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Carbon monoxide reduction of aqueous silver acetate McAndrew, Robert Thomson 1962

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CARBON .MONOXIDE REDUCTION OF AQUEOUS SILVER ACETATE  by  ROBERT THOMSON McANDREW B . S c , Queen's U n i v e r s i t y , 1957 M . S c , Queen s U n i v e r s i t y , I958 l;  A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHY  i n t h e Department of METALLURGY  We a c c e p t t h i s t h e s i s as conforming required standard  t o the  THE UNIVERSITY OF BRITISH COLUMBIA August,  1962  In presenting  t h i s thesis i n p a r t i a l f u l f i l m e n t of  the requirements f o r an advanced degree at the University of B r i t i s h Columbia, I agree that the Library s h a l l make i t f r e e l y available f o r reference and study.  I further agree that permission  f o r extensive copying of t h i s thesis f o r scholarly purposes may granted by the Head of my Department or by his  be  representatives.  It i s understood that copying or publication of t h i s thesis f o r f i n a n c i a l gain s h a l l not be allowed without my written permission.  Department of  Metallurgy  The University of B r i t i s h Columbia, Vancouver 8, Canada. Date  September 10,  1962  The University of B r i t i s h Columbia FACULTY OF GRADUATE STUDIES  PROGRAMME OF THE FINAL ORAL EXAMINATION FOR THE DEGREE OF DOCTOR OF PHILOSOPHY oof ROBERT THOMSON McANDREW B . S c , Queen's University 1957 M.Sc., Queen's University 1958 MONDAY, SEPTEMBER 10, I962 AT 10:00 A.M. IN ROOM 201; MINING BUILDING'  COMMITTEE IN CHARGE Chairman: F.H. SOWARD W.M..ARMSTRONG W.A. BRYCE D.L.G. JAMES  E. PETERS C.S. SAMIS E. TEGHTSOONIAN  External Examiner: W.K. WILMARTH, University of Southern C a l i f o r n i a , Los Angeles  CARBON MONOXIDE REDUCTION OF AQUEOUS SILVER ACETATE  GRADUATE STUDIES  ABSTRACT  F i e l d of Study: Metallurgy  The k i n e t i c s of the carbon monoxide reduction of s i l v e r perchlorate i n sodium acetate - acetic acid buffered aqueous .solution were studied between 60° and 110°C. by following the pressure decrease i n a closed system. The reduction occurs homogeneously i n the l i q u i d phase'by two p a r a l l e l reaction paths, one of which i s independent of pH. The second path i s favoured by increased pH and has both an acetateindependent and an acetate-dependent component. The observed k i n e t i c s are consistent with the formation of intermediate complexes by the i n s e r t i o n of a carbon monoxide molecule between a s i l v e r ion and a co-ordinated oxygen-donating base (e.g. OAc", ILpO) according to the following mechanism: Ag  Metallurgical Kinetics  E. Peters  Metallurgy of the Rarer Metals  E. Peters  Hydrometallurgy  '.  Nuclear Metallurgy.  Staff . . W.M.  Armstrong  Other Studies:  + OAc" -—* AgOAc (rapid equilibrium) 0 AgOAc + CO Ag-C-OAc (slow) 0 Ag-fi-OAc + Ag + H 0 2Ag + CO + HOAc + H (fast) 0 Ag + CO + H 0 ^==i Ag-C'-OH + H (rapid equilibrium) 0 Ag-C-OH + Ag ^ 2Ag + C0 + H (slow) 0 * Ag-C-OH + AgOAc 2Ag + COg + HOAc (slow) +  +  C'~S. ,Samis  M e t a l l u r g i c a l Thermodynamics.  S t a t i s t i c a l Mechanics  L.G.,-Harrison  Surface Chemistry  L.G.  Harrison  Chemical K i n e t i c s . . . J . H a l p e r n andCG.B..Porter Electronics  R.D.  Russell  +  2  %  D i f f e r e n t i a l Equations.,..-  +  F.M.C. Goodspeed  +  2  +  +  2  Silver-acetate' complexes are about a factor of three more reactive than hydrated s i l v e r ions i n the pHdependeht reaction. This enhanced r e a c t i v i t y i s a t t ributed to s t a b i l i z a t i o n by the basic acetate anion of the proton released i n the reduction process. The e f f e c t of increasing pH on the reduction rate i s much greater than the s p e c i f i c e f f e c t s associated with s i l v e r - a c e t a t e complexing.  PUBLICATIONS 1.  McAndrew, R.T., and Peters, E., "The Displacement of S i l v e r from Acid Solutionrby. Carbon Monoxide", XVIIIth Congress, International Union of Pure and Applied Chemistry, Montreal, 1961.  i ABSTRACT The kinetics of the carbon monoxide reduction of silver perchlorate in sodium acetate - acetic acid buffered aqueous solution were studied between 60 and 110°C by following the pressure decrease in a closed system. The reduction occurs homogeneously in the liquid phase by two parallel reaction paths, one of which is independent of pH. The second path is favoured by increased pH and has both an acetate-independent and an acetate-dependent component. The observed kinetics are consistent with the formation of intermediate complexes by the insertion of a carbon monoxide molecule between a silver ion and.a co-ordinated oxygen-donating base (e.g. OAc", H 0) according 2  to the following mechanism: Ag  +  + OAc ^=^ AgOAc -  AgOAc + C O — ^  (rapid equilibrium)  >AgiLoAc  0 Ag-fl-OAc + Ag+ + H 0  (slow)  >2Ag + C0 + HOAc + H  2  2  +  (fast)  p Ag  +  + CO + H 0 ^ 2  0 Ag-<3-0H + Ag  +  > Ag-d -OH + H  +  ^b_^2Ag + C0 + H 2  (rapid equilibrium) +  (slow)  Ag-3-OH + AgOAc — ^ > 2 A g + C0 + HOAc 2  (slow)  Silver-acetate complexes are about a factor of three more reactive than hydrated silver ions in the pH-dependent reaction.  This enhanced  reactivity is attributed to stabilization by the basic acetate anion of the proton released in the reduction process. The effect of increased pH on the reduction rate is much greater than the specific effects associated with silver-acetate complexing.  ACKNOWLEDGMENTS  I am deeply g r a t e f u l t o Dr. E. Peters for his i n s p i r i n g d i r e c t i o n of t h i s i n v e s t i g a t i o n .  Thanks are also extended t o A. M. Armstrong f o r her  thoughtful and constructive c r i t i c i s m during the preparation of the  manuscript.  F i n a n c i a l support from the National Research Council of Canada i n the form of grants i n a i d of research and a Studentship, and from the Consolidated Mining and Smelting Company i n the form of a Fellowship, i s greatly appreciated. I sincerely thank my wife f o r her encouragement and help throughout the period of study. Her e f f o r t s are responsible f o r converting my semi-legible scrawl t o readable typed copy.  iii TABLE OF CONTENTS Page I  INTRODUCTION I-l 1-2 1-3 1-4 I -5  II  1 2 5 6 10  EXPERIMENTAL I I -1 II-2 . II-3 II -k II-5 I I -6  III  General Thermodynamic Considerations of Carbon Monoxide Reactions i n Aqueous Solution Structure of Carbon Monoxide and i t s Compounds K i n e t i c s of Metal Ion Reduction by Carbon Monoxide in Aqueous Solution Object and Scope of the Present Investigation  Reactor System Pressure Measurement Temperature Control and Measurement Materials Chemical Analysis General Experimental Procedure  12 12 15 16 16 16  RESULTS AND DISCUSSION I I I -l Rate of Carbon Monoxide Reduction of S i l v e r (I) III-2 Chemistry and Stoichiometry of the Reaction IIIT3 E f f e c t of Carbon Monoxide Pressure III -k E f f e c t of Acetic Acid III-5 E f f e c t of Acetate Complexing III-6 Acid-Independent Reaction ; III-7 Acid-Dependent Reaction III-8 Acetate-Independent Reaction III-9 Reduction of Unbuffered S i l v e r Perchlorate III-10 "Best Value" Rate Parameters at 90°C ...' III-11 Proposed-Mechanism III-12 E f f e c t of Temperature  18 22 2k 2k 28 28 k2 ^9 52 62 67 73  IV  CONCLUSION  79  V  APPENDICES APPENDIX A APPENDIX B APPENDIX C  Method-of Estimating Rates from the Slope of Pressure-Time Records  8l  S o l u b i l i t y of Carbon Monoxide, Carbon Dioxide and Hydrogen i n Water  86  Silver-Acetate Complexing from E.M.F. Measurements  100  iv Page APPENDIX D  APPENDIX E  APPENDIX F  VI  Summary of Selected Experimental Data f o r the Reduction of S i l v e r ( I ) Solutions by Carbon Monoxide  105  Thermodynamic Calculations f o r the Oxidation of CO, E2, HCOOH and HCOO" i n Aqueous • Solutions at 25°C  119  Numerical Integration of Experimental Rate Law  122  REFERENCES  127  l VII  NOMENCLATURE USED IN RATE EXPRESSIONS  131  V  TABLES Page I  Reduction Rate of S i l v e r ( I ) by CO Under Various . Experimental Conditions  19  II  Reproducibility of Rate Measurements  III  Stoichiometry of Acetate-Buffered CO-Silver Perchlorate Reaction Summary of Intercepts and Slopes from R' vs [HOAc] i  IV  21  23  _  Plots at Various Degrees of Acetate Complexing  31 k6  V  Summary of the Dependence of S' on [Ag ] at 90°C  VI  Summary of Acetate-Independent  VII  Stoichiometry of Unbuffered CO-Silver Perchlorate Reaction at 90°C and 53 atm CO Summary of Experimental Rate Constants i n Unbuffered  56  S i l v e r Perchlorate Solutions at 90°C and 53 atm CO  59  VIII  +  Rates  53  IX  Summary of Rate Parameters f o r Equation 19 at 90°C  6k  X  "Best Value" Rate Parameters at 90°C  65  XI  Dependence of Reduction Rate on Temperature  75  B-I  S o l u b i l i t y of CO i n Water at 25 Atmospheres  89  B-II  E f f e c t of Pressure on S o l u b i l i t y of CO i n Water  B-III  S o l u b i l i t y of H  B-IV  S o l u b i l i t y of C0  2  i n Acetate Solutions at 90°C  96  B-V.  S o l u b i l i t y of C0  2  i n Water at One Atmosphere  98  C-I  E f f e c t of Temperature on Silver-Acetate Complexing  10k  D-I  E f f e c t of CO Pressure  105  D-II  E f f e c t of Acetic Acid  106  i n Water  2  D-III . Effect of Silver-Acetate Complexing D-IV  Rates Used i n Extrapolation t o Zero Acetate  91 95  106 112  vi Page D-V  Reduction of Unbuffered AgC10 Solutions  115  D-VI  Data Included i n Regression Analysis  117  D-VII  E f f e c t of Temperature  118  E-I  Standard Free Energy at 25°C  119  F-I  Comparison of Experimental and Calculated Pressure Records ...  126  4  vii FIGURES  1.  Page  Potential-pH Diagram f o r the Oxidation of CO, H , 2  HCOOH and HCOO" at 25°C  3  2.  Schematic Diagram of Reactor System  13  3.  Stainless Steel Reactor  Ik  k.  T y p i c a l Pressure-Time Records  20  5.  Dependence of Rate on CO Pressure  25  6.  Dependence of Rate on [HOAc] and. [H0Ac]-i  27  7.  Dependence of Rate on-[HOAc]-  at Various NaOAc Levels  29  8.  Dependence of Rate on [HOAc]  at Various AgC10 Levels  30  9.  Dependence of Acid-Independent Reaction on [Ag(I)] at Various NaOAc Levels Dependence of Acid-Independent Reaction on [NaOAc] at  10.  c  1  -1  4  33  Various Ag(I) Levels  3^  11.  Dependence of Acid-Independent Reaction on [Ag(I)][NaOAc]  35  12.  Dependence of Acid-Independent Reaction on [Ag(I)] and [NaOAc]; p l o t t e d according to equation 10, assuming [OAc ] = [NaOAc] .. Dependence of Acid-Independent Reaction on [Ag(I)] and-[OAc ]; plotted according to equation 10, assuming K = 3«7 M-i  39  Ik.  Dependence of Acid-Independent Reaction on [OAc ]  kO  15.  Dependence of Acid-Independent Reaction on [AgOAc]  ^1  16.  Dependence of Acid-Dependent Reaction on [Ag(I)] at Various NaOAc Levels  17.  Dependence of Acid-Dependent Reaction on [NaOAc] at Various Ag(I) Levels  kk  Dependence of Acid-Dependent Reaction on Acetate Complexing; p l o t t e d according to equation 15  ^7  Dependence of Acid-Dependent Reaction on Acetate Complexing; plotted according to equation 16  ^8  Dependence of Rate on [NaOAc] at Constant [HOAc]/[NaOAc]  50  -  13.  a  18. 19. 20.  37  -  -  [Ag(I)] and  viii Page 21.  Dependence of R [HOAc]/[AgOAc] on [OAc ] at Constant [Ag(I)] 1  -  and [HOAc]/[WaOAc]; p l o t t e d according to equation 2 1 22.  Dependence of Acetate-Independent Reaction on [Ag(I)]  23.  Reduction Rate of A g C 1 0  4  51 5^  i n Unbuffered Solution;  p l o t t e d according to equation 27  58 6l  2k.  Dependence of Rate on [Ag ] i n Unbuffered Solution  25.  Comparison of Experimental and Calculated Pressure Records  26.  Dependence of Rate on [HOAc]- at 6 0 , 8 0 , 9 0 and 110°C  7k  27.  Arrhenius-Plots for-Acid-Independent and Reactions  77  B-l  Measuring Burette System f o r Gas -.Solubility Determinations  B-2  S o l u b i l i t y of CO and-H  B-3  S o l u b i l i t y of C 0  C-l  Experimental C e l l f o r E.M.F. Measurements  +  ....  1  2  2  66  Acid-Dependent  i n Water from 2 5 to 225°C  i n Water at One Atmosphere  ....  87 90 99 101  CARBON'.'MONOXIDE:-'REDUCTION OF AQUEOUS SILVER ACETATE  I 1-1  INTRODUCTION  General Carbon monoxide, although i t i s a major by-product of many pyrometal-  l u r g i c a l processes, finds l i t t l e d i r e c t commercial use i n current m e t a l l u r g i c a l operations.  As a reducing agent carbon monoxide exerts a s l i g h t l y greater  p o t e n t i a l than hydrogen, which i s used commercially i n the production of copper, n i c k e l and cobalt powders from ammoniacal leach solutions (1,2), and therefore might be expected to f i n d s i m i l a r applications.  The fact that carbon monoxide  i s not used i n large-scale operations may be explained i n part by a general aversion to t h i s gas because of i t s t o x i c i t y and the perhaps mistaken concept that CO-reduction would lead to higher coats.  This l a t t e r objection i s not  necessarily v a l i d p a r t i c u l a r l y i n the v i c i n i t y of e l e c t r i c smelting plants where CO i s produced i n large volumes and available merely f o r the cost of collection.  What i s probably a more s i g n i f i c a n t reason f o r i t s l i m i t e d use i s  the general lack of d e t a i l e d information regarding p o t e n t i a l l y useful CO-metal reactions. The greatest use of CO i n the m e t a l l u r g i c a l industry has been the production of elemental n i c k e l by the Mond carbonyl process.  The process  con-  s i s t s e s s e n t i a l l y of reacting the gas with reduced n i c k e l to form a n i c k e l carbonyl which i s then thermally decomposed to give a high p u r i t y n i c k e l powder ( 3 ) . A modification of the Mond process was World War  used i n Germany during  II to produce n i c k e l and i r o n f o r powder metallurgy applications  (k).  The carbonyl process has also been adapted to the recovery of iron and n i c k e l from n i c k e l i f e r o u s l a t e r i t i c ores ( 5 ) .  /  Perhaps the greatest p o t e n t i a l use of CO i n processes  of m e t a l l u r g i c a l  importance l i e s i n the displacement of dissolved metals from hydrometallurgical  - 2 -  leach l i q u o r s .  A process has been developed (6,7) f °  r  the production of  elemental copper, n i c k e l or cobalt powders from ammoniacal solutions under autoclave conditions, reduction being accomplished at somewhat lower temperatures than corresponding hydrogen reactions. In the production of n i c k e l  r powder d i f f i c u l t i e s associated with carbonyl handling are a disadvantage.  In  the production of copper metal, however, CO i s p a r t i c u l a r l y useful because of i t s a b i l i t y to s t a b i l i z e the copper(I) species against hydrolysis thus preventing the p r e c i p i t a t i o n of cuprous oxide and providing a purer product from a wider range of solution compositions. A Russian process f o r the recovery of n i c k e l has been described (8) i n which n i c k e l concentrates are leached i n ammonia solution under oxygen pressure, and the r e s u l t i n g solution reduced with CO to form a carbonyl which i s then decomposed to y i e l d the n i c k e l product. 1-2  Thermodynamic Considerations of CO Reactions i n Acfueous Solution When CO acts as a reducing agent i n aqueous solution i t i s oxidized  to carbon dioxide, bicarbonate or carbonate, depending on the pH of the solu-  l tion.  The thermodynamic potentials of the corresponding couples at room  temperature  are depicted i n Figure 1 i n the form of a potential-pH diagram  using the International or Stockholm Convention electrode p o t e n t i a l s .  ( 9 ) f o r the sign of the -  The associated potentials of formic a c i d and  are also included together with that of the hydrogen electrode. ing the diagram, unit a c t i v i t y f o r a l l species except H  +  formate  In construct-  has been assumed.  The various reactions considered are l i s t e d i n Figure 1 together with a summary of the thermodynamic expressions used.  The d e t a i l s of the thermodynamic c a l -  culations are given i n Appendix E. It i s apparent from the diagram that CO i s a stronger reducing agent than H  2  by at least 0.1 v o l t s at a l l pH values.  Thus CO might be expected to  0  k  2  6  8  10  12  Ik  PH 1.  C0  + 2H + 2e  2.  HC0 - + 3H+ + 2e  3.  C0 = + 4H + 2e  IK  2H  5-  C0  2  + 2H + 2e  6.  C0  2  + H  7-  HCO3"  8.  C0 = + 3H + 2e  9-  HC00H  =  H  + 2e  + 2H  +  10.  C0  ll.  HCO3-  2  + H0 2  =  = H  +  2  -0.103 - O.O59 pH  E  =  0.128 - O.O89 PH  HCOO" = =  HCOO" •+ H 0 2  HCOO" + H 0 2  + HCOOH  +  + HCO3"  + C0  Figure 1.  0.435 - 0.118 pH =  E  HCOOH  =  +  +  =  E  CO + 2H 0  =  + 2e  H  E  E  3  =  2  2  +  +  CO + 2H 0  =  +  3  + 2e  2  =  3  +  CO + H 0  =  +  2  = 3  -O.O59 pH -O.I98 - 0.059 PH  E  =  -0.309 - 0.030 pH  E  =  -0.079 - 0.059 PH  E  =  0.227 - O.O89 PH  PH  =  -log K = 3.8  pH  =  -log K = 7.8  PH  =  -log K = 10.4  Potential-pH Diagram f o r the Oxidation of CO, H , HCOOH and HCOO- at 25°C 2  - k reduce several metal ions which are not reducible by H . 2  However, cobalt with  a standard reduction p o t e n t i a l of -0.277 volts (10) i s the least noble metal reported t o be reduced by CO (6) and i t i s also reduced by H  2  (1,2,6).  In both  cases reduction of ammonia complexes can be accomplished i n basic or s l i g h t l y a c i d i c solutions.  Other metals which have been produced by CO-reduction  include n i c k e l (6,8,11,12), bismuth (12), copper (6,13), s i l v e r (13,14), mercury (15) and gold (16).  Some metal species which are not reducible to metal by CO  can be reduced to a lower oxidation state  as i n the reduction of permanganate  to Mn0 i n a c i d or neutral solutions and t o Mn0 , which gives-Mn =  2  ++  4  portionation, i n basic solution (15,17)• solutions (18).  by dispro-  Chromate i s reduced to Cr 03 i n a c i d 2  The reduction p o t e n t i a l of CO can be increased by r a i s i n g the  CO p a r t i a l pressure, a hundred-fold increase being equivalent to 2 pH units or about 0.1 v o l t s at room temperature.  Higher temperatures increase the pressure  dependence of the CO p o t e n t i a l as w e l l as increasing the reaction rates. Figure 1 indicates that formic a c i d or the formate ion i s a stronger reducing agent than CO i n a c i d solution while the reverse i s true at higher pH values, the equipotential point occurring at about pH 7 at room temperature. Thus i t i s possible f o r CO to react with certain basic salt solutions to form metal formates.  Thermodynamically, the most favourable cases include the  hydroxides of the a l k a l i metals and the a l k a l i earths except beryllium (19)The reaction has been known since the o r i g i n a l synthesis of sodium formate from sodium hydroxide.in the mid-nineteenth century (20). The standard reduction p o t e n t i a l of the oxygen electrode i n a c i d solution i s 1.229 v o l t s (10) and thus the CO-reduction of H 0 t o form C 0 and 2  H  2  2  i s thermodynamically favourable. This i s the water-gas s h i f t reaction  f a m i l i a r i n the gas phase at elevated temperatures.  I t has also been observed  in basic aqueous solutions at temperatures greater than 150°C (21).  - 5 -  1-3  Structure  CO m o l e c u l e c o n t a i n s a t o t a l o f Ik  The  ls 2s 2p  atom ( i . e .  ls 2s 2p 2  2  o f C a r b o n M o n o x i d e a n d i t s Compounds  2  2  2  i n the ground  4  s t r u c t u r e as i s N , (MO)  structure  C[ls 2s 2p ] + 0[ls 2s 2p ] The  2  2  2  (AO);  MO  a n d wtr MOb  g i v e CO  2  C(2py)  and  0(2py).  i t s triple-bond structure.  2  2  4  of  C(2p ) x  C(2s) AO;  and  atomic  0(2p );  The the  C(2p ) x  wir r e p r e s e n t s two  Two  degenerate  pi-antibonding largely  of m e t a l c a r b o n y l complexes. (which corresponds  AO), a r e d o n a t e d t o a m e t a l a t o m a n d f o r m a s i g m a - b o n d i s formed*.  yKy  The b o n d s f o r m e d b y t h e  l o n e p a i r o f e l e c t r o n s i n t h e CO(x^T) o r b i t a l  c a r b o n y l complex  x^7  x  (vif) a r e a s s o c i a t e d w i t h w"[r a n d i t i s t h e s e o r b i t a l s w h i c h a r e  responsible f o r the s t a b i l i t y  to  CO[KK(z^ (yv) (xv) (wir) ]  r e p r e s e n t e d l a r g e l y by the  d e g e n e r a t e p i - b o n d i n g MOfe o f  by  r e p r e s e n t e d l a r g e l y b y t h e 0(2s)  MO  (i.e.  triply-bonded  (22)  i s described  >  4  carbon  I n terms o f the molecular  y y r e p r e s e n t s a s i g m a - a n t i b o n d i n g MO  i s a non-bonding  MO's  2  z ^ 7 o r b i t a l i s a non-bonding  orbital  from t h e oxygen atom  s t a t e ) and i s b e s t d e s c r i b e d as a  t r e a t m e n t t h e CO  2  8  w h i c h i s i s o e l e c t r o n i c w i t h CO.  2  orbital  i n the ground s t a t e ) and  e l e c t r o n s , 6 from the  S t a b i l i z a t i o n o f t h e complex  when a  occurs through p i -  b o n d i n g between o c c u p i e d m e t a l d - o r b i t a l s and t h e empty v i r p i - a n t i b o n d i n g orbitals  o f CO  which would  t h u s p r o v i d i n g a mechanism f o r t h e removal o f t h e excess  o t h e r w i s e b e p r e s e n t on t h e m e t a l a t o m . .The  donor c h a r a c t e r i s t i c s t a b l e complexes inability  *  o f CO  p i - a c c e p t o r and  charge sigma-  complement e a c h o t h e r and p e r m i t t h e f o r m a t i o n o f  e v e n t h o u g h CO  itself  t o form s t r o n g complexes  i s a p o o r d o n o r a s shown b y i t s  w i t h empty o r b i t a l a c c e p t o r s  (2k).  The l o n e p a i r o f e l e c t r o n s i n t h e C 0 ( z 9 ) MO, w h i c h c o r r e s p o n d s t o t h e 0(2s) AO, a p p a r e n t l y i s n o t d o n a t e d (23) b e c a u s e 0 i s much more e l e c t r o n e g a t i v e t h a n C. D o n a t i o n o f e l e c t r o n s i n t h e CO(wTT) MO's i s n o t o b s e r v e d e x p e r i mentally although t h i s type of donation i s observed w i t h c e r t a i n other l i g a n d s w h i c h a r e i s o e l e c t r o n i c w i t h CO ( e . g . a c e t y l e n e ) ( 2 3 ) .  - 6 -  Carbonyl compounds are formed with t r a n s i t i o n metals i n groups VI, VII and VIII of the periodic table ( 2 5 , 2 6 ) .  Carbon monoxide i s also absorbed by cuprous  chloride i n ammoniacal solution, by s i l v e r sulphate i n concentrated sulphuric acid, by mercuric acetate i n methyl alcohol, and by dry auric chloride ( 2 5 ) = In the Fischer-Tropsch synthesis  CO and H  2  react i n the presence of  a suitable metal catalyst t o y i e l d primary alcohols, o l e f i n s and saturated hydrocarbons, the r e l a t i v e y i e l d s depending oh the operating conditions ( 2 7 ) . When o l e f i n s are added to the CO-H gas mixtures, alcohols, aldehydes and 2  ketones are produced and the process i s known as the 0 X 0 or hydroformylation synthesis ( 2 8 ) .  The 0 X 0 synthesis requires a cobalt catalyst ( i n the metallic,  s a l t or carbonyl form) while catalysts f o r the Fischer-Tropsch synthesis include cobalt, n i c k e l , i r o n and other metals capable of forming metal carbonyls. The role of these catalysts i s associated with t h e i r a b i l i t y to form metal carbonyls and related complexes which act as intermediates i n the organic reactions ( 2 9 ) . The mechanisms f o r these and other syntheses involve the i n s e r t i o n of CO into metal-carbon and metal-oxygen bonds ( 2 9 ) .  I-k  K i n e t i c s of Metal Ion Reduction by Carbon Monoxide i n Aqueous Solution The rate of CO-reduction of both A g S 0 2  4  and C u S 0  4  has been reported  ( 1 5 ) t o be f i r s t - o r d e r i n CO p a r t i a l pressure up to at l e a s t 6 0 atm and secondorder i n dissolved metal.  The following mechanism was proposed t o account f o r  the observed k i n e t i c s :  Me  +  + CO ^ = = M e C 0  MeC0  +  Me C0 2  + Me + +  (rapid)  +  ^==> M e C 0  +  2  + H 0 2  s l Q W  + +  (a)  (rapid)  > 2Me + C 0  2  + 2H  (b) +  (I)  (c)  In the s i l v e r studies two series of measurements were made with i n i t i a l A g S 0 2  4  - 7-  concentrations between about 0.007  a n d  - 0.03 M_, one series being buffered with  O.65 M NH 0Ac* and the other unbuffered.  The rate law for the buffered series  4  between 70 and 110°C i s given as: -d[Ag(I)]/dt  =  6.02 x 104 [Ag(I)]  2  P  exp(-93OO/RT)  C Q  (M min" ) 1  and f o r the unbuffered series between 70 -d 150°C as: an  -d[Ag(I)]/dt  =  12.8x105 [Ag(I)]  2  P  exp(-l^,100/RT)  c o  ,(M l i n ' i )  Thus at 90°C the reaction rate i n the buffered system i s about 36 times faster than i n the unbuffered system. The reduction rate of CuS0 was measured i n d i l u t e 4  unbuffered  solutions, apparently to minimize corrosion and hydrolysis problems.  Also a  sheet of copper metal, etched t o give a high surface area, was required to obtain reproducible r e s u l t s .  The reported rate law as measured between 160  and 190°C has the form: -d[Cu(total)]/dt = 2.56 x 1 0  [Cu(total)]  1 3  P  2  exp(-33,500/RT) (M min" ) 1  c o  Another k i n e t i c study involving the CO-reduction of aqueous silver, amines i n basic solution has recently been reported (1^).  Measurable rates  were r e a d i l y obtained at atmospheric pressure and room temperature and were shown t o be consistent with the rate law: -d[C0]/dt  =  where L denotes an amine ligand. -d[C0]/dt  *  OAc  k  [C0][AgL ]/[HL+] +  exp  2  This rate law i s equivalent t o : =  k[C0][LAg0H]  i s used throughout the text t o denote the acetate r a d i c a l CH C00~. 3  - 8 The e x p e r i m e n t a l r a t e c o n s t a n t k  e X  p i s thus e q u i v a l e n t t o kK^K-^K^, where  the f i r s t i n s t a b i l i t y constant o f A g L ,  i s t h e b a s i c i t y constant o f t h e  +  2  amine a n d K  n  i s the a s s o c i a t i o n constant of A g L  w i t h 0H~.  +  mechanism was p r o p o s e d t o account f o r t h e observed AgL  + 2  + H 0 i=L^  LAgOH + H L  2  is  The f o l l o w i n g  kinetics:  (rapid equilibrium)  +  (a)  0 LAgOH + CO  k  > LAg-5-0H  0 LAg-C-OH + LAgOH  (rate-d.etermining)  > 2Ag + C0 ~ + 2HL  F o r t h e l i g a n d s NH , CH3NH > C H N H 3  k  e X  +  3  2  p were almost w h o l l y accounted  2  5  2  (b)  (fast)  (II)  (c)  and.(C H ) NH v a r i a t i o n s i n 2  5  2  f o r b y v a r i a t i o n s i n K^K^ w h i l e k K ^ remained  e s s e n t i a l l y independent o f t h e n a t u r e o f t h e amine and had a v a l u e o f about *10  5  M~  2  sec  - 1  a t 25°C . A t [ N H ] g r e a t e r t h a n about 0.02 M w i t h NH as the l i g a n d +  4  3  (14a). d e p a r t u r e s f r o m t h e above r a t e l a w were observed w h i c h were e x p l a i n e d by t h e s l o w i n g down o f r e a c t i o n 11(c), due t o l o w e r  [LAgOH], u n t i l  between 11(c) and t h e r e v e r s e o f 11(b) c o n t r o l l e d t h e r a t e . t h e s e f i n d i n g s i t was c o n c l u d e d  competition  On t h e b a s i s o f  (lk) t h a t h i g h pH r a t h e r t h a n s p e c i f i c com-  p l e x i n g e f f e c t s i s r e s p o n s i b l e f o r t h e h i g h r e a c t i v i t y o f A g ( I ) toward CO i n amine-buffered  aqueous s o l u t i o n s .  A d d i t i o n a l support f o r t h e n a t u r e o f t h e p r o p o s e d i n t e r m e d i a t e i n t h e r a t e - d e t e r m i n i n g s t e p 11(c) was drawn developed  f r o m a p a r a l l e l s t u d y (15)  by a n a l o g y from t h e mechanism  t o describe the CO-reduction  of Hg  + +  in  dilute perchloric acid, v i z : Hg  + +  0 Hg-C-0H + H  + CO + H 0  +  2  Hg-$-0H  +  Hg + H g  + +  >Hg + C0 >Kg£>  ++  2  + H  +  (fast)  (fast)  +  (a) (b) ( I I I ) (c)  - 9The observed k i n e t i c s at atmospheric pressure over the temperature range 26 to 5U°C were found to be consistent with the rate law: -d[CO]/dt  =  k[CO][Hg ] ++  with a AH* of lA.6 kcal/mole and a AS+ of -13  e.u.  The structure of the  proposed intermediate i n the rate-determining step III(a) i s analogous t o the stable methyl formate derivative AcO-Hg-C^-OCH which i s formed when CO reacts 3  with methanolic solutions of mercuric acetate under s i m i l a r conditions It was suggested (15) CO between H g  that the rate-determining step involves the i n s e r t i o n of  and a co-ordinated water molecule.  + +  (30).  Similar reactions involving  the i n s e r t i o n of CO into metal-oxygen bonds are important i n a number of metal (29).  carbonyl c a t a l y t i c reactions  The k i n e t i c s of the CO-reduction of Mn0 " (15),  as measured i n  4  aqueous solution at atmospheric pressure over the temperature range 28 to 50°C and the pH-range 1 t o 13,  are consistent with the rate law:  -d[C0]/dt  =  k[C0][Mn0 "] 4  with a AH* of 13 kcal/mole and a AS-t- of -17 catalyzed by A g  +  or H g  + +  (15)  a  n  e.u.  The reduction i s strongly  d i n d i l u t e perchloric acid solutions the rate  law i s :  -d[C0]/dt  where X represents A g e.u.; (15)  for Hg , ++  +  or H g . ++  =  k[C0][Mn0 -][X] 4  For Ag , AH-t- = 1.2  kcal/mole and AS-t- =  +  A H * = 6.k kcal/mole and AS* = -21  e.u.  .31  I t i s suggested  that the high c a t a l y t i c a c t i v i t y i n these reactions may be due to i n t e r -  mediates of the type Ag-8-0Mn0 and Hg-8-0Mn0 . +  3  3  The k i n e t i c s of the CO-reduction of B i ( S 0 ) 2  and -0.7  are described by the rate law  (12):  4  3  at pH's between O.k  - 10 -a[Bi(III)]/dt  =  8 . 2 x 1 0 6 [ B i ( I I I ) ] P ( [ H ] - - 0.24)exp(-23,OOO/RT) (M min" +  x  C 0  4 The high a c i d concentrations  are required to prevent hydrolysis.  suggested (12) that the active species i s B i O H . ++  decreases with increasing i n i t i a l  I t was  An induction period which  [ B i ( I I I ) ] indicates that the reaction i s  heterogeneous. The rate law f o r t h e CO-reduction of Ni(0Ac)  i n HOAc -NH 0Ac(0.5  2  4  M)  buffered solution at constant pH of 5.3 i s given (12) as; -d[Ni(II)]/dt  8 0 . 6 [Ni(II)] ( P  =  c o  -5.2) exp(-12,000/RT)  Other studies on the reduction of n i c k e l ( I I )  (M min-i)  amine sulphate complexes (11,21)  indicate that the rate increases with increasing pH.  The CO dependence may  r e f l e c t a thermodynamic influence on the rate.  1-5  Object and Scope of the Present Investigation  This t h e s i s embodies the r e s u l t s of a k i n e t i c study of the COreduction of aqueous AgOAc and AgC10  i n a c i d solution.  4  The work forms part  of a general i n v e s t i g a t i o n of the mechanisms by which CO displaces metals" from aqueous s a l t solutions. At the time the study was undertaken the only published  information  on the k i n e t i c s of the reaction was that of the CO-reduction of Ag S0 2  tions ( 1 3 ) -  4  solu-  At 90°C the rates measured i n NH 0Ac buffered solution were some 4  36 times greater than the corresponding I n i t i a l experiments i n the present  rates i n unbuffered  sulphate solution.  study revealed large pH e f f e c t s on the  reduction rate of AgC10 and AgOAc i n a c i d solution. 4  The previously proposed  mechanism (see mechanism I, Section 1-4), however, takes no account of pH effects. plexes  Also the possible v a r i a t i o n i n r e a c t i v i t y of d i f f e r e n t Ag(I) com-  (e.g. Ag , AgS0 , AgOAc, AgNH , Ag(NHs) , Ag(NH )0H, etc.) was not +  _  4  3  2  3  - 11 -  considered.  At the concentrations used i n the Ag S0 -NH OAc studies (13) i t 2  i s estimated,  4  4  from reported room temperature complexity  more than 60$ of the s i l v e r was  complexed with ammonia.  AgC10 , generally buffered with NaOAc and HOAc, was 4  constants  (31) that  The simpler system of  chosen to further elucidate  the mechanism by which CO displaces s i l v e r from aqueous solution. A preliminary study of the CO-reduction of Cu(OAc) and Cu(C10 ) 2  was also made.  4  The f i r s t step i n the reaction apparently i s the reduction of  Cu(II) to Cu(I) which forms a stable complex with CO and then hydrolyzes p r e c i p i t a t e a Cu 0 product. 2  and C u ( C 1 0 ) 4  2  I t was  demonstrated that both aqueous Cu(OAc)  2  2  The catalyzed reaction i s f i r s t - o r d e r i n CO and  zero-  and occurs homogeneously i n the l i q u i d phase i n a similar manner  2  to the CuS0 -and Cu(C10 ) -catalyzed 4  4  2  oxidation of H  2  by 0  2  (32).  The rate  increases with increasing [Cu(II)] and pH but the exact dependence was determined.  to  catalyze the oxidation of CO by 0 , measurable rates being  obtained above 125°C. order i n 0 ,  2  not  The CO-Cu(II) reaction i s currently•the subject of another study  i n these laboratories. No c a t a l y t i c a c t i v i t y , of aqueous AgC10 or AgOAc i n the C0-0 4  t i o n was  detected at temperatures to 150°C, although at 250°C Ag S0 2  reported (32) to catalyze the H -0 2  process.  2  4  2  reac-  is  reaction, possibly by a heterogeneous  - 12 -  II  II-1  EXPERIMENTAL  Reactor System Reduction experiments were conducted f o r the most part i n a small  (approximately- 120 ml) pressure vessel, the rate of reaction generally being followed by the decrease of pressure i n a closed system.  In some cases, how-  ever, rates were determined by analyzing periodic l i q u i d samples.  The pressure  vessel and associated f i t t i n g s were manufactured from stainless s t e e l (type 316) by Pressure-Products-Industries Inc., Hatboro, Pennsylvania 1  andxdesigned for: :  working: .pressures :.up. to, 7200,-psi.,. .A schematici.diagram, of the..reactor system i s 1  given i n Figure 2 and a section drawing of the reactor i s shown i n Figure 3Pressure-tight closure of the vessel was achieved by using a t o t a l l y enclosed stainless s t e e l F l e x i t e l l i c gasket having a Teflon and asbestos f i l l e r . Agitation was provided by mounting the reactor i n a v e r t i c a l p o s i t i o n on a shaker mechanism which when activated reciprocated h o r i z o n t a l l y at 275  oscil-  lations per minute with a l - l / 2 - i n stroke. Gas inlet- and outlet l i n e s of f l e x i b l e l / l 6 - i n o.d. stainless s t e e l (type 3^7)  c a p i l l a r y tubing, and a stainless s t e e l clad l / l 6 - i n o.d. i r o n -  constantan thermocouple were connected through the bottom of the reactor by means of Ermeto stainless s t e e l (type 17-4  PH) sleeve f i t t i n g s .  Also a short  l / 8 - i n o.d. l i q u i d sampling l i n e was connected through the bottom of the vessel.  II-2  Pressure Measurement The pressure of the system was measured with a Consolidated  Electrodynamics Corp. pressure transducer pick-up (type 4-311) mounted i n the gas i n l e t l i n e .  The transducer consisted e s s e n t i a l l y of strain-gauge windings  - 13 -  A  E  V  A B C D E SW  potentiometer recorder temp c o n t r o l l e r reactor CO cylinder selector switch  Figure 2 .  PT TH TP V! V 2  v  3  3  pressure transducer thermocouple thermister probe gas i n l e t valve gas outlet valve l i q u i d sampling valve  Schematic Diagram of Reactor System  - Ik -  1/8" Ermeto Set screw (8 at k^>°) Cover, 316-S.S. Thrust washer Cover nut, A-5136 s t e e l Gasket, 316 S.S. + Teflon + asbestos --Gasket, Teflon  Thermowell  Reactor body, 316 S.S.  1/8" Ermeto l / l 6 " Ermeto (k at 9O ) 0  Figure 3-  Stainless Steel Reactor  (full-scale)  - 15 connected i n a four-arm bridge c i r c u i t .  Pressure against a diaphragm displaced  the sensing element which changed the resistance of the two active arms and produced an e l e c t r i c a l output proportional t o the applied pressure. c i r c u i t was excited by a battery of dry c e l l s  The bridge  supplying up t o f i v e v o l t s at  10 ma, and the output was measured with a high p r e c i s i o n Leeds and Northrup potentiometer (No. 7552) or recorded on a 10 mv Brown Electronik strip-chart recorder.  Three transducers were available covering the ranges 0-150, 0-500  and 0-1000 psig, each of which had been calibrated at an e x c i t a t i o n voltage of 5.O v o l t s dc by the manufacturer and had a nominal 20 mv f u l l - s c a l e output. From the manufacturer's specifications i t i s estimated that the pressure measurements had an accuracy of better than 2$.  II-3  Temperature Control and Measurement The temperature of the experimental solution was regulated to ± 0.3 C ,o  using a Yellow Springs Instrument Co. Thermistemp Temperature Controller (Model 71) with a stainless s t e e l clad thermister probe (No. k-06) mounted i n a thermowell i n the reactor l i d .  The heating unit was a 600-watt external band heater  connected t o a variable transformer through the c o n t r o l l e r which was shunted by a four-ohm nichrome r e s i s t e r when i n the "off" p o s i t i o n .  Regulation was  made by adjustment of the transformer to give a heat input s u f f i c i e n t t o maint a i n the temperature at a value s l i g h t l y below the operating point when the regulator was i n the "off" position, and s l i g h t l y above the operating point when the regulator was i n the" "on" p o s i t i o n .  The top of the reactor was  wrapped with removable asbestos lagging t o reduce heat losses. Independent temperature measurements were made with an iron-constantan thermocouple sheathed i n stainless s t e e l inserted through the bottom of the reactor.  The thermocouple was calibrated at the melting point of indium  (156.4°C), the steam point and at room temperature by comparison with a  - 16 p r e c i s i o n mercury-in-glass thermometer graduated i n tenths of a degree.  II-k  Materials  A l l chemicals were of reagent grade and used without further p u r i f i cation.  Carbon monoxide  (99-5$ min) was obtained from the Matheson Co.  Analysis of several gas samples indicated that N , 2  less than 0.2$,  0  2  and H  2  impurities were  0.1$ and 0.002$ respectively.  -Experimental solutions were prepared by mixing'and d i l u t i n g aliquots of standard stock solutions.  AgC10 solutions were f i l t e r e d a f t e r preparation  and thus any trace amounts of C l AgCl.  4  -  or C103~  impurity were removed as insoluble  NaOAc and HOAc solutions were prepared with known weights of reagent  and were not standardized further.  II-5  Chemical Analysis S i l v e r solutions were analyzed by t i t r a t i o n i n d i l u t e n i t r i c a c i d  with standard ammonium thiocyanate using f e r r i c n i t r a t e as an indicator (33)Samples drawn during an experiment were f i l t e r e d before analysis. A l l pH measurements were made at room temperature with a Beckman pHmeter (Model G). Gas samples were generally analyzed i n a Beckman GC-1  chromatograph  using molecular sieve and s i l i c a g e l columns.  II-6  General Experimental  Procedure  In general, the experimental procedure was to pipette a known volume (90-100 mis) of solution into the reactor, seal and heat to the desired temperature with the shaking mechanism operating. had been achieved the shaking mechanism was  When temperature  control  stopped b r i e f l y while CO was added  - 17 to the desired pressure which usually required between 30 and 60 seconds.  The  pressure record indicated that CO equilibrium between the l i q u i d and gaseous phase was attained within two to four minutes a f t e r a g i t a t i o n had been resumed. Rates f o r reactions involving the reduction of AgC10 i n HOAc-NaOAc 4  buffered solutions were generally measured by a pressure-drop method i n which the decrease i n pressure of the closed reactor system was recorded. majority of the rate data was  The  obtained by measuring the slope of each pressure-  time record a f t e r CO saturation was complete (two to four minutes).  These  slope measurements were converted to rates expressed as M s e c i * using CO _  CO2  and  s o l u b i l i t y data and the estimated volumes of gas and l i q u i d i n the reactor.  Corrections were made to the concentrations of AgC104, HOAc, NaOAc and CO to take account of the small amount of reaction which occurred during the i n i t i a l CO saturation.  The derivation of the mathematical expression used to convert  the slope measurements to fundamental rate units i s given i n Appendix A. To augment meager published data (34) the s o l u b i l i t y of CO i n water was measured from room temperature to 220°C at pressures to 63 atm.  The  technique and results of these measurements are given i n Appendix B together •with some data on the s o l u b i l i t y of H tions.  2  i n water measured under similar condi-  Also included are a few measurements on the s o l u b i l i t y of C0  and HOAc-NaOAc solutions at 90°C and 2.6  *  2  i n water  atm.  A l l concentrations are expressed i n terms of l i t r e s measured at room temperature (20-25°C)  - 18 -  III  III-l  RESULTS AND DISCUSSION  Rate of Carbon Monoxide Reduction of S i l v e r ( I )  E a r l y experiments demonstrated that at low pH the reduction of AgC10 solutions was very slow even at temperatures atm.  of 175°C and CO pressures of 20  When the solutions were buffered with NaOAc and HOAc, however, the  rates were r e a d i l y measured under much milder conditions (e.g. 60°C and 5 atm CO) by the pressure-drop method.  Table I summarizes the rate of reaction  between CO and Ag(I) as measured over a wide range of experimental conditions. A set of t y p i c a l pressure-time records i s shown i n Figure k.  The  fast i n i t i a l decrease i n pressure corresponds to saturating the solution i n CO and the slower pressure decrease thereafter corresponds to chemical reaction exclusively.  The marked difference i n the rate of pressure drop i n the two  regions of the curve i s clear evidence that a f t e r i n i t i a l CO saturation the measured rate of CO consumption i s e f f e c t i v e l y independent  of mass transfer  of CO between the gas and l i q u i d phases. The rates measured from the i n i t i a l slopes of pressure-time records were estimated to have possible errors of at least t 10$ i n favourable cases and these errors may have r i s e n to ± 20$ or higher f o r faster or slower rates. The r e p r o d u c i b i l i t y of the rates, however, was generally better than ± 5$ shown i n Table I I . rate i s independent  a s  Results included i n the table also indicate that the of the surface area of stainless s t e e l and p r e c i p i t a t e d  s i l v e r i n contact with the solution.  The reaction therefore i s not hetero-  geneous i n nature but must occur homogeneously i n the l i q u i d phase. on the rate was detected when the C0 i n i t i a l slopes were measured was  2  No e f f e c t  concentration present at the time the  increased a factor of f i v e .  The presence of  small amounts of a i r (equivalent to about one atmosphere) was also without  4  - 19 -  TABLE I  Reduction Rate of S i l v e r ( I ) by CO Under Various Experimental Conditions  Temp °C  [Ag(I)]  175 90 90 60 60 80 90 90 110 110  .03 .05 1.0 .10 .10 .11  M  [NaOAc] M  [Acid] M  • 05 .002 .05c • 77 .05^ .06 .10 • 77 .28 .05 c  --  • 13 • 13 • 19 .66  .ok .2k  .Ok • 21  .11 .10  .18  c  d  d  d  d  d  d  co  R x 10s  atm  M s-i  20 53 53 5^ 5-3 5A 12.4 27.1 4.6 ^•5  0.36 0.58 12.0 O.95S 5.81S I8.58 4l.4§ 83.0 40.86 93-7  p  a  e  f  f  S  g  a  - R = -d[C0]/dt = -0.5d[Ag(I)]/dt = -0.5d[H ]/dt  b  - [Ag(I)]  +  pH  = s i l v e r analysis of periodic l i q u i d samples = pH determination of periodic l i q u i d samples  Pressure = pressure-drop method c  - Present as HC10  d  - Present as HOAc  e  - Average rate between 2 and 15 hours  f  - Average rate a f t e r 15 minutes  g  - I n i t i a l rate a f t e r CO saturation complete (2-4 minutes)  4  Method  13  [Ag(I)] PH pH Pressure Pressure Pressure Pressure Pressure Pressure Pressure  - 20 -  Expt No  [HOAc] M  • I n i t i a l Rate M sec-i  A  290  I.58  11.9 x  10-6  B  289  0.23  28.9 x  10-6  C  285  0.09  51.8 x IO"  6  0.5  0  5  10 Time  Figure k.  15  (min)  Typical Pressure-Time Records (0.235 M AgC10 ; 0.090 M NaOAc; 11.7 atm CO; 90°C) 4  - 21 -  TABLE I I  Reproducibility of Rate Measurements  •i-.. Expt Wo  Temp °C  139 14 0 l4l  161 l64  I65  a  b  265 264 266 267 268^ G  d  e  198 200 201  I n i t i a l Concentrations  [AgC10 ] [WaOAc][HOAc] M M M 4  R x 106 M s-  C0 atm  p  1  90 90 90  .100 .100 .100  .090 .090 .090  .180  .180 .180  26.6 26.7 26.3  21.4 20.2  90 90 90  .115 .115 .115  .045  .766 .766 .766  27.7 27.5 25.6  14.4  .045 .045  90 90 90 90 90  .144 .144 .144 .144  .195 .195 .195 .195 .195  • 195 • 195 .195 • 195 .195  5.04 4.78  110 110  .115 .115 .115  .195 • 195 .195  .023 .023 .023  4.49 4.53 4.49  110  .144  18.7 14.2 .12.7  5.16  19.4  5.10  20.9 20.3 I8.5  5.04 '  18.9 124. 111. 106.  R'x 10 M a-i s-  Average Deviation  s  1  i  • 755 .711  6.2 -0.3 -6.1  .520 .516 .496  1.8 1.0 -2.9  .804  3.75 4.09 4.02  3.87 3.75 27.6 24.6 23.7  -3.9 4.9 3.1 -0.8 -3.9 9.1 -2.8 -6.3  a  - Contained 0.1 g Ag p r e c i p i t a t e d during previous experiment plus 2.0 g s i l v e r sponge obtained from the Consolidated Mining and Smelting Co. L t d .  b  - Reactor thoroughly leached with HN0 p r i o r t o charging to remove a l l traces of previously p r e c i p i t a t e d Ag.  c  - A i r evacuated from solution i n reactor p r i o r t o heat-up.  d  - Added C 0 to 5.7 p s i p r i o r t o CO addition.  e  - Contained 3-3 g fine 316 S.S'. f i l i n g s .  f  - Contained 1.0 g Ag p r e c i p i t a t e d during previous experiments plus J.O g s i l v e r sponge obtained from the Consolidated Mining and Smelting Co. L t d .  3  2  - 22 -  e f f e c t on the rate.  Ill-2  Chemistry and Stoichiometry of the Reaction When dissolved s i l v e r undergoes reduction by carbon monoxide i n a c i d  solution the reaction products are s i l v e r metal and carbon dioxide  according  to the reaction:  (IV)  The Ag(I)-CO-C0  2  stoichiometry of reaction IV was examined at several tempera-  tures by analyzing l i q u i d and gas samples taken a f t e r about ^>0-6ofo reaction had occurred.  Solution composition  and experimental  r e l a t i v e l y fast reduction rates were selected. C0  2  conditions that gave  The concentrations  of CO and  were estimated from t h e i r p a r t i a l pressures using s o l u b i l i t y c o e f f i c i e n t s  given i n Appendix B, and assuming i d e a l gas laws.  The r e s u l t s of these  measurements, as summarized i n Table I I I , agree within 10$ with the s t o i c h i ometry represented by reaction IV*.  No H  2  was detected i n any gas sample.  Several f i l t e r e d l i q u i d samples of reaction products were tested for the presence of aldehydes and ketones using 2,4-dinitrophenylhydrozine  (35),  and f o r compounds containing a CH C0 group using the Iodoform test (36,37)3  Liquid samples were analyzed q u a l i t a t i v e l y i n a Beckman GC-2 chromatograph containing a Carbowax 1000 column using a thermal conductivity bridge detector, and also i n a Perkin-Elmer Vapor Fractometer (Model 15^C) containing a dinonyl phthalate column using a flame i o n i z a t i o n detector**.  No organic compounds  *  The m e t a l l i c s i l v e r p r e c i p i t a t e from a t y p i c a l experiment was confirmed by X-ray d i f f r a c t i o n . Mr. D.J. Rose performed the analysis and his a s s i s t ance i s g r a t e f u l l y acknowledged.  **  Thanks are extended t o Professors C.A. McDowell and J . Halpern of the U.B.C. Chemistry Department for allowing these analyses to be performed i n t h e i r laboratories.  TABLE I I I Stoichiometry o f Acetate-Buffered  Initial Expt No  Temp °C  [HOAc] M  Concentrations  ^co  CQ - S i l v e r P e r c h l o r a t e  [Ag(I)] Consumed M  [CO] Consumed M  Reaction  [co2]  A[Ag(I)]  A[C0 ] A [CO]  [NaOAc] M  [Ag(I)] M  0.222  0.115  5.44  O.O69  0.037  O.O38  1-9  1.0  atm  Produced •M  A[C0]  2  228  90  229  110  0.039  O.I95  0.115  4.42  O.O73  0.038  0.041  1.9  1.1  224  120  0.023  O.I95  0.115  4.08  0.064  0.035  0.036  1.8  1.0  225  120  0.023  0.195  0.115  4.08  0.063  0.035  0.039  1.8  1.1  226  120  0.222  0.115  4.08  0.061  0.031  0.035  2.0  1.1  - 2k except HOAc and NaOAc were detected. In the absence of Ag(I) no CO absorption was observed beyond the amount i n i t i a l l y required f o r saturation of the aqueous solutions.  The  presence of previously p r e c i p i t a t e d silver, metal had no e f f e c t on the CO absorption i n Ag(I)-free solutions.  In the absence of CO no reduction of  Ag(I) was detected i n HOAc-NaOAc solutions over a period of several hours at  90°C. These observations  are consistent with the absence of side reactions  a f f e c t i n g the CO-reduction of AgC10 i n NaOAc-HOAc buffered solutions and also 4  with the stoichiometry represented  Ill-3  in-reaction-IV.  E f f e c t of Carbon Monoxide Pressure Figure 5* shows that the reduction rate i s d i r e c t l y proportional t o  CO p a r t i a l pressure up t o at least 30 atm.  The CO p a r t i a l pressures were  estimated by subtracting the vapour pressure  of H 0 and the p a r t i a l  of a i r from the i n i t i a l saturation pressure.  2  pressure  Small corrections were also made  to take account of the reaction which occurred during CO saturation.  The  deviations from Henry's law f o r the s o l u b i l i t y of CO i n H 0 at 30 atm have 2  been.estimated to be only a few per cent (see Appendix B) and since the poss i b l e errors i n the i n i t i a l rate measurements were between ± 10 and 20$, the reduction rate i s proportional to the concentration of CO i n solution, within experimental error.  Figure 5 also indicates that the rate i s favoured by a  low HOAc/NaOAc r a t i o or by a high NaOAc concentration,  Ill-k  or both.  E f f e c t of Acetic A c i d In view of the large v a r i a t i o n i n rate with pH.(Table I) and HOAc/  NaOAc r a t i o (Figure 5) a series of experiments was conducted over a wide *  Data for rates shown i n Figures and- Tables throughout the text are tabulated i n Appendix.D.  - 25 -  Figure 5.  Dependence of Rate on CO Pressure (90°C)  - 26 -  range of HOAc concentrations  at constant amounts of added NaOAc and AgC10 . 4  The amount of free acetate i n i t i a l l y present was experiment and  [H ] was +  d i r e c t l y proportional to [HOAc] through the r e l a t i o n :  [H+]  =  K  ±  .{HOAc] LOAc-J  where Kj_ i s the i o n i z a t i o n constant of HOAc. experiments, as depicted i n Figure 6,  [HOAc]  -1  ( 1 )  '  v  The r e s u l t s of t h i s series of  indicate that the rate (reduced to unit  pressure*) i s inversely proportional to [HOAc]. vs  therefore i d e n t i c a l i n each  The non-zero intercept of R'  on the R'- axis indicates that there i s also an acid-independent  contribution to the o v e r a l l rate.  Therefore at a constant degree of s i l v e r -  acetate complexing ( i . e . constant amounts of added AgC10 and NaOAc) the 4  apparent rate law i s :  R  R'  where  R'  =  I' =  =  -d[C0]/dt  (2)  =  I ' P c o + S'P /[H0Ac]  (3)  =  R/Pco  (k)  =  I" + S* /[HOAc]  (5)  =  I' + D'  (6)  C0  !  t o t a l rate of CO consumption at u n i t pressure (M atm-i sec-i) acid-independent contribution to the t o t a l rate at unit pressure and constant complexing (M a t m i sec-i) -  *  Since i t has been demonstrated i n Section I I I - 3 that the rate i s d i r e c t l y proportional to the p a r t i a l pressure of CO, rates referred to i n the remainder of the text have generally been reduced to u n i t pressure and represented as R' ( i . e . • R = -d[CO]/Pco = / P c O M atm-i s e c - i ) . dt / 1  R  - 27 -  Figure 6.  Dependence of Rate on [HOAc] and [HOAc] (0.115 M AgC10 ; O.I95 M NaOAcj 90°C) 4  -1  - 28 -  D' =  acid-dependent contribution to the t o t a l rate at unit pressure and constant complexing (M atm  S' =  -1  sec ) - 1  acid-proportionality constant f o r the acid-dependent contribution to the t o t a l rate at unit pressure and constant complexing atm  (M  2  sec ).  -1  - 1  In terms of t h i s nomenclature I slope of R' vs [HOAc]  -1  and S  1  1  are respectively the intercept and  plots at constant Jcomplexing (e.g. Figures 6, 7  8).  a n  d  1  Since the experimental solutions were buffered, an inverse dependence of R*on [HOAc] at constant degree of complexing corresponds to an dependence on [H ]  or a d i r e c t dependence on [0H-]  +  t r i b u t i o n to the rate.  inverse  f o r the acid-dependent con-  This e f f e c t i s s i m i l a r to that reported for the  CO-  reduction of aqueous basic s i l v e r amine solutions at room temperature and atmospheric pressure (14).  In these amine solutions no pH-independent con-  t r i b u t i o n to the reduction rate was  Ill-5  observed.  E f f e c t of Acetate Complexing The e f f e c t on the rate of varying the amount of NaOAc added at  constant AgC104 addition i s shown i n Figure 7 while the e f f e c t of varying amount of AgC10 added at constant NaOAc i s shown i n Figure 8. 4  (I') and slopes (S')  The  the  intercepts  increase with increasing amounts of NaOAc and AgC10  4  indicating that the o v e r a l l rate i s favoured by complexing between s i l v e r and acetate.  Data of R  1  vs [HOAc]" at 90°C were obtained f o r f i f t e e n d i f f e r e n t 1  degrees of silver-acetate complexing as summarized i n Table IV.  Ill-6  Acid-Independent Reaction A plot of the intercept (I') from R' vs [HOAc]-i plots against  the  - 29 -  Figure 7.  Dependence of Rate on [HOAc] at Various NaOAc Levels (0.117 M AgC10 i n i t i a l l y ; 12 atm CO; 90°C) -1  4  - jo -  0  2  k  6  8  10  12  [HOAc]" (M-i) 1  Figure 8.  Dependence of Rate on [HOAc]" at Various AgC10 Levels (0.090 M NaOAc; 12 atm CO; 90°C) 1  4  - 31 TABLE IV  Summary of Intercepts and'Slopes from R' vs [HOAc]  -1  Plots  at Various Degrees of Acetate Complexing *  Wo of Expts  [Ag(I)] M  A  5  .048  .041  .0020  • 15  .017  B  4  .096  .041  .0039  • 31  .026  C  3  .01+5  .O89  .0040  .22  .058  D  3  .115  .043  .0050  • 32  .066  E  4  .090  .080  .0072  .40  .074  F  3  .044  .178  .OO78  .43  .081  G  4  .042  .215  .0090  .43  • 077  H  3  .230  .041  .0094  .50  .133  I  3  .114  .O87  .0099  .44  .178  J  4  .225  .082  .OI85  • 78  .384  K  3  .110  •173  .0190  • 70  .322  L  18  .110  .190  .0209  • 90  .320  M  3  .040  .670  .0268  .87  .249  W  3  .056  .554  .0310  .85  .275  0  5  .056  • 776  .0435  1.10  .340  *  [WaOAc] [Ag(I)][WaOAc] M M2  I'x 10 M,a-i s -  S'x 10 M2 a - i s - i  Series Code  6  s  1  Concentrations of Ag(I) and WaOAc are the averages of the corrected values of a l l experiments i n a p a r t i c u l a r R' vs [HOAc] series as tabulated i n Appendix D. -1  - 32 -  t o t a l [Ag(I)] at various NaOAc l e v e l s produces a group of non-linear curves passing through the o r i g i n , as shown i n Figure 9-  When I' i s p l o t t e d against  [NaOAc] at various Ag(I) l e v e l s a s i m i l a r set of curves i s obtained, as shown i n Figure 10.  Combining the two sets of curves by p l o t t i n g I' vs  [Ag(I)][NaOAc],  as i n Figure 11, produces a single curve which passes through the o r i g i n and approaches an asymptote with increasing [Ag(I)][NaOAc]. that the rate-determining step i n the acid-independent molecule and a complex between Ag  +  and  This evidence  suggests  reaction involves a  CO  OAc".  The t o t a l concentration of s i l v e r complexed with acetate can be approximated by an average value f o r the concentration of undissociated AgOAc molecules  ( i . e . [AgOAc]), which are the most abundant silver-acetate complexes  present i n the solutions studied*.  The acid-independent  rate law at unit pressure might then be expressed I'  where K  a  as:  =  k ' [AgOAc]  (7)  =  ki'KaUg+HOAc"]  (8)  x  i s the average s t a b i l i t y constant of AgOAc.  i n terms of t o t a l [Ag(I)], K  T'  contribution to the  Substituting f o r  [Ag ] +  and free [OAc ] gives**: -  a  -  v  [Ag(I)][QAc-]  , .  or upon rearrangement:  [Ag(i)3  *  1  =  +  1_  Based on a v a i l a b l e s t a b i l i t y constant data at room temperature (38) r a t i o of AgOAc to Ag(0Ac)g i s about 5:1 at 0.2 M NaOAc.  ** [AgOAc] = .-.[Ag ] = +  K [Ag ][0Ac"]  =  +  a  [Ag(I)] -  [ A g ( I ) j / ( l + K [0Ac"]) a  [Ag ] +  ( 1 0 )  the  - 33 -  [NaOAc] (M)  1.3  /  1.2 l.l  1.0 •9 .8 •7  •5 .4 •3 .2  .1  O  /  .6  -  /  / A i / -i /  1/ // A$  O-  0.2  Q-  1  :_ •/'• /  0.7 0.08  / •/ /  _  ©-  /  .  0.0k  /  o  /  /  )  •\ y/ is  i  1  .05  .1  .15  • 25  [Ag(I)] (M)  Figure 9'  Dependence of Acid-Independent Reaction on [Ag(I)] at Various NaOAc Levels (90°C)  - 34 -  0  .1  .2  .3  .4  .5  .6  .7  .8  .9  [NaOAc] (M)  Figure 10.  Dependence of Acid-Independent Reaction on [NaOAc at Various Ag(I) Levels (90°C)  - 35 -  0  .1  .2 [Ag(I)] [NaOAc]  Figure 11.  .3  .h  (M.2)  Dependence of Acid-Independent Reaction on [Ag(I)][NaOAc] (90°C)  .5  - 36 I f the proposed form of the acid-independent correct then a plot of [Ag(I)]/l' vs [OAc ]-  be l i n e a r with a slope of l/k!'K  contribution to the rate law i s 1  according t o equation 1 0 should  and an intercept of l / k ^ .  a  Figure 12 depicts  such a p l o t where [ 0 A c ~ ] has been approximated by [NaOAc], which i s a good estimate at high [NaOAc]. atm-i s e c  - 1  while K  From Figure 1 2 , k i ' has a value of about 3  has a value of about 2 M .  I f corrections are made f o r  -1  a  10-s  x  the amount of acetate e f f e c t i v e l y removed from solution through complexing with s i l v e r , assuming a value of 2 M-i f o r K , the intercept, and hence k i ' , a  remains e s s e n t i a l l y unchanged while the slope becomes f l a t t e r ( i . e . the value for K  a  increases). Using an i t e r a t i v e procedure of t h i s type  values of K  a  between about two and f i v e are obtained from the slope of [Ag(I)]/l' vs [ 0 A c ~ ] - i plots. E.M.F. measurements using c e l l s of the type:  AgOAc NaOAc  Ag  KN0  AgN0 NaN0  3  Saturated  Ag  3 3  have been made at room temperature (39) at i o n i c strength to about two. t h i s reported data, assuming complete d i s s o c i a t i o n of A g N 0  3  Using  and representing  the concentration of Ag(I) present as silver-acetate complexes by [AgOAc], values f o r the r a t i o [AgOAc]/[Ag ][OAc ] = K +  and 3.4 M  -1  -  a  were found t o l i e between 2.8  at i o n i c strengths between 0.1 and 0.9 increasing t o ^.0 M  infinite dilution.  -1  at  To evaluate the temperature c o e f f i c i e n t of K , further a  E.M.F. measurements were made during the present study at temperatures to 90°C and ionic strengths between 0.1 and 0.9 using c e l l s of the type:  Ag  AgC10  NaOAc  4  NaC10  8 .M  4  AgC10 NaC10  4  Ag  4  Complete d e t a i l s of these measurements and a summary of the r e s u l t s are given i n Appendix C.  Values of [AgOAc]/[Ag ][OAc"] = K +  a  assuming complete dissocia-  t i o n of A g C 1 0 were found, within experimental error, t o be e f f e c t i v e l y 4  F i g u r e 12.  Dependence o f Acid-Independent R e a c t i o n on [ A g ( I ) ] and [NaOAc]; p l o t t e d a c c o r d i n g t o e q u a t i o n 10, assuming [0Ac~] = [NaOAc] (90°C)  - 38 independent of temperature t o 90°C and to l i e between 3.0 and k.k M  -1  at ionic  strengths of 0.1 to 0.9. Using a value f o r K  of 3-7 M"  t o calculate the free acetate con-  1  a  centration [OAc-]* at 90°C, a plot of Ag(I)/l'vs [OAc-]" , as i n Figure 13, has 1  an intercept equivalent t o k i ' = 2.5 x 10-s atm-i sec-i and an average slope equivalent to K  = 3-7 M- .  The s l i g h t curvature of the plot i n Figure 13'  1  a  might be due to a small systematic error i n the estimation of [ 0 A c ~ ] , or t o higher order silver-acetate complexes (e.g. A g ( 0 A c " ) 2 ) a f f e c t i n g the rate at high'acetate l e v e l s .  A p l o t of I / [Ag ] vs [OAc ], as i n Figure Ik, exhibits 1  +  -  p o s i t i v e deviations from l i n e a r i t y with increasing [OAc"] and thus the l a t t e r explanation f o r the curvature i n Figure 13. Figure Ik i s equivalent to k^ K atm  -1  s e c " i assuming K  a  = 3•!  The average slope of  and y i e l d s a value f o r k ^  a  supports  of 2.5 x 1 0 5 -  M . -1  The dependence of the acid-independent  reaction on silver-acetate  complexing i s a l s o i l l u s t r a t e d i n Figure 15 where a plot of I' vs [AgOAc] i s l i n e a r , passing through the o r i g i n with a slope equivalent to kx' = 2.5 * 0.6 x 1 0 - 5 atm  -1  sec~i.  The acid-independent  contribution to the o v e r a l l rate  therefore has the form:  or  where  I  =  I' =  ki'' [AgOAc]  I  =  ki' P  =  ki[C0][AgOAc"]  =  k K [C0][Ag ][0Ac ]  c o  (7)  [AgOAc]  (11) (12)  +  1  the acid-independent  (13)  _  a  contribution to the o v e r a l l rate of CO con-  sumption (M s e c ) - 1  ki =  the apparent second-order s p e c i f i c rate constant f o r the a c i d independent reaction ( M  *  _1  sec ) - 1  For [OAc ], [AgOAc] and [Ag ] values using K Section III-7. -  +  a  = 3.7 M  _1  see Table V,  - 39 50  o  .5  io  15  20  25  [OAc ]  35  ^o  45  50  (M )  -  Figure 13.  30 _1  Dependence of Acid-Independent Reaction on [Ag(I)] and [OAc ]; plotted according to equation 10, assuming K = 3.7 M (90°C) -  _1  a  - ko -  [OAc"]  Figure lk.  (M)  Dependence of Acid-Independent Reaction on [OAc ] (90°C) -  - kl -  Figure 15.  Dependence of Acid-Independent Reaction on [AgOAc] (90°C)  - 42 Using a value for the s o l u b i l i t y of CO i n water of 6.9 x 10 - 4 M/atm at 90°C (see Appendix B) the experimental bimolecular rate constant kj. i n equations 12 and 13 has a value of 0.04  Ill-7  ± .01 M-i s e c "  1  at 90°C.  Acid-Dependent Reaction The v a r i a t i o n of the acid-dependent  reaction rate with i n i t i a l [Ag(I)]  and [NaOAc], as summarized i n Table IV (Section III-6), i s depicted i n Figures 16 and 17.  The slope of the S' v s [ A g ( I ) ] curves at constant [NaOAc] (Figure  16) increases with increasing [Ag(I)] and thus indicates that the  acid-dependent  reaction has a higher than f i r s t - o r d e r dependence on [Ag ]. +  Plots of S' vs [NaOAc] (Figure 17) and I' vs [NaOAc] (Figure 10) at constant [Ag(I)] have the same general shape increasing [NaOAc],  i n that the slopes decrease with  This indicates that S' has approximately the same depend-  ence as I' on [OAc ], which has been shown to be f i r s t - o r d e r . -  been demonstrated from R' vs [HOAc] r  that the acid-dependent [0Ac~],  -1  I t has also  plots at constant degree of complexing  reaction i s inversely proportional to [H ] or [HOAc]/ +  The slopes, S' , of R  vs [H0Ac]-i plots therefore have a f i r s t - o r d e r  1  [0Ac~] factor incorporated i n them from the a c i d dependence. A simple rate law, consistent with the shape of plots i n Figures 16 and 17, to describe the acid-dependent  reaction might involve a  second-order  dependence on uncomplexed s i l v e r and an inverse dependence on [H ] or [HOAc]/ +  [OAc ]. -  In t h i s case the slopes of R' vs [HOAc]  -1  plots at constant degree of  complexing would be given by: S'  =  kg [Ag ] [0Ac-] 1  +  (14)  2  I f equation 14 accurately represented the v a r i a t i o n of S' then S'/[Ag ] [OAc ] +  should be constant f o r a l l degrees of complexing.  2  The experimental data, as  -  - 43 -  F i g u r e 16.  Dependence o f Acid-Dependent R e a c t i o n a t V a r i o u s NaOAc L e v e l s (90°C)  on  [Ag(I)]  - kh  -  [NaOAc] (M)  Figure 17.  Dependence of Acid-Dependent Reaction on [NaOAc] at Various Ag(I) Levels (90°C)  - k  -  5  summarized i n Table V, i n d i c a t e s t h a t such i s n o t t h e case, b u t r a t h e r S'/ [ A g ] [ O A c ] i n c r e a s e s w i t h i n c r e a s i n g [OAc~] o r d e c r e a s i n g +  -  v a r i a t i o n i s t a k e n as evidence  [Ag ],  This  +  f o r a t l e a s t a two-term r a t e l a w t o d e s c r i b e  t h e acid-dependent r e a c t i o n , e.g.:  or  S*  =  k ' [Ag+]2[0Ac-] + k "[Ag ]2[0Ac-]2  (15)  S1  =  k ' [Ag ] [OAc-] + k "[Ag ][OAc-]  (16)  +  2  3  +  2  +  2  4  I f e i t h e r e q u a t i o n 15 o r 16 r e p r e s e n t s t h e c o n t r i b u t i o n t o t h e r a t e f o r t h e acid-dependent t e r m t h e n p l o t s o f S " / [ A g ] [ O A c ] v s [OAc~] o r [ A g ] r  +  2  -  +  _ 1  should  be l i n e a r w i t h i n t e r c e p t s o f k '- and s l o p e s o f ~k " o r k " r e s p e c t i v e l y . F i g u r e 2  3  4  18 i n d i c a t e s t h a t S ' / [ A g ] [ O A c ] v s [OAc ] i s l i n e a r as p r e d i c t e d by e q u a t i o n +  2  _  -  15 w i t h a s l o p e o f 3 . 2 ± 0 . 8 x 10~ x 1 0 5 M" -  1  atm  3  M"  a t m ~ i s e c - i and an i n t e r c e p t o f 7 ± 7  2  sec" .  - 1  1  F i g u r e 19 shows t h a t a p l o t o f S ' / [ A g ] [ O A c ] v s [Ag" "]" i s n o t +  l i n e a r , except perhaps a t low [ A g ] +  inadequate  - 1  2  -  1  1  v a l u e s , i n d i c a t i n g t h a t e q u a t i o n 16 i s  t o d e s c r i b e t h e e x p e r i m e n t a l o b s e r v a t i o n s over t h e e n t i r e c o n c e n t r a -  t i o n range i n v e s t i g a t e d .  Increasing p o s i t i v e deviations a t higher  v a l u e s , which a l s o correspond a c t i v i t y o f Ag(0Ac)  2  t o h i g h e r [OAc"],•might be e v i d e n c e  r e l a t i v e t o AgOAc*.  S e c t i o n s I I I - 8 and I I I - 9 ,  [Ag ] +  _ 1  f o r greater  Subsequent measurements p r e s e n t e d i n  however, f a v o u r e q u a t i o n 15 t o d e s c r i b e t h e a c i d -  dependent c o n t r i b u t i o n t o t h e r a t e l a w . Thus t h e r a t e l a w f o r t h e a c i d dependent r e a c t i o n can be e x p r e s s e d a s :  D  *  '  "  k  =  •  k  ='  , [ A 8 + ] 2  [ A g + ] 2  liiiT  +  lisT  +  vibrio*,-! [  A  B  +  ]  ™ { i ^ T  (IT) <> 18  Based on a v a i l a b l e s t a b i l i t y c o n s t a n t d a t a a t room temperature ( 3 8 ) , t h e r a t i o o f AgOAc t o Ag(0Ac) i s about 5 : 1 a t 0 . 2 M NaOAc. 2  TABLE V  Summary of the Dependence of S'on [ A g ] * at 90°C +  *-* Series Code B H D J E A I L K C F G N  0  M  S'x M  2  10s  a - i s-  .026 .133 .066 .384 .074 .017 .178 .320 .322 .058 .081 .077 .275 .340 .249  1  [Ag(I)] M  .096 .230 • 115 .225 .090 .048 .114 .110 .110 .045 .044 .042 .056 .056 .040  [NaOAc] . M  .041  .041 .043 .082 .080  .041  .087 .190 .173 .089 .178 .215 .550 .780 .670  [AgOAc] M  [Ag ] M  [OAc-] M  .010 .018 .011 .034 .017 .OO56 .022 .039 .036 .010 .016 .017 .036  .086 .212 .103 .191 .063 .042 .092 .071 .073 .035 .027 .024 .019 .015 .012  .031 .023 .031  .041  .028  +  Values of [Ag ] and [AgOAc] were estimated using Ka = 3-7 See Appendix D f o r Series Code references  .048  .073 .035 .065 .151 .136 .078 .161 • 197 .513 • 739 .642  M"  [Ag*]" M" 1  11.6 4.7 9.7 5.3 15.9 23.6 10.9 14.1 13.7 28.6 36.4 4l.2 51.3 66.7  84.4  1  S'x 104 [Ag ] l0Ac-] M-i a ' l s - i +  2  1.1 1.3 2.1 2.2 2.6 2.7 3-3 4.3 4.4 6.1 7.0 7-0  14.8  20.4 27.4  - 4  Figure 18.  7  -  Dependence of Acid-Dependent Reaction on Acetate Complexing; p l o t t e d according to equation 15 (90°C)  Figure 19.  Dependence of Acid-Dependent Reaction on Acetate Complexingj p l o t t e d according to equation 16 (90°C)  - k9  -  The o v e r a l l rate law including both the acid-independent  and the acid-dependent  reactions then becomes:  R' = ' [AgOAc ] + k ' [ A g ] | ^ l | . +  k l  R  2  2  = k [CO] [AgOAc"] + x  +  3  A  katCOltAg+FJI^j +  k  From Figure 18, k ' = 7 ± 7 x 10-5 M"  1  2  = 9.0*  2.1 x IO" M 4  _1  atm  -1  sec . - 1  3  [ 0] C  2  sec  - 1  A  +  atm" s e c " and k ' = k "/K 1  1  3  Using a value of 6.9 x I O  90°C.  From Section III-5 k  Ill-8  Acetate-Independent Reaction  x  and 1.3  (20)  [Ag ] [AiOAc"]-[| |l|  s o l u b i l i t y of CO i n water (see Appendix B) y i e l d s values f o r k equation 20 of 0.1 ± 0.1 M-  (19)  + k ' [ g ][AibAc"]-[|^  - 0.3 M  -2  sec  has a value of 0.0k ± .01 M  - 1  _1  In equation 19 the k ' -term i s acetate-independent 2  - 4  3  a  M/atm f o r the  and k i n  2  3  respectively, at sec . - 1  since [HOAc]/  [OAc ] i s simply a measure of [ H ] . T O v e r i f y the form of the acetate-  +  independent contribution to the rate law several series of experiments were conducted at various low WaOAc l e v e l s using constant ratios.  [Ag(I)] and HOAc/WaOAc  Extrapolation of each series t o zero [NaOAc] y i e l d s a measure of the  reduction rate i n the absence of acetate complexing.  Several p l o t s of t h i s  type are depicted i n Figure 20. Although the t o t a l [Ag(I)] and HOAc/NaOAc r a t i o were held constant during each series, [Ag ] and [HOAc]/[OAc ] (or [H ]) v a r i e d with each experi+  -  +  ment due to d i f f e r i n g degrees of silver-acetate complexing at the various acetate l e v e l s used.  To minimize these effects  plots of R'[HOAc]/[AgOAc]  vs [OAc ] calculated using Ka = 3'7 M-i were extrapolated t o zero [OAc ] as -  shown i n Figure 21.  -  Rearrangement of equation 19 according to equation  21,  - 50 -  Figure 20.  Dependence of Rate on [NaOAc] at Constant and [HOAc]/[NaOAc] (90°C)  [Ag(I)]  - 51 -  0  .02  .06  .Ok  [OAc ] -  Figure 21.  .08  .10  .12  (M)  Dependence of R [HOAc]/[AgOAc] on [ 0 A c ~ ] at Constant [Ag(I)] and [HOAc)/[NaOAc]; plotted according to equation 21 (90°C) 1  - 52 -  indicates that the intercepts of plots of t h i s l a t t e r type are equivalent to k ' [Ag ]/K +  2  a  while the slopes are equivalent to average values of k i ' [HOAc]/ Table VI summarizes the intercepts, ^R' [HOAc]/[AgOAc])  [OAc-] + k ' [ A g ] . +  3  ,  Q  obtained from ten R"[HOAc]/[AgOAc] vs [OAc ] plots including those shown i n -  Figure 21.  Since the intercepts are equivalent to k ' [Ag ]/K  a plot of  +  2  (R [HOAc ]/[AgOAc ] ^  vs [Ag ] should be l i n e a r  1,  passing through the o r i g i n  +  0  with a slope of k '/K . 2  a  Such a p l o t , as depicted i n Figure 22,  a  indicates that  the data, although e x h i b i t i n g a high degree of scatter (due presumably to the inherent d i f f i c u l t y of measuring slow rates by the pressure-drop  method), i s  consistent with a l i n e a r r e l a t i o n s h i p having a slope of 8 ± 4 x I O Thus the acetate-independent expressed  =  or  where  R R  -  °  ~  =  k  ,  [Ag+f [H J  K  +  -1  [CO][Ag F[OAc-]  [CO][Ag ]  +  1:  K 2  THOACT  +  =  y  K  K  2  K  L  ' •  K  2 f  25)  — [ I T —  (  O  ;  i s the i o n i z a t i o n constant of acetic a c i d . From the slope of Figure 22, using a value of 3.7  30 i 15 x 10-5 M = 0.4  t 0.2  -1  atm  -1  s e c , and using a value of 6.9 - 1  M-2, s e c - i at 90°C.  to be 7 * 7 x 10-5 M  Ill-9  sec .  -1  contribution to the o v e r a l l rate law can be  ytAg^OAc-] [HOAc]  °  2  atm  as: V  k  - 5  Reduction  _1  atm  -1  sec  of Unbuffered  - 1  In Section IJI-7, k ' 2  and 0.1  ± 0.1 M  -2  M 10  x  for K ,  _1  2  =  HOTOLQQ,  were estimated  s e c - respectively. 1  S i l v e r Perchlorate  A more d i r e c t estimation of the acetate-independent  contribution to  the o v e r a l l rate can be made from rate measurements i n unbuffered solutions.  2  M/atm  -4  and k  k'  a  AgC10  4  The perchlorate ion i s among the weakest complexing ligands f o r  metal ions i n aqueous s o l u t i o n (40).  Thus i n a solution of  AgC10  4  i n the  - 53 -  TABLE VI  Summary of:Acetate-Independent• Rates At Various HOAc/NaOAc,ratios and Ag(I).Levels at 90°C  Series* Code  No of Expts  [Ag(I)] M  [HOAc] [NaOAc]  / i[HOAc] \ R  \  [AgOAc]/o  M a-i s-i  x 10  *  P  6  .050  1  4.6  Q  3  .100  1  6.8  R  2  .123  1  9-6  S  2  • 235  1  13.7  T  5  .100  5  6.0  U  2  .064.  8.7  4.4  V  2  .130  8.7  18.0  W  2  .163  8.7  19.4  X  3  .233  8.7  12.8  Y  4  .270  8.7  18.4  See Appendix D f o r Series Code references  s  - ^  -  ,[Ag(I)]  Figure 2 2 .  (M)  Dependence of Acetate-Independent Reaction on [Ag(I)] ( 9 0 ° C )  - 55 -  absence of other complexing agents. the s i l v e r can be considered to exist exclusively as the aquo-complex.  Unfortunately the rate of CO consumption i n  such a system i s too slow to be e a s i l y measured using the method.  pressure-drop  By measuring the pH of periodic l i q u i d samples, however, the CO  consumption can be calculated using the stoichiometry represented by reaction IV ( i . e . -2d[C0]/dt  =  d[H ]/dt): +  2Ag + + CO + H 2 0  >2Ag + C0 2  + 2H +  (IV)  Four such experiments were conducted at varying [AgClO.4] using constant CO pressures of about 55 atm at 90°C. consumption of A g  and the production of R"*  +  1  The stoichiometry between the  as shown i n Table VII i s con-  sistent with reaction IV. In Section III-8 i t was  shown that the acetate-independent  reaction  can be represented by a rate law of the form:  V  =  (22)  k '-Ki[Ag+]2/[H ] +  2  The duration of each experiment reported i n the present section was  equivalent  to about 15$ reduction and hence [Ag ] as w e l l as the CO pressure remained +  e s s e n t i a l l y constant during each experiment.  Equation 22 i s thus equivalent  to: d[H+]/dt  =  v2d[C0]/dt  =  k "/[H ]  The e f f e c t of CO pressure was not investigated and i s assumed to be order.  •*  (25)  +  2  first-  Since [Ag ] remained constant during each experiment i t i s not +  Concentrations of H were calculated from pH measurements made with a Beckman pH-meter (Model G) calibrated against mixtures of standard HC104 and AgC104 solutions using glass and saturated calomel electrodes with an a d d i t i o n a l saturated potassium chloride and f i l t e r paper s a l t bridge. +  - 56 -  TABLE VII  Stoichiometry of Unbuffered CO - S i l v e r 'Perchlorate Reaction 1  Expt No  [AgC10 ] (M) Final Initial  at"9'0°C and 55 atm CO  [H ] (M) Initial  Final  Produced  A[Ag ] A[H j  .0081  .0081  0.9  +  4  Consumed  +  +  262  .0530  .0458  .0072  7 x 10"  260  '.1040  .O897  .0143  1 x  10-5  • 0135  .0135  1.1  261  .211  .179  .032  3 x  10-5  • 0339  .0339  0.9  263  1.060  .850  .210  1 x 10-  .204  .204  1.0  6  4  - 57 necessary to assume a second-order solution.  The rate constant k '" 2  k "'  =  2  [Ag ] dependence i n unbuffered AgClC>4 +  i n equation 23 i s thus equivalent to: 2k 'Ki[Ag+] P  (2k)  n  2  C0  Integration of equation 23 from zero to t y i e l d s :  [H ] +  - [H ]|  2  =  +  2k "*t  But rapid i n i t i a l reaction quickly makes [ H ] +  to [H ]. +  (25)  2  0  n e g l i g i b l y small with respect  Taking logarithms of equation 25 gives: log [H+]  =  0.5 log t + log k "'  (26)  pH  =  -0.5 log t - log k "'  (27)  2  2  Plots of pH vs log t at constant [Ag ] and Pco should be l i n e a r with slopes of +  -0.5-  Figure 23 shows that t h i s prediction i s e s s e n t i a l l y correct  with s l i g h t  deviations from l i n e a r i t y being explained by the decrease i n [Ag ] and the +  simultaneous formation of small amounts of C0  2  during each experiment.  Values of k '" can be obtained from Figure 23 using the d i f f e r e n t i a l 2  logarithmic form of equation 23:  k 2  -  =  M f t  *(log[H+]) d(log t )  v  '  A summary of k "' values obtained i n t h i s manner, assuming d(log[H ])/d(log t ) +  2  = 0.5,  i s given i n Table VIII. The apparent dependence of k '"  on [Ag ] can also be found from +  2  Figure 23 by p l o t t i n g pH values, obtained at equal times f o r each of the four experiments, against log [Ag+].  That the slope of such a plot should be  -O.5  times the apparent order of the [Ag ] dependence follows from equations 28 and +  2k  where at constant t , Pco, and d(log[H ])/d(log t ) : +  - 58 -  3.6  A  I 2.0  I  I  I  I  I  I  I  I  I  I  .2  .4  .6  .8  3.0  .2  .4  .6  .8  4.0  I  I  .2  .4  log tjme  Figure 23.  Reduction Rate of AgC10 i n Unbuffered Solution; plotted according to equation 27 (53 atm; 90°C) 4  TABLE VIII Summary of Experimental Rate Constants i n Unbuffered S i l v e r Perchlorate Solutions at 90°C and 53 atm CO  t = 1000 sec log [Ag ]  [Ag ] M +  [H ] M +  pH  +  x 10  *  k "'* M s"  t = 10,000 sec k " ** a- s - i  2  2  1  x 10  3  [H ] M +  2 1  pH  x  9  k * M sm  2  2  0.051  2.71  2.65  2.24  2.6 x 10-9  9-5  2.14  0.097  2.99  2.40  3.98  7.9 x 10-9  7-9  1.88  13.2  8.7 x IO"  0.195  1.24  2.11  7.76  3.0 x 10-8  7-5  1.55  28.2  4.0 x IO"  O.96  I.98  1.33  46.8  > K 2  -  [H ] • t  k"  =  0.5 k " ' / [ A g ]  2  +  2  6  11.3  O.79  d(log[H ]) „ J ) ° ' according to equation 28 d(log t ) +  n  L  +  2  1.1 x 10"  2  P  c o  o  Q  according to equation 34  162.  Mean k" a-i s - i  x 10  x 10  2  1  103  7.25  k " ** a-i s-i  2.6 x IO"  1.3 x IO"  9  2  9  9-5  9-5  9  8.7  8.3  8  9-9  8.7  6  13.3  12.3  9  - 60 -  k '" 2  =  constant x' [ H ]  =  constant x-[Ag+]  +  .". pH =  (28)  2  (24)  n  -0„5n l o g [Ag ] + constant  (29)  +  Two plots of t h i s type prepared using the data summarized i n Table VIII are presented i n Figure 2k, and indicate that the rate of CO-reduction of unbuffered A g C 1 0  solution i s second-order i n [Ag ], +  4  In constructing Figure  2k no account was taken of the possible e f f e c t on ks'" of ionic strength which varied from about 0.05 t o 1.0 throughout the experimental series (but remained e s s e n t i a l l y constant during each i n d i v i d u a l experiment). Rate measurements of the CO-reduction of unbuffered A g C 1 0 assuming a f i r s t - o r d e r dependence on Pco,  a  r  e  4  solutions,  consistent with a one-term  experimental rate law of the form:  .R  Q  or R'. 0  Equation 32  =  -d[C0]/dt  =  0.5  d[H+]/dt  (30)  =  k " [ A g + ] 2 P /[H+]  (31)  =  k "[Ag ]2/[H ]  (32)  2  C0  +  +  2  . i s of the same form as equation 22, thus supporting the experi-  mental rate law developed t o represent the rate of CO-reduction of acetatebuffered A g C 1 0  solution.  4  From a comparison of equations 32, 22 and 2k i t i s  evident that:  k" 2  =  k ''K  k"  =  0.5 k " ' / [ A g + P P  2  2  (33)  ±  2  (34) ,  CO  Average values of k " calculated using equation 34 are included i n Table VIII 2  and l i e i n the range 8.3 - 12.3 x 1 0 - 9 atm-i s e c - i .  The 40$ v a r i a t i o n i n k " 2  may be due l a r g e l y to ionic strength effects p a r t i c u l a r l y with the solution O.96 M i n A g C 1 0 . 4  The average of the other three values i s 8.8 x 1 0 - 9 atm-i  s e c - i and may be compared with a k ' o f 3 . O i l . 5 x 1 0 - M 4  2  _1  atm  -1  sec-i  Figure 2k.  Dependence of Rate on [Ag ] i n Unbuffered Solution (53 stm; 90°C) +  - 62 -  estimated i n Section III-8 f o r the acetate-independent reaction i n buffered solutions.  Thus from equation 33 K i should have a value of about 3 x I O  M  - 5  to make the rate measurements i n the buffered and unbuffered systems consistent . The above result i s i n agreement with accepted values f o r Kj_ taking into account the e f f e c t of ionic strength on the i o n i z a t i o n of weak acids i n aqueous solution.  For example the i o n i z a t i o n constant of HOAc i n NaCl solution  at 25°C varies from 1.75  x  IO  to a maximum of 3-32  - 5  (kl).  strength of zero and 0.5  x 10-5 between an ionic  The e f f e c t of temperature  K i f o r HOAc i n water to pass through a maximum of I.76 and f a l l to I.55 x 10-  5  at 60°C (kl).  causes the value of  x 10  -5  at about 25°C  Extrapolation of t h i s data to 90°C  y i e l d s a value of about 1.2 x 1 0 .  Thus i t appears from a comparison of the  -5  rates of CO-reduction of AgC10 i n buffered and unbuffered solutions that the 4  ionic strength rather than the temperature  has a greater e f f e c t on the i o n i z a -  t i o n of acetic a c i d i n HOAc-NaOAc solutions at 90°C. The results of the rate measurements made i n buffered and unbuffered solutions are i n good agreement considering p a r t i c u l a r l y the widely different methods used to obtain the measurements and the approximately hundred-fold difference i n o v e r a l l rate.  "Best Value" Rate Parameters at 90°C  Ill-10  The CO-reduction of AgC10 i n acetate-buffered solution i s consist4  ent with a rate law represented by equations 19 and 20:  y  +] ^oAc"]|P^l|  R' =  k ' [AgOAc ] + k ' [ A g ] 2 ^ A g l j  R  k [C0][Ai0Ac"] + k [ C 0 ] [ A g ] | 0 A c 4 + k [CO ] [Ag ] [ A i o A c " ] - ^ ! ! LHOAc J IHOAc J  =  +  x  2  +  1  2  +  [Ag  [A  +  a  3  ( 1 9 )  (20)  - 63 Values for the rate constants at 90°C i n equation 19, found by graphical analysis i n Sections III-6, III-7 and III-8 are summarized i n Table IX*  Also  included are the results of least square regression analyses on equation 19 using various sets of experimental data.  The regression c o e f f i c i e n t s were  evaluated on an IBM 1620 d i g i t a l computer using a s l i g h t l y modified S3-4 s t a t i s t i c a l analysis program available from the program l i b r a r y of the U.B.C. Computing Centre.  Unfortunately the program considers the experimental rate  data t o have equal possible errors on an absolute rather than a r e l a t i v e basis and thus the higher values of R'are given undue weight. From consideration of the various sets of rate parameters, derived by graphical and s t a t i s t i c a l methods and summarized i n Table IX, a set of "best value" parameters at 90°C f o r equations 19 and 20 have been selected and presented i n Table X.  The r e l i a b i l i t y of the acid-independent parameter i s  estimated to be ± 25$ while the r e l i a b i l i t y of each of the two acid-dependent parameters i s estimated to be ± 50$. As a further v e r i f i c a t i o n of the proposed rate law, equation 19 was integrated numerically using the "best value" rate constants from Table X and compared with the pressure records of several extended reduction experiments. Figure 25 shows the results f o r f i v e experiments, two of which represent about 50$ reduction of AgC10 . 4  Experiments  of longer duration were not made because  of the p o s s i b i l i t y of minute gas leaks producing erroneous results f o r extended measurements of slow rates. i s given i n Appendix F.  A sample calculation f o r the numerical integration  The calculated pressure-time curves are i n good agree-  ment with the experimental curves and thus equation 19, developed from measurements of i n i t i a l rates, i s adequate to describe the CO-reduction of AgC10 i n 4  acetate-buffered solutions f o r at least 50$ reaction.  *  The rate constants i n equation 20 can be obtained by dividing the corresponding constants i n eqution 19 by aQQ = 6.9 x 1 0 - 4 M/atm.  - 6k  -  TABLE IX Summary of Rate Parameters f o r Equation 1  at  9  90°C  Rate Parameters ki'x 105 a" s-  No of Expts  Varied [HOAc] at constant [NaOAc] and [AgC10 ] ". Series A-0  67  1  2.5  1  k 'x 104 M-i a - i s " 2  k 'x 10* M-i a - s - i 3  1  1  Figures Ik, 15  ± 0.6 0.7  67  Reference  ±  0.7  9.0 ± 2.1  Figure 18  5-7  Regression Analysis  4  3.0  67  Varied [NaOAc] at constant [AgC10 ] and [HOAc]/[NaOAc] Series P-Y  51  0.9  3-0  —  --  ± 1.5  4  A l l Experiments  Figure 22  31  2.3  2-5  3-0  Regression Analysis  119  3-0  l.k  5-8  Regression Analysis  Equation 19 R' = k ' [AgOAc] + k ' [ A g ] +  x  *  2  2  + kg' [Ag ] [AgOAc +  ]-[§|f^  To obtain rate parameters i n terms of [CO] divide values i n Table IX by  a  = 6.9 x IO- M/atm. 4  65  TABLE X  "Best Value" Rate Parameters at 9Q°C  Equation 20  Equation 19  = 2.7  X  10-5 a - i  k ' = 2.1 x 10" M-  a-  k ' = 6.2 x 10"  a-i  4  1  2  3  4  M-i  = 0.04  s-i  1  s"  s-  k  2  = 0.3  M-2  S"  k  3  = 0.9  M-2  s-i  1  s-i  M-i  Reliability  ± 25$  1  ± 50$  1  ± 50$  Equation 19  R' = kx' [AgOAc] + k ' 2  [Ag+]2j|A^|  + kg' [Ag ] [AgOAc ]|^£l-| +  Equation 20 OAc-] R = kitCO][AgOAc] + k 2 [ C 0 ] [ A g j [ ~ ] + k [CO ] [Ag ] [AgOAc ]• .HOAc ] [HOAc J +  2  0 A c  +  3  o> 0\  Time (min) Figure 25.  Comparison of Experimental and Calculated Pressure Records ( I n i t i a l Conditions: 0;115 M AgC10 ; O.I95M.NaOAc; 5.4 atm CO; 90°C) 4  - 67 III-11  Proposed Mechanism :  The k i n e t i c s of the CO-reduction of acetate-buffered AgC10 solutions 4  are described by the experimental rate law: R  =  k [CO] [AgOAc"] + k [ C O ] [ A g ] j P i ^ + katCO] [ A g + ] [ A i O A 7 ] l | £ | ^ +  x  2  2  (20)  In terms of [H ] equation 20 may be re-written as: +  R  =  k  kJCOHAiOAc-] +  2 K i  [ C  °^f  ] g  +  ^ K± [C0]^llf^l  (35)  3  A mechanism which i s " consistent with equation 35  c  a  be represented  n  by the following scheme: -. Ag  + OAc  AgOAc  0  +  -  v, CO  k  AgOAc  (rapid equilibrium)  >Ag-JJ-0Ac  a  Ag-C-OAc + Ag+ + H 0  f  a  s  t  2  Ag  +  (b)  > 2Ag + C0  + CO + H a O - J & s Ag-C"-0H S-H-OH + H  Ag-^-OH + A g  +  (a)  k  > 2Ag + C0  b  2  + HOAc + H  (c)  +  (rapid equilibrium)  +  + H  Ag-C-OH + AgOAc —££->2Ag + C0  2  (d) (e)  +  + HOAc  2  (V)  (f)  The rate law derived from t h i s sequence corresponds t o :  R =  =  k [CO][AgT0A7] a  k ^ M ^ p  +  which i s i d e n t i c a l with equation 35 i f k  + W "  Ag^ [AgOAc]  00  a  = k , kbKc = k K i , k K x  2  c  c  ^  = k K i and 3  i f [H 0] i s incorporated i n K . 2  c  Support f o r the nature of the proposed intermediates i s drawn from studies on the CO-reduction of H g  + +  Ag(I) amines i n basic solution (Ik),  i n d i l u t e HCIO4 (15), and the H g  of Mn0 ~ i n both a c i d and basic solution (15)• 4  + +  the reduction of  and A g catalyzed reductions +  Analogous  intermediates were  proposed i n these cases with the evidence f o r (HgJj-0H) being p a r t i c u l a r l y +  - 68 -  strong because of the existence of the stable methyl formate derivative AcO-Hg J!-OCH formed when CO reacts with methanolic 3  (30).  solutions of Hg(OAc)  2  The oxidation of CO i n aqueous solution i s apparently f a c i l i t a t e d by  the presence of an oxygen-donating base (e.g. OH , -  OAc",  H 0, 2  Mn0 "). 4  The reduction of s i l v e r acetate by a pH-independent mechanism i s s i m i l a r to the reduction of H g  + +  and Mn0 ~ (see mechanism I I I , Section 1-4), 4  while the pH-dependent contribution i s s i m i l a r to the reduction of s i l v e r amines i n basic s o l u t i o n (see mechanism I I , Section 1-4).  In the l a t t e r case  the formation of the LAgJ!-OH complex (where L denotes an amine ligand) i s rate-determining as evidenced by a f i r s t - o r d e r dependence on [Ag(I)] and Apparently at [NH ] greater than 0.02 +  4  M (i.e.  lower pH) with NH  ligand the decomposition of the complex to form f i n a l products  3  [OH ]. -  as the amine  i s retarded to"  the extent that the rate i s determined by competition between t h i s step and a back-reaction to form i n i t i a l reactants.  In the reduction of Ag(I) i n a c i d  solution the rate of complex decomposition to products r e l a t i v e to backreaction to form i n i t i a l reactants must be decreased even further u n t i l the complex formation becomes a pre-equilibrium and the rate i s determined exclusively by attack of another Ag(I) species (e,.g. A g complex to form f i n a l  +  or AgOAc) on the  products.  A hydrolyzed s i l v e r ( I ) species (e.g. AgOH) i s not required to explain the observed k i n e t i c s i n a c i d s o l u t i o n . and OH  -  i s given as about 10  4  The association constant  of  Ag  +  at room temperature i n d i l u t e s o l u t i o n (42).  Thus while LAgOH i s the predominant s i l v e r species i n basic amine solutions (assuming  f o r AgL  +  i s of the same order as f o r Ag ) +  AgOH i s present  only  i n trace amounts i n a c i d s o l u t i o n . In the proposed mechanism f o r reduction of Ag(I) i n acid solution reaction V(d) i s indistinguishable from Ag  +  0 + CO + 0 H " ^ = ^ Ag-C-OH  - 69 preceded by the water d i s s o c i a t i o n equilibrium : H 0 ^=£- H 2  +  + 0H-  .TnepH e f f e c t arises i n an equilibrium preceding the rate-controlling step and therefore i t s exact nature cannot be determined  from the present k i n e t i c  study. The mechanism proposed t o describe the CO-reduction  of AgC10 i s 4  probably the simplest which i s consistent with the present k i n e t i c study i n both unbuffered and acetate-buffered solutions. Further contributions of silver-acetate complexes can be proposed but i t i s d i f f i c u l t t o assess t h e i r v a l i d i t y on an experimental basis.  For example Ag(0Ac) might attack the 2  intermediate complex formed i n V(d) by a reaction p a r a l l e l to V(e) and V ( f ) , viz: 0 Ag-C'-OH + Ag(OAc)  ^d > 2Ag + C 0  2  + HOAc + OAc"  2  (V)(g)  Also AgOAc might form an intermediate by an equilibrium analogous t o V(d), v i z : , ^ c ^  AgOAc + CO + H 0  0 (AcOAg-C'-OH) " + H  K  2  (V)(h)  +  This intermediate could be attacked through processes p a r a l l e l to reactions V(e), V(f) and V(g), v i z : ( AcOAgil-OH)" + A g — e _ +  Ag  k  > 2  + C0  2  + HOAc  (V){i)  0 . ( AcOAg-fi-OH)" + AgOAc — ^ £ _ > 2Ag + C 0  + HOAc + OAc"  2  (V)(j)  0 (Ac0Ag-6-0H)" + Ag(OAc) __ JL_ k  2  >  2Ag + C 0 + HOAc +20Ac"  (V)(k)  2  I f the association constants of AgOAc and Ag(OAc) from A g and OAc are +  -  2  represented by K  x  and K  2  r e s p e c t i v e l y and [H 0] i s incorporated i n each of K  and K ' , the acid-dependent c  2  c  reduction rate, considering a l l contributions from  - 70 reaction V(d) to V(k), can be represented by:  D = [C0][Ag ]2 [H ] +  +  (k K * b  c  + (k K Ki + k K 'Ki) c  c  e  c  [0Ac~] +  (k K K  + kfK 'K^^OAc"] . " + kgKc'KiKatOAc-] ) (37) d  c  2  c  3  The sum of the constants of l i k e terms i n equation 37 can be defined by new parameters C  n  (n = 1,2,3,4) and the expression reduced to equation J>8 which  represents a power series i n [OAc ] t o describe the k i n e t i c s of the pH-dependent -  CO-reduction of acetate-buffered AgC104 solutions. D  = [H ] +  n  =  2 i  CjOAc-f"  (38)  1  By a suitable choice of parameters equation J>Q can be f i t t e d t o almost  any  set of rate measurements, be they good or bad, with a high degree of p r e c i s i o n . Graphical analyses of the measurements made i n the present investigation are consistent with a rate expression t o describe the pH-dependent reaction that involves only the i n i t i a l two terms i n equation 38 and reactions V(d), V(e) and V(f) give r i s e to a rate law of t h i s form. Similar considerations t o those presented above can be applied to the acid-independent  reduction process and again i t i s concluded that the rate  law a r i s i n g from reaction V(b) i s of the simplest form consistent with experimental observation. Many ions with a d  1 0  structure, including Ag , C u , A u and H g , +  +  +  ++  tend t o form l i n e a r l y co-ordinated complexes'* and thus i t i s u n l i k e l y that Ag(0Ac)  2  i s capable of reacting without p r i o r d i s s o c i a t i o n t o form-an i n t e r -  mediate complex s i m i l a r to those described by reactions V(d) and V(h). For  *  I t has been suggested (43) that t h i s e f f e c t i s due t o hybridization of d and s o r b i t a l s which remove charge from the region between the metal ion and i t s ligands and thus favours l i n e a r co-ordination for those &10 ions with s u f f i c i e n t l y low d-s separations.  2  - 71 similar reasons i t i s doubtful that the Ag(0Ac)3 complex proposed to help explain some e q u i l i b r i a studies i n aqueous solutions (hk) plays a role i n the reduction mechanism, i f i n fact the'complex e x i s t s .  In the interpretation of  another study on e q u i l i b r i a i n aqueous s i l v e r acetate solutions an Ag OAc  +  2  complex has been postulated (^5).  The existence of such a complex r e f l e c t s  an a f f i n i t y of AgOAc f o r another s i l v e r ion. Because of the higher b a s i c i t y of water compared t o HOAc, which i s analogous t o AgOAc, t h i s second s i l v e r ion should strongly prefer t o remain i n the simple hydrated form, and therefore the existence of s i g n i f i c a n t concentrations of Ag OAc  i s doubted.  +  2  The amount of silver, present i n the various acetate complexes has been approximated by an average [AgOAc] calculated from a mean association constant ( K ) . a  I f the r e a c t i v i t y of AgOAc and Ag(OAc)  2  were s i g n i f i c a n t l y  different deviations from the proposed rate law would be expected p a r t i c u l a r l y at high [OAc ]. . Reference to Figure 18,- Section III-7, i n which S'/ [Ag ] [OAc~] -  +  2  i s plotted against [OAc ], shows that no deviation i n the acid-dependent rate -  expression i s apparent at [OAc~] up to about 0.8 M, where AgOAc and Ag(0Ac)  2  should be present i n approximately equal concentrations, based on published room temperature complexity constants.(38). AgOAc and Ag(0Ac)  2  I t i s therefore concluded that  are approximately equally reactive i n the pH-dependent  reduction process. Reference to Figure lk, Section III-6, i n which l ' / [ A g ] +  i s plotted against [OAc ], gives some evidence that Ag(0Ac) -  2  may be more  reactive than AgOAc toward direct attack by CO. S i l v e r carbonyl intermediates have been proposed t o explain the experimentally observed rate of CO-reduction of d i l u t e Ag S04 solutions 2  buffered with O.65 M NH 0Ac- (13). 4  At the concentrations used the majority of  the s i l v e r was complexed with ammonia.  The e f f e c t of pH was not investigated  and thus a f i r s t - o r d e r dependence on [CO] and an apparent second-order dependence on [Ag(I)] was taken as evidence f o r the rate-determining step involving  - 72 the reaction of H 0 with A g C 0 2  Section 1-4).  2  + +  formed i n a pre-equilibrium^(see mechanism I,  More recent studies (14) have shown that at [NH ] greater than +  4  about 0.02 M the reduction of Ag(NH3)0H approaches second-order i n both 0 +  [Ag(I)] and [OH ]. -  In t h i s l a t t e r case a mechanism involving NH Ag-C -0H as an l  3  intermediate was proposed.  The existence of a s i l v e r carbonyl complex i s s t i l l  f e a s i b l e p a r t i c u l a r l y since analogous stable Cu(I) carbonyl complexes are known (25).  Thus the observed k i n e t i c s f o r the CO-reduction  of unbuffered AgC10  4  might be explained by a mechanism s i m i l a r t o : Ag  +  + CO ^---^ AgC0  AgC0  A rate law developed [AgC0 ] +  solution.  +  + OH"  +  > products  from t h i s scheme using the steady-state approximation  for  successfully described the experimental observations i n unbuffered When acetate contributions were considered, however, a s a t i s f a c t o r y  simple expression could not be developed.  I t was therefore concluded that  such a mechanism i s not responsible f o r the CO-reduction  of Ag(I) i n a c i d solu-  tion . S i l v e r carbonyl complexes formed i n p r e - e q u i l i b r i a might take part i n the formation of Ag-C^-OH i n V(d), e.g.:  Ag  +  + CO ;==^ AgC0  AgC0  +  +  + H 0 ^ = ^ AgJ-OH + H  +  2  This implies that a CO molecule f i r s t co-ordinates with an A g 0  +  ion before  reacting further with an H 0 molecule to form the Ag-£-0H intermediate. 2  K i n e t i c a l l y , such a process i s indistinguishable from the d i r e c t i n s e r t i o n of CO between A g and a co-ordinated H 0 +  2  molecule.  S i l v e r hydride complexes, which are postulated to be active i n t e r mediates i n H -reduction processes i n aqueous solution (46,47), apparently do 2  - 73 not influence the k i n e t i c s of corresponding CO-reduction processes. mechanism by which each gas reacts with Ag(I) i s s p e c i f i c ; H  2  The  i s activated by  dissociation, while CO is' oxidized by the transfer of an oxygen atom from a donor-base (e.g. H 0, OAc-).  In both cases basic ligands increase the metal  2  ion r e a c t i v i t y through s t a b i l i z a t i o n of protons released i n the reduction processes. I t i s u n l i k e l y that nucleation or growth of s i l v e r c r y s t a l s influence the reduction k i n e t i c s .  Nucleation generally involves a high-order dependence  on metal ion concentration, f o r example, the disproportionation of Cu(I) i s tenth-order  i n [Cu(I)] (48), and growth rates f o r most metals are also fast as  evidenced by low overvoltages  required f o r electrodeposition (49). Trace  amounts of p r e c i p i t a t e d s i l v e r from previous experiments were usually present i n the reactor and served to minimize possible nucleation e f f e c t s .  Ill-12  E f f e c t of Temperature The e f f e c t of temperature on the rate of CO-reduction of AgC10 i n 4  acetate-buffered  solution was determined at 60, 80, 90 and 110°C  the rate of CO consumption at a constant number of HOAc concentrations.  by measuring  degree of acetate complexing at a  The r e s u l t s of these measurements are shown  as R' vs [H0Ac]-i plots i n Figure 26 and summarized i n Table XI.  In Section  III-4 the o v e r a l l reaction was shown t o be made up of an acid-independent and an-acid-dependent component, i . e . : R' = =  I ' + S/[H0Ac]  Analysis of the intercepts of R' vs [HOAc] I'  (6)  I' + D'  -1  (5) plots at 90°C indicated that:  = ki''[AgOAc]  (7)  - 74 -  30  0  -  5  1°  15  [H0Ac]-i  Figure 26.  20  25  30  (M-i)  Dependence of Rate on [H0Ac]-i at 60, 80, 90 and 110°C  TABLE XI Dependence of Reduction Rate on Temperature  Temp °C  103/T  [NaOAc] M  [AgOAc] M  I ' X 106 M a" s-  S' X 106 a s  ki'x 106 a-i s - i  ° _I  [Ag(I)] M  60  3.00  .099  .134  .028  0.1k  0.029  5.0  6.70  80  2.83  .110  .190  • 039  0.64  0.164  16.3  5.21  7 22  90  2.76  .110  .190  .039  0.90  0.320  23.1  5-36  7 51  110  2.61  .105  .185  .037  3-2  1.04  86.5  5.94  6 02  K  1  1  M  2  - 1  - 1  log k '  I n i t i a l Conditions: 60°C _  0.100 M AgC10 ;  0.135 M NaOAc;  5 atm CO  80, 90, 110 C _  0.115 M AgC10 ;  O.I95 M NaOAc;  5 atm CO  P  4  4  x  logj s 1  - 76 while analysis of the slopes indicated that: S' = k ' [Ag+] [OAc-] + kg' [Ag+][AgOAc][OAc"]  (15)  2  2  The e f f e c t of acetate complexing  on the reduction rate was studied  only at 90°C and thus i t was not possible to calculate the dependence of k ' and k ' 2  acid-dependent  3  individually.  An average a c t i v a t i o n energy f o r the  reaction was estimated, however, from the slope of a log S' vs  l/T plot at a constant degree of complexing. 27 f o r 80, 90 and 110°C kcal/mole.  temperature  -Such a plot i s given i n Figure  and y i e l d s an average a c t i v a t i o n energy of 17 ± 3  Also included i n Figure 27 i s a plot of log I'/[AgOAc] vs l/T at  60, 80, 90 and 110°C  which gives the temperature  dependence of k i ' and y i e l d s  an a c t i v a t i o n energy of 15 ± 2 kcal/mole f o r the acid-independent  reaction.  The experimental a c t i v a t i o n energies incorporate the enthalpies of a l l e q u i l i b r i a preceding the rate-determining step.  In the acid-independent  reaction these include the heat of dissolution of CO and the heat of formation of AgOAc, while i n the acid-dependent  reaction the heat of formation of KOAc  and the intermediate complex, Ag-C^-OH, must also be considered.  The heats of  CO d i s s o l u t i o n and AgOAc formation are small (see Appendices B and C) as i s the heat of formation of HOAc  (e.g. ..-0.1 kcal/mole at 25°C (hh)).  No  information i s available f o r the heat of formation of the intermediate complex.  I f OH- rather than H 0 2  i s involved i n the formation, the experimental  a c t i v a t i o n energy f o r the acid-dependent d i s s o c i a t i o n of H 0 2  reaction includes the heat of  which i s about Ik kcal/mole i n d i l u t e aqueous solution  Combining the experimental a c t i v a t i o n energies with the "best value" rate parameters at 90°C l i s t e d i n Table X (Section III-10), assuming equal a c t i v a t i o n energies f o r each of the acid-dependent lowing Arrhenius expressions f o r the temperature  reactions, gives the f o l -  dependence of the rate  Figure 2 7 .  Arrhenius Plots f o r Acid-Independent and Acid-Dependent Reactions  - 78 parameters over the range investigated: ki  =  107.6  k  2  =  IO -  k  3  =  10 -  9  ± 1.3 7  10  ±  2  2  expt-(15 ± 2)  103/RT]  - ° exp[-(17 ± 3) 103/RT]  * - ° exp[-(17 ± 3) 2  M-I  M-2  103/RT] M"  2  (39)  sec-i sec-i  (kO)  (4l)  sec-i  The experimental a c t i v a t i o n entropies corresponding to the frequency factors i n the above expressions are: A S * = -l4 - 9 e.u. 3  ASf = -26  - 6 e.u.,  A S 2 * = -16  * 9 e.u.  and  (based on a standard state of one mole per l i t r e ) .  These experimental a c t i v a t i o n entropies also incorporate contributions from e q u i l i b r i a preceding the rate-determining step.  In the a c i d -  independent reaction these contributions may be small and the true a c t i v a t i o n entropy may  correspond to that found from equation 39-  abnormally low for a bimolecular reaction (50)  and may  This'value i s reflect steric  hindrance  f o r the i n s e r t i o n of a CO molecule i n t o the silver-oxygen bond. A l t e r n a t i v e l y the reactive s i l v e r species may not be undissociated AgOAc molecules but rather one present i n very low concentration (e.g. ion p a i r s ) . explanation may  A further  involve a large solvent ordering e f f e c t i n the formation of  the activated complex. between undissociated  Such an effect i s not generally expected f o r reaction molecules.  - 79 -  IV  CONCLUSION  The CO-reduction of s i l v e r perchlorate i n a c i d solution i s described by the o v e r a l l reaction: 2Ag  + CO + H 0  +  > 2Ag + C0 + 2H  2  +  2  In acetate-buffered solution the reaction proceeds homogeneously i n the l i q u i d phase by two p a r a l l e l routes, one of which i s pH-independent and the other, pH-dependent.  The pH-dependent route i s favoured by increased pH and i s made  up of both an acetate-independent  and acetate-dependent contribution.  The observed k i n e t i c s are consistent with the following mechanism: Ag  +  + 0 A c - ^ = ^ AgOAc  AgOAc + CO — 0 Ag-C-OAc + A g  K  Ag  +  > AgJi-OAc  A  +  (rapid equilibrium)  + H0 2  (slow)  >2Ag + C 0  0 ^ Ag-d-OH + H  + CO + H 0 v. 0 2  Ag-C-OH + A g — 2 A g +  + C0  2  + H  k  2  + HOAc + H  +  (fast)  (rapid equilibrium)  +  AgJLoH + AgOAc — i ^ _ > 2 A g + C 0  2  +  (slow)  + HOAc  (slow)  In the pH-dependent reaction the r e a c t i v i t y of silver-acetate complexes i s about a f a c t o r of three greater than the r e a c t i v i t y of simple hydrated s i l v e r ions.  This enhanced r e a c t i v i t y i s a t t r i b u t e d t o s t a b i l i z a t i o n  by the basic acetate anion of the proton released i n the reduction process. Buffering s i l v e r perchlorate solutions with sodium acetate and acetic a c i d increases the reduction rate by ( i ) increasing the pH, ( i i ) i n creasing the r e a c t i v i t y of s i l v e r ions i n the pH-dependent reaction through complexing and ( i i i ) providing an alternate pH-independent route for reduction. The e f f e c t of increased pH i s much greater than the s p e c i f i c e f f e c t s of  -  silver-acetate  80  -  complexing.  The CO-reduction of s i l v e r ( I ) i n a c i d solution i s consistent with the formation of intermediate complexes by the i n s e r t i o n of a CO molecule between a s i l v e r ion and co-ordinated oxygen-donating base.  - 81 APPENDIX A METHOD OF ESTIMATING RATES FROM THE SLOPE OF PRESSURE-TIME RECORDS The slope of a t o t a l pressure vs time record (e.g. Figure k}  Section  I I I - l ) can be converted to the rate of CO consumption i n fundamental units (e.g.  M s e c i ) from a knowledge of ( i ) the g a s - l i q u i d r a t i o i n the reactor, _  ( i i ) the s o l u b i l i t y of CO and C 0 under the experimental conditions and ( i i i ) 2  the  stoichiometric relationship between the consumption of CO and the produc-  t i o n of C0 .  The mathematical expression used f o r t h i s conversion i s given  2  by equation A-1:  -d[CO] dt  =  f(F  -dP dt  T  ^  +  aC0)(F + QCCOgA ( a c o - aco) J  r  A  [  2  where -d[CO]/dt  =  rate of CO consumption (M s e c ~ i ) *  dPip/dt  =  slope of a pressure-time record (atm sec- )  C0  =  s o l u b i l i t y of CO (M/atm)*  2  =  s o l u b i l i t y of C 0  F  =  a g a s - l i q u i d volume factor (M atm)  a  a  and  C0  = where  1  2  (M/atm)*  (V Ai)(1000/RT) g  Vg  =  gas volume (mis) measured at experimental temperature  VT_  =  l i q u i d volume (mis) measured at room temperature  R  =  universal gas constant (82.05 mis atm mole-i deg-i)  T  =  experimental temperature (°K)  Concentrations are expressed i n terms of l i t r e s of solution measured at room temperature (20-25°C).  _ - , N  >  - 82 -  The derivation of equation A - l i s given below. A p a r t i c u l a r gas (e.g. CO or C0 ) present i n the reactor w i l l be 2  d i s t r i b u t e d between the gas and l i q u i d phase i n f i x e d proportion dependent on the r a t i o of gas and l i q u i d volumes and on the s o l u b i l i t y of the gas i n the l i q u i d phase.  Assuming that the i d e a l gas law applies with s u f f i c i e n t  accuracy at the temperature and pressure of i n t e r e s t , the concentration of a gas i n the gas phase i s given by:  [gas]g  =  ^ £ (moles per V i mis of solution)  =  P x g (moles per l i t r e of solution) V l RT  =  F x P  V  1  0  0  0  where P = p a r t i a l pressure of gas (atm). The concentration of a gas i n the l i q u i d phase i s given by Henry's law: [gas]j_  =  a x P (moles per l i t r e of solution)  Thus the t o t a l concentration of a p a r t i c u l a r gas present i n the reactor i s [gas]  P(F + a)  =  From the stoichiometry of the Ag(I)-CO reaction i n a c i d solution, viz : 2Ag(I) + CO + H 0  > 2Ag + C0  2  + 2H " 4  2  the quantity of CO consumed equals the quantity of C0 produced. 2  change i n concentration of each gas i s given by: X  =  A[gas]  = AP  C 0  =  C Q  AP  (F  + C0) a  (F + « C 0 ) 2  (A-2) Thus the  - 83 The observed pressure change during a reduction experiment i s the difference between the decrease i n CO pressure and the increase i n C0  2  pressure, i . e . :  APr  =  AP  - AP o C  C 0  X  _  (F + QT'CO)  y(  2  X  (F + a  C0 ) 2  ( a C0 - a CO) (F + ©C!0)(F + OC C0 ) 2  2  or A [CO]  = A P / ( F + <* Co)(F + <* C0 ) T  2  (Qco " Qco) 2  Thus the rate of CO consumption i s given by:  -d[Cpi d t  =  -dPr  a  /(F +  . dt  V  (  a  C 0  ) ( F + oc C 0 ) \  C0  2  2  ( A  _  1 }  " cc CO)  The values f o r the s o l u b i l i t y coefficients used i n equation A - l were obtained from measurements described i n Appendix B and are summarized i n Table A-I.  TABLE A-I S o l u b i l i t y of CO and C0  Temp °C  60 80 90 110  i n Water at 60, 80, 90 and 110°C  2  CC c  104 M/atm o  x  7.0 6.8 6.9 7-3  a™  x 104 M/atm  163-5 122.5 IO5.O  84.0  - 8k -  The gas volume i n each experiment was estimated by subtraction from the  t o t a l volume of the reactor system, the i n i t i a l l y added volume of solution,  corrected to the experimental temperature using the v a r i a t i o n i n density of water with temperature (51)-  The volume of the reactor system was determined  at each experimental temperature from the sum of the gas volume measured from several observed decreases i n pressure due to CO saturation when the reactor was about 90$ f i l l e d with water, plus the volume of water present i n the reactor during each determination.  Reactor volumes determined i n t h i s way  were estimated to be accurate to ± 0„3 mis i n a t o t a l volume of about 120 mis (e.g.  118.9 at 60°, 119.5 at 80°, 120.0 at 90°, 120.1 at 110° and 120.7 at  120°C). Corrections to the i n i t i a l concentrations of AgC10 , NaOAc, HOAc and 4  CO to take account of the small amount of reaction which occurred during CO saturation were made on the basis of the stoichiometry of reaction A-3 using the  i n i t i a l rates calculated from equation A - l . 2Ag(I) + CO + H 0  + 2NaOAc  2  S» 2Ag + C0  2  + 2H0Ac + 2Na  +  (A-3)  Sample Calculation Expt No 285 (Figure k, Section I I I - l , Curve C; see also Appendix D-III-I) I n i t i a l Conditions:  0.235 M AgC10 ; 4  12.8 atm CO; I n i t i a l Slope  (dP /dt) T  Gas S o l u b i l i t y at 90°C  a  1  C Q  C  0  2  T o t a l volume of reactor at 90°C 2  =  90°C a  Volume of solution added (Vl)  Density of H 0  0.090 M HOAc;  -5.76 x 1 0 - 3 tm s e c " a f t e r 2-l/k  = a  0.090 M NaOAc;  =  =  6.9 x 10"  =  105.0 x IO" M/atm  4  M/atm 4  101.5 mis =  120.0 mis  O.9653 g/ml at 90°C;  O.9982 g/ml at 20°C  min  - 8 5  V  K  =  F  =  120.0 - 101.5 x 0.9982/0.9655 , , = 82.05 x 363 1  x  101.5 -d [CO] dt it  =  =  7  0  0  15.1 mis  i+9.9 x 10- M/atm 4  - 3 / (49.9 + 6.9) IO' (49-9 + 105-0) !Q- \ ^ (105.0 - 6.9) 10y 4  6  x  1 0  J %1  =  0  =  4  4  5I.8 x 10-  6  M sec-i  Using t h i s rate value an estimate can be made of the amount of AgC104, NaOAc and CO consumed and the HOAc produced at the point the slope was measured. -0.5 A[AgC10 ] 4  =  -A[C0]  =  -0.5 A [NaOAc]  =  2-1/4 x 60 x 5I.8 x 10"  6  =  0.5 A [HOAc]  =  0.007 M  .The concentrations at the point the rate was measured are therefore estimated to be [Ag(I)]  =  0.235 - 0.014  [NaOAc]  =  0.090 - 0.014 =  O.O76 M  [HOAc]  =  0.090 + 0.014  0.104 M  P  =  12.8 - 0.007/(49.9 + 6.9) 10-  c o  =  =  0.221 M  4  =  11.6 atm  - 86 -  APPENDIX B  SOLUBILITY OF CARBON MONOXIDE, CARBON DIOXIDE AND HYDROGEN IN WATER Data f o r the s o l u b i l i t y of carbon monoxide i n water are available (52) at atmospheric pressure and temperatures to 100°C while carbon dioxide s o l u b i l i t y data are available (53) to 700 atm and 120°C.  As an i n i t i a l phase  of the present investigation data on the s o l u b i l i t y of carbon monoxide i n water were extended to 63 atm and 220°C using the previously described reactor system (Section II-1).  The s o l u b i l i t y of hydrogen i n water was also measured  at about 25 atm and temperatures to 225°C.  A value f o r the s o l u b i l i t y of C 0  2  in water and the s a l t e f f e c t of sodium acetate - acetic acid mixtures were determined at ^0°C and 2.6 atm.  Experimental The gas outlet l i n e inside the reactor was bent so that l i q u i d samples could be drawn through the l/l6-±n o.d. c a p i l l a r y tubing and collected over mercury i n the ^O-ml water-jacketed burette shown i n Figure B - l . ent  Suffici-  water was charged to the reactor to leave an i n i t i a l gas volume of about  10 mis at experimental temperature.  The head space and water were degassed  under vacuum or by b o i l i n g the solution and steam flushing at atmospheric pressure.  A f t e r degassing  the outlet valves were closed, the shaking mechanism  activated and the solution heated to the desired temperature before introducing carbon monoxide or hydrogen.  When equilibrium had been attained (approximately  5 min), the sampling l i n e was flushed with a few mis of solution and a 10 to 30-ml l i q u i d sample c o l l e c t e d i n the burette while the shaking mechanism was stopped.  The burette valves were closed and the excess gas flushed from solu-  t i o n by rapidly r a i s i n g and lowering the mercury l e v e l .  After equilibrium had  - 87 -  from reactor  Y  l / l 6 " o.d. c a p i l l a r y tubing  -J—3-way microvalve  Water manometer-  K5S  S3  - Thermometer  50-ml burette  Levelling bottle  -Water jacket  —  Figure B - l .  Mercury  Measuring Burette System f o r Gas S o l u b i l i t y  Determinations  - 88 been established i n the burette the volumes of l i q u i d and gas were measured at the temperature of the water jacket and at atmospheric pressure as determined by balancing the two legs of a water manometer.  Several samples could  be drawn from a single charge. Corrections were made f o r the residual amount of gas remaining i n the burette solution and f o r the vapour pressure of water both i n the burette and the reactor.  The pressure i n the reactor system was measured with a  Consolidated Electrodynamics 0-1000 psig pressure transducer (Type 4-311). The carbon monoxide was c.p. reagent grade (99*5$ min) supplied by the Matheson Co. and was used without further p u r i f i c a t i o n .  Hydrogen was of  commercial grade (99-8$ min) supplied by the Canadian Liquid A i r Co. and was used without further p u r i f i c a t i o n .  Degassed d i s t i l l e d water was used i n a l l  determinations.  Results (a)  S o l u b i l i t y of CO i n Water Data obtained f o r the s o l u b i l i t y of CO i n water at 25 atm from room  temperature to 220°C are summarized i n Table B-I and shown i n Figure B-2. Published data available to 100°C at atmospheric pressure (52) when extrapolated to 25 atm agree with the present measurements to about 50°C.  At higher  temperatures the published values are too low by about 4$ at 75°C and about df at 100°C. 0  The s o l u b i l i t y c o e f f i c i e n t of CO i n water passes through a minimum between 50° and 100°C, a c h a r a c t e r i s t i c exhibited by other permanent gases (5*0The e f f e c t on the CO s o l u b i l i t y of pressure to 65 atm at 4l.5°C and to 40 atm at l40°C i s summarized i n Table B-II.  Deviations from Henry's law  - 89 TABLE B-I S o l u b i l i t y * of CO i n Water at 25 Atmospheres  S25  a x 10*  mls/g  M/atm  .561 .564 • 554  10.00 10.05 9.87 9.76 9.34 9.71  24.7  .567 .568 •565 .560 • 53.2 • 559 .562 .528 .522  41.5 41.5  21.4 27.2  • 390 .489  .455 .450  8.10 8.02  76.5  25.9  • 387  • 383  6.82  100 100  24.8  .403 • 394  .406 • 394  7.24 7.01  120 120  25.O  .431 .425  .431 .424  7.67 7-55  iko iko  24.6 25.4 27.0 27.5 24.5 24.8  A75 .480 .522 .516  .484  8.61  180 180  24.1  Temp °C  s mls/g  PCO atm  24 .4  25.2 25.2 25.5 25.6 25A 25.6 25A  2k. k 2k.6 2k. 9  25.4 26.1 26.2 26.4 27.1  24.9  25.O  25.1  .548  .524 .545 • 552 • 530 .529  .472  9.84  9.43 9.43  .484  8.40 8.64  .470  8.36  • 5^5 .545  • 555  9-89 9.76  .624 .638  .647  25.1  .635  11.5 11.3  200 200  23.4 24.2  .718 • 731  .768 • 757  13.7 13.5  220 220 220 220 220 220  25.5 25.9 26.2 26.5 26.6 27.1  .940 .928 .909 .981 • 907' • 983  .920 .898 .869 .924 .852 .905  16.4 16.0 15.5 16.5 15.2 16.1  i4o iko  160 160  .548  * S = mis gas (measured at S.T.P.) per gram H 0 S = mis gas (S.T.P.) per gram H 0 corrected to 25 atm assuming Henry's lawa = moles gas per l i t r e H 0 (measured at 20°C) per atmosphere of gas = S x I.78I x IO" (M/atm) 2  2 5  2  2  3  2 5  - 90 -  Figure B-2-  Solubility of CO and H  2  in Water from 25 to 225°C  - 91 -  '  TABLE B-II  E f f e c t of Pressure on S o l u b i l i t y * of CO i n Water  Temp °C  PCO atm  s mls/g  41.5  11A 11.7 'i4.o 21.4 27,2 34.1 41.4 48 :o 56.9 64.3 64.3  .209 .216 .251 .390 .489 .601 .725 .830 • 973 1.079 1.096  8.17 8.26 7-97 8.10 8.02  8.6 13.6 19.0 20.7  .177 .274 .367 .400 .475 .480 .522 .516 .625 .635 .727  9.15 9.00 8.59 8.61 8.61 8.40  i4o  24.6  25.4 27.O  27.5 33-5 34.3' 39-5  a  x 10 M/atm  7.84  7.82 7.70 7.61 7.48  7-59  8.64  8.36 8.31 8.26 8.21  * S = mis gas (measured at S.T.P.) per gram H 0 2  CC = moles gas per l i t r e H 0 atmosphere of gas 2  = S  x I.781 x 10-  3  2 5  (measured at 20°C) per  (M/atm)  4  - 92 were observed above about JO atm.  Sample Calculations The s o l u b i l i t y (S) i n mis gas (measured at S.T.P.) per gram of water was calculated using equations B-1 and B-2.  At room temperature the correction  for the amount of gas remaining i n solution i n the burette was made on the basis of Henry's law by subtracting the p a r t i a l pressure of the gas i n the burette from the experimental reactor gas pressure. At higher temperatures the correction was made by adding the volume of gas remaining i n the burette solution, as estimated from room temperature s o l u b i l i t y data, to the measured gas volume.  The two methods give i d e n t i c a l . r e s u l t s and were used i n t h e i r  respective temperature regions merely t o f a c i l i t a t e computation of the data. The r e p r o d u c i b i l i t y of duplicate determinations was better than the estimated maximum possible error of 3 to ^ (i)  f o r both CO and H  2  measurements.  At room temperature:  (B-1)  S  where  temperature of water jacket (°C)  T P  =  atmospheric pressure (mm Hg)  =  vapour pressure of water (mm Hg) at T  =  gas volume (mis) at T and P  =  water volume (mis) at T  d  =  density of water at T  S  =  gas s o l u b i l i t y mis (S.T.P.) per g water at the experimental  P  w  V D  V  w  temperature ( T ) and pressure ( P ) reduced by the p a r t i a l a  pressure of gas i n the burette.  a  - 93 Example;  of Carbon Monoxide i n Water at 24.9 ± 0.3°C ( T )  Solubility  a  V = 15.1 ± 0.1 mis  d = O.9977 g/ml  V = 24.4 ± 0.1 mis  P = 762.O ± 0 . 5 mm  g  w  T  22.2 * 0.1°C  =  P  Measured CO pressure (P )  =  Corrected CO pressure  =  a  S  Possible error  20.1 ± 0.2  mm  ± O.5 atm  26.6 - 762.0 - 20.1 760  25.6  =  O.56O mis CO (S.T.P.) per g water at 25.6 'AV2  =  =  atm  15.1 \ (762.0- 20.l\ f__273__ 24.4 x 0.9977/ I 76O j I22.2 + 273y  = =  26.6  w  AV — — — V  AP« A(P - PV) AT + — — — — + P - Pwi T + 275  W  +  0.1 15.1  +  J4-  24.4  +  0.1 295.2  0.7 741.9  +  26.6  +  AT  +  + —  w  atm CO a  273  Ta + 0-3 279.9  100  = 3-1$ (ii)  At elevated temperature:  v  g  + s .a.Vw (T + d«V, w x  273)/273\ /P - P^ 760  \T  273 + 273  (B-2)  where  T, P, P , Vg, V  and  Si =  gas s o l u b i l i t y mis (S.T.P.) per g water at T and 1 atm  S  gas s o l u b i l i t y mis (S.T.P.) per g water at experimental  W  =  w  and d have the same meaning as previously  temperature ( T ) and pressure (Pa) a  Example:  Solubility  of carbon monoxide i n water at 200.0 ± 0.3°C (T_)  V  g  =  12.5 ± 0.1 mis  d = O.997O g/ml  V  w  =  I5.8 ± 0.1 mis  P  S  x  =  0.0214 mis (S.T.P.)/g  L  Measured CO pressure (P ) a  w  =  =  752.8 t 0.5 mm 24.0 ± 0.2 mm  23.4 ± 0.7 atm  - 94 -  S = /12.5 + (0.02l4)(0.9970)(l5.8)(298.2/273)\ /752.8 - 24.o\ / 273 \ ^ 15.8 x 0.9970 y \^ 760 ) ^298.2) =  O.718  mis (S.T.P.) per g water at 23.4 + A(P  + w  r> . Possible error =  AV  M  ——  -rr^-  VV  V  g  - Pw) +  —r  AT  AP  —  P - P  w  T + 273  w  r  1  S o l u b i l i t y of H  +  ATa  +  \  •———•  P  T a  a +" 273 J  _PVL + 00 \ 23.4 473 J  1 0 0  i n Water  2  Data obtained f o r the s o l u b i l i t y of H temperature  a  +  = ( o- + + Q-7 + Q-i \12.5 15.8 728.8 298.2  (b)  atm CO  2  i n water at 25 atm from room  to 225°C are summarized i n Table B-III and shown i n Figure  B-2.  These results are i n excellent agreement with other published values (55,56,  57). Figure B-2 indicates that CO and H about 50°C.  At higher temperatures  temperatures  the reverse i s true.  The s o l u b i l i t y of H (c)  S o l u b i l i t y of C0  2  H  2  2  are equally soluble i n water at  i s the more soluble while at lower  i n water obeys Henry's law to about 100 atm  2  (55).  i n Water and Acetate Solutions  The s o l u b i l i t y of C0  2  i n water and acetate solutions was  at 90°C from the observed pressure drop due to C0 quantity of solution i n the reactor.  2  determined  saturation of a measured  The results of these measurements as  summarized i n Table B-IV indicate that 1:1  solution mixtures of NaOAc and HOAc  up to 2 M have l i t t l e e f f e c t on the s o l u b i l i t y of C0 . 2  The value f o r the C0  2  s o l u b i l i t y obtained i n t h i s way agrees with a value interpolated from published data (58) at 25 atm, i f a 10$ deviation from Henry's law i s assumed. published data at 15°C  Other  (59) indicate that such a deviation i s reasonable.  TABLE B-III  S o l u b i l i t y * ofH  Temp °C  C0 atm  g  i n Water  S mls/g  P  S 5 2  mls/g  a x 10 M/atm 4  24.9 25.2 25.4  .434 .438 .447  • 435 .434 .440  7.75 7-73 7.84  50 50 50  24.7 24.6 24.6  .414 .414 .402  .420 .421 .409  7.47 7A9 7.28  100 100 100  24.3 24.4 24.8  .440 .445 .444  .452 .456 .448  8.05 8.12 7-97  125 125 125 125  25.3 25.9 26.4 25.8  • 535 .547 • 551 .542  .529 .528 .521 .526  9.4i 9.40 9.28 9-58  150 150 150  25.7 25.8 24.8  .626 .634 .615  .609 .616 .619  10.8 11.0 11.0  175 175  24.6 25.1  .742 .743  • 753 .739  13.4 13.2  200 200  25.8 25.5  .900 .908  .871 .889  15.5 15.8  225 225  25.3 24.9  1.101 1.075  1.089 1.081  19.4 19.3  28.1 • 28.1 28.2  * S = mis/gas (measured a t S T.P.) p e r gram H 0 2  S  2 5  = mis gas (S.T.P.) per gram H Q corrected to 25 atm assuming Henry's law 2  a = moles gas p e r l i t r e H 0 (measured a t 20°C) p e r atmosphere o f gas 2  = S  2 5  x I.78I x 10-3 (M/atm)  - 96 -  TABLE B-IV  [NaOAc]  2  C0  2  [HOAc]  P  i n Acetate Solutions at 90°C  s mls/g  Si mls/g  a  x 10  2  M/atm  M  atm  --  —  2.61  .605  .232  1.02  —  —  2.66  .627  .236  1.05  M  *  S o l u b i l i t y * of C0  0.2  0.2  2.57  .606  .236  1.05  0.2  0.2  2.60  ,6l4  .236  1.05  2.0  2.0  2.56  .604  .238  1.06  S = mis gas (measured at S.T.P.) per gram solution S i = mis gas (S.T.P.) per gram solution corrected to one atmosphere assuming Henry s law 1  a, = moles gas per l i t r e of solution (measured at 20°C) per atmosphere = S  x  x 4.454 x 10-2 (M/atm)  - 97 -  ./ The measured absorption c o e f f i c i e n t f o r C0 other data f o r the s o l u b i l i t y of C0  2  2  at 90°C together with  i n water at one atmosphere and tempera-  tures to 60°C as included i n a recent review (53) are summarized i n Table B-V and are shown i n Figure B-3-  - 98 -  TABLE B-V S o l u b i l i t y of C0  2  i n Water at One Atmosphere  C0  Temp  2  Solubility  °C  g/100 g *  20  .172  3.91  25  .149  3-38  30  .131  2.97  ho  .105  2.38  50  .087  1.98  60  .072  1.64  90  M/atm  1.05  —  *  C0  **  a, = moles gas per l i t r e H 0 (measured at 20°C) per atmosphere of gas  2  s o l u b i l i t y i n grams C0 per 100 g H 0 from reference 53 2  2  = 0.227(g C0 /l00 g H 0) 2  2  2  - 99 -  Figure B-3.  S o l u b i l i t y of C 0  2  i n Water at One Atmosphere  - 100  APPENDIX C  SILVER-ACETATE COMPLEXING FROM E.M.F. MEASUREMENTS E q u i l i b r i a i n s i l v e r acetate solutions have been extensively studied at room temperature by s o l u b i l i t y measurements (1+5,60,61,62,63,64,65) and by E.M.F. measurements (39>^>66»).  The e f f e c t of temperatures to 90°C on s i l v e r -  acetate complexing at ionic strengths between 0.1 and 0.9 has been estimated during the present investigation from E.M.F. measurements on c e l l s of the type:  Ag  AgC10 NaOAc  NaC10 8 M  4  4  AgC10 NaC10  Ag  4  4  Experimental A l l c e l l solutions contained 0.0104 M AgC10 of either NaOAc or NaC10  4  4  and appropriate amounts  i n equal concentrations to maintain constant  ionic  strength i n both sides of each c e l l . S i l v e r f o i l electrodes, approximately 4 x 4-cm with a 5were conditioned by e l e c t r o l y z i n g approximately M AgC10 with HC10 . 4  4  cm  lug,  solution a c i d i f i e d  Each electrode was f i r s t treated as an anode to expose a fresh  surface and then as a cathode to produce a coherent s i l v e r deposit.  Current  densities of about 0.05 amps/cm gave suitable deposits within one to two 2  minutes.  The e l e c t r o l y s i s conditions were not c r i t i c a l .  The electrodes were  then formed into c y l i n d r i c a l shapes, washed and stored i n a c i d i f i e d  distilled  water u n t i l required. The experimental c e l l , depicted i n Figure C - l , consisted of two 22-mm i . d . pyrex solution compartments joined by about 15 cm of 6-mm  i.d.  - 101 -  - - S i l v e r electrode • — Helium flushing tube  F i l t e r paper plug  8 M NaC10 Salt Bridge 4  Figure C - l .  Experimental C e l l f o r E.M.F. Measurements (approximately half-scale)  - 102 -  pyrex tubing and had a t o t a l volume of about 75 mis.  A c a p i l l a r y tube f o r  gas flushing was attached near the bottom of each solution compartment. F i l t e r paper plugs were used to separate the 8 M N a C 1 0  4  compartment and solution contact was made i n the plugs.  salt-bridge from each The c e l l was about  80$ immersed i n a. s t i r r e d , e l e c t r i c a l l y heated two-litre bath of ethylene glycol.  The temperature was controlled to ± 0.1°C using a Yellow Springs  Instrument Co. Thermistemp Temperature C o n t r o l l e r (Model 71) with a stainless s t e e l - c l a d thermister probe (No. k-06) .  E.M.F. measurements were made with a  high p r e c i s i o n Leeds and Northrup potentiometer (No. 7552). The general procedure consisted of p i p e t t i n g s u f f i c i e n t 8 M N a C 1 0  4  (previously prepared with degassed d i s t i l l e d water and stored i n a stoppered f l a s k ) t o bring the solution l e v e l i n the s a l t bridge tubing to within about two cm of the bottom of each solution compartment. plugs were inserted approximately  After the f i l t e r paper  25 mis of experimental solution were  added to each solution compartment and the electrodes secured h to 5 cm above the plugs with rubber bungs.  The vessel was mounted i n the g l y c o l bath at  room temperature and the solution compartments were flushed with a slow flow of helium f o r about f i v e minutes.  After flushing the gas outlet l i n e s were  clamped and a s l i g h t p o s i t i v e pressure of helium maintained over the solutions . At room temperature the E.M.F. became constant within 2 t o k hours and showed l i t t l e deviation f o r periods up to 1 2 hours.  A f t e r the room  temperature value had been determined the g l y c o l bath was heated and the E.M.F. measured at successively higher temperatures, s u f f i c i e n t time being allowed to a t t a i n equilibrium as evidenced by constant cell-voltage readings.  The E.M.F. of c e l l s allowed to cool from higher temperatures over-  night generally reproduced the room temperature value within 1 0 $ .  - 103 Results The following assumptions are made: 1.  The difference i n l i q u i d junction potentials i s n e g l i g i b l e .  2.  A c t i v i t y c o e f f i c i e n t s of A g are equal i n a l l solutions of equal ionic +  strength. 3.  AgC10 , WaC10 and WaOAc are completely dissociated.  4.  The average silver-acetate complex i s represented by the formula AgOAc,  4  4  On t h i s basis the E.M.F. of each c e l l , using the International or Stockholm sign convention (9), i s given by:  2.30 RT  =  where  c e l l E.M.F.  -nF  1  0  g  0.0104 [Ag+]  (c-l) k  E  =  (volts)  R  = u n i v e r s a l gas constant (I.987 c a l m o l e i  F  =  Faraday constant (23.06 kcal/equivalent)  T  =  absolute temperature (°K)  n  =  number of v o l t equivalents (= unity)  [Ag ]  =  s i l v e r ion concentration i n the acetate solution compartment (M)  0.0104  =  s i l v e r ion concentration i n the perchlorate solution compartment (M)  _  +  deg-i)  The concentration r a t i o (K ) f o r the formation of AgOAc a  K,  i s given by:  LAg+JlOAc"]  a  The results of measurements made on s i x c e l l s are summarized i n Table C-I.  K  a  estimated using equations C - l and C-2 i s e s s e n t i a l l y independent  of ionic strength between 0.1 and 0.9 and temperatures t o 90°C and has an average value of 3-7 * 0.7 M" . 1  - 10k -  TABLE C-I  E f f e c t of Temperature on Silver-Acetate Using 0 . 0 1 0 4 M AgC10  Cell No  [NaOAc] M  Temp °C  EMF mv  1  .10  23  9.0  2  .20 .20 .20 .20  24  3  Complexing  4  [Ag ] +  M  M-l  ± 1.0  • 0073  4.4  25 50 90  12.0 12.5 15.0 17.0  ± * ± ±  1.0 1.0 1.0 1.5  .OO65 .0064 .0061 .0060  3-1 3-2 3.6 3-7  .50 .50 •50  25 35 50  27.0 ± 1.0 2 9 . 0 ± 1.5 3 1 . 5 * 1-5  .0036 • 0035 .0034  3.8 4.0 4.2  4  .50  23  24.0 ± 1.0  .0041  3.1  5  .50 • 50 .50  24  23.5 ± 1.0 26.0 ± 1.0 33-5 * 1.5  .0042  25 90  .0038 • 0037  3.0 3.5 3-7  6  • 90 .90 .90 • 90  35 50 90  35.0 39.0 45.0 50.0  .0027 .0024 .0021 .0021  3.2 3.7 4.4 4.4  K  = a  24  0.0104 - [Ag ] [Ag ]([NaOAc]-[Ag ]) +  +  +  ± 1.0 ± 1.0 ± 1.5 ±2.0  - 105 APPENDIX D SUMMARY OF SELECTED EXPERIMENTAL DATA FOR THE REDUCTION OF SILVER(I) SOLUTIONS BY CARBON MONOXIDE  I. (a)  (b)  EFFECT OF CO PRESSURE AT 90°C (Figure 5) I n i t i a l Conditions:  4  Expt No  Pco  234 233 231 232 2l4 230 207 236 235 237 238  1.67 1.69 3.10 4.17 • 5-17 5.22 5.24 9.28 12.8 13.3 29.8  I n i t i a l Conditions:  11.4 19.4 25.6 27.5 27.7  b  161 a  b  42.0  43.3 54.4 132.  4  162 163 165?  6  7.2 5.0 13.6 20.0 27.1 23.7 23.7  0.115 M AgC10 ; PCO atm  O.I95 M NaOAc; R* x 10 M s-i  atm  Expt No  l64  * **  0.115 MAgC10 ;  0.045 M NaOAc;  O.O78 M HOAc R'**x I O M - i s-l 6  a  4.31 2.96 4.39 4.79 5.24 4". 35 4.34 4.57 3.38 4.09 4.43  O.766 M HOAc  R* x 106 M. s - i  R***x 106 M a-i s - i  5.8 10.0 12.7  .509 .515 .496 .516 .520  14.2  14.4  - Reactor thoroughly cleaned with HNO3 p r i o r to charging. - Contained O.lg Ag p r e c i p i t a t e d i n previous experiment plus 2.0 g s i l v e r sponge obtained from Consolidated Mining and Smelting Co. Ltd.  R = -d[C0]/dt R' = R/P C0  - 106  EFFECT OF ACETIC ACID AT 90° C (Figure 6)  II.  I n i t i a l Conditions:  III. A.  -  0.115  O . I 9 5 M NaOAc  M AgC10 ; 4  Expt No  [HOAc] M  [HOAc] M"  209 212 208 210 238 237 235 236 230 207 214 232 231 233 234 213 206 211  .585 .585 .160 .160 .0978 .O898 .0886 .0868 .0832 .0830 .0828 .0818 .0805 .0786 .0786 .0624 .0460 .0334  1.7 1-7 6.3 6.3 10.2 11.1 11.3 11.5 12.0 12.0 12.1 12.2 12.4 12.7 12.7 16.0 21.7 29.9  R' x 1 0 6 M a s."  PCO atm  -1  1  -  O.O5O M AgC10 j 4  Expt No  Ug(I)] M  A-180 A-181 A-191 A-192 A-193  .048 .042 .046 .046 .049  .043 .037 .041 .041 .043  Average:  .048  .041  [NaOAc] M  I'  = 0.15  [HOAc]" M-i  1  4.4 18.9 10.6 5A 1.2  S' = 0.017  X  10-6  8 , - TablelV)  0.045 M NaOAc PCO atm  R' x 1 0 p M a s  28.0 27.9 28.1 28.1 27.8  .152 • 559 .418 .304 .106  1 1)  1  1.52 1.84 3.61 3.89 4.43 4.09 3.39 4.53 4.54 4.34 5.24 4-79 4.39 2.96 4.31 6.88 7-73 10.1  5.17 5.24 5.31 5.31 29.8 13.3 12.8 9.28 5.22 5.24 5.17 4.17 3.10 1.69 1.67 5.24 5.17 5.17  EFFECT OF SILVER-ACETATE COMPLEXING AT 90°C (Figures 6 , 7 , I n i t i a l Conditions:  1  M atm  x 10-6 M  2  - 1  atm  sec -1  1  sec"  1  - 1  _ 1  - 107 B.  I n i t i a l Conditions: Expt No B-lk2 -B-I4-9  B-152 B-155 Average:  0.100 M AgC10 ;  [Ag(I)] M  [NaOAc] M  [HOAc]" M-i  • 097  .096 .O96 .095  .0k2 .01+1 .041 .040  4.39 4.36 1.30 1.30  .096  .04l  I'  =  S' =  C.  I n i t i a l Conditions: Expt No  0.045 M NaOAc  4  0.31  X  10 "  0.026 x I O  - 6  C-308 C-307 C-306  .046 .045 .044  .090 .089 .087  Average:  .045  .O89 I'  I n i t i a l Conditions: Expt No  [Ag(I)3 M  2  1  .64 4.41 10.7  .116 .115 .114  .044 .043 .042  Average:  .115  .043  I' =  1  secT  s  1  _ 1  .230 .504 .853  1  M2 atm" s e c "  6  R' x 10 M a" s  12.4 12.2 12.1  6  D-299 D-298 D-297  co  p  _1  = 0.22 x IO" Matm"  [NaOAc] M  .420 .368 .368 •.347  - 1  atm  4  1  - 1  ;M atm'" s e c  [HOAc]  0.117 M AgC10 ;  1  0.090 M NaOAc  S' = O.O58 x IO"  D.  R' x 106 M a " .'s"  27.2 27.6 27.6 27.7  -1  4  [NaOAc] M  P  M atm - s e c  6  0.047 M AgC10 ;  [Ag(I)] M  C0 atm  1  1  1  0.045 M NaOAc [HOAc] M-  _1  C0 atm  R x 10 M a s  12.1 12.1 12.4  • 552 .619 1.00"  P  1  .63 4.41 10.7'  0.32 x IO" M atm-i sec"! 6  S' = 0.066 x 1 0 - 6 M a t m - s e c 2  1  - 1  1  - 1  s - 1  - 108 E.  I n i t i a l Conditions: Expt No  0.090  0.100 M AgC10 ; 4  [NaOAc] M  [Ag(I)] M  E-159 E-140 E-l4l E-151  .094 .094 .089 .095  .084 .084 .079 .O85  Average:  .090  .080  M NaOAc  [H0Ac]-i M-i  PCO atm  26.6 26.7 26.5 27.4  5.38 5.38  5.24  2.20  I'' = 0.40 x I O  - 6  Matm" s e c " 1  F.  I n i t i a l Conditions: Expt No  0.047 M AgC10 ; [NaOAc] M  [HOAc]" M"  F-505 F-504 F-505  .045 .044 .043  .179 • 177 .176  • 63 4.39 10.6  Average:  .044  .178  S' =  G.  1  [Ag(I)] M  = 0.43 x  I n i t i a l Conditions: Expt No  [Ag(I)] M  G-I89  .044  G-190  .043  Average:  .042  .215  .033 .o4i  6  12.2 12.3 12.2  1  sec"  2  [HOAc]M"  .467 • 799 1.27  M NaOAc C0 atm  1  P  27.4 28.2 27.7 28.4  16.1 4.3 2.1 1.6  I' = 0.43 x IO" M a t m 6  M  1  1  1  6  1  1  0.225  S' = 0.077 x 10"  R'x 106 M a- s-  atm" sec _ i  10-6 M  4  .208 .215 .219 .218  G-155 G-182  atm  IO" M atm"  O.O5O M AgC10 ;  [NaOAc] M  1  Pco  1  1  0.081 x  .804 • 755 .711 .566  0.180 M NaOAc  4  ]:'  1  1  S' = 0.074 x IO" M2 atm" sec" 6  R'x 106 M a - i s"  2  -1  sec-i  atm"  1  sec"  1  R'x 106 M a-i s-i  1.73 .701 • 5^9 .602  - 109 H.  I n i t i a l Conditions: Expt No  0.235 M AgC10 ;  0.045 M NaOAc  4  [NaOAc] M  [Ag(I)] M  H-293 H-292 H-291  .233 .231 .229  .0U3 .041 .039  Average:  .230  .041  =  I  S' =  PCO atm  M-i [HOAc]-1  11.6 11.8 11.5  • 63 4.37 10.4  O.5O x IO" M atm" s e c " 6  0.133 x IO"  1  6  M  2  atm"  I n i t i a l Conditions: Expt No  J.  0.117 M AgC10 ; [NaOAc] M  M  1-296 1-295 1-294  .115 .114 .110  .088 .087 .085  Average:  .114  .O87  Expt No  [HOAc]" M"  1  C0 atm  R'x 1 0 M a" s-  11.8 12.0 12.2  .56 1.11 2.32  • 63 4.37 10.2  =  0 . 4 4 x 1 0 - 6 M atm" s e c " i  S  =  O.I78 X  S  1  1  1  10-6  M2  0.235 M AgC10 ; [NaOAc] M  J-290 J-283 J-289 J-285  .231 .230 .227 .221  .087  Average:  .225  .082  a  t  - i  m  sec"  1  0.090 M NaOAc  4  [Ag(I)] M  •56 1.21 1.83  1  p  1  I  I n i t i a l Conditions:  1  O.O9OM NaOAc  4  [Ag(I)]  1  -1  1  sec I.  R'x 106 M a- s"  [HOAc] M"  _1  .62 1.26 4.30 9-60  .O85  .082  .O76 I  =  O.78 x 1 0 " M atm" s e c  S  =  0.384 x 10 "  s  6  1  R'x 1 0 M a s"  11.6 11.9 11.9 11.6  1.02 1.34 2.43 4.47  - 1  •1  M atm" 2  C0 atm P  1  sec  - 1  6  - 1  1  - 110 -  K.  I n i t i a l Conditions: Expt No  K-302 K-301 K-300  Average:  0.117  [Ag(I)] M  M AgC10 ;  0.180 M NaOAc  4  [NaOAc] M  .114 .111 .105  • 177 .174 .168  .110  • 173  [HOAc]"  C0 atm P  1  M-I  12.1 12.0 11-7  .63 4.32 9.76  0.70 x IO" M.atm s e c  I' =  6  -1  L.  I n i t i a l Conditions: Expt No  L-209 L-212, L-208 L-210 L-238 L-237 L-235 L-236 L-230 L-207 L-214 L-232 L-231 L-233 L-234 L-213 L-206 L-211  Average:  2  0.115  M AgC10 ;  [NaOAc] M  [HOAc]" M-i  .113 .113 .111 .111 .096 .103 .104 .106 .110 .110 .110 .112 .114 .114 .108 .108 .105  • 193 .193 .191 .191 • 175 .183 .184 .186 .190 .190 .190 .191 .192 .194 .194 .188 .188 .185  1.7 1-7 6.3 6.3 10.2 11.1 11.3 11.5 12.0 12.0 12.1 12.2 12.4 12.7 12.7 16.0 21.7 29.9  .110  .190  ]:'  = 0.90 x 10"  S' - 0.32 x  10-6  - 1  s"  .92 2.06 3.87  1  - 1  O.I95 M NaOAc  4  [Ag(I)] M  .ill  Ma  s  - 1  S = 0.322 x 10-6 M atm" s e c 1  R' x 10  6  C0 atm  R'x 106 M a s-i  5.17 5.24 5-31 5-31 29.8 13.3 12.8 9.28 5.22 5.24 5.17 4.17 3.10 I.69 1.67  1.52  1  p  5.24  5-17 5.17  M atm"s e c MS atm" s e c 1  1  - 1  - 1  - 1  1.84  3.61 3.89 4.43 4.09 3.39 4.53 4.54 4.34 5.24 M 9  4.39 2.96 4.31 6.88 7-73 10.1  1  - Ill -  M.  I n i t i a l Conditions: Expt No  0.04-7 M AgC10 ; [NaOAc] M  [Ag(I)] M  M-511 M-310 M-309  .043 • 039 .036  .672 .670 .665  Average:  .040  .670  I' = S  N.  I n i t i a l Conditions: Expt No  0.  1  =  [HOAc]"  • 63 4.32 9.87  6  0.249 x IO"  6  O.O58 M AgC10 ;  M  •51 2.57 8.78  • 055  Average:  .O56  • 55^ I  = O.85 x 10 "  S  = 0.275 x 1 0 ~ M  M atm 6  O.O58 M AgCIO ;  I n i t i a l Conditions:  M  .52 1.03 2.57 5.10 8.76  .055  • 777 • 777 .776 .776 • 775  Average:  .O56  .776  .O56 .O56  sec"  atm  2  [HOAc]M-i  .057 • 057  1  1  -1  [NaOAc] M  0-250 0-252 0-251 0-253 0-249  S  1  1  sec  -1  PCO atm  R"x 106 M,a-i,s-i  5-51 5.44 5-57  1.00 I.58 5.24  1  - 1  O.778 M NaOAc  4  [Ag(I)]  sec"  -1  1  6  1.04 1.94 3.34  1  [HOAc]" M"  •555 • 55^ • 553  12.1 12.4 12.4  atm" s e c "  [NaOAc] M  .057  .O56  2  R'x 1 0 M a" s-i  O.556 M NaOAc  4  [Ag(I)]  M  C0 atm  p  1  O.87 x 1 0 " M a t m  N-254 N-255 N-256  Expt No  0.676 M NaOAc  4  1  I  =  1 . 1 0 x 1 0 " M atm" s e c - i  S  =  0.540 X 1 0  6  _ s  C0 atm  R'x 1 0 M a" s-  5.24 5-57 5.51 5.51 5.37  1.24 I.58 1.95 2.92 3.92  P  1  M  2  atm" s e c " 1  1  S  1  1  - 112 REACTION RATES USED IN EXTRAPOLATION TO ZERO ACETATE AT 90°C (Figures 20, 21, 22; Table VI).  IV.  P.  0.050 M AgC10 ;  I n i t i a l Conditions:  Expt No  •184 -182 -I85  -183 -186 P-181  4  [Ag(I)] M  [NaOAc] M  [HOAc] M  .OH-3  .262 .215 .172 .126 .083 • 037  .278 .235 .188 .144 .097 .053  .04l .043 .04l  .044 .042  p  Q.  I n i t i a l Conditions:  [AgOAT] [OAc"] R'xlOs M M M a s"  C0 atm r  - 1  .0201 .0173 .0156 .0123 .0094 .0046  28.2 28.2 28.2 27.9 28.0 27.9  .[HOAc] LAgOAc J  [HOAc]/[NaOAc] = 1  .245 .198 .156 .114  .074 .032  = 4.6 x IO" M a t m 6  0.100 M AgC10 ; 4  R  M a  1  - 1  s"  1  0  6  1  10.8 9-53 9.15 7.85 6.71 5.54  • 777 .702 • 759 .671 .650 .476 -1  , [HOAci  sec~i  [HOAc]/[NaOAc] = 1  .[HOAc] xlOe [AgOAc] [OAc"] R' x l O [AgOAc] M M a-i s - i M a-i s - i .M R  Expt No  Q-137 Q-138 Q-136  [Ag(I)] M  .082  .O85 .O85  [NaOAc] M  .207 .165 .120  R  ,  I n i t i a l Conditions:  Expt No  R-287 R-288  [Ag(I)] M  .115 .118  T A W T )  0.123  [NaOAc] M  .102 .029  p  co  atm  12.7 12.8  (  [AgOAc J J o  • 175 .137 .098  1.73 1.53 I.58  13.1 10.5 10.5  = 6.8 x I O - M a t m - s e c -  4  ^..[HOAc] \ K  0  .0322 .0285 .0225  M AgC10 ;  [HOAc] M  .O87 .017  6  C0 atm  P  26.4 26.4 26.5  .243 .195 .150 (  R.  [HOAc] M  [HOAc]/[NaOAc] = 1  [HOAc] 6 [AgOAc] [OAc-] R' xlO© [AgOAcJ M M M a-i s - i M a-i s - i 1  n  .0223 .0051  .065 .012  = 9.6 x IO- M a t m 6  11.4 10.0  2.48  1-75 -1  sec  - 1  Q  - 113 -  S.  I n i t i a l Conditions:  Expt No S-285 S-286  [Ag(I)] M .225 .226  0.235 M AgC10 ; 4  [NaOAc] M  [HOAc] M  .076 .013  .104 .032  ,[HOAc] [AgOAc ] [OAc ] R'xlO LAgOTSff M M M a - i s'" M a ~ i s  C0 atm P  R  0  6  -  .0313 .OO59  »  .045 .007  = 15.7 x l O - e  6 10  _ : L  1  11.6 11-9  ( 'S)  (  [HOAc]/[NaOAc] = 1  14.9  4.47 2.59  M  i4.o  atm- s e c -  .  *„  T.  Initial'Conditions:  Expt No T-150 T-154T-151 T-lV? T-l42  [Ag(I)] M. .094 .095 .095 .096 .097  0.100 M AgC10 ; 4  [HOAc] •Pco atm M  [NaOAc] M .129 .108  .681 .569 .455 .229 .228  .O85  .041 .042  U.  I n i t i a l Conditions:  Expt ' [Ag(I)] No M U-276 U-277  .065 .065  [AgOAc] [OAc"] R'xlO M M M a s-i 6  j^  .0259 .0228 .0187 .0099 .0102  .090 .065  .785 •595  (  C0 atm P  4.83 4.90  „,.[H0Ac]\ [AgOAc J J  - 1  1  [AgOAc] [OAc ] R' xlO© M M M a s-  •or [HOAc ] ^  -  - 1  .0139 .0078  .076 .038  = 4.4 x IO" M a t m 6  q  1  [HOAc]/[NaOAc] = 8.7  M a-i s-i-  1  52.2 18.0  • 570 .55^  N  K  Xl06  1  16.0 14.4 13.5 8.5 9-4  .611 •575 • 555 .368 .419  6  4  [HOAc] M  .103 .O85 .066 .031 .032  , iHOAc] LAgOAcJ M a s-  = 6.0 x IO" M atm" s e c "  0.064 M AgC10 j  [NaOAc] M  R  - 1  27.5 27.5 27.4 27.6 27.2  ^ R'  [HOAc]/[NaOAc] = 5  -1  sec  - 1  c  -114 -  V.  0.130 M AgC104;  I n i t i a l Conditions:  [HOAc]/[NaOAc] = 8.7 [HOAc]  Expt No  V-281 V-280  [Ag(I)] M  .126 .127  [NaOAc] M  [HOAc] •M  12.1 12.3  1.314 .221  .146 .022  [AgOAc] [OAc-] M M  C0 atm P  rlP^ 1 ]  ( R' i \ [AgOAc J J  - 18.0  x  10-6  6  1  .110 .015  .0362 .0066  ^  R xlO [AgOAcJ M a " s" M a " i s " i 1  R  1  1.19  43.3 21.8  .649  M atm"  sec"  1  1  0  W.  0.163  I n i t i a l Conditions:  Expt  Ug(I)]  [NaOAc]  No  M  M  W-278 W-279  .161 .162  4  [HOAc] M  • 073 .024  X.  [AgOAcV  0.235  I n i t i a l Conditions:  -  M  4.94 4.68  / , [HOAc] \ R  [AgOAc] [OAc ]  Pco atm  .655 .219  ^  [HOAc]/[NaOAc]= 8.7  M AgC10 ;  .0246 .OO87  _  nq  M  R' x l O M a-i s - i M a-i s - i  .049 .015  .u x 10-6  M  6  30.1 22.9  1.13 •.908  a t T n  -i  R  e  n  - i  Q  M AgC10 ; 4  [HOAc]/[NaOAc]=  8.7  p,-[HOAc] , R [AgOAc] [OAc ] R"xl06 LAgOA-c]XlCP M a-i s - i M a-i s - i M M r l A B  Expt No  X-283 X-282 X-284  [Ag(I)] [NaOAc] M M  .230 .232 .232  [HOAc] M  .085  • 793 • 399 .319  .042  • 053  (  E  ' & )  Pco atm  11.9 11.8 12.1  0  -  .0355 .0184 .0146  .050 .024 .018  1.34 .903 .890  12.8 x 10~6 M atm  -1  sec  29.9 19.6 19.4 -1  - 115 Y.  I n i t i a l Conditions:  Expt No Y-270 Y-274 Y-273 Y-272  0.270 M AgC10 ; 4  [HOAc] M  [Ag(I)] [NaOAc] M M .267 .268 .269 .268  .092  .829 .415 .167 .084  .046  .018 .0083  (, R  V  V. P  H  •p r  [HOAc]/[NaOAc] = 8.7 *'[HOAc] [AgOAc] M s-i a-i X  co  [AgOAc] [OAc-] R' x l O M.a" . s - i M M 6  atm  1  .0419 .0218 .0087 .0041  5.05 4.88 4.92 4.97  -\ [AgOAc J /  [HOAc]  = 18.4  -  .050 .024 .009 .004  X  10 "  2.27 1.55 1.18 .969  M atm  6  -1  45.1 29.5 22.8 19.8  sec-.1  Q  CHANGES DURING REDUCTION OF UNBUFFERED AgC10 SOLUTIONS AT 90°C AND 53 ATM CO (Figures 23 and 24; Tables VII:and-VIII) 4  Experiment No. 262  Time sec  [Ag+] M  0 180 390 78O 1,800 3,600 6,300 9,000 12,780  .0530  .0458  [H ] x 103 M +  pH  ~  5.2 3.15 2.92 2.71 2.55 2.32 2.27 2.14 2.09  ^0.007 0.7 1.2 2.0 2.8 4.7 5-4 7.2 8.1  log t  2.26 2.59 2.89 3.26 3.56 3.80 3.96 4.11  The pH i s defined i n terms of hydrogen ion concentration rather than a c t i v i t y since the pH-meter was calibrated against mixtures of standard HC10 and AgC10 solutions. 4  4  L  0  6  - 116 -  Experiment No. 260  Time sec  0 120  [Ag ] M +  .104  ~  240  480 990 1,920 4,740  10,000  [H+] x 10  PH  .090  5.0 3.05 2.77 2.61 2.41 2.26 2.05 1.87  3  log t  M  ~  0.01 0.9 1-7 2-5 3-9 5-5 8.9 13.5  2.08 2.38 2.68 3.00 3.28 3.68 4.00  Experiment No. 26l  Time sec  [Ag ] M  0 180 300 525 1,020 1,810 3,600 6,310 9,000 11,700  .211  [H ] x 10 M  +  +  PH  ~ 4.5 2.62 2.51 2.30 2.14  1.96 1.80 1.67 1.63 1.47  .179  3  ^ 0.03 2.4 3-1 5.0 7.3 11.0 15.9 22.4 23.4 33-9  log t  2.26 2.48  2.72 3.01 3.26 3.56 3.80 3.96 4.07  Experiment No. 265 Time sec  [Ag+] M  pH  0 180 360 660 1,140 2,040  1.06  ^4.0 1.67 1-57 1.43 1.50 1.10 0.95 0.90 0.83 " 0.69  3,840  5,640 8,340  11,040  •85  [H+] x 103 M  —' 0.1 21 27 37 50 79 112 126 148  204  log t  2.26 2.56 2.82 5.06 3.31 5.58 3.75 5.92 4.04  - 117 VI.  EXPERIMENTAL DATA INCLUDED IN LEAST SQUARE REGRESSION ANALYSIS, SECTION IH-IO(TABLEIX) BUT NOT USED IN GRAPHICAL ANALYSES, AT 9Q°C  (a)  I n i t i a l Conditions:  Expt No  265 264 266 267 268  [Ag(I)] M  a  '  b  c  a b c d  d  -  (b)  .142 .141 .141 .142 .142  0.144 M AgC10 ; 4  [NaOAc] M  .192 .191 .191 .192 .192  0.195 M NaOAc;  [HOAc] M  .197 .198 .197 • 197 •197  PCO atm  •  5.04  3-75 4.09 4.02 3.87 3-75  2  0.111 M HOAc  Expt No  [Ag(I)] M  [NaOAc] M  [HOAc] M  239 • 240  .068 .068 .070 .069 .133 .132 .129 .278 .044 .069  .220 • 330  • 113 .114 .112 • 113 .113 .115 .118 .118 .112 .113  24l  242 243 244 245 246  247 248  I n i t i a l Conditions:  Expt No  [Ag(l)] M  162 163  .114 .110 .111 .110 .110  I65  5.04  4.78  R' x 106 M a-i s-i  A i revacuated from solution i n reactor p r i o r to heat-up Added C0 to 5.7 p s i p r i o r to CO addition Contained 3-3 g fine 316 S.S. f i l i n g s Contained 1.0 g Ag p r e c i p i t a t e d during previous experiments plus ).0 g Ag sponge  I n i t i a l Conditions:  (c)  5.16 5.10  O.I95 M HOAc  164 161  .044  .109 .043 .108 .216 .104 .110 .110  0.115 M AgC10 ; 4  [NaOAc] M .044 .040  .041 .040 .040  C0 atm P  5.22 5.25 5.45 5.17 5.11 5.15 5.11 5.10 5.27 5-30  0.045 M NaOAc;  [HOAc] M  .767 .771 .770 • 771 • 771  Pco atm  11.4 19.4 25.6 27.5 27.7  R'x 106 M a-i s-i  2.22 2.97 .785 1-57 1-57 3.16 6.34 6.53 • 955 1.43  O.766 M HOAc R'x 106 M a-i s-i  .509 .515 .496 .516 .520  - 118 -  VII. (a)  EFFECT OF TEMPERATURE. (Figure 26; Table XI) I n i t i a l Conditions:  Expt Wo  60°C, 0.100 M AgC10 ;  [Ag(I)]  [NaOAc]  [H0Ac]-i  M  M  M-I  .135 .134 .134  1.3 11.0 21.5  223 222 221  .100 .099 .099 .  Average:  .099  0.135 M NaOAc  4  '  R'x 10s  Pco  atm  M a-i s - i  5.58 5.51 5.30  .177 .468  .720  .134 I' = 0.14 x 10-6 M atm-i sec-i  s' = 0.029 x  (b)  I n i t i a l Conditions : Expt No  [Ag(I)]  220 219 218 217 216 215 Average:  80°C;  M2 atm  10-6  4  [HOAc]-1  M  M-i  • 113 .112 .111 .110 .110 .107  • 193 .192 .192 .191 .190 .187  1-7 6.3 12.4 17.2 25.2 32.6  .110  .191  M  sec  - 1  O.I95 M NaOAc  0.115 M AgC10 ;  [NaOAc]  _ 1  R'x 106 M a-i s-i  Pco atm  5.4l 5.47  0.93 1.79  5.38 5.45 5.4.1  3.43 3-95 6.76  5.48  2.64  I' = 0.64 x 1 0 - 6 M atm-i.sec-i  S' = O.I65 I O M atm" s e c " x  (c)  I n i t i a l Conditions:  Expt No  205 199 204 203 202  197  201 200  198 Average:  110°C;  [Ag(I)]  [NaOAc]  M  M  .112  .192 .189 .186 .183 .183 .182 .182 .181 .180  .109 .105 .105 .105  .102 .101 .100  .099 .105  - 6  2  1  0.115 M AgC10 ;  1  O.I95 M NaOAc  4  Pco  [H0Ac]-i M-I  atm  1-7 6.2 11.6 20.0 20.0  4.56 4.56 4.56 4.49 4.49 4.65 4.49 4.53 4.49  19.7 27.6  27.2 26.1  .I85  I' = 5.2 X  10-6  S' = 1.04 x  10-6  M atm" s e c 1  - 1  M2 atm- s e c 1  - 1  R' x 106 M.a--L S"  5 .01 9 • 72 15 .6 20 • 9 20 • 9 21 .6 23 • 7 24 .6 27 .6  1  - 119 APPENDIX E THERMODYNAMICS OF THE OXIDATION OF CO, H , HCOOH AND HCOO" 2  IN AQUEOUS SOLUTION AT 25°C (Figure 1)  The thermodynamics of the reactions considered i n Figure 1 (Section 1-2) are summarized i n terms of the free energy data (10) i n Table E-I using the equations E - l and E-2.  The standard states are chosen as unit molarity  for dissolved species and one atmosphere f o r gaseous species. national or Stockholm Convention  The Inter-  ( 9 ) i s used f o r the sign of electrode poten-  tials .  TABLE E-I Standard Free Energy at 25°C (kcal/mole)  -32.808 -94.260 -126.22  C0(g) C0 (g) C0 =(aq) HC0 "(aq) HCOO-(aq) HCOOH(aq) H (aq) H (g) HaO(l) 2  3  -140.31  3  -80.0 -85.I  0.0 0.0 -56.690  +  2  AG° = -nFE° = -2.30 RT l o g K E  = E° -  -30 RT  2  nF  l  o  g  Q  (E-l) /  E  _  2  )  where AG°  = standard Gibbs free energy f o r the electrode reaction  E°  = standard electrode p o t e n t i a l (volt/mole)  E  = electrode p o t e n t i a l (volt/mole)  n  = number of v o l t equivalents  (kcal/mole)  - 120 -  •  (1)  F  = Faraday constant  R  = u n i v e r s a l gas constant  T  = a b s o l u t e t e m p e r a t u r e (°K)  K  = thermodynamic e q u i l i b r i u m  Q  = activity  =  -F  constant  +  2  0  = 4.762 k c a l / m o l e ; = -0.103 -  E° = 0.103 v o l t s  0.059pH - 0.030  log(aco/aco ) 2  H C 0 - ( a q ) .+ 3 H ( a q ) + 2e = 0 0 ( g ) + 2H 0(1) +  3  2  0  E  = -5.88 k c a l / m o l e ; = 0.128 -  0.089PH  E° = 0.128 v o l t s - 0.030 l o g ( a  c o  /a  H C  o -) 3  C 0 = ( a q ) + 4 H ( a q ) + 2e = 0 0 ( g ) + 2H 0(1) +  3  2  AG°  = -19.97 k c a l / m o l e ;  E° = 0.433 v o l t s  E . = 0.433 - 0 . 1 l 8 p H - 0.030  log(000/8^03 =)  2 H ( a q ) + 2e = H ( g ) +  2  AG°  =0;  E, =  (5)  deg-  quotient  2  AG  (4)  - 1  C 0 ( g ) + 2 H ( a q ) + 2e = C 0 ( g ) + H 0 ( 1 )  E  (3)  (I.987 c a l m o l e  0.059 ( v o l t s / e q u i v a l e n t ) a t 25°C  AG  (2)  (23.06 k c a l / e q u i v a l e n t )  E° = 0  -0.059pH -  0.030 l o g ( a  H 2  )  C 0 ( g ) + 2 H ( a q ) = 2e = HCOOH(aq) +  2  A G ° =9.16 k c a l / m o l e ;  E° =  -O.I99  volts  E : = -0.199 - 0.059PH - 0.030 l o g ( a 0 0 H / a c 0 H C  (6)  2  C 0 ( g ) + H ( a q ) + 2e = HC00"(aq) +  2  AG" E  = 14.26 k c a l / m o l e ;  E° = -O.3O9 v o l t s  = -0.309 - 0.030PH - 0.030 l o g ( a  H C O  o-/aco  2  - 121 -  HC0 (aq) + 2H (aq) + 2e = HCOO-(aq.) + H 0(1) _  +  3  2  AG" =3.62 kcal/mole; E  = _o„079  E° = -O.O79 volts  - 0.059PH - 0.030  log(a  HC00  -/a  HC03  C0 = (aq.) + 3H (aq.) + 2e = HC00 (aq) + H 0(1) +  _  3  2  AG° = -10.47 kcal/mole; E  = 0.227 - 0.089PH-0.030 log(aHCOO-/aco ") 3  HCOOH(aq)  = H (aq.) + HC00 (aq) +  _  AG° =5.1 kcal/mole; pH = 3.7  +  l o g K = -3.7  l o g ( HC00"/aHC00H) a  = H (aq) +, HC0 "(aq)  C0 (g) + H 0(1) 2  +  2  3  AG° = 10.64 kcal/mole;  pH = 7-8 + l o g ( a HC0 -(ao_) 3  E° = 0.227 v o l t s  HC03  l o g K = -7.8  -/a ) C02  = H (aq) + C 0 = (aq) +  3  AG° = 14.09 kcal/mole;  l o g K = -10.3  pH = 10.3 + log^ogz/ancog-)  -)  - 122  APPENDIX F  NUMERICAL INTEGRATION OF EXPERIMENTAL RATE LAW The rate of CO-reduction  (Figure 2 5 , Section 1 1 1 - 1 0 )  of s i l v e r perchlorate i n sodium acetate -  acetic a c i d buffered solution i s given by equation 1 9 : R' =  ( k l  [AgOAc ] + k ' [ g + P - [ | | l j + kg' [Ag+] A  2  [AgOAc-]{jjgjgj.  ( 1 9 )  A method f o r numerically integrating equation 1 9 i s outlined i n the present section.  From equation 1 9 ,  -d[CO]_ dt  p  =  c o  R  .  (M sec-i)  (F-l)  From Appendix A* : -d[CO]  dPr  =  dt  Let  /(F + a 0 ) ( F + a C  \ ( «C0  d t  rate f a c t o r (R.F.)  .  -dP  T  =  (  _  PCO  "dt -  F  + a  CO)(F  RI  o .  )\  c o  1  - C0) a  2  +aC0 ) 2  (M sec-i)  ( M  /  (A-1)  a t m )  .  ( v  p  ,  R  i F  CO  - ^CO  R.F.  \ R«.  (atm  sec-1)  (F-3)  i where and  PQQ APco  =  i n i t i a l CO pressure  =• decrease i n CO pressure i n time t  From Appendix A, the amount of CO consumed i n time t i s given by:  X *  - APco (F +«co)  = APT  See Appendix A f o r d e f i n i t i o n of terms.  x  r  - f  () m  ( +) fJ  - 123 -  AP  =  C 0  AP  '*'  (atm)  B T  ^ "CO  (F-5)  Substitution i n F-3 gives:  -dPrp  / P c o " APT R.F./(F  dt  V  Co)  (atm s e c - i )  R'  R.F.  (F-6)  QQ can be calculated by the method outlined i n Appendix  R.F. and F + A.  +a  PQQ and APIJ can be obtained from an experimental pressure vs time record. R' can be evaluated i n the following manner, using equation 19 and the  stoichiometry of the o v e r a l l reduction reaction represented by equation F-7:  2Ag(I)  + CO + H 0 + 2NaOAc  > 2Ag + C0  2  2  + 2H0Ac + 2Na  +  (F-7)  Since X represents the amount of CO consumed i n time t , then, i f [ ]^ denotes i n i t i a l concentration,  [ A g C l O J i - 2X  =  [AgOAc ] + [Ag ]  (F-8)  [NaOAc]± - 2X  =  [AgOAc] + [OAc ]  (F-9)  [HOAcJi  =  [HOAc]  (F-10)  where X i s given by F-k.  + 2X  +  -  [AgOAc] can a l s o be expressed i n terms of the  association constant, K , according to equation F - l l : a  [AgOAc]  =  K [Ag+][OAc-]  (F-ll)  a  Substitution- f o r [Ag ] and [OAc ] from F-8 and F-9 gives: +  -  [AgOAc]  where  =  b - (b  b  =  0.5  c  =  ( [AgC10 ]i [NaOAc ]i - 2X  ( [ A g C l O ^ i + [NaOAc ] 4  ±  2  - c)2"  + l / K - kX) a  ([AgC10 ]i 4  + [NaOAc ]± )••+ kX)  (F-12)  - 124 i n terms of F - 8 , F - 9 , F-10 and F-12 using an experi-  Thus R can be evaluated 1  mental set of rate parameters. A l l the terms on the right-hand  side of F-6 are now available and  F-7 can be integrated from zero t o t : ^0-A^R.F./(F °c  -dPT dt~  +  60 =  x  l  f(APT)  ^t  m  ^0  l  j , r  -AP ^F./(F T  (  a  t  m  s  e  c  _  i  a )y,  +  }  ( p s i  C0  - .  min  (psi min-i)  1)(F  (F-l4)  =  t  =  -/ ~o  (f ( A P T ) )  _1  dP  T  (min)  (F-15)  The time ( i n min) required f o r a decrease i n t o t a l pressure  ( i n p s i ) can be  calculated from the area under a (f (ABj)) ""h/s APrpj plot-. The i n t e g r a l i n F - 1 5 was evaluated numerically on an IBM 1 6 2 0 d i g i t a l computer using the trapezoidal rule at 1 . 0 p s i i n t e r v a l s . Sample Calculations Expt Nos 2 0 7 and 2 3 0 (Figure 2 5 , Section 111-10; see also Appendix D-III-L) I n i t i a l Conditions:  O.I95 M NaOAc;  0 . 1 1 5 M AgC104 ; 7 9 p s i CO;  Gas S o l u b i l i t y at 90°C  OQ  O.O78MHOAC;  90°C  = 6 . 9 x 10- M/atm 4  0  aco = IO5.O x 10-4 M/atm 2  Volume of solution added (V]_) = T o t a l volume of reactor at 90°C = Density Vg =  of H 0 2  15)  TAPT dt  o  k  0 )  =  O.9655  9 0 . 0 mis 120.0 mis  g/ml at 90°C;  120.0 - 9 0 . 0 x O . 9 9 8 2 / O . 9 6 5 5  =  O . 9 9 8 2 g/ml at 20°C 2 6 . 9 mis  - 125 -  F  R.F.  =  !g Vi  1  0  0  26.9 90.0  0  x  =  RT  =  100.3 x 10"  =  (F + a c o ) ( ( co a  2  +  1000 82.05 x 363  M/atm  4  F  X  a  C0 ) 2  - co) a  =  (100.3 + 6.9)10~ (100.5 + 105.0)10(105.0 - 6.9)10-4  =  0.0224  4  Values f o r ( f ( A P T ) )  _1  4  M/atm and calculated time computed from the above  data at 1-psi i n t e r v a l s up to about 50$ reaction using the "best value" rate constants* f o r equation 19 from Table X, are given i n Table F-I. Also included are the time values taken from the experimental records f o r experiments 207 and 230.  *  k' =  2.7 x 10- atm-i s e c - i ;  k' =  6.2 x 10-4 M  x  3  5  _1  sec-i  k ' = 2.1 x 10" M-i s e c - i ; 4  2  - 126 TABLE F-I COMPARISON OF EXPERIMENTAL AND CALCULATED PRESSURE RECORDS (Figure 25, Section III-10)  ,  APT* 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14  15 16 17 18 19 20  (f(APrrj))  _1  Calculated  min p s i - i  0.67 0.75 O.85 O.95 1.07 1.21 1.37 1.55 1.76 2.00 2.28 2.60 2.97  3.41  3.92 4.53 5.26 6.12 7.16 8.43 9.98  0 0.7 1-5 2.4 3.4 4.6 5.9 7.3 9.0 10.9 13.0 15.4 18.2 21.4 25.1 29.3 34.2 39.9 46.5  54.3 65.5  Time  (min) Experimental  No 207  No 230  0 1.0 1.8 2.8 3-7  0 0.9 1.9 2.7 3-8  5-6  6.0  8.7  10.3  12. C 15.8  13.7  ---  --  --  19.0  --  27.0 30.0 34.5 37-2 45.7 50.2 57-8  A t o t a l pressure decrease of 20 p s i i s equivalent to about 50$ reaction.  -  VI  127  -  REFERENCES  1.  Forward, F.A., Trans. Can. Inst. Min. and Met., 5_6, 363 (1953).  2.  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Orgel, L.E., "An Introduction t o Transition-Metal Chemistry", John Wiley and Sons, Inc., New York, i960, p 135.  24.  Orgel, L.E., International Conference on Co-ordination Chemistry, Special Publication No. 13, Chem. S o c , London, 1959, p 93.  25.  Sidgwick, N.V., "The Chemical Elements and Their Compounds", v o l I, Oxford University Press, London, 1950, P 5^7-  26.  Cable, J.W. and Sheline, R.K., Chem. Rev., 5_6, 1 (1956).  27.  Anderson, R.B., i n "Catalysis", v o l IV, edited by Emmett, P.H., Reinhold Publishing Corp., New York, I956, p 29.  28.  Wender, I., Sternberg, H.W. and Orchin, M., i n "Catalysis", v o l V, edited by Emmett, P.H., Reinhold Publishing Corp., New York, 1957, p 73.  29.  Sternberg, H.W. and Wender, I., International Conference on Co-ordination  r  Chemistry, Special Publication No. 13, Chem. S o c , London, 1959, P 3530.  Halpern, J . and Kettle, S.F.A.,•Chem. and Ind., 668 (1961).  51.  Bjerrum, J . , "Metal Ammine Formation i n Aqueous Solution", P. 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Chem.,  5_6, 1090 (1952).  - 131 VII  NOMENCLATURE USED IN RATE EXPRESSIONS  Rate Functions R  = -d[CO]/dt = rate of CO consumption  R' = I  R/Pco  (M sec-i)  rate of CO consumption at unit pressure (M atm - l  =  sec-i)  = acid-independent rate (M sec-i)  I' = acid-independent rate at unit pressure (M atm-i sec-i) = acid-dependent rate (M sec-i)  D  D' = acid-dependent rate at unit pressure (M atm i _  S  sec ) - 1  = acid-proportionality constant f o r acid-dependent  rate (M  S' = acid-proportionality constant at unit pressure (M  2  2  atm'- i  sec-i) sec-i)  Ro = acetate-independent rate (M s e c i ) -  Ho' = acetate-independent rate at unit pressure (M a t m  -1  sec  Rate Constants k, k 3 =  experimental rate constants i n terms of [CO]  k ', k 3 =  experimental rate constants i n terms of PQQ  2  ki',  !  2  k ", k '"  =  rate constants f o r reaction i n unbuffered solution (Section I I I - 9 )  k ", k "  =  rate constants f o r acid-dependent reaction (equations 15 and  2  2  3  4  16)  ka, kb, k , k^, k , kf, kg = rate constants f o r proposed mechanism (reaction V) c  e  Cn(n=l,2,3,4) = [OAc ]-power series c o e f f i c i e n t s i n eqn 38 f o r acid-dependent _  react.  Equilibrium Constants = i o n i z a t i o n constant of acetic a c i d (M) K  = association constant f o r average silver-acetate complex (AgOAc) from Ag+ and OAc- (M~i) K i , K = association constants f o r AgOAc and Ag(OAc) from A g and OAc (eqn 37) a  +  2  K  -  2  = formation constant, incorporating [ H 0 ] , f o r intermediate complex from Ag , CO and H 2 0 i n proposed mechanism (reaction V(d)) 2  c  +  K'  = formation constant, incorporating [ H 0 ] , f o r intermediate complex from AgOAc, CO and H 2 0 i n proposed mechanism (reaction V(h)) 2  c  a  C0  a  C 0  =  2  CO s o l u b i l i t y c o e f f i c i e n t (M/atm) =  C02  s o l u b i l i t y c o e f f i c i e n t (M/atm)  

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