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A kinetic study of the catalytic activation of molecular hydrogen by silver amine complexes Milne, John Buchanan 1960

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A KINETIC STUDY GF THE CATALYTIC ACTIVATION OF MOLECULAR HYDROGEN BY SILVER AMINE COMPLEXES  by JOHN BUCHANAN MILNE B. A., University of British Columbia, 1956  A THESIS SUBMITTED IN PfiRTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF SCIENCE i n the Department of CHEMISTRY  We accept this thesis as conforming to the required standard  THE UNIVERSITY OF BRITISH COLUMBIA April, I960  In p r e s e n t i n g  t h i s t h e s i s i n p a r t i a l f u l f i l m e n t of  the r e q u i r e m e n t s f o r an advanced degree at the  University  o f B r i t i s h Columbia, I agree t h a t the L i b r a r y s h a l l make it  freely  a v a i l a b l e f o r r e f e r e n c e and  agree t h a t p e r m i s s i o n f o r e x t e n s i v e f o r s c h o l a r l y purposes may  study.  I further  c o p y i n g of t h i s  be g r a n t e d by t h e Head of  Department or by h i s r e p r e s e n t a t i v e s .  g a i n s h a l l not be a l l o w e d w i t h o u t my w r i t t e n  The U n i v e r s i t y of B r i t i s h Columbia, Vancouver 3, Canada.  my  I t i s understood  t h a t copying or p u b l i c a t i o n of t h i s t h e s i s f o r  Department of  thesis  financial  permission.  (ii) ABSTRACT  The kinetics of the reduction of silver amine complexes in aqueous solutions were investigated and found to be second order overall, the rate being proportional to silver complex and hydrogen concentrations.  These  systems were studied under conditions of essentially complete complexing and therefore the rates were independent of amine concentration.  The rates were  also shown to be independent of amine perchlorate concentration and thus of pH within a limited range.  Enthalpies and entropies of activation were  determined for each system and an attempt was made to correlate kinetic data with information on complex stability constants and amine basicities.  The  most prominent trend i n the results was the inverse dependence of rate on complex stability constant.  Two mechanisms are proposed both involving  heterolytic cleavage of the hydrogen molecule.  In the f i r s t mechanism, the  proton released i n the rate determining step i s taken up by the basic ligand directly.  In the second mechanism, a water molecule replaces the amine  ligand and acts as the proton acceptor.  The strength of the silver-ligand  bond and the d i f f i c u l t y of ligand replacement by water account for the inverse dependence of rate on complex stability for each mechanism respectively.  Arguments are presented to support both mechanisms.  In general dibasic amine complexes activated hydrogen more readily and displayed a more negative entropy of activation than do the monoamine complexes.  These observations are attributed to the presence of a free basic  group in the ligand aiding the cleavage of the hydrogen molecule i n the rate determining step.  The proximity of the second basic group to the central  silver atom also appears to be important.  These effects and the possible  role of the free basic group in the ligand are discussed,,  ACKNOWLEDGEMENTS  The author wishes to express his gratitude for the advice, help and encouragement given by Dr. J. Halpern for his inspiring direction of the research reported i n this thesis.  The author also wishes to express  his appreciation to Dr. W. A. Bryce for his constructive criticism of the manuscript during i t s preparation.  (iv) TABLE OF CONTENTS Page INTRODUCTION  1  Heterogeneous Catalysis  1  Homogeneous Catalysis  3  EXPERIMENTAL  10  Materials  10  Analysis  10  Procedure  11  RESULTS AND DISCUSSION  H  CONCLUSION  38  REFERENCES  41  (v) TABLES  Table No.  Title  Page  I  EFFECT OF COMPLEXING ON ACTIVITY  II  RATES OF REACTION WITH SILVER TERTIARY DIAMINE COMPLEXES  III  RATES OF REACTION WITH SILVER - DIAMINE COMPLEXES  IV  RATES OF REACTION WITH SILVER - SIMPLE MONOAMINE COMPLEXES  9  20 21  22  V  RATES OF REACTION WITH SILVER - AMINOACID COMPLEXES  24  VI  RATES OF REACTION WITH MISCELLANEOUS SILVER COMPLEXES  25  VII  SUMMARY OF KINETIC DATA, STABILITY CONSTANTS AND BASICITIES RELATIONSHIP BETWEEN RATE OF ACTIVATION AND A FUNCTION OF AMINE BASICITY JND COMPLEX STABILITY  VIII  31 33  (vi) FIGURES  Figure No. 1 2  Title Apparatus for Experiments from 50° C, to 85° C. at Atmospheric Pressure  13  Typical Rate Plots for the Reduction of Silver Complexes of Methyl-, Ethyl-, Propylamine, Ethylenediamine and 1,3-Diaminopropane  15  3  Typical Rate Plots for the Reduction of Silver Complexes of N,N-Dimethylethylenediamine, N,N'-Dimethylpiperazine, Piperazine, Triethylenediamine and N,N,N«,N'-Tetramethylethylenediamine 16  U  Typical Rate Plots for the Reduction of Silver Complexes of Triethylamine, Triethanolamine and Diethylamine  17  Typical Rate Plots for the Reduction ofSilver Complexes of Pyridine and Ammonia  18  Typical Rate Plots for the Reduction of Silver Complexes of Aminoacids  19  Dependence of Rate on Hydrogen Partial Pressure at 65° C. for N,N-Dimethylethylenediamine Complex  27  Dependence of Rate on Hydrogen Partial Pressure at 70° C. for Ethylenediamine Complex  28  5 6 7 8  1  INTRODUCTION  Molecular hydrogen i s a relatively inert substance as indicated by the high endothermicity of i t s uncatalysed reactions.  The chain i n i t -  iating step for the hydrogenation of ethylene;  H  2  + C H 2  4  >- H  +  C H 2  5  has an activation energy of 60-70 ke&U/nole ( l ) , (2) The unreactive #  nature of this molecule has been attributed to the high bond energy of 103 kcal./mole and closed shell electronic configuration.  However, on the sur-  face of a catalyst an alternative reaction path requiring less energy i s provided.  Thus, while the apparent activation energy f o r the homogeneous  hydrogenation of ethylene i s 43 kcal./mole, the reaction takes place on a nickel surface with an apparent activation energy of only 11 kcal./mole (3). The catalytic activation of hydrogen may be conveniently divided into two types; activation taking place on the surface of a solid and activation occurring homogeneously i n solution.  HETEROGENEOUS CATALYSIS Heterogeneous catalytic processes are widely known and extensively used.  However, the mechanism by which hydrogen reacts on a surface with a  lower energy requirement than i n the gas phase remains to be adequately  2 explained.  The activation of hydrogen i s considered to occur through homo»  l y t i c cleavage of the H-H bond with simultaneous formation of covalent bonds with the catalysts  Both crystal dimensions and electronic character of the  metal have been employed i n attempts to explain the heterogeneous activation of hydrogen but with only limited success. Some experimental support for the relationship between activity and crystal dimensions and geometry has been provided by studies on the hydrogenation of ethylene (3) and benzene (4)* (5) on surfaces of transition metals.  A semiquantitative treatment of the variation of activation energy  for chemisorption of hydrogen on carbon with lattice parameters by Sherman and Eyring (6) has lent theoretical support to this approach.  However,  several objections to this point of view have been raised; (a) metals such as Cu and Zn have optimum lattice spacing, yet show no catalytic activity and (b) catalysis on alloy surfaces bears no relationship to geometric factors (7),  (8), (9)* At present crystal dimensions appear to be of  secondary significance and of greater importance i n catalytic activation i s the electronic character of the solid (10). The ability of metals to catalyse hydrogen reactions has been related to incomplete f i l l i n g of the metal d-bands (9).  This i s demonstrated  by the catalytic activity of the transition metals and the reduction i n activity caused by alloying with strong electron-donating metals such as those of group IB where electrons are donated into the d-bands of the transi t i o n metal (9).  Similarly these metal catalysts are poisoned by electron-  donating compounds such as the sulfides.  The paramagnetic susceptibility  of palladium decreases upon absorption of hydrogen indicating f i l l i n g of metal d-bands (11).  This indicates that hydrogen i s activated through the  3 formation of an actual bond involving donation of electrons from the hydrogen into the metal d-bands. Dowden has given a theoretical treatment to the electronic approach (12).  He points out the proportionality between catalytic activity and the  energy density of electron levels at the Fermi surface and the electronic work function of the transition metals.  Some experimental evidence has been  presented to support these ideas but at present quantitative application of the theory i s d i f f i c u l t .  HOMOGENEOUS CATALYSIS Calvin i n 1938 was the f i r s t to demonstrate that hydrogen could be activated homogeneously (13) and since that time numerous similar systems have been discovered and studied (14), (15), (16).  These systems are of  particular interest i n the study of the catalytic process because of their simplicity compared with the heterogeneous systems. Calvin showed that cupric acetate and benzoquinone could be reduced homogeneously i n the presence of cuprous acetate i n quinoline at temperatures of about 100° C.  He reported the process to be between f i r s t  and second order dependent i n Cu(I) concentration and i t was proposed that the active species was a dimer of Gu(l)  (13), (17).  However, further study  of this system has shown the Cu(l) dependence to be exactly second order (18),  (19) and this along with other evidence (20) indicates that a single  termolecular step involving homolytic splitting of hydrogen i s rate determining: 2Gu(l)  + Hg  ^-  2CuH  4 Cuprous acetate has also been shown to activate hydrogen at moderate temperatures i n dodecylamine and pyridine (21) but i n these solvents f i r s t order dependence on Cu(I) concentration was observed.  The reason for the difference  in order of dependence on Cu(l) concentration remains to be adequately explained.  While cupric acetate was not noted to activate hydrogen in these  solvents, this salt was reduced to Cu^O in aqueous solution (22).  Similarly  cupric perchlorate was observed to catalyse reduction of substrates such as dichromate and iodate i n aqueous solution (23). found to be inactive i n aqueous medium (24).  However cuprous salts were  Both Cu(l) and Cu(II)  heptanoates were found to activate hydrogen in heptanoic acid, biphenyl and octadecane (20) at 125° C.  In a l l solvents except quinoline the kinetics  were f i r s t order i n the active copper species and the mechanisms proposed involve the formation of hydride intermediates with heterolytic splitting of hydrogen? Cu(ll)  +  H  2  Cu(l)  +  H  2  CuH >-  GuH  +  +  H  +  +  H  +  The nature of the copper salt and i t s state of complexing affect the rate of hydrogen activation.  This has been attributed to the difference  in basicity of the ligand or anion but other evidence (16) and the results of the present work indicate that the stability of the activating species as defined by the degree to which the d-orbitals are f i l l e d by electrons from the ligand i s of prime importance. acid and ethylenediamlne  Thus ethylenediaminetetra-acetic  inhibit the rate of reaction i n quinoline (25) and  cuprion (2j2'-bisquinoline) appears to have the same effect on the catalytic activity of Cu(l) i n heptanoic acid (20).  A summary of previous work on  5 the effect of complexing Cu(Il) i n aqueous solution and information on complex stability are given i n Table I. The activation of hydrogen by Ag(I) salts has been studied i n aqueous solution (26), pyridine and dodecylamine (21) and heptanoic acid (27). F i r s t order dependence on Ag(I) concentration was observed i n a l l four solvents suggesting a mechanism similar to that for Cu(l) i n pyridine and heptanoic acid involving heterolytic splitting of the hydrogen molecule: Ag  *  +  H  AgH  +  H  +  In aqueous solution a second activation path kinetically second order i n Ag(I) concentration was also observed for which a mechanism similar to that for cuprous acetate i n quinoline has been proposed.  In this mechanism the  hydrogen i s s p l i t homolytically: 2Ag  +  +  Hg  ^  2AgH  +  Where the silver i s activating hydrogen i n a basic ligand environment the bimolecular mechanism i s preferred. Both Hg(l) and Hg(Il) activate hydrogen i n aqueous solution (28), (29) and Hg(ll) undergoes reduction by hydrogen i n both heptanoic acid and biphenyl (27).  These reactions were a l l kinetically f i r s t order i n the mer-  cury salt and the mechanisms proposed are similar to those for Cu(Il) and Ag(I) although thermodynamic considerations indicate that the formation of an hydride intermediate i s unfavourable  Eg*+  +  Hgg*  +  Hg  (30):  Hg  +  2H  +  2Hg  +  2H  +  6 The effects of completing Cu(ll) (31), Hg(ll) (32), (23) and Ag(l) (33) have been studied in aqueous solution and varying rates noted* This work i s summarised i n Table I with data on complex stability.  The  presence of a proton-accepting ligand i n the heterolytic mechanism would tend to accelerate the rate of activation but i n several instances marked slowing of the rate i s observed even with strong basic ligands.  These ef-  fects may be explained on the basis of an activated complex of the typer  n  *" r:~M 1  s-  z +  - X « H " S+  H  where X = ligand, M = metal cation, n = total no. of ligands and Z = total cation charge.  In a complex of this type i t i s apparent that the rate w i l l  depend directly on the strength of the M-H the strength of the M-X bond.  and X-H bonds and inversely on  The strengths of a l l these bonds w i l l depend  on the character of the ligand, X.  Donation of electrons from the ligand  into the d-orbitals of the metal will reduce i t s ability to form the  M-H  bond, a process analogous to poisoning i n heterogeneous catalysis, while increasing the basicity of the ligand should accelerate the reaction by strengthening the X-H bond.  This explains the results of complexing with  Ag(I), Cu(Il), and Hg(II)) where increasing complex stability reduces the rate of reaction while with the weak complexes, increasing ligand basicity accelerates the rate. Several other systems have been noted to activate hydrogen homogeneously or cause deuterium exchange with the solvent* (34) and chlororhodate  Chloropalladate (II)  (II) (35) complexes both activate hydrogen as shown  by the reduction of Fe(III) substrate i n their presence.  The kinetics are  f i r s t order i n each complex. Dicobaltoctacarbonyl has also been demonstrated  7 to activate hydrogen homogeneously (36).  Hydrogen i s apparently s p l i t  homolytically through the mechanism: Cog(GO)  +  8  Hg  *~  -  2HCo(G0)  4  Similarly cobaltous cyanide solutions catalyse the reduction of substrates such as cinnamic acid homogeneously (37).  The species (GN)^Co-GN-Co(CN)^  has been proposed as the catalyst by Winfield (38).  Ethylene platinous  chloride has been observed to catalyse addition of hydrogen to thee ethylene double bond homogeneously at high ethylene concentrations and temperatures below 0° C. (39).  I t has been suggested that the reduction takes place  through the sequence of steps: (PtGl C H^) 2  2  2  2EtCl (C H^) 2  2  2  +  C H^  *  2H2  <  2  > ^  2PtCl (C H ) 2  2  4  (PtCljjCgH^g  2  + 2CgH  6  The oxyenion MnO£ has also been shown to cause hydrogen oxidation (4O) involving a bimolecular rate determining step of the type: MhO£  +  Hg  >-  HMnO~  +  H  +  However, a solution of Ag(I) and MnO^ obeys kinetics of the form: rate = k [H ] [Ag ] [mil] +  2  :  <  a)  B n d  V^  0  ( 4 2 )  » ^ >»  exchange  3  reactions have been observed i n basic medium and the reaction was found to be f i r s t order i n NHg and OH" concentrations respectively.  However, no  reduction of dissolved substrate, CrO^, was observed i n the aqueous solution. The homogeneous activation process discussed above i s unique i n that i t occurs at temperatures f a r below that required for uncatalysed hydrogenations. •.  Severs! other metal cations including C a , Mg  ^4.  » Zn  , Mn  ++  4+  , Ni  , Gd  ++  ,  8 Pb  , Al  , Cr  , TI  , Ce  and CrG^ have been studied up to tempera-  tures of 150° G. and found to be inactive. The purpose of the work on which this thesis i s based was to study further the effect of completing on the rate of hydrogen activation by Ag(l). A series of silver-amine complexes was chosen f o r study since information about the stability of these complexes i s readily available.  I t was hoped  that such a study would provide more information about the configuration of the activated complex and reveal the effects of ligand basicity and cony plex stability on the rate of catalytic activation.  9  TABLE I Effect of Completing on Activity Complex (a)  HgS0  Stability Constant (to) 22  4  Relative Activity 1.8 1.0  Hg~ 2.7 x 10  HgAc2  8  4  x 10"  4  x 1G"  2  HgPr  2  HgCl  1.7 x l ©  1 3  2  2.5 x 1G"  HgBr  1.2 x 1 0  15  2  1.7 x l G "  Hgten)^  2.6 x 1 0  23  1.G x 1G"  AgAc  4.4  Ag(en) Ag  +  3  3  25  7  1  —  Ag(CN)J  6 x 10  CuBag  —  150  GuPi^  —  150  CuAe  4 x 1G  CUSO4  2.2 x 10  GuCl£  15  2  3  80  5 x 1G  2  2  inactive  2 0  120  2  2  2.5  G  _ ++ Gu  MO  CuGlg  4  Cu(en )2  1 x 10  6.5  1 x 1G  15  < 0.5 0.1  (a) Ligand designations: Bo* = butyrate, Pr~ = propionate, Ac" = acetate, Gl" = glycinate, en = ethylenediamine. (b) Kn  = [MXn] / [M]  [xf  (44).  10  EXPERIMENTAL  MATERIALS Silver perchlorate used i n this investigation was a G. F. Smith Reagent grade product. NjNjNSN'-tetramethylethylenediamine, N,N'-dimethylpiperazine, N,N-dimethylethylenediamine and 1,3-diaminopropane were products of K&K Laboratories.  D i s t i l l a t i o n of these products had no effect on reaction  rate and hence they were a l l used without further purification. amine, BDH Reagent grade, was redistilled before use.  Triethyl-  The Houdry Process  Corporation triethylenediamine was used directly, since the rate of Ag(I) reduction was unaffected by using recrystalllzed amine. A l l other amines were Eastman Kodak White Label and except for methyl-, ethyl- and propylamine, were redistilled before use. Reagent grade. Liquid Air Go.  A l l other chemicals were Baker and Adamson  Hydrogen and nitrogen gases were supplied by the Canadian Distilled water was used i n the preparation of a l l solutions.  In the D 0 experiment, recovered D 0 (88.855) was employed, 2  2  ANALYSIS The normality of a l l the liquid amines was determined by direct titration using 0.1N HGIO^ or by back titration using 0.1N NaOH and HCIO^ solutions. tures.  These amines were used directly i n making up the reaction mix-  Aqueous solutions of the amino-acids of known normality were made up  gravimetrically and were used to prepare reaction mixtures. Other solid amines were made up i n aqueous solution and employed to prepare reaction  11 mixtures i n this form. The silver analysis was done by thiocyanate titration using ferric indicator.  To avoid interference due to silver metal and insoluble amine  salts the sample was filtered while warm, then diluted and acidified.  PROCEDURE For rates that could be determined at atmospheric pressure and in the temperature range 50° G. to 85° C,, the glass apparatus depicted i n Figure I. was used.  A presaturator f i l l e d with aqueous solution of the same  ionic strength as the reaction mixture was employed to prevent volume changes i n the mixture.  The whole apparatus was immersed i n a constant temperature  bath thermoatated to within - 0.03° G.  A 250 ml reaction mixture was made up?  from stock solutions and placed i n the reaction vessel.  Nitrogen was run into  the mixture, entering the reaction vessel through a glass sinter at the bottom, for periods of up to an hour.  Samples were taken during this time to establish  the stability of the mixture.  Hydrogen was then introduced and, when the  mixture became saturated with hydrogen, samples were taken at appropriate intervals and analysed for silver content. The uncertainty of the reaction starting time was of the order of a minute during which time the solution was becoming saturated with hydrogen.  A high rate of hydrogen flow was used to  insure complete saturation of the reaction mixture and to provide adequate mixing. For rates of reaction requiring more severe conditions an autoclave was used.  The apparatus used was of Parr manufacture (series 4500) with a  stainless steel reaction vessel provided with a stirrer, gas inlet, sampling  12 tube, pressure gauge and thermowell, surrounded by a 250 watt electric heating mantle controlled by a rheostat.  A glass liner was used In a l l the  present work. Fine temperature control was achieved by use of an auxiliary 5 watt heater placed directly i n the reaction mixture through the thermowell. I t was controlled by a Thermistemp Temperature Controller (model 71) and probe placed i n the reaction mixture through as opening originally provided for a cooling c o i l .  This arrangement gave temperature control of * 0.3° C.  A 500 ml reaction mixture was made up from stock solutions and placed i n the reaction vessel.  Nitrogen gas was run into the mixture at  atmospheric pressure, the vessel was sealed and brought to temperature. Samples were taken over a one hour period to establish the stability of the solution.  One of two procedures was then used to introduce the hydrogen.  With the less volatile amines the solution was flushed with hydrogen and then brought to the desired pressure. With the volatile amines the desired hydrogen partial pressure was established directly and was determined by correcting the total gauge pressure for the contribution made by the nitrogen. Samples were then taken at appropriate times and analysed for silver content* The stirrer was rotated at 600 rpm.  I t has previously been established that  these reactions are independent of stirring rate,,  A. B. C. D. E. F i g . 1.  Gas Outlet Reaction Mixture Sintered Glass Plate Presaturator Solution Gas Inlet  Apparatus f o r Experiments from 50° at Atmospheric Pressure.  u  RESULTS AND DISCUSSION  The stoichiometry of the reduction of silver amine complexes by hydrogen reported i n this thesis i s probably the same as that reported by Webster (26) i n the absence of substrate such as Gr^O^t 2Ag  +  +  H2  >-  2kg  +  2H  +  It was not possible to follow the reaction using CrgO^ i n the basic medium in which the reactions reported here were studied due to the preferential reduction of Ag ». Webster observed that at high temperatures and i n basic + +  solutions the reaction was predominantly f i r s t order i n Ag and i n the perchlorate system f i r s t order dependence on (26), (33 )w  concentration was observed  Similar results are reported i n this work,  The disappearance of Ag  +  at constant  pressure for a l l the silver  amine complexes studied i n the present investigation obeyed f i r s t order kinetics as i s shown by the linear plots given i n Figures 2 - 6 inclusive. This dependence was verified by the fact that the same rate constant was obtained for two different silver concentrations i n the study of each of several complexes.  Results are given i n Tables II - VI inclusive.  In  several instances, notably diethylamine, triethylamine, ammonia, ethylenediamine, 1,3-diaminopropane, ^ - a l a n i n e , methyl-, ethyl- and propylamine, an increase to higher order of Ag proceeded.  dependence was observed as the reaction  This was attributed to autocatalysis an the surface of the  precipitated silver metal.  The onset of autocatalysis could normally be  delayed by buffering the reaction mixture at a lower pH,  In those instances  15  0  1200  2400  3600:'  Time-Seconds Figure 2 Typical Rate Plots for the Reduction of Silver Complexes of Methyl-, Ethyl-, Propylamine, Ethylenediamine and 1,3-Diaminepropane  16 -2.2  A •  o  V o -2.4  Amine Run No. N,N-Dimethylethylenediamine 4f 2f N, N'-Dimethylpiperazine Piperazine  5e  Triethylenedi amine  lg  N,N,N«,N»-Tetramethylethylenediamine  3a  -2.6  -2.8 LOG Ag  +  -3.0  -3.2 2400  1200  Time-seconds Figure 3  Typical Rate Plots for the Reduction of Silver Complexes of N,N-Dimethylethylenediamine, N,N -Dimethylpiperazine, Piperazine, Triethylenediamine and N,N,N ,N'-Tetramethylethylenediamine 1  1  Typical Rate Plots for the Reduction of Silver Complexes of Triethylamine, Triethanol amine and Diethylamine  18  O •  900  1800  Amine Pyridine Ammonia  2700  Run No. 18d 8d  3000  Time-seconds Figure 5 Typical Rate Plots for the Reduction of Silver Complexes of Pyridine and Ammonia  Typical Rate Plots for theReduction of Silver Complexes of Aminoacids  20 Table II Rates of Reaction with Silver-Tertiary Diamine Complexes  Initial AgC10  Eg  4  M.xl0  Triethylenedi amine  2.7 5.5 5.5 5.5 5.5 5.5 5.5 5.5 5.5  0.803 0.803 0.803 0.803 0.803 0.755 0.691 0.879 0.848  0.18 0.18 0.13 0.13 0.09 0.09 0.09 0.09 0.09  0.060 0.060 0.022 0.135 0.030 0.030 0.030 0.030 0.030  60 60 60 60 60 65 70 50 55  N,N«-Dimethylpiperazine  2.7 5.5 5.5 5.5 5.5 5.5 5.5 5.5  0.803 0.803 Gv803 0.803 0.755 0.691 0;879 0.848  0.09 0.09 0.13 0.13 0.09 0.09 0.09 0.09  0.030 0.030 0.045 0.022 0.030 0.030 0.030 0.030  60 60 60 60  5.5 5.5 5.5 5.5 5.5 5.5 5.5  0.691 0.691 0.691 0.755 0.803  0.13 0.18 0.13 0.13 0.13 0.13 0.13  0.045 0.060 0.022 0.045 0.045 0.045 0.045  70 70 70  5.5 5.5 5.5  0.532 0.532 0.532 0.429 0i620 0.691 0.755 1.74 2.75 3.75  0.18  0.060 0.080 0.030 0.060 0.060 0.060 0.060 0.060 0.060 0.060  80 80 80 85 75 70  Amine Name  N,N,N ,N -Tetramethylethylenedi amine f  f  N,N-Dimethyle thylenedi amine  3  5.5  5.5 5.5 5.5 5.5 5.5  5.5  0.848 0.879  Amine M,  Amine Perchlorate M.  Pressure atm.  0.24  0il8 0.18 0.18 0.18 0.18 0.18 0.18 0.18  T°C  65  70  50  55  65  60 55 50  65  65  65 65  h  sec 0.83 0.87 0.81 0.82 0.87 1.22 1.86 0.42 0,62  No. la lb lc Id le If lg lh li  0.40 0.40 0.43 0.43 0.57 0.81 0.17 0.27  2a 2b 2c 2d 2e 2f 2g 2h  1.48 1.44 1.48 1.02 0.72 0.50 0.34  3a 3b 3c 3d 3e 3f 3g  0.72 0.71 0.70 0.95 0.47 0.30 0.22 0.18 0.21 0.18  4a 4b  4c 4d 4e  4f 4g 4h 4i  21 Table III Rates of Reaction with Silver-Pi amine Complexes  Initial AgClO^  T°C  ;-i M -1 sec  No.  Amine Name  M.xlQ  Piperazine  5.5 5.5 5.5 5.5 5.5 5.5 5.5  9.53 9.53 9.53 9.75 9.69 9.62 9.43  0.075 0.150 0.075 0.075 0.075 0.075 0.075  0.150 0.300 0.300 0.150 0.150 0.150 0.150  80 80 80 65 70 75 85  0.097 0.093 0.100 0.031 0.051 0,069 0.135  5a 5b 5c 5d 5e 5f 5g  Ethylenediamine  5.5 5.5 5.5 5.5 5.5 5.5 5.5 5.5 5.5  10.69 10.69 10.69 7.69 5.09 10.75 10.62 10.53 10.43  0.180 0.270 0.180 0.180 0.180 0.180 0.180 0.180 0.180  0.060 0.090 0.180 0.180 0.180 0.180 0.180 0.180 0.180  70 70 70 70 70 65 75 80 85  0.017 0.017 0.016 0.016 0.015 0.010 0.039 0.059  6a 6b 6c 6d 6e 6f 6g 6h 6i  5^5 5.5 5.5 5.5 5.5 5.5 5.5  10.69 10.69 10.69 10.62 10.53 10.43 10.31  0.180 0.270 0.180 0.180 0.180 0.180 0.180  0.180 0.270 0.360 0.360 0.360 0.360 0.360  70 70 70 75 80 85 90  0.0082 0.0090 0.0088 0.014 0.020 0.029 0.045  7a 7b 7c 7d 7e 7f 7g  1,3-diaminopropane  3  Amine  Amine Perchlorate  Pressure atnu  Hi  0.024  22 Table 17 Rates of Reaction with Silver-Simple Monoamine Complexes Initial AgClO^  H  2  Amine Perehlorate M.  T°C  sec 0.0034  M.X10  Pressure atau  Amine  5.5 5.5 5.5 5.5 5.5 5.5 5.5  31.31 31.31 31.31 31.43 31.17 31.00 30.85  0.045 0.090 0.045 0.045 0.045 0.045 0.045  0.360 0.720 0.720 0.360 0.360 0.360 0.360  90 90 90 85 95 100 104  0.0034 0.0021 0.0061 0.0096 0.0174  8a 8b 8c 8d 8e 8f 8g  Methylamine  5.5 5.5 5.5 5.5 5.5 5.5 5.5  16.43 16.43 16.43 16.62 16.53 16.31 16.17  0.045 0.090 0.045 0.045 0.045 0.045 0.045  0.360 0.720 0.720 0.360 0.360 0.360 0.360  85 85 85 75 80 90 95  0.054 0.051 0.051 0.022 0.041 0.086 0.121  9a 9b 9c 9d 9e 9f 9g  Ethylamine  5.5 5.5 5.5 5.5 5.5 5.5 5.5  16.43 16.43 16.43 16.69 16.62 16.53 16.31  0.090 0.180 0.090 0.090 0.090 0.090 0.090  0.360 0.720 0.720 0.720 0.720 0.720 0.720  85 85 85 70 75 80 90  0.015 0.016 0.016  10a  5.5 5.5 5.5 5.5 5.5 5.5 5.5  16.62 16.62 16.62 16.80 16.75 16.69 16.53  0.045 0.090 0.045 0.045 0.045 0.045 0.045  0.360 0.720 0.720 0.360 0.360 0.360 0.360  75 75 75  60  11.0 11.0 11.0 11.0 11.0 11.0 11.0  10.69 10.69 10.69 10.80 10.75 10.62 10.53  0.180 0.270 0.180 0.180 0.180 0.180 0.180  0.180 0.270 0.360 0.180 0.180 0.180 0.180  Amine Name Ammonia  Propylamine  Diethylamine  3  0.0035  Mo.  0.0059 0.0095 0.022  10b 10c lOd lOe lOf lOg  65 70 80  0.0085 0.0084 0.0085 0.0020 0.0032 0.0048 0.0154  11a lib 11c lid lie llf Hg  70 70 70 60 65 75 80  0.044 0.044 0.041 0.019 0.027 0.067 0.114  12a 12b 12c 12d 12e 12f 12g  0.0040  23 Table IV (continued)  Initial AgClO^ Amine Heme Triethylamine  M.xl0  3  11.0 11.0 11.0 11.0 11.0. 11.0 11. G11.0*  Pressure atm.  Amine  7.92 7.92 7.92 7.96 7.94 7.90 7.88 7.90  0.180 0.270 0.270 0.270 0.270 0.270 0.270 0.270  t - rate determined i n 88.8$ D 0 2  M,  Amine Perchlorate M.  T°C  0.360 0.540 0.270 0.540 0.540 0.540  40 40 40 30  0.540  45  0.540  i  k  35 45  50  r  -i  sec  0.053 0.049 0.053 0.017 0.028 0.079 0.130 0.079  No. 13a 13b 13c 13d 13e 13f 13g 13h  24 Table V Rates of Reaction with Silver-Aminoacid Complexes  Initial AgClO^ Amino-Acid Name  M.X10  Glycine  5.5 5.5 5.5 5.5 5.5 5.5 11.0 11.0 11.0 5.5 5.5 5.5 5.5 5.5  o<-Alanine  ^3 -Alanine  3  Pressure atnu  Amine -acid Mi  9.53 9.80 9.75 9.69 9.62 9.80 9.80 9.80 9.80 '9.8G0 19.80 19;80 19.80  0.107 0.214 0.107 0.107 0.107 0.107 0.107 0.107 0.107 0.107 0.107 0.107 0.107 0.107  0.107 0.214 0.107 0.107 0.107 0.107 0.107 0.240 0.190 0.130 0.090 0.090  5.5 5.5 5.5 5.5 5.5 5.5  9.62 9.62 9.75 9.69 9.53 9.43  0.120 0.240 0il20 0.120 0.120 0.120  0.120 0.240 0.120 0.120 0.120 0.120  5.5 5.5 5.5 5.5 5.5 5.5  9.62 9.62 9.75 9.69 9.53 9.43  0.120 0.240 0.120 0.120 0.120 0.120  0.120 0.240 0.120 0.120 0.120 0.120  9.53  -l -l sec k  NaOH M,  0.065  0.050  M  80 80  60 65  70 75 60 60  60  60  60  60 60 60 75 75  65  70 80 85 75 75  65  70 80 85  No,  0.104 0.106 0.0176 0.030 0.048 0.076 0.0178 0.161 0.127 0.050 0.0084 0.0086 0.0055 0.0051  14a 14b 14c 14d 14e 14f 14g l4h 14i 143 14k 141 14m 14n  0.0305 0.0309 0.0173 0.0250 0.0491 0.0699  15a 15b 15c 15d 15e 15f  0.040 0.039 0.015 0.029 0.051 0.075  16a 16b 16c I6d I6e I6f  25 Table VI Rates of Reaction with Miscellaneous Silver Complexes Initial AgC10 4  Amine Name  M.xlG  3  *2 Pressure atm.  Amine  Amine Perchlorate  k  T°G  Triethanolamine  5.5 5.5 5.5 5.5 5.5 5.5 5.5  0.691 0.691 0.691 0.803 0.755 0.620 0.532  0.180 0.225 0.180 0.180 0.180 0.180 0.180  0.720 0.900 0.360 0.720 0.720 0.720 0.720  70 70 70 60 65 75 80  Pyridine  5.5 5.5 5.5 5.5 5.5 5.5 5.5  9.53 9.53 9.53 9.43 9.31 9.17 9.00  0.270 0.540 0.270 0.270 0.270 0.270 0.270  0.270 0.540 0.135 0.135 0.135 0.135 0.135  30 80 80 85 90 95 100  i  sec  No.  0.47 0.45 0.49 0.19 0.30 0.68 0.96  17a 17b 17c 17d 17e 17f 17g  0.019 0.018 0.020 0.029 0.043 0.060 0.093  18a 18b 18c 18d 18e 18f 18g  26 i n which antocatalysis interfered to a large extent, reproducible rates could be determined from the f i r s t 5% of reaction. As i n the acid perchlorate medium, the rate was directly dependent on hydrogen concentration as i s shown i n Figures 7 and 8.  At different  hydrogen pressures a constant k' was also obtained for glycine complexes (Table V),  Thus, the overall rate law may be expressed ast  dfe! = -£  iL&g!3  dt  dt  =  "k'fAg+JpgOC ^  =  kJAg+^E,]  where p ^ i s the hydrogen pressure and <X i s Henry's constant. region of pressures investigated, Henry's law i s obeyed.  In the  The values taken  for hydrogen solubility at different temperatures were those of Weibe and Gaddy (46),  I t was assumed that the hydrogen solubilities i n the solutions  used i n the present study were, within the limits of experimental error, the same as for pure water. These kinetics suggest a mechanism of the types + A g  . « AgH  +  . +  ^ . + Ag  S l 0 w  >  fast  AgH  +  H  „+ >- H  +  2Ag  +  which has been justified on thermodynamic grounds (30). Except for triethylamine, a l l experiments were made under conditions i n which the silver was completely complexed by the amine. Limitations due to the insolubility of triethylamine i n water allowed only 95% complexLng i n this case but, since the rate with this amine i s 1000 times that for uncomplexed silver, the error i n true rate for triethylamine complex i s small.  Figure 7 Dependence of Rate on Hydrogen Partial Pressure at 65° C. for N,N-Dimethylethylenediamine Complex  H  2  Pressure-ATM. Figure 8  Dependence of Rate on Hydrogen Partial Pressure at 70° C. for Ethylenediamine Complex  29 That the rate i s independent of amine concentration i s shown i n Tables II-VI.  A l l the reaction mixtures were buffered with an amine-amine per-  chlorate system to prevent large changes i n pH. The rate was shown to be independent of amine perchlorate concentration i n a limited range. shown i n Tables II-VI.  This i s  It i s consequently independent of pH within the fiH  change caused by the protons liberated i n the complete reduction of the Ag . +  The agreement between rates obtained i n the glass apparatus and those determined i n the autoclave i s shown In the results for N,N-dimethylethylenediamine (Table II). The activity of the silver-glycinate complex was studied i n solutions containing excess base (NaOH) and excess glycine (Table V, runs 14g-14n inclusive).  The reaction mixtures containing excess base were  unstable, silver being reduced i n the absence of hydrogen.  However, by sub-  tracting this background rate from the observed rate with hydrogen, semiquantitative rate constants could be determined. with increasing OH  The large increase i n rate  concentration has been observed before for silver ethyl-  enediamine and mercury (II) ethylenediamine complexes (26), (32) and may be explained i n one of two ways. The OH" may participate directly i n the rate determining step: +Ag(H NCH C00~) 2  2  + OH"  2  + Eg  AgH + HgO  + 2H2NCH2G00"  or an equilibrium of the type observed with Co amine complexes by Pearson and Basolo (47) may be responsible: + 00GCH IH AgNH CH C00 2  2  2  2  + OH  *-  +" OOCCHjNEjAgNNHCBgCOO  +  H0 2  30 where "NHCHgCOO" i s a much stronger base than EjNCBgCOO".  The decrease i n  rate i n the region of constant total glycine plus glycinate concentration with decreasing glycinate concentration i s anomalous i n as much as sufficient glycinate i s present to complex f u l l y the silver i n a l l cases and the rate has been shown to be independent of glycinate concentration i n this region. This effect remains to be explained. A 10,000-fold variation i n rate for the silver complexes studied —Z. —1 _ i + was observed, ranging from 3.3 x 10 l.moles sec. for Ag(NH ) to 1.8A 3  £  Lmoles'^sec."^ for the triethylenediamine complex; at 70°C. A complete -  resume' of thermodynamic values, complex stability constants, amine basicities and rates i s given i n Table VII.  In general the differences i n rates  are reflected primarily in the activation energies.  The activation entropies  are normal for the simple bimolecular mechanism suggested, although with the exception of triethanolamine, pyridine and glycine, those for the dibasic ligands are somewhat more negative (-7 to -19 eu.) than those for the monoamines (7 to -6 eu.).  These differences suggest differences i n the rate  determining step for the dibasic amines and the monobasic amines. Table VII contains evidence of two different trends: (a) a decrease i n catalytic activity with increasing complex stability and decreasing amine basicity and (b) a generally greater activity for the dibasic amines than the monobasic.  These two trends w i l l be discussed separately i n this order.  On the basis of the rate determining step:  +  HjNAgNR-j + Eg  ^  AgH  + R^NH  +  + R^N  31  Table VII Stnnmary of Kinetle Data Stability Constants and Basicities Amine , » 4H* B a s i c i t y Stability C o n s t a n t s ^ kcal. log K log K mole Pa  AS* eu. log k(70 )  u ;  Amine Name  K  Ammonia Methylamine Ethylamine Propylamine Diethylamine Triethylamine  x  9.25  2  3.78 3.53 3.93  3.37 3.15 3.37  10.72 10.61 10.58 10.98 10.77  -  -  -  7.15 6.68 7.30 7.39 6.36 4.76  28.3 22.3 21.6 23.3 20.5 18.8  +7.6 -2.8 -7.1 -1.5 -5.4 -4.8  -3.48 -1.91 -2.42 -2.26 -1.33 -0.13  -6.7  -1.80  Ethylenediamine 1,3-Diaminopropane Piperazine N,N-Dimethylethylenedi amine N,N,N',N'-Tetramethylethylenedi amine N,N'-Dimethylpiperazine Triethylenedi amine  10.18  4.62  2.92  7.54  20.8  10.72 9.81  5.77 3.32  --  -  18.9 16.8  -13 -16  -2.00 -1.32  -  -  17.1  -11  -0.49  -  -  15.1  -14  0.17  Glycine ex -Alanine /? -Alanine  9.78 9.87 10.19  3.51 3.64  7.90 5.45  2.30 2.04  9.53  -  9.30  -  8.30 8.19  1.57  Triethanolamine Pyridine  (a) K a  (b) K  x  = [ AIH*) , [AH ]  -  -  -  -  -7.2. -8.1  -0.07 0.26  -  17.9 17.1  3.38 3.54  6.89 7.18  -  21.5 16.2 18.3  -2.5 -19 -13  -1.33 -1.61 -1.63  1.34 2.18  3.64 4.22  17.5 19.3  -9.7 -12  -0.36 -2.09  -  Values taken from (45)  +  =  [AgNR^]  , K  2  =  [Ag ][NR "] +  3  [Ag(NR3) '] , /3  2  2  = 1 2 Values taken from K  K  [Ag NR ][NR ]  ^  +  3  3  corrected where necessary to 25° G., using temperature coefficient for Ag(NH C H )2 2  2  5  =  d log % / d t  =  d log Kg/dt  = -0.016  (48)  32: an activated complex of the type:  + R3N - Ag - NR3 %- H - H£+ may be postulated for the monoamines (R may represent an alkyl group or H). On the basis of such a configuration the rate of hydrogen activation by the complex should depend directly on the strength of the Ag-H and H-NR^ bonds and inversely on the strength of the Ag-NR^ bond. pK  a  Therefore, since the  i s a measure of the strength of the H-NR3 bond and log Kg a measure of  the strength of the Ag-NR^ bond being cleaved, one might expect a correlation between the rate of activation of hydrogen and a function involving these parameters of the type: log k  CX  (pK  a  - log Kg)  A comparison of log k at 70° C. and the function (pK - log Kg) for simple a  monoamines i s given in Table VIII and the dependence i s clearly shown. Pyridine and triethanolamine are monoamines but their characters are d i f ferent from that of the simple monoamine and they should be considered i n their own homologous series. The diamines follow a similar pattern (Table VIII) but because for most of these amines only information on the f i r s t stability constant i s known, log K^ i s used.  The relative correlation between the f i r s t  stability constant and the overall stability constant i s good and the use of log K^ should not detract from the relationship drawn. For the dibasic amines i t was f e l t appropriate to use the logarithm of the f i r s t basicity  Table T i l l Relationship between Rate of Activation and a Function of Amine Basicity and Complex Stability  Monoamine Name  Function (pK - log Kg) ft  Ammonia Ethylamine Methylamine Diethylamine Triethylamine  5.47 6.68 7.19 7.68 8.61  Diamine Name  Function (pK - log K-^)  1,3-Diaminopropane Ethylenediamine Piperazine Triethylenediamine  a  4.95 5.56 6.49 7.41  log k -3.48 -2.42 -1.91 -1.33 -0.13  log k -2.00 -1.80 -1.32 0.26  34 constant in as much as the ligand accepts one proton i n the rate determining step. (pK  a  Table VIII shows the relationship between log k and  - Log K ) . x  In the dibasic amines (including amino-acid anions) the presence of the second basic group may increase the rate.  A comparison of the rates  of ethylamine and ethylenediamine complexes and of diethylamine and piperazine complexes shows that the rate for the diamines i s greater. * Whether this i s due to ligand basicity and complex stability effects or the presence of the second basic group i s not clear. second basic group appears to be important.  The proximity of the  The logarithm of the stability  constant for the Gu(II) glycine complex i s greater than that for the Cu(Il) -alanine complex.  This relationship i s l i k e l y to hold for the Ag(l)  complexes as well and i f this i s the case the faster rate with the glycine complex may be attributed to the proximity of the carboxy1 group. The tendency for chelation i n silver complexes i s small.  In the  solutions studied the silver complex would carry two monocoordinated dibasic  + ligands e.g. HgNCHgCHgNHgAgHgNCHgCHgNHg.  A rate determining step involving  direct acceptance of the proton by the non-bonded basic group i n the ligand suggests an activated complex of the form:  +  HgNCRgC^NHgAgBgNCEj  The greater negative entropy change noted above for the diamine complexes may be indicative of an activated complex different from that suggested for the monoamine complexes.  The acceptance of the released proton at a different  site i n the activated complex for the diamine mechanism may require a greater  35 degree of solvation for these complexes than for the monoamine complexes. This greater degree of 'freezing out' of solvent molecules for the diamine complexes may be the cause of the greater negative entropy change. A second configuration deserving consideration i s : CHgCHgm,  NEg - Ag - HgNCHgGHgNHg S- H - H S+ However, i f the formation of the Ag-H bond i s dependent upon vacant dorbitals i n the silver as has been suggested earlier (16), coordination with the second amino group of one ligand would be expected to hinder this process and this type of activated complex would be unlikely. A l l previous discussion has depended upon a mechanism i n which the ligand accepts the proton i n the rate determining step directly. However, certain general considerations and evidence would suggest a mechanism in which water acts as the proton acceptor or as a bridge between the proton and the ligand.  This would give rise to an activated complex of the  type:  + R3N - Ag  H 1 0-H - H£ +  S-H A similar configuration may be postulated for the dibasic complexes. In as much as the electrons responsible for the basicity of an amine would be donated to the silver i n the formation of the covalent bond i n the complex  36 i t would appear d i f f i c u l t for the ligand amine to accept a proton i n the rate determining step.  If this were true, i t seems reasonable that water  as a ligand may act more readily as the proton acceptor i n that i t has two coordinative positions.  Evidence which would favour such a mechanism i s  the possible proton accepting role of OH", CO", acetate and propionate suggested by Korinek to explain the acceleration of hydrogen activation noted with Hg(II)-ethylenediamine complexes when these anions were added (32). Support for this point of view i s found i n the measurement of rates of proton transfer for methylamines and ammonia by NMR techniques. Meiboom and co-workers by observing the broadening of the water and amine proton signals have been able to distinguish and measure the rates of the two separate paths of proton transfer; one employing an intervening water molecule end one by direct transfers + R-jNH R NH 3  +  +  l  k  +  0-H H  +  NRj NR  3  —  =>^  R^N R^N  + +  + H NR^ H-0 H  + H NR +  3  They have observed that the second path i s preferred by the methylamines (kj/k^  =  1.4, 14 and  >10  for methylamine, dimethylamine and trimethyl-  amine respectively) but not by ammonia (^/^  *  0^085) (49).  The above considerations explain the observed variation i n rates with different complexes.  A water mediated mechanism seems plausible for  the alky!amine complexes at least and the observed inverse variation of rate with complex stability constant may be due to the difficulty of replace-  37 merit of the ligand by a water molecule i n the activated complex.  The  abnormally low rate constant for ammonia complex may be the result of direct proton transfer to the ligand i n preference to the water mediated mechanism. Such a direct transfer would account for the anomalously high activation energy and positive entropy change for this complex i n that the ligand removal would not be compensated for by replacement with a water molecule. In order to determine whether the direct transfer mechanism was i n operation for a l l the complexes or whether a water mediated mechanism was used, an experiment was made i n DgO with the triethylamine complex which should activate hydrogen by the water mediated mechanism i f these ideas are correct.  It would be expected that i f a water mediated mechanism  were i n operation the difference i n basicity of water and DgO would be reflected i n the rate.  No such difference i n rate was observed (Table IV)  which would appear to substantiate the mechanism suggested earlier of direct proton transfer to the ligand.  However, i n as much as i t i s not  clear whether the difference i n basicity of the solvents i s the only effect, this evidence i s not conclusive.  38 CONCLUSION  The rate of activation of hydrogen by silver amine complexes was observed to be f i r s t order dependent on both complex and hydrogen concentration.  This i s i n agreement with Webster's observations on the silver  system at high temperatures and i n the presence of basic ligands.  Hetero-  l y t i c splitting of the hydrogen molecule i s suggested in the rate determining step (26), (33).  The rate of activation was seen to vary inversely with  complex stability i n a manner similar to that reported for Cu(II), Hg(ll) and some silver complexes (31), (32), (33).  Two possible mechanisms can be  used to explain these results; the f i r s t involves direct transfer of the proton released i n the rate determining steps Hj  + Ag(R N) —*-R N 3  2  3  - Ag - NR^—=>- AgH f-H  + R NH 3  +  +  R^  - Es-  In the other water acts as the main proton acceptor i n a l l cases except that of the ammonia eomplexs '  +  ?  •  Hj + Ag(R N) + B^O-^RjN - Ag - 0-H 3  2  g-E - Hg  N R ^ R N + AgH + R NH 3  3  3  + IL^O  +  The f i r s t mechanism finds support i n the reasonable correlation drawn between the rate of reaction and the function (pK - logK^) for the ft  simple monoamines. Since amine basicity should be a measure of the strength of the incipient H - NR bond and the complex stability a measure of the 3  strength of the Ag - NR bond, the reaction rate should increase with i n 3  creasing magnitude of this function.  In addition the rate of activation by  silver - triethylamine complex i n D 0 and HjO i s the same suggesting that 2  39 solvent molecules are not involved i n the rate determining step and that the proton i s accepted directly by the ligand amine.  However, this evidence  i s not conclusive since i t i s not clear whether the use of Dg  0  ^  Place of  HgO would introduce only a solvent basicity difference, A water mediated mechanism appears plausible i n that the ability of a ligand such as an amine with i t s one coordinative position bonding to silver would be hindered from accepting a proton,  A water ligand having one  free coordinative position would be able to accept the proton i n the rate determining step.  The inverse dependence of rate on complex stability would  reflect the d i f f i c u l t y of replacing the amine ligand by a water molecule, A recent study of proton transfer between amines i n aqueous medium indicates two mechanismsi the f i r s t , preferred by the alkylamlnes, involves proton transfer through an intervening water molecule.  The second, used by ammonia,  involves transfer of the proton directly between ammonia molecules.  The  anomalously low rate observed f o r the silver-ammonia complex (slower than for the water complexed species) compared to those rates observed for the other monoamines may be explained by the d i f f i c u l t y i n effecting a direct proton transfer to the ammonia ligand i n the rate determining step.  The  abnormally high activation energy and positive entropy change for the ammonia complex appears reasonable i n that, unlike the alkylamine complexes, release of the proton accepting ligand would be uncompensated by coordination of a water molecule. The dibasic complexes displayed a greater activity and were observed to have a greater negative entropy change than the monoamine complexes.  This suggests an accelerating effect due to the presence of the  40 second basic group possibly by accepting the proton released or by coordinating with the silver i n the position vacated by the ligand i n the rate determining step.  proton-accepting  The participation of the free basic  group in the activation i s further substantiated by the fact that the rate decreases as the second basic group becomes more distant from the silver atom i n the complex.  A water mediated mechanism may occur with these com-  plexes as well. It would appear reasonable that the presence of a basic ligand i s responsible for the generally faster rates observed with silver amine complexes than the water coordinated silver species.  Unlike similar Cu(Il)  and Hg(ll) complexes, donation of electrons from the ligand into the silver d-orbitals does not have a great effect on the formation of a bond between hydrogen and the activating species.  41  REFERENCES  1.  Rice, F. 0. and Herzfeld, K.F., J. Am,Chem. Soc.,  j>6, 284 (1934).  2.  Varnerin, R. E. and Dooling, J. S., J. Am. Chem. Soc. 28, 1119 (1956).  3.  Beeck, 0., Rev. Mod. Phys., 12, 61 (1945).  4.  Balandin, A. A., Zeit.fur physik. Chem., 2B, 289 (1929), ibid.3B. 167 (1929).  5.  Beeck, G. and Ritchie, A. W., Disc. Farad. S o c , 8, 159 (1950).  6.  Sherman, A. and Eyring, H., J.Am.Chem.Soc., 5Jt» 2661 (1932).  7.  Reynolds, P. W., J.Chem.Soc, 265 (1950).  8.  Dowden, D. A. and Reynolds, P. W., Disc.Farad.Soc., 8, 184 (1950).  9.  Couper, A. and Eley, D. 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