange, but which decomposes only slowly to metallic copper. Another system i n which hydrogen i s activated homogeneously and which has attracted much interest i n recent years i s the "0X0"  or hydroformylation  synthesis (28,29), i n which an o l e f i n reacts with hydrogen and carbon monoxide in the presence of metallic cobalt.  Simple hydrogenation reactions have also  been observed under these conditions (30,31,32).  Studies i n this system (32)  have shown that Hg i s apparently activated by reaction with dicobalt octacarbonyl, i . e . Co (C0) + H 2  8  2  —2HCo(C0)  4  •-  [l]  while the observed products result from the subsequent reaction of HCo(CO)^ with the o l e f i n .  These reactions are homogeneous and proceed i n a variety  of organic solvents. A number of further examples of homogeneous hydrogenation reactions i n organic liquids, which have been reported recently, involve the activation of hydrogen by silver salts (53*34) and by ethylene platinous chloride  (35).  HOMOGENEOUS HYDROGENATION REACTIONS IN AQUEOUS SOLUTIONS The f i r s t established homogeneous hydrogenation reaction i n aqueous solution  i s the reduction by hydrogen of cupric acetate to cuprous oxide  reported i n 1953 by Dakers and Halpern (3°»37,38).  9  This reaction was  originally observed by Ipatieff and his co-workers (2) i n their early metal displacement experiments.  However, i t s homogeneous nature was not noted;,  although i t was known to proceed rapidly under r e l a t i v e l y mild conditions without the addition of any extraneous catalyst. By kinetic studies, Dakers and Halpern showed this reaction to be f i r s t order i n the concentrations of both dissolved hydrogen and of cupric acetate, and independent of solid surfaces i n contact with the solution. The reaction was postulated to proceed through the following mechanism, the f i r s t step i n the sequence being rate-determining: k CuAc + H >~ CuAc «H slow 2  g  2  CuAc °H + CuAc  2  + HO  Cu 0 + 4HAc  [2]  - - [3]  fast  d  The rate-controlling step i s a bimolecular reaction between a cupric acetate molecule and a hydrogen molecule to form an intermediate complex whose subsequent fast reaction with cupric acetate and water yields the observed product, CUgO.  Of special interest i s the conclusion that only one  metal atom i s involved i n the hydrogen activation process. The reaction sequence implies that cupric acetate should also be capable of functioning as a homogeneous catalyst f o r reactions with hydrogen of other substrates which are thermodynamically  more readily reducible than  cupric acetate i t s e l f , but which do not react with hydrogen i n the absence of a catalyst for kinetic reasons.  By analogy with early observations on  cuprous acetate, such a substrate might be expected to react preferentially with the complex formed i n the f i r s t step (equation [2])  to regenerate the  cupric acetate molecule, i . e . CuACgoHg + 2A  CuAc + 2B + 2H*  [4]  2  where A represents an equivalent of the substrate and.B; an equivalent of i t s reduction product. Confirming this reasoning, i t was demonstrated ..early i n the research described i n t h i s thesis that substrates such as CrgCy, I0^~, C.e ' l  11  a n d  were reduced by Hg i n the presence of cupric acetate according to this  mechanism.  The reduction of CTgOy" (because of the low concentration which  could be accurately determined) proved a particularly convenient reaction f o r studying the kinetics of the hydrogen activation process, and permitted the accurate determination, of very low rates of activation. The object of the research on which this thesis i s based, was to investigate the kinetics of the homogeneous activation of hydrogen by a number of cupric salts.  In particular, i t was hoped to achieve some i n -  formation on the following points: 1.  Whether other cupric salts share with cupric acetate the a b i l i t y  to activate hydrogen homogeneously. 2.  How the catalytic a c t i v i t i e s of different cupric salts and  com~  plexes are related to the structure and configuration of the catalyst species. 3.  The detailed mechanism of the hydrogen activation process i n these  systems, and i t s relation to the mechanism of heterogeneous hydrogenation catalysis. Since the work described i n this thesis was undertaken, a number of other homogeneous hydrogenation reactions have come to l i g h t , partly as a result of this work.  In particular, mercuric and mercurous salts have  -8been shown to activate hydrogen homogeneously (59,40,41), the kinetics of the activating process being analogous to that observed f o r cupric s a l t s  9  i . e . - f i r s t order i n the concentration of hydrogen and f i r s t order i n the concentration of Hg Ag  +  and Hgg  •  On the other hand, the kinetics of the  catalyzed hydrogenation of CrgO^~ i n aqueous solution were found to be  f i r s t order i n hydrogen and second order i n the.concentration of Ag  +  (42,45)»  The homogeneous reduction of MnO^~* by hydrogen i n acid solutions i s f i r s t order i n MnO^  and Hg, but i n the presence of Ag , a kinetic contribution +  of the form k[H ][Ag ][Mn0 ~] was also noted +  o  (45).  -9-  BXPEEIMENTAL 1 • \  Glycine and ethylenediamine were of Eastman white label grade; other chemicals were of reagent grade.  all  Hydrogen, and nitrogen gas were supplied  i n cylinders by the Canadian Liquid A i r Company, and were used without further purification.  D i s t i l l e d water was used i n the preparation of a l l solutions.  The reaction studies were conducted i n a one-gallon stainless steel autoclave manufactured by Autoclave Engineers, Inc.  The autoclave was  equipped with a motor-driven propellor type s t i r r e r mounted on a v e r t i c a l shaft passing through a pressure gland.  Auxiliary f i t t i n g s included a thermo-  couple well, a sampling tube, and gas connections through the l i d .  For  solutions corrosive to stainless steel, a titanium l i n e r and f i t t i n g s were used.  The autoclave was heated with a gas ring burner and the temperature  was controlled to - 0.3 C. by a Leeds and Northrop Micromax controlling B  recorder. The experimental procedure comprised flushing the autoclave with nitrogen and heating to temperature.  When the desired temperature was reached,  a number of samples were taken while the solution (usually 3 l i t e r s ) remained under steam or under a slight partial pressure of nitrogen. sired pressure of hydrogen was introduced. to follow the course of the reaction.  Then the de-  Samples were taken periodically  When required, a sintered stainless  steel f i l t e r was placed on the end of the sampling, tube so that the samples of solution removed for analysis were free of precipitated products. Samples of solution were analyzed either f o r the reactant or f o r the products of reaction. reduction of Ox^i^",  Most of the studies were concerned with the  the concentration of the l a t t e r usually being determined  spectrophotometrically at i t s 350 m IX absorption peak with a Beckman DU  -10* Spectrophotometer after suitable d i l u t i o n .  Copper was determined either  spectrophotometrically  (as the ammine complex at 610 mjn (36)) or electro-  lytically.  or Ce  When 10 ^  were measured volumetrically.  were used as substrates, their concentrations A known amount of ferrous sulphate was added  to the sample and the excess was back t i t r a t e d with a standard eerie sulphate or potassium permanganate solution.  The same procedure was applied to deter-  mine Cr 0^~ i n chloride solutions, as the chloride complexes of copper, which 2  absorb light at 350 mju,, interfered with the spectrophotometric method. pH measurements were made with a Beckman Model H-2 A.C. pH meter.  See Appendix A f o r details of absorption spectra of solutions.  RESULTS AND DISCUSSION  A.  CATALYTIC ACTIVATION; OF MOLECULAR HYDROGEN BY CUPRIC PERCHLORATE Among the cupric salts that were examined for homogeneous catalytic  a c t i v i t y i n hydrogenation reactions was cupric perchlorate, which i s generally considered to be completely dissociated when dissolved i n aqueous perchloric acid solution.  Some typical rate plots for the reduction of dichromate by  hydrogen i n the presence of cupric perchlorate are shown i n F i g . 1.  At  constant temperature and hydrogen pressure, the concentration of Cr^O^" i s seen to decrease l i n e a r l y with time, corresponding to zero order kinetic behaviour.  The slopes of the zero order plots are independent of the  i n i t i a l Cr 0^~ concentration.  The stoichiometry of the reaction i s  2  represented  by Cr 0 2  = 7  • 3H  + 8H  2  Other substrates, including 10^  ^  +  and Ce  2Cr  + + +  + 7H 0  [5]  2  , were also found to react with  hydrogen i n the presence of cupric perchlorate, as follows: I 0 ~ + 2{  + H  3  i.H  C e ^ *  >  +  — ^  2  £,I Ce  + + + +  2  H  + 3H 0 2  +  [6] [7]  or i n general Ox + n H  2  + q H  +  >•  r Red. + s HgO  [8]  In each case the kinetics of the reduction were zero order with respect to the substrate.  No change i n the concentration of copper could  be detected during the reactions. The reaction rates were estimated from the slopes of rate plots such as those shown i n F i g . 1,  after introducing corrections f o r contributions due  to side reactions not involving hydrogen.  This contribution, which i s  F i g . 1. Typical- Rate plots f o r the Cupric Ferchiorate Catalyzed Reduction of Dichromate by Hydrogen. 0.10 M./L. Cu(C10i,.) ; I 1 0 ° C ; 20 atm. H . 2  2  probably due to corrosion of the reaction vessel, was estimated from the rate of reduction of C r 0 ^ while the reaction mixture was under nitrogen.  This  -  2  correction rarely exceeded 5$ of the rate of reaction with hydrogen.  Duplicate  rate measurements generally reproduced to within - 5$ i n the case of CrgO^ and 10, 5  and to within - 10% with Ce  .  -  The rates for different substrates  can best be compared by expressing them i n terms of the equivalent rates of consumption of hydrogen* calculated from the stoichiometric r e l a t i o n of equation [8], i . e . -d[H J/dt 2  =  -nd[0x],/dt  [9]  Equivalent rate plots for the reaction of hydrogen with various substrates, i . e . Cr 0^"*, 10^ , and Ce 2  Cu  are shown i n F i g . 2.  At constant  concentration, the slopes of these plots are seen to be essentially  independent of the nature of the substrate, within the l i m i t s of experimental uncertainties involved i n their comparison.  This observation, coupled with  the zero order reaction kinetics, strongly suggests that the substrate does not participate i n the rate-determining  species  step, and that t h i s step i s the  same for a l l the substrates. Further indication of this i s provided by F i g . 3 i i n which the rates are plotted as functions of the cupric perchlorate concentration, and points for CrgOy", 10^ , and Ce line.  the  are a l l seen to l i e on a single straight  The fact that -dtHgj/dt i s d i r e c t l y proportional to [ C u ] i s con++  sistent with the suggested catalytic role of the l a t t e r . The reaction of Cr 0^~ proved to be the most convenient one 2  on  experimental grounds, and to y i e l d the most precise measurements of the rate of activation of hydrogen. remaining kinetic variables. rate of this reaction.  It was therefore selected to elucidate the Figs. 4 and 5 depict the effect of !,, on the  In the range from 0 to 25 atm.,  -d[H ]/dt i s seen to 0  -14-  i.o h  Figure 2.- Comparison of the Reaction Rates f o r Various Substrates. 110°C; 20 Atm. H„.  Fig. 3. Dependence of the Reaction Rate on Cupric Perchlorate Concentration; 110°C; , 20 atm. H . 2  ^16-  TIME  -  MINUTES  F i g . 4. Reduction of Dichromate i n the Presence of Cupric Perchlorate at d i f f e r e n t Hydrogen P a r t i a l Pressures. 0.10 M./L. C u ( G 1 0 O ; iio°c. 2  -17-  Fig. 5. Dependence of the Reaction Rate on the Hydrogen Partial Pressure. 0.10 M./L. CuCClO^a; 110 C. C  be directly proportional to the hydrogen partial pressure.  The residual rate  of disappearance of CrgO^"" i n the absence of hydrogen, emdent i n Fig. 4, was subtracted from the slopes of the other rate plots to evaluate -d[Hg]/dt. The rate of reaction can thus be represented by -d[H ]/dt = k'tCu^lPg 2  =  kCCu^Klg]  [10]  where [Hg] i s the concentration of hydrogen i n solution, since the system obeys Henry's law (44,45) i n the pressure range under consideration. k' and k are related by Henry's constant, a, denoting the s o l u b i l i t y of Hg, I060  k*  =  a k  [ll]  The s o l u b i l i t y of hydrogen i n water varies with temperature,, approximately as shown by the plot i n F i g . 6, which i s based on the data of Wiebe and Gaddy (44) and Pray, Schweichert, and JUinnich (45).  Values of  a used i n the calculation of k were taken from this plot, and are l i s t e d i n the legend of F i g . 6. At 100°C, the experimental value of k f o r the cupric perchlorate-  —1  —5 catalyzed reaction between hydrogen and dichromate i s 4.0 x 10  l i t e r mole  -1 sec.  . Rate measurements were made at seven different temperatures, ranging  from 80 to 140«C.  The slope of the resulting plot of log k/T vs. l/fr, ±  shown i n F i g . 7, corresponds to an apparent enthalpy of activation, AH , of + 24.6 - 1.5  * kilocalories per mole.  The rate constant for the cupric perchlorate catalyzed reaction This large uncertainty i n the enthalpy of activation i s primarily due to an i s thus given by uncertainty i n the variation of the hydrogen s o l u b i l i t y with temperature. See F i g . 6.  -19-  Fig. 6 . The Solubility of Hydrogen i n Water as a Function of Temperature. Based on Published Data ( 4 5 , 4 6 ) .  Fig. 7 . Plot of Log(k/T) vs. 1/T for the Cupric Perchloric Catalyzed Reaction between Hydrogen and Dichromate. 0*10 M./L. Cu (ClOiJ;,; 20 atm. H . 2  rnp 24,600 k = ^ - ' e * R T  ' e  _ UJI R  *  [l2J  +  The entropy of activation, AS^ , i . e . -13«1 - 4.0 e.u., i s normal for a simple bimolecular reaction i n solution (46,47). In calculating the values of k, i t was assumed that the solutions (which contained 0.1 M./L. Gu(G10 ) , 0.01 M./L. HC10 , and 0.0005 M./L. 4  2  4  NagCrgO^) were s u f f i c i e n t l y dilute so that a had the same value as that for pure water, i . e . that given i n F i g . 6. In Table I are l i s t e d the rates of solutions of differing salt concentration.  GT^OJ"  reduction i n several  The rate of reaction i s seen to  be essentially independent of NaClO^ concentration from 0 to 1.0 M./L. Additions of 0.38 M./L. Ca(C10 ) or ZnCciO^g were also without effect 4  2  while 0.25 M./L. Al{C10^)^ depressed the rate s l i g h t l y .  The general ionic  strength effect, i f any, i s apparently small, consistent with the fact that one of the reactants, i . e . , hydrogen, i s uncharged.  The absence of a  specific perchlorate salt effect supports the conclusion that the catalytic a c t i v i t y , observed i n this system, i s due to the uncomplexed cupric ion. Direct evidence for the homogeneous character of the reaction i s provided by the results of an experiment (W-3, Table l), i n which 30 g. of stainless steel powder, of similar composition to the autoclave i t s e l f , , was added.  Although the area of this powder was at least twice that of the  autoclave surface normally exposed to the solution, no increase i n the reaction rate could be detected. Varying the s t i r r i n g efficiency, either by reducing the impeller velocity to 600 r.p.m. from its.normal value of 900 r.p.m. (expt. W-2) or by introducing a baffle into the autoclave (expt. W-l) also had no effect on the rate.  This indicates, as do other features of the kinetics, that under the  -22-  TABLE I  RATES OF REACTION BETWEEN HYDROGEN AND DICHROMATE IN CUPRIC PERCHLORATE SOLUTIONS OF VARYING COMPOSITION  0.1 M./L. Cu(C10 ) J 4  2  0.010 M./L. HC10 ; 4  0.0005 M./L. N a ^ C r ^ *  20 atm. H ; 2  110»C.  Expt. No.  Salt Concentration M./L.  T-10  0.00  T-13  -dtH^/dt x IO  k x IO  7  Mole l i t e r " Sec." 1  1  L. Mole"  1.55  9.5  0.25 NaC10  1.55  9.5  T-14  0.50  "  1.50  9.2  T-15  0.75  "  1.53  9.3  T-12  1.00  "  1.48  9.0  T-16  0.38 Ca(C10 )  2  1.50  9.2  T-17  0.38 Zn(C10 )  2  1.53  9.3  T-18  0.25 A l ( C 1 0 )  5  1.39  8.5  W-l  a  0.00  1.56  9.5  W-2  b  0.00  1.55  9.3  0.00  1.49  9.1  w-3  4  4  C  4  4  Baffle added to increase s t i r r i n g efficiency. S t i r r e r velocity decreased from 900 to 600 R.P.M. 3 0 g. of stainless steel powder added.  1  5  Sec."  1  -23conditions of these experiments, the reaction was not limited by the rate of solution of hydrogen. The dependence of the reaction on perchloric acid concentration i s shown i n F i g . 8.  The observed decrease i n the rate of reduction of dichromate  with increasing acid concentration i s believed to be due to a competing r e action involving activated hydrogen but not dichromate. competing reaction was not d e f i n i t e l y established.  The nature of this  The absence of any  chloride i n the reduced solutions suggests that there i s no reduction of perchlorate.  The linear r e l a t i o n between l/k and HCIO^ concentration, shown  i n F i g . 8, can be interpreted i n terms of a competing reaction, involving which i s essentially a reversal of the rate-controlling step, i . e . k H + Cu ^ CuE + H + +  1  +  +  H, +  [13]  -1 The simplest mechanism for the subsequent reaction of CuH* consistent with the observed zero, order kinetics with respect to dichromate,  + involves the formation of Cu  as a further intermediate, through the following  sequence of steps:  , K g  CuH + C u +  >•  + +  2Cu + H +  .[14]  +  followed by _ + C r 0 ~ + 14H  6Cu  2  fast *~  7  2Cr  w  + 6Cu  + 7HgO [15 ]  Application of the steady state treatment (-d[CuH ]/dt = 0) to the above +  sequence leads to the following rate law:  -d[Hj/dt= 2  — [Cu^J  1  +  '  —  [16]  (k^/kgHH*]  This implies, i n agreement with the results which have been reported,  first  order kinetics with respect to [Cu**] and [Hg] at low H* concentrations,  as  well as a linear dependence of the reciprocal of the rate on [H*] at constant  F i g . 8. Dependence of the Reaction Rate on the Perchloric Acid Concentration. 0.10 M./L. Cu(C10i ) ; 20 atm. H / f  2  2  -25* [Cu*"*"] and [Hg].  The further suggestion that at high H  kinetic dependence on [Cu received experimental  +  concentrations the  ] would approach second order, has also recently  support  (48).  The intermediate proposed here, CuH*, i s similar to that which has been postulated for the homogeneous activation of hydrogen by cuprous acetate i n quinoline (21,22,24,27).  In this system the kinetics demonstrate that two  cuprous ions are involved i n the rate-determining step which has been depicted as (Cu ),, + H  ^  1  or  2CU  1  + H  2  >-  2  <Sto\)  [17]  2  2Cu 'H  [18]  I  Further evidence i n support of the above mechanism w i l l be  presented  later.  B.  CATALYTIC ACTIVATION OF MOLECULAR HYDROGEN BY CUPRIC ACETATE A similar kinetic study was made of the reduction of dichromate by  hydrogen i n the presence of dissolved cupric acetate, which, unlike cupric perchlorate, i s known to be largely undissooiated i n aqueous solution (49). This was of particular interest i n view of the e a r l i e r studies (36,37,58) on the kinetics of the hydrogenation of cupric acetate i t s e l f .  As i n the case  of cupric perchlorate, the reaction was found to be zero order k i n e t i c a l l y i n CrgO^ (Fig. 9), f i r s t order i n Hg (Fig. 10) and nearly f i r s t order i n -  CuACg (Figs. 11 and 12).  Thus the reaction rate i s given by the expression  - 5d[Cr 0 ]/dt =  2  7  = -d[Hg]/dt = ktCuACgHHg)  [19]  The relation between the reaction rate and the cupric acetate concentration i s seen to be not quite linear (Fig. l l ) , k showing a tendency to f a l l off s l i g h t l y with increasing cupric acetate.  Thus at 100°C. the  -26-  TIME MINUTES  F i g . 9. T y p i c a l Rate Plots f o r the Cupric Acetate Catalyzed Reduction-of Dichromate by Hydrogen. 0.10 M./L. CuAc ; 100°C. 13.6 atm. H . 2  2  -27-  TIME  - MINUTES  F i g . 10. Rate Plots f o r the Reduction of Dichromate at Different Cupric Acetate Concentrations. 0.25 M./L. NaAc; 0.50 M./L. HAcj 100°C.j 13.6 atm. H . 2  F i g . 11. Dependence of the Reaction Rate on the Cupric Acetate Concentration. 0.25 M./L. NaAc; 0.50 M./L. HAc; 100°C; 13.6 atm. H . 2  Fig. 1 2 . Dependence of the Rate of the Cupric Acetate Catalyzed Reaction on the Hydrogen P a r t i a l Pressure. 0 . 1 0 M./L. CuAc ; 0 . 2 5 M./L. NaAc; 0 . 5 0 M./L. HAc; 1 0 0 ° C . 2  value of k, extrapolated to zero [CuACgl i s 4»7 x 10  1. mole.  -  sec."* ,  while at the highest cupric acetate concentration investigated, i . e . 0.25 M./L., i t i s only 4.05 x 10~^ or about 14$ lower.  The significance of this  decrease w i l l be discussed later. The rate constant, k, for the cupric acetate catalyzed reduction of dichromate, i s thus seen to be nearly 120 times that for the corresponding cupric perchlorate catalyzed reaction. the correspondingly lower value of AH  This increased rate i s reflected i n (calculated fromi the linear plot of  log (k/T) vs. l/T i n F i g . 13) > i . e . 23»0 - 1.5 kilocalories per mole. corresponding entropy of activation, A %  The  , i . e . -8.0 - 4»0 e.u., i s i n  reasonable agreement with that reported for the perchlorate system and  may  be considered normal for a simple bimolecular reaction i n solution (46,47). Using the Eyring rate equation, the rate constant may be expressed as:  k =  ^  ' e  -23,000 RT  -^0 ° e R  [203  In most of these experiments the solutions contained, i n addition to sodium dichromate and cupric acetate, 0.25 M./L. 0.5 M./L.  of acetic acid.  of sodium acetate and  The pH was thus buffered at a value of about 4.5  and changed very l i t t l e during the course of the reaction.  That these values  are not c r i t i c a l i s indicated by the results i n Table II which show that wide variations i n the acetate and acetic acid concentrations do not affect the rate significantly.  However, where sufficient perchloric acid was added to  lower the pH of the solution below 4, the rate decreased appreciably, as shown in F i g . 14, presumably as a result of the dissociation of the more active cupric acetate complexes, i . e . CuAc + H  +  2  and  0uAc + H* +  > —>-  CuAc + HAc  [2l]  Cu "* + HAc  [22]  +  4  -31-  F i g . 13". Plot of Log (k/T) vs. l / T f o r the Cupric Acetate Catalysed Reduction of Dichromate by Hydrogen. 0.05 M./L. CuAc j 1 3 . 6 atm. H . 2  2  TABLE I I BATES OF REACTION BETWEEN HYDROGEN AND DICHROMATE IN CUPRIC ACETATE SOLUTIONS OP VARYING COMPOSITION  0.10 M./L. CuAc ; 2  100»C.; 13.6 atm. H  Solution Composition  g  Reaction Rate  [NaAc]  [HAc]  M./L.  M./L.  H-l*  0.010 0.010 0.010 0.010 0.010 0.010 0.010 0.010  0.25 0.25 0.25 0.25 0.25 0.25 0.25 0.25  0.50 0.50 0.50 0.50 0.50 0.50 0.50 0.50  4.82 5.04 4.92 4.76 4.78 4.71 4.66 4.66  4.49 4.69 4.57 4.43 4.45 4.38 4.33 4.34  C-4 ©-3 C-2 C-5  0.005 0.010 0.015 0.020  0.25 0.25 0.25 0.25  0.50 0.50 0.50 0.50  5.13 4.92 4.62 4.76  4.77 4.57 4.30 4.43  F-5 F-4 F-3 F-2 F-l G-2 G-3  0.010 0.010 0.010 0.010 0.010 0.010 0.010 0.010  0.00 0.05 0.15 0.25 0.50 0.25 0.25 0.25  0.50 0.50 0.50 0.50 0.50 0.50 0.75 1.00  4.61 4.71 4.74 4.71 4.79 4.66 4.75 4.63  4.29 4.39 4.42 4.39 4.46 4.33 4.42 4.30  Expt. [Na Cr 0 ] 2  No. A-4 B-4 C-3 D-3 E-3 F-2  2  M./L.  7  S t i r r i n g velocity reduced to 600 r.p.m. velocity was 900 r.p.m.  -dtHgl/dt x IO mole l . "  1  6  k x IO , -1 -1 1. mole. sec. 5  sec.""  1  In a l l other experiments s t i r r i n g  -33-  F i g . 14. Dependence of the Reaction Rate i n Acetate Solutions on pH. 0.05 M./L. Cu(II); 1.0 M./L. T o t a l Acetate; 100°C.| 20 atm. H . pH measured at Room Temperature. 2  pH F i g . 15. E f f e c t of pH on the Relative Concentrations of Cupric Species i n Acetate Solutions at Room Temperature. 0.05 M./L. Cu(Il); 1.0 M./L. T o t a l Acetate. Based on Published ComplexityConstants (49,50,51,52).  -34-  The shape of the curve, depicting pH dependence of the reaction rate (Fig. 1 4 ) , i s qualitatively similar to that for the variation of total concentration of complexed Cu(ll) ( i . e . , CuAc,, + CuAc ) plotted i n F i g . 15, +  u t i l i z i n g the known room temperature complexity constants for cupric acetate (49,50,51,52).  This supports the view that the catalytic a c t i v i t y of cupric  acetate solutions i s primarily due to complexes such as CuACg and CuAc . +  A comparison of the curves i n Figs. 14 and 15 suggests that the two complexes have similar catalytic a c t i v i t i e s . It i s probable that the observed slight decrease i n k with i n — creasing cupric acetate concentration arises from a. secondary effect such as the lowering of the s o l u b i l i t y of hydrogen..  Similar decreases i n rate  were observed when inert salts were added to the solution, as shown i n F i g . 16.  In most cases for which data are available, the variation of k with  salt concentration was found to be i n f a i r agreement with the lowering of the s o l u b i l i t y of hydrogen ( 5 3 ) •  The anomalous increase i n reaction rate  due to sodium perchlorate may be due to a slight salting-in effect, i . e . an increased hydrogen s o l u b i l i t y , sometimes observed for large negative ions ( 5 4 ) . The kinetics of the cupric acetate-catalyzed reduction of dichromate by hydrogen, reported above, are i n good quantitative agreement with those found e a r l i e r for the reduction by hydrogen of cupric acetate i t s e l f to cuprous oxide  (37,38|39).  A similar reduction of cupric  acetate was observed i f the reduction of. dichromate was allowed to proceed to completion, as shown i n F i g . 1 7 .  The f i r s t order rate constant f o r this  reduction agreed with the value obtained previously for this reaction (39) and with the rate constant for the activation of hydrogen calculated from the rate of the preceding dichromate reduction. As i n the case of the perchlorate system, the pattern of kinetic  -35-  NaC10  -O  NaN0  -A  4  3  Na S0 2  - •  4  Ca(N0 ) 3  MgSO*.  0  2  -O -O  0.5  1.0 NORMALITY  1.5  OF SALT  Fig. 16. Effect of Various- Salts on the Rate of the Cupric Acetate Catalyzed Reaction. 0.10 M./L. CuAc ; 0.25 M./L. NaAc; 0.50 M./L. HAc; 100°C; 13.6 atm. H . 2  2  -36-  F i g . 17. Typical Rate Plots f o r the Consecutive Reduction of Dichromate and Cupric Acetate by Hydrogen. 100°C; 13.6 atm. H . 2  behaviour suggests that the rate i s determined by a bimolecular process involving one cupric acetate complex and one hydrogen molecule.  The large  difference i n the rate constants for the two systems, ( i . e . a factor of 120) reflecting the higher catalytic a c t i v i t y of the cupric acetate, may  be  explained by assigning to the acetate ligand the role of a proton acceptor i n the rate determining step, i . e . CuAc* + H  g  a  ^ > Qw  CuH* + HAc  [23 J  Some measure of further support would thus appear to be provided for the suggestion that CuH* i s the activated intermediate i n these systems.  C.  EFFECT OF COMPLEXING ON THE HOMOGENEOUS ACTIVATION OF MOLECULAR HYDROGEN BY CUPRIC SALTS In view of the fact that the catalytic a c t i v i t y of the undissooiated  cupric acetate complex was found to.be about 120 times greater than that of the uncomplexed cupric ion, i t seemed of interest to examine the effect of other complexing ligands.  Therefore, kinetic studies on the catalytic  activation of hydrogen by a number of organic and Inorganic cupric complexes were made.  In each case, the reduction of dichromate was used as a measure  of the rate.  Cupric Carboxylate Complexes Measurements of the. rate of reduction of dichromate by hydrogen i n the presence of various cupric carboxylate salts are reported i n Table I I I . As with cupric acetate, the kinetics were found to be zero order i n dichromate.  The reaction rates could generally be estimated to within - 5 $ p (  except i n solutions containing formate,, where the uncertainty was much greater, because the rate of reaction of dichromate with formate was of the same order  -38TABLE III  RATES OP REACTION BETWEEN HYDROGEN AND DICHROMATE IN SOLUTIONS OP VARIOUS CUPRIC CARBOXYLASE SALTS 0.0005 M./L. Na Cr 0 ; 2  Expt. N o  2  *  Pr-3 Pr-8 Pr-5 Pr-4 Pr-6 Pr-7  0.05 CuPr 0.05 ll £ 0.05 II 0.05 it 0.05 n 0.05 ti  Bu-1 Bu-2 Fo-1 Po-2  2  100°C.  -d[H I/dt x 10**  Solution Composition M o / L  20 atm. H ;  7  Mole L i t e r - S e c . "  °  0.0 NaPr 0.25 " 0.50 " 1.00 " 0.25 " 0.25  k x 10^  2  1  L.Mole^Sec."  0.5 HPr 0.5 " 0.5 " 0.5 " 0.2 " 2.0" "  4-38 4.38 4.83 4.50 4.44 5.49  5.75 5.83 6.33 5.94 5.90 7.13  0.026 CuBu 0.052 ii ^  0.25 NaBu 0.5 HBu 0.5 " 0.25 "  2.31 4.59  5.82 5.86  0.10 CuPo0.10 ii 2  0.25 NaPo 0.25 HPo 0.25 " 0.25 "  2.94 2.88  1.8 1.8  0.39  2.8 2.2 1.1 2.5  0  Su-4 0.01 Su-10 0.01 Ma-2 0.01 Fu-1 0.002  CuSu *' CuMa CuFu  11  0.025 Na Su 0.05 " 0.13 Na.Ma 0.03 Na^Pu  0.25 H„Su 0.60 • 0.015 H„Ma 0.13 H_Pu  0.33  0.16 0.08  The following abbreviations are used i n this thesis: Pr - propionate; Fu - fumarate;  Bu - butyrate;  Ma - maleate;  ethylenediamine.  Po - formate;  Ac - acetate;  Su - succinate;  Gl - glycine or glycinate;  EDA -  1  as that with hydrogen. The results for cupric propionate (CuPTg) parallel those obtained e a r l i e r for cupric acetate, suggesting that the mechanism of activation of hydrogen i n the two systems i s essentially the same.  Pig. 18 shows that the  rate i s nearly proportional to the cupric propionate concentration, a slight tendency being apparent for the rate constant, k, to decrease at higher cupric propionate concentrations.  A similar effect i n the cupric acetate system  was attributed to a lowering of the s o l u b i l i t y of hydrogen.  The rate i s  seen to be essentially independent of sodium propionate or propionic acid over a wide range.  By analogy with the cupric acetate system i t i s concluded  that the active species i n the present case are probably CuPrg and CuPr . +  At 1 0 0 C , the catalytic a c t i v i t y of cupric propionate i n B  propionate-buffered  solutions (k sir 6.0 x 10 ^ l i t e r mole  1  sec. ^) was found  to be about 25$ higher than that determined for cupric acetate (k •= x 10~^ l i t e r mole" s e c . ) . 1  - 1  4.7  Rate measurements for cupric propionate at  temperatures between 80" and 120°C. gave a good linear plot of log (k/T) vs. l/T  (Pig. 19) from whose slope an apparent activation enthalpy, AH*,  of  - 1.5 k i l o c a l o r i e s per mole and an activation entropy, AS *, of -7.5 C  e.u. were calculated.  23.0 4.0  These values are i n excellent agreement with those  for cupric acetate. Two experiments with cupric butyrate i n butyrate-buffered  solutions  yielded rates identical with those f o r cupric propionate (see Table I I I ) . Measurements of the rate of activation of hydrogen i n formatebuffered solutions were subject to large uncertainties because of arising from side reactions between formate and dichromate.  complications  The rate at  100»C. (k = 1.8 x 10"^ l i t e r mole" sec." ) was about one-third that for 1  1  cupric acetate and the higher carboxylate complexes.  This may r e f l e c t the  -40-  F i g . 18. Dependence of the Reaction Rate on the Cupric Propionate Concentration. 0.25 M./L. NaPr; 0.50 M./L. HPr; 100°C.j 20 atm. H . . . . a  -4.0  2.4  2.6  2.5  l/T X 10  1  2.7 3  )•  F i g . 19. Plots of Log (k/T) vs. l / T f o r the Cupric Propionate and Cupric Sulphate Catalyzed Reactions Between Dichromate and Hydrogen. 20 atm. H . 2  -42mm  fact that complexing between Fo  \ 'fr  and Cu  i s incomplete, consistent with the  higher dissociation constant of formic acid.  Similar considerations apply-  to cupric fumarate and cupric maleate. The catalytic a c t i v i t y of succinate-buffered cupric solutions (k -2: 2.5 x 10~^ l i t e r mole cupric acetate.  -1  sec.  -1  at 100»C.) was about one half that for  It i s probable that the cupric succinate complex has a  chelate structure. Cupric Sulphate A typical rate plot for the reduction of dichromate i n the presence of cupric sulphate i s shown i n F i g . 20. summarized i n Figs. 21 and 22.  The results for cupric sulphate are  It i s seen that the addition of NagS0  found to increase the catalytic a c t i v i t y of Cu complex formation.  was  4  , presumably as a result of  It has been suggested that only the f i r s t stage of  association, i . e . Cu** + S 0  = 4  CuS0  [24]  4  i s important (55,56), and that higher cupric sulphate complexes are not formed to an appreciable extent.  The failure of the curve i n F i g . 21 to  level off suggests that complexing i s incomplete at S0 "~ concentrations as 4  high as 0.5  M./L.  Consistent with the above interpretation, i t was found that the experimental results could be f i t t e d by a relation of the form: k [Cu**l + k [CuSO ] .  k  =  tCu(ll)]  ~  t  2 5 1  where k^ and kg represent the specific rate constants f o r the activation of Hg by Cu  and CuS0 respectively; 4  i s known from measurements i n per-  chlorate solutions and has a value of 0.93 x 10  1A  - l _ i l i t e r mole sec. at  -43-  100  F i g . 20. T y p i c a l Rate Plots f o r the Reduction of C r 0 by H i n the Presence of Various Cupric Complexes. 20 atm. H , =  2  7  z  2  1. 2. 3. 4.  Ethylenediamine (0.05 M./L. C u ( l l ) , 0.20 M./L. EDA, pH 4.0, 130°C, Chloride (0.10 M./L. Cu(II)j 1.0 M./L. CI"; 110°C. Sulphate (0.10 M./L. Cu(II); 0.10 M./L. 3 0 ^ ; 110°C. Glycine"(0.05 M./L. C u ( I l ) ; 0.05 M./L. G l ; pH 3.5; 130°C.  Fig. 21, Dependence of the Reaction Rate on SO^" Concentration. 0.10 M. Cu(Il)j 20 atm. H ; 110 C. P  2  F i g . 22-. .. Dependence of the Reaction Rate i n Sulphate Solution on Cu(Il) Concentration. 0.25 M./L. S04=; 0.10 M./L. HSO^", 20 atm. H ; 110 C. 9  2  -46110«C. (Table i ) . Prom equation [24], [Cu**] and [CuSO^J are related by [CuSOj [Cu^itSO^l while  =  [Cu**] + [CuS0 ] 4  K  [26]  = [Cu(ll)I  [27]  The best correspondence between a curve f o r k, calculated from equations [25], [26], and [27], and the experimental points (shown i n Pig. \ 21J was obtained using the following values:  kg = 6.5 x 10  sec."* at 110°C. (i.e. about 7 times as great as k ^ 1  mole"" . 1  —4  l i t e r mole  —1  and K = 6.7 l i t e r  The excellent f i t lends support to these values.  However, some  uncertainty arises through neglect of a c t i v i t y coefficients and of the effect of variable salt concentration on the s o l u b i l i t y of hydrogen.  It  i s of interest that the above value of K, derived from the kinetic results at 1 1 0 C , i s of the same order as the room temperature value ( zr 4.5 B  liter  mole"" f o r ionic strength of unity) estimated by a spectrophotbmetric 1  method (56). A plot of log (k/T) vs. l/T f o r CuSQ^ based on reaction rate measurements at temperatures between 90  B  and 130°  i s shown i n Pig. 18.  The apparent activation enthalpy of 27.3 kilocalories per mole, obtained from the slope of this plot, presumably reflects the effect of temperature on the relative proportions of the cupric sulphate complex and the cupric ion present, as well as on their specific rate constants. Cupric Chloride Complexes A typical rate plot f o r the reduction of dichromate by hydrogen i n the presence of cupric chloride i s shown i n F i g . 20.  The results f o r  cupric chloride solutions, summarized i n Fig. 23» indicate that C l ~  -47-  F i g . 23. Depsndence of the Rate of the Cupric Chloride Catalyzed Reaction on the Chloride Concentration. 0.10 M./L. C u C C l O O a ; 2 0 a t m « 2> H0°C. H  -48increases the catalytic a c t i v i t y i n a manner attributable to the formation of cupric chloride complexes. 1 M./L.  It i s probable that i n the region above  C l " , where the rate levels o f f , the cupric ion i s f u l l y  plexed, i . e . exists as CuCl^ .  com-  Hence i t seems reasonable to attribute the  -  measured value of k i n this region to this complex*  This value (2.5 x  10~^ l i t e r mole"" sec." ) i s about 2.7 times as great as for the uncomplexed 1  Cu  1  ion. At constant Cl~" concentrations the rate was found to be d i r e c t l y  proportional to the total cupric concentration (Fig. 24) confirming that the catalytic a c t i v i t y i s associated with the cupric species. Cupric Ethylenediamine Complexes .The extent of complexing between Cu H  and EDA i s a function of the  concentration, i . e . CHg— NHg  ^NH —CH 2  CHg——NHj  2  +  CH  2  NHg  NH  2  4 f  CHg  Cu  + 2  [28] C H g — NH^  [All  The variation of k with pH, shown i n F i g . 25, thus probably reflects a variation i n the extent of complexing of the cupric ion. At low pH ( <C 3» measured at room temperature) dissociation of the complex i s apparently complete and the observed value of k approaches that f o r  Spectrophotometrically determined values of the complexity constants of cupric chloride at room temperature  (57) suggest that higher C l ~ concentrations  would be required to achieve complete complexing.  However, i t has been  pointed out that these values probably increase with temperature (58).  -49-  -50-  F i g . 25. Solutions.  E f f e c t of pH on the Reaction Rate i n Cupric Ethylenediamine.  the uncomplexed Cu  ion.  Increasing the pH (by decreasing the amount of added  HCIO^) results i n a lowering of the rate, presumably reflecting the lower catalytic a c t i v i t y of the C U ( E D A )  + + 2  complex [A].  -4 levels o f f at a value ( k ; t 0.6 x 10  -1 l i t e r mole  At high pH ( > 5) the rate -1 sec.  \ at 130°C.; which i s  about one-eighth that for the uncomplexed cupric ion at the same temperature. Room temperature values of the complexity constants (59) suggest that i n this region complexing i s essentially complete, a l l the Cu(ll) being present as the CU(EDA)2  complex;  hence i t seems reasonable to associate the measured value  of k with this complex. This interpretation i s consistent with the results shown i n Pig. 26. In a series of experiments i n which the Cu(ll) concentration was increased while the total EDA concentration was held constant, the rate (and the apparent value of k) rose sharply as the ratio [Cu(ll)l/[EDA], approached the value 0.5» Cupric Glycinate Complexes Figs. 27 and 28 depict the effect of glycine on the rate of catalytic activation of hydrogen by Cu(ll) i n solutions of different pH. In basic solutions complexing of Cu** with glycine (involving formation of the CuGlg chelate complex) i s essentially complete (59). However, with decreasing pH i t i s probable that the complex undergoes dissociation, passing through a number of intermediate stages including the following*  F i g . 26. Dependence of the Reaction Rate i n Ethylenediamine Solutions on the Cupric Salt Concentration. 0.20 M./L. EDA; 20 Atm. H ; 1 3 0 ° C ; pH 6 (Room Temperature). 2  -53-  Fig. 27. Effect of pH on the Reaction Rate i n Cupric Glycinate Solutions.. 0.05 M../L. Cu(C10*)2i 0.20 M./L. Glj 20 Atm. H j 130°C. 2  -54-  Fig. 28. Dependence of the Rate of the Cupric Glycinate Catalyzed Reaction on the Glycine Concentration at Several pH's. 0.05 M./L. Cu(C10^) j 20 Atm; H;. 130°C. 2  2  -55C H „ — NH_ 2  0 = C—  0— G = 0  2  \  0  Cu  /  0  •  I °  CE,  + 4 H  rNHg-— CH  Cu  r /  /  CH„  N8L  0 = C—0  2  rC = 0  [C]  [B] NH, ^---  :i,t  Cu  + 2  CH  [293  2  0 = C — OH  Consistent with this i t was found that at low pH (~ l ) the value of k approaches that f o r the undissooiated C u  + +  ion (Fig. 2 7 ) and i s  essentially independent of the Gl concentration (Fig. 2 8 ) . In the region of high pH ( > 3.5), k decreases with increasing pH (Fig. 2 7 ) and with increasing glycine concentration (Fig. 2 8 ) approaching a value  /  -4 - 1 - 1 \ 1.6 x 10 l i t e r mole sec. at 130»C.J about one-third that i  of the uncomplexed Cu  ion at the same temperature.  presumably r e f l e c t s the catalytic a c t i v i t y of the CuGl  This value, which 2  complex [B],  (equation [ 2 9 J ) i s subject to considerable uncertainty because of the competing side reactions between dichromate and glycine (see F i g . 2 0 ) . In the region of intermediate pH ( l - 3*5) the measured value of "Hfr*  k i s greater than that f o r Cu  and increases with the Gl concentration,  presumably reflecting the presence of cupric carboxylate complexes such as [C] (equation [ 2 9 ] ) '  In common, with the other cupric carboxylate com-  plexes discussed earlier, the catalytic a c t i v i t y of this complex i s . apparently greater than that of the uncomplexed cupric ion. Summary of the Effect of Complexing  Complex forms of copper i n aqueous solutions apparently activate  -56hydrogen by essentially the same mechanism as the simple hydrated cupric ion, but with widely d i f f e r i n g a c t i v i t i e s .  The data obtained i n these studies  have been used to compile Table IV i n which the various complexes, f o r which values of k have been determined, are arranged i n order of decreasing catalytic a c t i v i t y .  The l i s t e d values of the relative catalytic a c t i v i t i e s  (that of the uncomplexed cupric ion being taken as unity) are only approximate since they are based on rate measurements at different  temperatures  (ranging from IOC-G to 130*C.) and of varying precision. Of particular interest i s the fact that the negative ions enhance the catalytic a c t i v i t y of Cu i . e . Bu~, Pr~;>  Ac"* >  SO^" >  i n order of increasing basicity,  Cl~ >  c  l°4~*  This supports the suggestion,  made previously f o r Ac"", that their promoting action results from their role as proton acceptors i n the rate determining step, i . e . CuX + H  CuH + HX  +  +  2  [30]  The depressing effect of chelating reagents containing ammane groups i s not clearly understood.  The effect i s thought to be primarily  due to co-ordination of the cupric ion with nitrogen, which i s less electronegative than oxygen or chlorine and, therefore, through electron donation ++ to the Cu  ion decreases the electron accepting power of the l a t t e r .  role of the electron accepting power ( i . e . electron a f f i n i t y ) of the C u  The + +  ion i n the hydrogen activation process w i l l be discussed l a t e r . D.  HOMOGENEOUS ACTIVATION OF HYDROGEN BY SALTS OF OTHER METALS  In addition to the cupric salts, salts of certain other metals were tested f o r their a b i l i t y to activate molecular hydrogen homogeneously i n aqueous solutions.  In most cases, dichromate was used as an indicator.  The following salts were tested:  -5717  TABLE  RELATIVE CATALYTIC ACTIVITIES OF VARIOUS CUPRIC COMPLEXES  —1 -1 k - l i t e r mole sec.  Complex CuBu  2  5.84 x 10-5 at 100*C.  CuPr  2  5.95 x IO"  * Catalytic A c t i v i t y 150  ti  II  150  CuACg  4.70 x 10'5 it  II  120  Cu  3.88 x IO"  it  II  1  CuSO. 4  6.5  x IO"  4  2.5  x IO"  4  CuCl  = 4  Cu _  - H r  Cu CuGl  1.6 t  2  CU(EDA)  2  +  +  Relative to that of Cu  5  5  9.55 x IO"  5  5.35 x IO"  4  i.o x io6  x IO"  5  4  at 110»C.  7  n  II  2.7  it  II  1  at 130»C. II  II  II  II  1 < 0.5 0.1  which i s taken as unity at each temperature.  NiAc  MgAc  2  Ni(C10 )  MgS0  4  4  Mg(C10 ) 4  CaAc  ZnAc  2  2  2  Ga(C10 ) 4  2  HgAc  2  Zn(C10 )  2  4  NiS0  4  ZnS0  CoAc  2  CdAc 2  HgS0  2  4  4  AgAc  2  A1(C10 )  5  Gr(ci0 )  5  4  U0 Ac  MnACg  2  2  4  Of these, only the mercuric and s i l v e r salts showed measurable reactivity towards hydrogen at temperatures up to 140°C. and pressures of 15*6  atmospheres.  E.  MECHANISM OP HYDROGEN ACTIVATION In the above discussions of the experimental results, repeated  references have been made to the occurrence of an "activated intermediate". While the nature of this intermediate complex has not been determined  by  direct observation, the kinetic evidence f o r i t s formation i s considered f a i r l y conclusive.  Of the various possible intermediates for which  reasonable 'classical' structures can be written, the most l i k e l y i s believed to be CuH , whose formation i s depicted by equation [30]. +  This suggestion  i s supported by a consideration of the energetics of the processes leading to the formation of the possible alternative intermediates.  ** The available thermochemical. data  lead to the following expressions'  for the endothermicities of the alternative rate-determining processes:  * Reference has already been made to the results of subsequent kinetic studies of these reactions.  See pages 7 and 8.  The following standard enthalpies.of formation (AH ) have been used: B  f  H* (aq.), 0;  H<> (g), 52.1 k i l o c a l s . ,  Cu** (aq.), 15.4 k i l o c a l s . (60)j  He (g), 0;  Cu» (g), 81.5 k i l o c a l s . ;  Cu* (aq.), 17.1 k i l o c a l s . ( 6 l ) .  -59-  Cu  + +  + H  where S  Cu** + H  g  Cu** + H  2  >- Cu + H +  •  »» Cuo + 2H  >" CuH + H +  2  9  H  +  [32]  ; AH° = 66 kcal. -  ; AH» = 54 kcal. - D +_ + (S +  +  [31]  + H° ; AH = 54 kcal. - S „  +  Cu  H  Qu  [33)  - 3^+)  represents the hydration energy of the gaseous species,' x, 'and'  x  g represents the dissociation energy of the gaseous molecule, AB. Of these p o s s i b i l i t i e s , reactions [3IJ  and [32I appear to be  :  energetically inconsistent with the observed activation energy of 24 kcal. since S^ and  s  a  C u  g » corresponding to the hydration energies of the gaseous  atoms, can hardly be very large. be evaluated very accurately.  The energetics of reaction [33)  While CuH  cannot  appears to have been detected  spectroscopically (62) i t s dissociation energy i s not known; however, the value i s not l i k e l y to be smaller than that of the isoelectronic molecule NiH, i . e . 60 kcal. (9). than S  CuH  At the same time, while S ^ * i s probably larger  + , the difference i s not l i k e l y to exceed the 30-40 kcal. which  would be required to make the endothermicity of the process greater than the observed activation energy.  Thus, on energetic grounds, the occurrence  of CuH as the activated intermediate as a result of the process represented +  by equation [33]  appears the most probable.  It i s unlikely that the formation of Hg equation  [3li)»  or of C u H  ++ 2  ( i n place of H° + H i n  +  ( i n place of CuH + H +  +  +  i n equation  [33J} i n  aqueous solution would be energetically favoured, because of the high solvation energy of the proton. The reaction represented by equation [33]  can be depicted schemat-  i c a l l y by means of a potential energy diagram such as that shown i n F i g . 29» 11  Repulsion between the reacting species, Cu  and H , i s reflected i n the 2  increasing potential energy as the Cu - H distance decreases.  The  activated complex corresponding to the cross-over point of the two curves  -60=  CS  Pi  w  Cu ~ H  DISTANCE  F i g . 29. Schematic P o t e n t i a l Energy Diagram f o r the Reaction Cu + .H; — CuH + H . + +  +  2  +  -61probably has a configuration approximating [Cu  —- H  H ].  A  further decrease i n the Cu - H distance from this point results i n the formation of a CuH molecule, accompanied by a decrease i n potential energy, +  the resulting proton being released to an adjacent base.  The entire pro-  cess of hydrogen activation can therefore be represented by the expression Cu** + H  2  H" -T-  [Cu**  This mechanism implies that a H  H*]*  CuH* + H*  displacement or, essentially, an  electron displacement from the. hydrogen molecule to the cupric ion, occurs during the activation process.  A prerequisite requirement for this i s  presumably that the catalyst species, i . e . Cu  , must have a high electron  affinity. Certain electron displacement phenomena that occur i n solution, such as the formation of co-ordinated complex molecules, have been related to the electron a f f i n i t i e s of the corresponding gaseous ions.  Thus i n a  series of metal chelate complexes, a linear relation was observed between the log of the s t a b i l i t y constants and the ionization potential for the removal of the l a s t electron i n the formation of the corresponding gaseous ion (63)0  Where an electron displacement mechanism i s responsible f o r the  homogeneous activation of hydrogen,, some degree of correlation might similarly be expected between catalytic a c t i v i t y and the corresponding gaseous electron a f f i n i t i e s of the active metal ions. view, i t might be pointed out that Cu ++ Hg  +  In support of this  and other metal ions such as Cu ,  +  , Hg , and Ag  [34]  which have been found to activate hydrogen homogeneously  i n solution are characterized by high electron a f f i n i t i e s i n the gaseous state, compared with analogous inactive ions. Also consistent with the suggestion that the activation process involves electron transfer from the hydrogen molecule to the catalyst, i s  '  /  the observation that other electron transfer reactions of the c a t a l y t i c a l l y active ions are also notably fast, as indicated by studies of the isotopic electron exchange between Hgg  ++  and Hg** (64) and of the catalytic influence  of Cu** on various oxidation reactions (65,66).  The promoting influence  of negative ions on the catalytic a c t i v i t y of Cu** also parallels the well known general catalytic influence of complexing negative ions such as OH  and CI  on the Fe  - Fe'  and other isotopic electron exchange re-  actions (68), an effect that has been attributed to their bridging action, or effectiveness i n lowering the electrostatic repulsion which constitutes a barrier to the mutual approach of two exchanging ions (68,69).  This  may  be an alternative interpretation of the influence of these ions, to that which has already been considered. Eyring and Smith (70) have pointed out that the fusion of more than one electronically saturated molecule into an activated complex normally involves electron promotion;  hence a lowering of the activation  energy i s to be expected when the activated complex i s co-ordinated with a suitable electron acceptor.  The mechanism proposed here, which attributes  the catalytic activation of hydrogen to the electron accepting  properties  of the metal ion catalysts, reflects the same principle. The interpretation proposed here has some important features i n common with that originally suggested by Dowden (20,21,71) for the ordinary heterogeneous catalytic a c t i v i t y of metals, i n which the electronic work function of the metal was introduced i n a manner analogous to the electron a f f i n i t y of the metal ion i n the present case.  CONCLUSIONS  The a b i l i t y to activate molecular hydrogen homogeneously i n aqueous solutions was found to be a general property of the dissolved cupric salts that were studied.  The catalytic a c t i v i t y of the cupric ion was  found to be strongly affected by complexing, the a c t i v i t i e s of the cupric complexes decreasing i n the order: sulphate >  butyrate, propionate > acetate  chloride ^> water ( i . e . the uncomplexed cupric ion) }  glycine,  ethylenediamine). The species CuH i s postulated to occur as the reactive i n t e r +  mediate i n the stepwise process leading to the overall reduction of the substrate. (1)  This suggestion i s supported by the following considerations:  The dependence of the dichromate reduction rate on the H  +  con-  centration can be explained most readily i n terms of competition from a back reaction, i f CuH and H +  +  are the immediate products of the rate con-  t r o l l i n g step. (2)  The promoting influence of complexing anions which follows the  same order as their b a s i c i t i e s finds explanation i f they are assigned the role of proton acceptors i n the rate determining step. (j)  Within the l i m i t s of endothermicity consistent with the observed  activation energy of the reaction, CuH appears energetically more plausible +  than the other intermediates which have been suggested. (4)  A similar intermediate, i . e . Cu^'H, has been postulated to  explain the homogeneous activation of hydrogen by cuprous acetate i n quinoline (21,22,24,27). The mechanism proposed here, entails the displacement of a hydride ion from EL to the C u  ++  ion, both electrons i n the Cu  +  H bond being con-  -64tributed by the hydrogen molecule.  In this context, the role of Cu  appears to be that of an electron-acceptor suggesting that i t s catalytic activity may be related to the presence of low-lying unoccupied orbitals into which the Hg electrons can enter.  The lowering of the catalytic  activity of Cu** on chelation with glycine and ethylenediamine (Table IV) may reflect the fact that in the complex these orbitals are used in bonding. The heterolytic splitting of the Hg molecule in the activation process requires only one metal atom 'site , as evidenced by the first 1  order kinetics with respect to the cupric salt.  The 'dual site  1  requirement of the process applies only i n the sense that a base molecule such as H_0 or Ac  may be necessary as a second site to absorb the proton.  -65RSFERENCES  1.  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Dowden, D.A., J.Chem.Soc, 242 (1950)  APPENDIX A  ABSORPTION SPECTRA OF SODIUM DICHROMATE AND CHROMIC ACETATE IN AQUEOUS SOLUTIONS  Dichromate was normally determined i n solution samples by a spectrophotometric method based on the absorption peak at 350rn^ushown i n F i g . 30. Except i n chloride solutions where another method was used, cupric salts do not generally absorb l i g h t i n t h i s region.  Beer's law i s obeyed at  -4 a l l concentrations up to 5 x 10 moles per l i t e r CrgO^ , the maximum that was determined without further d i l u t i o n .  Dichromate concentrations as low  as 5 x 10~^ moles per l i t e r were detectable by this method. The absorption spectrum of dissolved chromic acetate, shown i n F i g . 30, was found to be identical with that of reduced dichromate solutions of similar total chromium concentration, allowance being made for absorption by cupric acetate i n the red region of the spectrum.  This was accepted as  evidence that the reduction product of CrgO^ i n these solutions was C r ( l l l ) . -  From Fig. 30, CrAc^ i s seen to absorb l i g h t i n the 350m/<region of the spectrum.  This absorption was found to be general for chromic salts  i n aqueous solution, their optical density averaging about 1.7$ of that for similar concentrations of dichromate.  A small correction was made to  eliminate systematic errors from this source.  -70-  WAVELENGTH - w\/A-  Figure 30. Absorption Spectra of Sodium Dichromate and Chromic Acetate Solutions. (Determined with a Beckman DU Spectrophotometer).  -71APPENDIX B  SUMMARY OP SELECTED EXPERIMENTAL DATA  (a)  Effect of I n i t i a l Dichromate Concentration on the Rate of the Cupric Perchlorate Catalyzed Reduction of Dichromate by Hydrogen (Fig. l ) . 0.10 M./L. Cu(C10,) , 0.01 M./L. HC10,; 110»C; o  Expt, No.  -d[Cr_0 ~]/dt x I O -1 -1 mole l i t e r sec.  Initial [Cr 0 "J M./L.  8  20 atm. H  -d[Hul/dt x 1 0 -1 -1 mole l i t e r sec. 8  k x 10 - l •1 l i t e r mole s e c ' 5  S-7-R  0.0010  5.25  15.7  9.48  S-5  0.00025  5-37  16.1  9.63  S-6  0.00075  5.30  15.9  9.55  S-8  0.00050  5.37  16.1  9.75  (b)  Comparison of the Reaction Rates f o r Various Substrates;  Effect of  Cupric Perchlorate Concentration (Figs . 2 and 3)-  M./L.  HC10 ; 4  110»C;  Expt. [Cu(C10 ) J M./L. No.  0.01  20 atm. Hg  Substrate ^  -d[X]/dt  -d[HU]/dt  x 10 - mole  x 10 -mole , .. -1 -1 liter sec  8  liter sec. _ 1  U-l  0.20  Cr 0  ?  U-2  0.27  Cr 0  ?  U-3  0.36  Cr 0  ?  U-4  0.51  Cr 0  ?  U-5  0.10  Cr 0  ?  2  2  2  2  2  -1  k x 10  l i t e r mole"" -1 sec  1  10.3  3.10  9.43  13.8  4.13  9.32  =  18.0  5.41  =  27.2  8.15  9.80  1.57  9.60  =  =  =  5.23  .-  5  9.07  (b) contd.  -72-  X-l  0.11  I 0  3~  X-2  0.22  I 0  3~  X-3  0.32  I  X-4  0.20  X-5  0.10  (c)  O  1.59  8.80  12.9  3.22  8.90  18.6  4.64  8.78  2.72  8.18  1.31  7.77  6.37  f  54.5 _ ++++ Ce  .  26.2  Effect of Hydrogen Pressure on the Reaction Rate .(Figs. 4 and 5). 0.10 M./L. Cu(C10 ) , 0.01 M./L. HC10 ; 110»C. 4  Expt. P No 2 * .. Atm. w  2  4  - d [ C r 0 l / d t x 10 —1 —1 mole l i t e r s e c . =  o  -d[Hj/dt x 1 0 —1 —1 mole l i t e r " sec."  8  -  k x 10 —1 —1 l i t e r mole" s e c .  8  7  -  5  -  U-5  20  5.23  15.7  9.60  U-6  25  6.50  19.5  9.48  U-7  15  3-87  11.6 .  9.43  U-8  10  2.60  7.8  9.48  U-9  5  1.23  3.7  8.88  (d)  Effect of Temperature on the Reaction Rate (Fig. 7)« Cu(C10 ) , 0.01 M./L. HC10 ; 4  2  4  Expt. Temp. -d[Cr 0 j / d t =  Nn  or  1  7  ft  '  x 10 -mole liter" s e c 1  - 1  1 ° M./L.  20 atm. Hg  -d[Hj/dt.  k x 10  x 10 -mole  l i t e r mole  R  liter sec. - 1  0 ,  - 1  sec.  1  5  —1  l o g (k/T)  —1  l i t e r mole sec.  1  °K  V-l  130  31.6  94.8  53-5  6.123  ¥-2  140  73.3  220  111  6".430  V-3  120  13.3  40.0  23.8  7.783  1  (d) contd.  -73-  V-4  110  5.08  V-5  100  2.05  V-6  90  V-7  80,  (e)  15.5  9.18  7.452  6.15  3.88  7.017  0.83  2.47  1.61  8.648  0.33  1.00  0.67  8.289  Effect of Perchloric Acid on the Reaction Rate (Fig. 8). 0.10 M./L. Cu(C10 ) ; 4  2  20 atm. Hg  Expt. HC10. -d[Cr 0 J / d t x 1 0 No X X M./L. mole l i t e r s e c . =  7  -  S-16  -  8  -d[H ]/dt x 1 0 ^X X mole l i t e r sec. &  k x IO ' 5  X X l i t e r mole s e c . -  0.25  3.10  9.30  5.72  S-17 0.50  2.22  6.66  4.07  S-18 0.75  1.88  5.65  3.44  S-19  1.00  1.37  4.10  2.51  T-5  0.10  4.28  12.8  7.76  T-6  0.05  4.72  14.2  8.67  TN-7  0.025  5.13  15.4  9.33  T-8  0.017  5.32  16.0  9.64  T-9  0.013  5.15  15.4  9.25  T-10  0.010  5.17  15.5  9.45  T-ll  0.004  5.23  15.7  9.42  -74(f)  Effect of I n i t i a l Dichromate Concentration on the Rate of the Cupric Acetate Catalyzed Reduction of Dichromate by Hydrogen (Fig. 9). 0.10 M./L. CuAc ;  0.25 M./L. NaAc;  2  0.50 M./L. HAc;  13.6 atm. H ; 2  100«C.  Expt. No.  Initial [Cr 0 ] =  2  7  -d[Cr.O "]/dt x 10 -1 -1 mole l i t e r sec.  6  -d[H ]/dt x 10  k x 10  6  2  mole l i t e r ^ s e c . " *  1  5  l i t e r mole^sec."  1  M./L. C-4  0.005  1.67  5.00  4.65  C-5  0.020  1.55  4.64'  4.34  C-6  0.015  1.52  4.55  4.20  C-7  0.010  1.56  4.67  4.34  is)  Effect of Cupric Acetate Concentration on the Reaction Rate (Figs. 10 and l l )  0.25 M./L. NaAc;  0.50 M./L.HAc; 1 3 . 6 atm. H ;. 2  -d[Cr 0 ] / d t x 10 -1 -1 mole l i t e r sec.  mole l i t e r ^ s e c . " "  0.198  2.89  8.66  4.08  E-2  0.150  2.20  6.60  4.10  E-3  0.102  1.58  4.74  4.34  E-4  0.051  0.81  2.44  4.47  K-2  0.248  3-50  Expt. No.  [CuAc ] M./LT  E-l  =  6  -d[H ]/dt x 10  k x 10  6  2  10.5  100»C.  1  5  l i t e r mole^sec."  3.94  1  -75(h)  Effect of Hydrogen Pressure on the Cupric Acetate Catalyzed Reaction Rate (Pig. 12).  0.25 M./L. NaAc;  0.50 M./L. HAc;  0.10 M./L. CuAc ; 9  100»C. Eixpt. No.  \  -d[Cr 0 "]/dt x 10 -1 mole l i t e r " " sec . 2  ?  1  6  -d[H ]/dt x 10 . , .. -1 -1 mole l i t e r sec.  k x 10  6  l i t e r mole ''"sec.  -1  D-l  27.2  3.08  9.24  4.27  D-2  20.4  2.34  7.03  4.35  B-3  13.6  1.54  4.64  4.31  D-4  7.1  0.81  2.44  4.30  (i)  5  Effect of Temperature on the Cupric Acetate Catalyzed Reaction Rate (Pig. 13).  0.10 M./L. CuAc ; 2  0.25 M./L. NaAc;  0.50 M./L. HAc;  13.6 atm. Hg  l i t e r "^sec."  •-d[H ]/dt x 10 -1 mole l i t e r sec.-1  22.2  66.7  54.2  4.129  Expt. Temp, -d[Cr 0 ~]/dt 5 ' op. No. x 10-mole o  1  5  k x IO  5  -1 l i t e r mole sec.-1  l o g (k/T) -1 l i t e r mole sec.-1 <™--l °K  A-l  130.5  A-2  120.0  8.78  26.3  22.8  ?.763  A-3  110.2  4.03  12.1  10.8.  5.452  A-4  100.5  1.57  4.70  " 4.38  5.070  A-5  90.3  0.65  1.93  1.87  6.712  A-6  80.0  0.22  0.66  0.65  6.266  B-l  140  39.8  B-2  130  16.4  B-3  120.0  8.95  B>-4  100.0  1.53  95-3  4.364  49.2  41.1  4.008  26.8  23.2  5.771  119  4.59  4.27  F.057  (i) contd.  -76-  B-5  80.3  0.23  0.70  0.69  6.292  B-6  90.5  0.67  2.00  1.94  6.728  B-7  109.5  3.42  9.2  5.380  (3)  Effect of pH on the Reaction Rate i n Acetate Solutions (Fig. 14).  10.3  0.05 M./L. CuAc ;; 2  pH  Expt. [HCi0 l or No. [NaOH] M./L. 4  Ac-1  2  =  k x 10  7  2  7  mole l i t e r " -1 sec.  l i t e r sec.  -  100»C.  -d[H J/dt x 10  -d[Cr 0 }/dt x 10 mole  0.25 NaOH  Ac-2  0.90 HAc; 20 atm. H ;  5  l i t e r mole" sec."  1  1  4.25  13.0  38.9  4.68  3-39  11.3  34.0  4.11  2.05  2.32  Ao-4  0.10 HC10. 4 0.05 HC10 4  2.90  8.03  Ac-5  0.15 HC10  4  1.25  0.28  0.84  0.10  Ac-6  0.11 HC10  1.70  0.90  2.68  0.34  Ac-7  0.025 HC10 3.17  10.31  30.9  3-70  Ac-8  0.075 HC10 2.55  5.00  15.0  1.86  (k)  Effect of Various Salts on the Rate of the Cupric Acetate Catalyzed  Ac-3  4  4  4  Reaction (Fig. 16). CuAc  2  (Q series);  6.96  0.88  24.1  3.12  0.05 M./L. CuACg (M,N Series);  0.25 M./L. NaAc;  1  0.10 M./L.  0.50 M./L. HAc; 13.6 atm. H ; g  100»C. Expt. No.  M-l M-2  Salt M./L.  0.20 CaAc  -d[Cr 0 ] / d t x 1 0 , -1 -1 mole l i t e r sec. =  2  6  -d[nj/dt x 10 k x 10 , .. -1 -1 ... , -1 -1 mole l i t e r sec. l i t e r mole sec. 6  5  n  0.822  2.47  4.33  0.857  2.57  4.67  (k) contd.  -77-  N-2-B  0.20 CrAc^  0.773  2.32  4.29  N-5-B  0.20 NiAc  0.755  2.26  4.20  N-4  0.20 NaC10  4  0.836  2.51  4.60  N-5  0.20 Na S0  4  0.690  2.07  3.87  N-6  0.75 NaAc  0.772  2.31  4.31  Q-l  1.50 NaC10  1.77  5.30  4.81  Q-2  1.00 NaC10  1.77  5.30  4.81  0-3  0.50 NaC10  1.70  5.10  4.64  1.64  4.93  4.45  Q-4  2  2  4  4  4  -  Q-5  0.25 Na S0  4  1.44  4.33  3.93  Q-6  0.50 Na S0  4  1.27  3.80  3.45  Q-7  0.75 Na S0  4  1.13  3.39  3.06  Q-8  1.50 NaNO^  1.58  4.75  4.40  Q-9  1.00 NaN0  1.58  4.74  4.39  Q-10  0.50 NaN0  1.59  4.77  4.40  1.60  4.81  4.38  1.45  4.35  4.04  1.55  4.67  4.28  1.53  4.60  4.26  1.60  4.81  4.45  Q-ll  2  2  2  5  5  -  Q-12  0.75 Ca(N0 )  2  Q-13  0.50 Ga(N0 )  2  Q-14  0.25 Ca(N0 )  2  0-15  3  5  5  -  Q-16  0.75 MgS0  1.09  3.25  3.00  0-17  0.50 MgS0  1.24  3.72  3.41  Q-18  0.25 MgS0  1.42  4.26  3.94  0-19  1.00 KNO^  1.59  4.78  4.40  Q-20  1.30  3.90  3.85  Q-21  0.90 NH.N0 4 3 0.33 Cr(N0 )  }  Q-22  0.50 Mg(N0 )  2  4  4  4  7  5  negligible 1.60  4.80  -  4.32  -78-  (k) contd. Q-23  0.50 Ca(C10 )  2  1.52  4.58  4.17  Q-24  0.50 Mg(C10 )  2  1.62  4.85  4.44  Q-25  0.25 Ni(ci0 )  2  1.56  4.69  4.29  Q-26  0,50 (NH ) S0  4  1.10  3.30  3.05  Q-27  0.50 NiS0  4  1.09  3.29  2.98  Q-28  0.50 ZnS0  4  0.99  2.96  2.72  Q-29  0.50 NH Ac  1.17  3.52  3.28  Q-30  0.50 CaACg  1.42  4.25  3.85  (l)  4  4  4  4  2  4  Reaction Rate Measurements from a Typical Experiment Involving the Consecutive Reduction of Dichromate and Cupric Acetate by Hydrogen (Expt. H-2. F i g . 17). I n i t i a l CuACgj  0,01 M./L. I n i t i a l N a g C r ^ ;  0.25 M./L. NaAc;  0.50 M./L. HAc;  0.10 M./L.  13.6 atm. H ; 2  100»C. Substrate Cr 0 ~ 2  -d[Cr 0 ]/dt - mole l i t e r ' =s  2  1  7  -d[CuAc ]/[CuAc ]dt - sec." "  X1  2  2  2  k - l i t e r mole" sec." 1  1  1.60 x IO"  1  ~  4.44 x 10~  2  6  1.07 1.07 x x 10 10" :  4  5.35 x 10  4.77 x 10>  -d[H ]/[CuAc ldt - sec. 2  sec"  CuAc  7  5  4.98 x 10~  5  (m)  Effect of Cupric Propionate Concentration on the Reaction Rate. 0.25 M./L. NaPr;  Expt. [CuPr ]i No. M./L. 2  0.50 HPr; 20 atm. H ; 100»C.  -d[Cr 0 "]/dt x 10~ - 1 - 1 mole l i t e r sec.  6  7  -d[H 3/dt x I O -1 -1 mole l i t e r sec.  k x 10  6  l i t e r mole^sec.  Pr-1  0.10  2.80  8.40  5.38  Pr-8  0.05  1.46  4.39  5.45  Pr-9  0.15  4.22  12.6  5.18  Pr-10  0.20  5.24  15-7  4.88  Pr-13  0.05  1.50  (n)  -1  5.30  4.51  Effect of Temperature on the Rate of the Cupric Propionate Catalyzed Reaction (Fig. 19). HPr;  8 0  0.05 M./L. CuPr ; 2  0.25 M./L. NaPr;  0.50 M./L.  20 atm. Hg.  Expt. Temp. -dtCrgO^ 1/dt No. ' x 10 -mole . .. -1 -1 l i t e r sec.  Pr-8  5  6  100.0  1.46  Pr-11 120.6  8.58  Pr-12 110.0  3.51  Pr-13 100.0  1.50  Pr-14  90.2  Pr-15  80.5  -d[H ]/dt x 1 0 -1 mole l i t e r -1 sec.  4.39  6  k x 10  log (k/T)  5  l i t e r mole -1 sec.  1  l i t e r mole"* -1 -1 sec. °K  1  5.45  5.164  25.7  28.1  5.854  10.6  12.5  5.513  4.51  5.30  5.153  0.60  1.78  2.28  6.799  0.23  0.69  0.92  "6.415  -80(o)  Effect of Sulphate Concentration on the Reaction Rate (Fig. 21). 0.10 M./L. C u ( l l ) ;  20 atm. H ; 110«C.  Expt. [S0 ~]  -d[Cr 0 ~]/dt x 10  M./L.  mole l i t e r s e c . " *  4  2  7  7  -1  1  -d[B ]/dt x 10  k x 10  7  2  mole l i t e r s e c . 1  l i t e r mole sec.  - 1  -1  0.30  2.55  7.65  4.58  0.30  2.56  7.68  4.58  S0 -3 0.28  2.43  7.32  4.36  S0 -4 - 0.18  2.00  6.00  3.58  S0 -5  a  0.08  1.38  4.14  2.47  S0 -6  a  0.03  0.85  .2.54  1.52  S0 -7  a  0.0  0.50  1.48  0.95  S0 -9  b  0.27  2.38  7.15  4.27  SQ -10 0.23  2.15  6.45  3.85  so -n  1.73  5.20  3.11  S0.-12 0.026 4  0.77  2.31  1.29  S0 -13 0.08  1.50  4.50  2.66  S0 -14 0.49  2.83  8.51  5.03  S0 -15° 0.39  2.67  8.00  4.76  S0,-18 0.26 4  2.09  6.25  3.70  so -i  a  S0 -2  a  4  4  a  4  a  4  4  4  4  4  b  4  0.13  b  4  b  S  4  C  4  4  a  a - Ionic Strength Constant at 1.0; C - Ionic Strength  ^ 1.0.  b - [ S 0 ] + [HS0 ~] = 0.35; =  4  4  4  -81(p)  Effect of Temperature on the Rate of the Cupric Sulphate Catalyzed Reaction (Pig. 19). M./L. H S0 J g  Expt. No.  Temp. *C.  -d[Cr 0„ ]/dt 7 x 10—mole , ., -1 -1 l i t e r sec. =  2.09  130.0  13.45  SO.-20 89.8  0.23  4  4  A  (q)  4  0.20 M./L. MgSO^; 0.05  20 atm. Hg.  4  S0 -18 110.0 S0 -19  0.10 M./L. CuS0 »  -d[H J / d t  x 10 ^*1 mole l i t e r -1 sec.  7  k x 10  4  —1 l i t e r mole -1 sec.  3.70  7.985  22.6.  ?.750  0.44  7.083  6.25 40.4  log (k/T) —1 l i t e r mole" -1 -1 sec. °K  0.69  Effect of Cu(ll) Concentration on the Rate of the Cupric Sulphate Catalyzed Reaction (Fig. 22). 0.25 M./L. S0  4  ; 0.10 M./L. HSO^;  20 atm. Hgj 110°C.  [Cu(ll)l -d[Cr 0 "J/dt x I O  Expt. No.  -d[H J / d t  x 10 -1 -1 mole l i t e r sec.  7  7  k x 10 -1 -1 l i t e r mole sec. 4  M./L.  , , !. -1 -1 mole l i t e r sec.  S0 r-16  0.314  5.40  16.2  3.14  S0 -17  0.207  3.78  11.3  3.48  S0 -18  0.103  2.09 .  4  4  4  (r)  3.70  6.25  Effect of Chloride Concentration on the Reaction Rate (Fig. 23). 0.10 M./L. C u ( l l ) ;  Expt, [Cl"J  20 atm. H, ; 2  - d [ C r 0 ] / d t x 10 =  2  -d[H J / d t  x 10 -1 -1 mole l i t e r sec.  k x 10 -1 -1 l i t e r mole sec.  7  4  M./L.  Cl-30  1.00  1.44  4.32  2.57  Cl-31 0.80  1.31  3.93  2.33  o  e  liter^sec.  7  No.  m  i  7  110°C.  1  (r) contd.  -82-  Cl-32  0.60  1.33  3.99  2.37  Cl-33  0.40  1.06  3.18  1.93  Cl-34  0.20  0.93  2.78  1.66  Cl-35  0.21  0.95  2.84  1.67  Cl-39  O.30  1.12  3-36  1.96  Cl-40  0.10  0.83  2.49  1.50  Cl-44  0.20  0.79  2.37  1.42  Cl-47  1.40  1.36  4.08  2.45  (s)  Effect of Cu(ll) Concentration on the Cupric Chloride Catalyzed Reaction Rate (Fig. 24).  Expt. [ C l " l [Cu(ll)] No.  Cl-25 Cl-26 Cl-27 Cl-28  M./L.  M./L,  20 atm. Hgj  -d[Cr 0 ~l/dt 2  „  7  7  „ ,  x 1 0 . mole , ., -1 -1 l i t e r sec. n A  110 C. a  -d[H J/dt x 10 2  ,.4.-l  mole l i t e r -1 sec. T  7  k x 10 ...  4  ,-1  l i t e r mole -1 sec.  -  0  0.07  0.20  it  0.020  0.31  0.93  2.80  n  0.041  0.58  1.75  2.62  n  0.062  0.74  2.23  2.20  1.00  Cl-29  II  0.080  0.96  2.87  2.18  Cl-30  II  0.103  1.44  4.32  2.57  Cl-42  0.20  0.390  2.83  8.50  1.33  0.209  1.63  4.90  1.43  II  0.102  0.79  2.37  1.42  II  0.067  0.53  1.60  1.46  Cl-46  II  0.051  0.46  1.38  1.67  Cl-35  0.21  0.104  0.95  2.84  1.67  01^43 Cl-44 Cl-45  it  (s) contd.  -83-  Cl-36  0.42  0.205  1.94  5.82  1.73  Cl-37  0.63  0.306  2.78  8.35  1.67  Cl-38  0.84  0.408  3.79  11.37  1.70  (t)  Effect of pH on the Reaction Rate i n Cupric Ethylenediamine Solutions (Pig.  25).  0.05 M./L. Cu(C10 ) ; 4  2  0.20 M./L. EDA;  20 atm. BL;  130«C.  Expt. [HC10 ] No. M./L. 4  pH  -d[Cr  0 ~]/dt 7  x 10-mole T4. " I "I l i t e r sec.  -d[B 1/dt x 10 -1 -1 mole l i t e r sec. 7  k x 10  4  l i t e r mole "'"sec.""  1  ED-2  0.38  2.5  1.37  4.11  4.79  ED-3  0.36  3.5  1.35  4.05  4.76  ED-4  0.26  4.1  0.72  2.17  2.56  ED-5  0.20  5.0  0.15  0.45  0.54  ED-6  0.19  6.1  0.16  0.49  0.59  ED-7  0.31  3.9  0.82  2.45  2.87  ED-7-R 0.31  3.9  0.78  2.35  2.72  ED-8  0.23  4.6  0.34  1.02  1.20  ED-9  0*40  1.6  1.39  4.17  4.91  ED-13  0.18  6.1  0.19  0.57  0.66  -84(u)  Effect of Cu(ll) Concentration on the Reaction Rate i n Ethylenediamine Solutions (Pig. 26).  0.20 M./L. EDA; pH 6 (Room Temp.); 20 atm. H ; g  1300C, Expt.  [Cu(C10 ) ] [HC10 1 -d(Cr 0 l / d t =  4  2  4  *  M./L.  M./L.  ED-10  0.025  ED-11  7  -dtHgl/dt x 1 0  k x 10  7  x 10'-mole , .. -1 -1 l i t e r sec.  mole l i t e r " " -1 sec.  0.27  0.105  0.315  0.72  0.095  0.00  0.781  2.34  1-33  ED-12  0.073  0.08  0.232  0.696  0.54  ED-13  0.049  0.18  0.189  0.567  0.66  EIW.7  0.000  0.36  .0.008  0.024  -  H o  4  l i t e r ..mole"* -1 sec.  1  1  »  (v)  Effect of pH on the Reaction Rate i n Cupric Glycinate Solutions (Fig. 0-.05 M./L. Cu(C10 ) ; 0.20 M./L. G l ; 20 atm.-H ; 130°C.  27).  4  pH  Expt. [NaOH] or No. [HC10.] M./L  2  2  -d(Cr 0 ]/dt =  2  7  x 10 -mole -1 -1 l i t e r . sec. 7  4  -d[H J/dt x 1 0 2  mole l i t e r sec.  1  7  k x 10  4  l i t e r mole" -1 sec.  Gl-14  0.10 NaOH  5.1  0.50  1.49  1.66  Sl-15  0.08  4.1  0.89  2.68  2.97  Gl-17  0.02  "  3.1  3»I7  9.51  Gl-20  0.08 "  4.2  0.56  1.69  1.92  Gl-21  0.06 "  3.7  1.25  3.75  4.31  Gl-22  0.04  11  3.4  2.28  6.84  7.55  Gl-23  0.02  11  3.1  3.21  9.63  Gl-24  -  2.9  3.90  11.7  12.7  2.2  4.98  14.9  16.4  11  G1^26 0.10 HC10  4  10.6  10.5  1  (v) contd.  —85™* 1.8  4.18  12.5  13.7  II  2.1  4.98  14.9  15.9  0.20  H.  1.5  3.23  9.7  Gl-30  0.30  II  1.0  1.47  4.41  4.9  G1-3I  0.50  II  0.55  0.68  2.04  2.26  Gl-32  0.03  II  2.6  3.90  11.7  12.5  Gl-33  0.05  II  2.45  4.57  13.7  15.1  Gl-41  0.05 NaOH  3.5  1.56  4.68  5.4  Gl-46  0.10  6.2  0.51  1.53  1.8  Gl-27  0.15 HC10. 4  Gl-28  0.10  Gl-29  (w)  fl  10.7  Effect of Glycine Concentration on the Reaction Rate at Several pi's  (Pig. 28). Expt. [Gill No. M./L.  0.05 M./L. Cu(C10 ) ; 4  4  2  [NaOHj-  0.04 HC10  Gl-12 0.10  0.05  4  "  G3,-13 0.10 0.09 NaOH  7  x 10 -mole i .. -1 -1 l i t e r sec. 7  M./L.  Gl-4 0.05  20 atm. H ; 130»G.  pH -d[Cr 0 "]/dt  [HC10 1 or  2  2  -d[H l/dt x I O 2  mole liter""' -1 sec.  11  7  k x 10  l i t e r mole"* -1 sec.  1  2.0  2.47  2.2  3.40  5.1  0.33  0.98  1.09  5.1  0.50  1.49  1.66  2.1  4.98  7.41 10.2  8.51 11.2  Gl-14 0.20  0.10 "  Gl-28 0.20  0.10 HC10  Gl-30 0.20  0.30  "  1.0  1.47  4.41  . 4.89  Gl-35  0.10  "  1.2  1.04  3.12  3-49  Gl-37 0.05  0.15  "  1.1  1.07  3.21  3-78  Gl-38. 0.10  0.20  "  1.1  1.21  3.63  4.39  Gl-39 0.05  0.03 NaOH  3.5  1.49  4.47  5.10  Gl-40 0.10  0.04  3.6-  1.56  4.68  5.42  0  11  4  14.9  4  15.9  (w) contd) Sl-41 V-l  a  -86-  0.20  0.05 NaOH  0  0.01 HC10  4  a- E a r l i e r Experiment.  3.5  1.56  2.0  3.16  See page 12,  <  4.68  5.36  9.48  5.35  

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