UBC Theses and Dissertations

UBC Theses Logo

UBC Theses and Dissertations

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

You don't seem to have a PDF reader installed, try download the pdf

Item Metadata

Download

Media
[if-you-see-this-DO-NOT-CLICK]
UBC_1962_A1 M2 C2.pdf [ 7.08MB ]
[if-you-see-this-DO-NOT-CLICK]
[if-you-see-this-DO-NOT-CLICK]
Metadata
JSON: 1.0105706.json
JSON-LD: 1.0105706+ld.json
RDF/XML (Pretty): 1.0105706.xml
RDF/JSON: 1.0105706+rdf.json
Turtle: 1.0105706+rdf-turtle.txt
N-Triples: 1.0105706+rdf-ntriples.txt
Original Record: 1.0105706 +original-record.json
Full Text
1.0105706.txt
Citation
1.0105706.ris

Full Text

CARBON .MONOXIDE REDUCTION OF AQUEOUS SILVER ACETATE by ROBERT THOMSON McANDREW B.Sc, Queen's Un i v e r s i t y , 1957 M.Sc, Queen l;s U n i v e r s i t y , I958 A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHY i n the Department of METALLURGY We accept t h i s t h e s i s as conforming to the required standard THE UNIVERSITY OF BRITISH COLUMBIA August, 1962 In presenting this thesis in p a r t i a l fulfilment of the requirements for an advanced degree at the University of British Columbia, I agree that the Library shall make i t freely available for reference and study. I further agree that permission for extensive copying of this thesis for scholarly purposes may be granted by the Head of my Department or by his representatives. It i s understood that copying or publication of this thesis for financial gain shall not be allowed without my written permission. Department of Metallurgy  The University of British Columbia, Vancouver 8, Canada. Date September 10, 1962 The University of Br 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.Sc, Queen's University 1957 M.Sc., Queen's University 1958 MONDAY, SEPTEMBER 10, I962 AT 10:00 A.M. IN ROOM 201; MINING BUILDING' External Examiner: W.K. WILMARTH, University of Southern California, Los Angeles COMMITTEE IN CHARGE Chairman: F.H. SOWARD W.M..ARMSTRONG W.A. BRYCE D.L.G. JAMES E. PETERS C.S. SAMIS E. TEGHTSOONIAN CARBON MONOXIDE REDUCTION OF AQUEOUS SILVER ACETATE 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 l i q u i d phase'by two parallel reaction paths, one of which i s 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", ILpO) according to the following mechanism: Ag + + OAc" -—* AgOAc (rapid equilibrium) 0 AgOAc + CO Ag-C-OAc (slow) 0 Ag-fi-OAc + Ag + + H20 % 2Ag + CO + HOAc + H + (fast) 0 + Ag + + CO + H20 ^==i Ag-C'-OH + H (rapid equilibrium) 0 Ag-C-OH + Ag + ^ 2Ag + C0 2 + H + (slow) 0 * Ag-C-OH + AgOAc 2Ag + COg + HOAc (slow) Silver-acetate' complexes are about a factor of three more reactive than hydrated silver ions in the pH- dependeht reaction. This enhanced reactivity is att ributed to stabilization by the basic acetate anion of the proton released in the reduction process. The effect of increasing pH on the reduction rate i s much greater than the specific effects associated with silver-acetate complexing. GRADUATE STUDIES Field of Study: Metallurgy Metallurgical Thermodynamics. C'~S. ,Samis Metallurgical Kinetics E. Peters Metallurgy of the Rarer Metals E. Peters Hydrometallurgy '. Staff Nuclear Metallurgy. . . W.M. Armstrong Other Studies: 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 Differential Equations.,..- F.M.C. Goodspeed PUBLICATIONS 1. McAndrew, R.T., and Peters, E., "The Displacement of Silver 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 inter mediate complexes by the insertion of a carbon monoxide molecule between a silver ion and.a co-ordinated oxygen-donating base (e.g. OAc", H20) according to the following mechanism: Ag + + OAc -^=^ AgOAc (rapid equilibrium) AgOAc + CO—^ >AgiLoAc (slow) 0 Ag-fl-OAc + Ag+ + H20 >2Ag + C02 + HOAc + H + (fast) p Ag + + CO + H20 ^  > Ag-d -OH + H + (rapid equilibrium) 0 Ag-<3-0H + Ag + ^b_^2Ag + C02 + H + (slow) Ag-3-OH + AgOAc —^>2Ag + C02 + HOAc (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 grateful to Dr. E. Peters for his inspiring direction of this investigation. Thanks are also extended to A. M. Armstrong for her thoughtful and constructive criticism during the preparation of the manuscript. Financial support from the National Research Council of Canada in the form of grants in aid of research and a Studentship, and from the Consolidated Mining and Smelting Company in the form of a Fellowship, is greatly appreciated. I sincerely thank my wife for her encouragement and help throughout the period of study. Her efforts are responsible for converting my semi-legible scrawl to readable typed copy. i i i TABLE OF CONTENTS Page I INTRODUCTION I - l General 1 1-2 Thermodynamic Considerations of Carbon Monoxide Reactions in Aqueous Solution 2 1-3 Structure of Carbon Monoxide and i t s Compounds 5 1-4 Kinetics of Metal Ion Reduction by Carbon Monoxide in Aqueous Solution 6 I- 5 Object and Scope of the Present Investigation 10 II EXPERIMENTAL II- 1 Reactor System 12 II-2 Pressure Measurement 12 . II-3 Temperature Control and Measurement 15 II -k Materials 16 II-5 Chemical Analysis 16 II- 6 General Experimental Procedure 16 III RESULTS AND DISCUSSION III- l Rate of Carbon Monoxide Reduction of Silver (I) 18 III-2 Chemistry and Stoichiometry of the Reaction 22 IIIT3 Effect of Carbon Monoxide Pressure 2k III -k Effect of Acetic Acid 2k III-5 Effect of Acetate Complexing 28 III-6 Acid-Independent Reaction ; 28 III-7 Acid-Dependent Reaction k2 III-8 Acetate-Independent Reaction ^9 III-9 Reduction of Unbuffered Silver Perchlorate 52 III-10 "Best Value" Rate Parameters at 90°C ...' 62 III-11 Proposed-Mechanism 67 III-12 Effect of Temperature 73 IV CONCLUSION 79 V APPENDICES APPENDIX A Method-of Estimating Rates from the Slope of Pressure-Time Records 8l APPENDIX B Solubility of Carbon Monoxide, Carbon Dioxide and Hydrogen in Water 86 APPENDIX C Silver-Acetate Complexing from E.M.F. Measurements 100 iv Page APPENDIX D Summary of Selected Experimental Data for the Reduction of Silver(I) Solutions by Carbon Monoxide 105 APPENDIX E Thermodynamic Calculations for the Oxidation of CO, E2, HCOOH and HCOO" in Aqueous • Solutions at 25°C 119 APPENDIX F Numerical Integration of Experimental Rate Law 122 VI REFERENCES VII NOMENCLATURE USED IN RATE EXPRESSIONS 127 l 131 V TABLES Page I Reduction Rate of Silver(I) by CO Under Various . Experimental Conditions 19 II Reproducibility of Rate Measurements 21 III Stoichiometry of Acetate-Buffered CO-Silver Perchlorate Reaction 23 IV Summary of Intercepts and Slopes from R' vs [HOAc] _i Plots at Various Degrees of Acetate Complexing 31 V Summary of the Dependence of S' on [Ag+] at 90°C k6 VI Summary of Acetate-Independent Rates 53 VII Stoichiometry of Unbuffered CO-Silver Perchlorate Reaction at 90°C and 53 atm CO 56 VIII Summary of Experimental Rate Constants in Unbuffered Silver Perchlorate Solutions at 90°C and 53 atm CO 59 IX Summary of Rate Parameters for 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 Solubility of CO in Water at 25 Atmospheres 89 B-II Effect of Pressure on Solubility of CO i n Water 91 B-III Solubility of H 2 in Water 95 B-IV Solubility of C02 in Acetate Solutions at 90°C 96 B-V. Solubility of C02 in Water at One Atmosphere 98 C-I Effect of Temperature on Silver-Acetate Complexing 10k D-I Effect of CO Pressure 105 D-II Effect of Acetic Acid 106 D-III . Effect of Silver-Acetate Complexing 106 D-IV Rates Used in Extrapolation to Zero Acetate 112 v i Page D-V Reduction of Unbuffered AgC104 Solutions 115 D-VI Data Included in Regression Analysis 117 D-VII Effect of Temperature 118 E-I Standard Free Energy at 25°C 119 F-I Comparison of Experimental and Calculated Pressure Records ... 126 v i i FIGURES Page 1. Potential-pH Diagram for the Oxidation of CO, H2, HCOOH and HCOO" at 25°C 3 2. Schematic Diagram of Reactor System 13 3. Stainless Steel Reactor Ik k. Typical Pressure-Time Records 20 5. Dependence of Rate on COcPressure 25 6. Dependence of Rate on [HOAc] and. [H0Ac]-i 27 7. Dependence of Rate on-[HOAc]-1 at Various NaOAc Levels 29 8. Dependence of Rate on [HOAc] - 1 at Various AgC104 Levels 30 9. Dependence of Acid-Independent Reaction on [Ag(I)] at Various NaOAc Levels 33 10. Dependence of Acid-Independent Reaction on [NaOAc] at 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]; plotted according to equation 10, assuming [OAc-] = [NaOAc] .. 37 13. Dependence of Acid-Independent Reaction on [Ag(I)] and-[OAc-]; plotted according to equation 10, assuming K a = 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 18. Dependence of Acid-Dependent Reaction on Acetate Complexing; plotted according to equation 15 7^ 19. Dependence of Acid-Dependent Reaction on Acetate Complexing; plotted according to equation 16 8^ 20. Dependence of Rate on [NaOAc] at Constant [Ag(I)] and [HOAc]/[NaOAc] 50 v i i i Page 2 1 . Dependence of R1[HOAc]/[AgOAc] on [OAc-] at Constant [Ag(I)] and [HOAc]/[WaOAc]; plotted according to equation 21 51 2 2 . Dependence of Acetate-Independent Reaction on [Ag(I)] 5^ 2 3 . Reduction Rate of A g C 1 0 4 in Unbuffered Solution; plotted according to equation 27 58 2k. Dependence of Rate on [Ag+] in Unbuffered Solution 6 l 2 5 . Comparison of Experimental and Calculated Pressure Records .... 66 2 6 . Dependence of Rate on [HOAc]-1 at 6 0 , 8 0 , 90 and 110°C 7k 2 7 . Arrhenius-Plots for-Acid-Independent and Acid-Dependent Reactions 77 B-l Measuring Burette System for Gas -.Solubility Determinations .... 87 B-2 Solubility of CO and-H2 in Water from 25 to 225°C 90 B-3 Solubility of C 0 2 in Water at One Atmosphere 99 C-l Experimental C e l l for E.M.F. Measurements 101 CARBON'.'MONOXIDE:-'REDUCTION OF AQUEOUS SILVER ACETATE I INTRODUCTION 1-1 General Carbon monoxide, although i t is a major by-product of many pyrometal- lurgical processes, finds l i t t l e direct commercial use in current metallurgical operations. As a reducing agent carbon monoxide exerts a slightly greater potential than hydrogen, which is used commercially in the production of copper, nickel and cobalt powders from ammoniacal leach solutions (1,2), and therefore might be expected to find similar applications. The fact that carbon monoxide is not used in large-scale operations may be explained in part by a general aversion to this gas because of i t s toxicity and the perhaps mistaken concept that CO-reduction would lead to higher coats. This latter objection is not necessarily vali d particularly in the v i c i n i t y of electric smelting plants where CO is produced in large volumes and available merely for the cost of collection. What is probably a more significant reason for i t s limited use is the general lack of detailed information regarding potentially useful CO-metal reactions. The greatest use of CO in the metallurgical industry has been the production of elemental nickel by the Mond carbonyl process. The process con sists essentially of reacting the gas with reduced nickel to form a nickel carbonyl which is then thermally decomposed to give a high purity nickel powder (3). A modification of the Mond process was used in Germany during World War II to produce nickel and iron for powder metallurgy applications (k). The carbonyl process has also been adapted to the recovery of iron and nickel from nickeliferous l a t e r i t i c ores (5). / Perhaps the greatest potential use of CO in processes of metallurgical importance l i e s in the displacement of dissolved metals from hydrometallurgical - 2 - leach liquors. A process has been developed (6,7) f° r the production of elemental copper, nickel or cobalt powders from ammoniacal solutions under autoclave conditions, reduction being accomplished at somewhat lower tempera tures than corresponding hydrogen reactions. In the production of nickel 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 is particularly useful because of i t s a b i l i t y to stabilize the copper(I) species against hydrolysis thus pre venting the precipitation of cuprous oxide and providing a purer product from a wider range of solution compositions. A Russian process for the recovery of nickel has been described (8) in which nickel concentrates are leached in ammonia solution under oxygen pressure, and the resulting solution reduced with CO to form a carbonyl which is then decomposed to yield the nickel product. 1-2 Thermodynamic Considerations of CO Reactions in Acfueous Solution When CO acts as a reducing agent in aqueous solution i t is 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 in Figure 1 in the form of a potential-pH diagram using the International or Stockholm Convention ( 9 ) for the sign of the - electrode potentials. The associated potentials of formic acid and formate are also included together with that of the hydrogen electrode. In construct ing the diagram, unit activity for a l l species except H + has been assumed. The various reactions considered are l i s t e d in Figure 1 together with a summary of the thermodynamic expressions used. The details of the thermodynamic cal culations are given in Appendix E. It is apparent from the diagram that CO is a stronger reducing agent than H 2 by at least 0.1 volts at a l l pH values. Thus CO might be expected to 0 2 k 6 8 10 12 Ik PH 1. C0 2 + 2H + + 2e = CO + H20 E = -0.103 - O.O59 pH 2. HC03- + 3H+ + 2e = CO + 2H20 E = 0.128 - O.O89 PH 3. C03= + 4H+ + 2e = CO + 2H20 E 0.435 - 0.118 pH IK 2H + + 2e = H 2 E = -O.O59 pH 5- C0 2 + 2H + + 2e = HCOOH E -O.I98 - 0.059 PH 6. C0 2 + H + + 2e = HCOO" E = -0.309 - 0.030 pH 7- H C O 3 " + 2H + + 2e = HCOO" •+ H20 E = -0.079 - 0.059 PH 8. C03= + 3H+ + 2e = HCOO" + H20 E = 0.227 - O.O89 PH 9- HC00H = H + + HCOO- PH = -log K = 3.8 10. C0 2 + H20 = H + + H C O 3 " pH = -log K = 7.8 l l . H C O 3 - = H + + C0 3 = PH = -log K = 10.4 Figure 1. Potential-pH Diagram for the Oxidation of CO, H2, HCOOH and HCOO- at 25°C - k - reduce several metal ions which are not reducible by H2. However, cobalt with a standard reduction potential of -0.277 volts (10) is the least noble metal reported to 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 in basic or slightly acidic solutions. Other metals which have been produced by CO-reduction include nickel (6,8,11,12), bismuth (12), copper (6,13), silver (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 in the reduction of permanganate to Mn02 in acid or neutral solutions and to Mn04=, which gives-Mn + + by dispro portionation, i n basic solution (15,17)• Chromate i s reduced to Cr 203 in acid solutions (18). The reduction potential of CO can be increased by raising the CO parti a l pressure, a hundred-fold increase being equivalent to 2 pH units or about 0.1 volts at room temperature. Higher temperatures increase the pressure dependence of the CO potential as well as increasing the reaction rates. Figure 1 indicates that formic acid or the formate ion is a stronger reducing agent than CO in acid solution while the reverse is true at higher pH values, the equipotential point occurring at about pH 7 at room temperature. Thus i t is possible for 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 original synthesis of sodium formate from sodium hydroxide.in the mid-nineteenth century (20). The standard reduction potential of the oxygen electrode in acid solution is 1.229 volts (10) and thus the CO-reduction of H20 to form C0 2 and H 2 is thermodynamically favourable. This is the water-gas shift reaction familiar in the gas phase at elevated temperatures. It has also been observed in basic aqueous solutions at temperatures greater than 150°C (21). - 5 - 1-3 S t r u c t u r e o f C a r b o n Monoxide a n d i t s Compounds The 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 e l e c t r o n s , 6 f r o m t h e c a r b o n atom ( i . e . ls 2 2 s 2 2 p 2 i n t h e g r o u n d s t a t e ) and 8 f r o m t h e oxygen atom ( i . e . ls 2 2 s 2 2 p 4 i n t h e g r o u n d 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 i p l y - b o n d e d s t r u c t u r e as i s N 2 , 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. I n terms o f t h e m o l e c u l a r o r b i t a l (MO) t r e a t m e n t t h e CO s t r u c t u r e i s d e s c r i b e d (22) by C[ls 2 2 s 2 2 p 2 ] + 0[ls 2 2 s 2 2 p 4 ] > C O [ K K ( z^ 2 ( y v ) 2 ( x v ) 2 ( w i r ) 4 ] The z ^ 7 o r b i t a l i s a non-bonding MO r e p r e s e n t e d l a r g e l y by t h e 0(2s) a t o m i c o r b i t a l (AO); 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 o f C(2s) and 0(2p x ); x^7 i s a n o n-bonding MO r e p r e s e n t e d l a r g e l y by t h e C(2px) AO; wir r e p r e s e n t s two d e g e n e r a t e p i - b o n d i n g MOfe o f C(2py) and 0(2py). The bonds formed by t h e yKy and wtr MOb g i v e CO i t s t r i p l e - b o n d s t r u c t u r e . Two d e g e n e r a t e p i - a n t i b o n d i n g MO's (vif) a r e a s s o c i a t e d w i t h w"[r and 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 l a r g e l y r e s p o n s i b l e f o r t h e s t a b i l i t y o f m e t a l c a r b o n y l complexes. 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 CO(x^T) o r b i t a l ( w h i c h c o r r e s p o n d s t o t h e C(2px) AO), a r e d o n a t e d t o a m e t a l atom and f o r m a sigma-bond when a c a r b o n y l complex i s formed*. S t a b i l i z a t i o n o f t h e complex o c c u r s t h r o u g h 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 o r b i t a l s o f CO t h u s p r o v i d i n g a mechanism f o r t h e r e m o v a l o f t h e e x c e s s charge w h i c h w o u l d o t h e r w i s e be p r e s e n t on t h e m e t a l atom. .The p i - a c c e p t o r and sigma- donor c h a r a c t e r i s t i c o f CO 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 s t a b l e complexes even t h o u g h CO i t s e l f i s a p o o r donor as shown by i t s i n a b i l i t y t o f o r m s t r o n g complexes 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) because 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  m e n t a l l y a l t h o u g h t h i s t y p e o f d o n a t i o n i s o b s e r v e d w i t h c e r t a i n o t h e r 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 transition metals in groups VI, VII and VIII of the periodic table ( 2 5 , 2 6 ) . Carbon monoxide is also absorbed by cuprous chloride in ammoniacal solution, by silver sulphate in concentrated sulphuric acid, by mercuric acetate in methyl alcohol, and by dry auric chloride (25)= In the Fischer-Tropsch synthesis CO and H 2 react in the presence of a suitable metal catalyst to yield primary alcohols, olefins and saturated hydrocarbons, the relative yields depending oh the operating conditions ( 2 7 ) . When olefins are added to the CO-H2 gas mixtures, alcohols, aldehydes and ketones are produced and the process is known as the 0 X 0 or hydroformylation synthesis ( 2 8 ) . The 0 X 0 synthesis requires a cobalt catalyst (in the metallic, salt or carbonyl form) while catalysts for the Fischer-Tropsch synthesis include cobalt, nickel, iron and other metals capable of forming metal carbonyls. The role of these catalysts is associated with their a b i l i t y to form metal carbonyls and related complexes which act as intermediates in the organic reactions ( 2 9 ) . The mechanisms for these and other syntheses involve the insertion of CO into metal-carbon and metal-oxygen bonds ( 2 9 ) . I-k Kinetics of Metal Ion Reduction by Carbon Monoxide in Aqueous Solution The rate of CO-reduction of both Ag 2 S 0 4 and C u S 0 4 has been reported (15) to be first-order in CO parti a l pressure up to at least 60 atm and second- order in dissolved metal. The following mechanism was proposed to account for the observed kinetics: Me+ + CO ^ = = M e C 0 + (rapid) (a) M e C 0 + + Me+ =^=> Me 2 C 0 + + (rapid) (b) (I) Me 2 C 0 + + + H 2 0 s l Q W > 2Me + C 0 2 + 2 H + (c) In the silver studies two series of measurements were made with i n i t i a l Ag 2 S 0 4 - 7 - concentrations between about 0.007 a n d- 0.03 M_, one series being buffered with O.65 M NH40Ac* and the other unbuffered. The rate law for the buffered series between 70 and 110°C is given as: -d[Ag(I)]/dt = 6 . 0 2 x 1 0 4 [Ag(I)] 2 P C Q exp(-93OO/RT) (M min"1) and for the unbuffered series between 70 a n-d 150°C as: -d[Ag(I)]/dt = 12.8x105 [Ag(I)] 2 P c o exp(-l^,100/RT) ,(M l i n ' i ) Thus at 90°C the reaction rate in the buffered system is about 36 times faster than in the unbuffered system. The reduction rate of CuS04 was measured in dilute unbuffered solutions, apparently to minimize corrosion and hydrolysis problems. Also a sheet of copper metal, etched to give a high surface area, was required to obtain reproducible results. The reported rate law as measured between 160 and 190°C has the form: -d[Cu(total)]/dt = 2.56 x 10 1 3 [Cu(total)] 2 P c o exp(-33,500/RT) (M min"1) Another kinetic study involving the CO-reduction of aqueous silver, amines in basic solution has recently been reported (1^). Measurable rates were readily obtained at atmospheric pressure and room temperature and were shown to be consistent with the rate law: -d[C0]/dt = k e x p[C0][AgL 2 +]/[HL+] where L denotes an amine ligand. This rate law i s equivalent to: -d[C0]/dt = k[C0][LAg0H] * OAc is used throughout the text to denote the acetate radical CH3C00~. - 8 - The experimental r a t e constant k e X p i s thus e q u i v a l e n t t o kK^K-^K^, where i s the f i r s t i n s t a b i l i t y constant of A g L 2 + , i s the b a s i c i t y constant of the amine and K n i s the a s s o c i a t i o n constant of AgL + w i t h 0H~. The f o l l o w i n g mechanism was proposed t o account f o r the observed k i n e t i c s : A g L 2 + + H 20 i=L^ LAgOH + HL + ( r a p i d e q u i l i b r i u m ) (a) 0 LAgOH + CO k > LAg-5-0H (rate-d.etermining) (b) ( I I ) 0 LAg-C-OH + LAgOH > 2Ag + C0 3~ + 2HL+ ( f a s t ) (c) For the l i g a n d s NH 3, CH3NH2> C 2H 5NH 2 and.(C 2H 5) 2NH v a r i a t i o n s i n k e X p were almost w h o l l y accounted f o r by v a r i a t i o n s i n K^K^ wh i l e kK^ remained e s s e n t i a l l y independent of the nature of the amine and had a value of about *105 M~2 s e c - 1 at 25°C . At [NH 4 +] gre a t e r than about 0.02 M w i t h NH 3 as the ligand (14a). departures from the above rat e law were observed which were e x p l a i n e d by the slowing down of r e a c t i o n 11(c), due t o lower [LAgOH], u n t i l competition between 11(c) and the reverse of 11(b) c o n t r o l l e d the r a t e . On the b a s i s of these f i n d i n g s i t was concluded (lk) t h a t h i g h pH r a t h e r than 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 the high r e a c t i v i t y of Ag(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 the nature of the proposed intermediate i n the rate-determining step 11(c) was drawn by analogy from the mechanism developed from a p a r a l l e l study (15) t o describe the CO-reduction of H g + + i n d i l u t e p e r c h l o r i c a c i d , v i z : 0 H g + + + CO + H 20 Hg-C-0H+ + H + (a) Hg-$-0H+ >Hg + C0 2 + H + ( f a s t ) (b) ( I I I ) Hg + H g + + >Kg£>++ ( f a s t ) (c) - 9 - The observed kinetics 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 in the rate-determining step III(a) is analogous to the stable methyl formate derivative AcO-Hg-C^-OCH3 which i s formed when CO reacts with methanolic solutions of mercuric acetate under similar conditions (30). It was suggested (15) that the rate-determining step involves the insertion of CO between Hg + + and a co-ordinated water molecule. Similar reactions involving the insertion of CO into metal-oxygen bonds are important in a number of metal carbonyl catalytic reactions (29). The kinetics of the CO-reduction of Mn04" (15), as measured in aqueous solution at atmospheric pressure over the temperature range 28 to 50°C and the pH-range 1 to 13, are consistent with the rate law: -d[C0]/dt = k[C0][Mn04"] with a AH* of 13 kcal/mole and a AS-t- of -17 e.u. The reduction i s strongly catalyzed by Ag + or Hg + + (15) a n d in dilute perchloric acid solutions the rate law i s : -d[C0]/dt = k[C0][Mn0 4-][X] where X represents Ag + or Hg + +. For Ag +, AH-t- = 1.2 kcal/mole and AS-t- = .31 e.u.; for Hg + +, A H * = 6.k kcal/mole and AS* = -21 e.u. It is suggested (15) that the high catalytic activity in these reactions may be due to inter mediates of the type Ag-8-0Mn03 and Hg-8-0Mn03+. The kinetics of the CO-reduction of B i 2 ( S 0 4 ) 3 at pH's between O.k and -0.7 are described by the rate law (12): - 10 - -a[Bi(III)]/dt = 8.2x106 [ B i ( I I I ) ] P C 0 ( [ H + ] - x - 0.24)exp(-23,OOO/RT) (M min" 4 The high acid concentrations are required to prevent hydrolysis. It was suggested (12) that the active species is BiOH + +. An induction period which decreases with increasing i n i t i a l [Bi(III)] indicates that the reaction is heterogeneous. The rate law f o r t h e CO-reduction of Ni(0Ac) 2 in HOAc -NH40Ac(0.5 M) buffered solution at constant pH of 5.3 is given (12) as; -d[Ni(II)]/dt = 80.6 [Ni(II)] ( P c o -5.2) exp(-12,000/RT) (M min-i) Other studies on the reduction of nickel(II) amine sulphate complexes (11,21) indicate that the rate increases with increasing pH. The CO dependence may reflect a thermodynamic influence on the rate. 1-5 Object and Scope of the Present Investigation This thesis embodies the results of a kinetic study of the CO- reduction of aqueous AgOAc and AgC104 in acid solution. The work forms part of a general investigation of the mechanisms by which CO displaces metals" from aqueous salt solutions. At the time the study was undertaken the only published information on the kinetics of the reaction was that of the CO-reduction of Ag 2S0 4 solu tions (13)- At 90°C the rates measured in NH40Ac buffered solution were some 36 times greater than the corresponding rates in unbuffered sulphate solution. I n i t i a l experiments in the present study revealed large pH effects on the reduction rate of AgC104 and AgOAc in acid solution. The previously proposed mechanism (see mechanism I, Section 1-4), however, takes no account of pH effects. Also the possible variation in reactivity of different Ag(I) com plexes (e.g. Ag +, AgS0 4 _, AgOAc, AgNH3, Ag(NHs)2, Ag(NH3)0H, etc.) was not - 11 - considered. At the concentrations used in the Ag2S04-NH4OAc studies (13) i t is estimated, from reported room temperature complexity constants (31) that more than 60$ of the silv e r was complexed with ammonia. The simpler system of AgC104, generally buffered with NaOAc and HOAc, was chosen to further elucidate the mechanism by which CO displaces silver from aqueous solution. A preliminary study of the CO-reduction of Cu(OAc)2 and Cu(C10 4) 2 was also made. The f i r s t step in the reaction apparently is the reduction of Cu(II) to Cu(I) which forms a stable complex with CO and then hydrolyzes to precipitate a Cu 20 product. It was demonstrated that both aqueous Cu(OAc)2 and Cu(C10 4) 2 catalyze the oxidation of CO by 0 2, measurable rates being obtained above 125°C. The catalyzed reaction is first-order in CO and zero- order in 0 2, and occurs homogeneously in the l i q u i d phase in a similar manner to the CuS04 -and Cu(C10 4) 2-catalyzed oxidation of H 2 by 0 2 (32). The rate increases with increasing [Cu(II)] and pH but the exact dependence was not determined. The CO-Cu(II) reaction is currently•the subject of another study in these laboratories. No catalytic activity, of aqueous AgC104 or AgOAc in the C0-02 reac tion was detected at temperatures to 150°C, although at 250°C Ag 2S0 4 is reported (32) to catalyze the H 2-0 2 reaction, possibly by a heterogeneous process. - 12 - II EXPERIMENTAL II-1 Reactor System Reduction experiments were conducted for the most part in a small (approximately- 120 ml) pressure vessel, the rate of reaction generally being followed by the decrease of pressure in 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 fitti n g s were manufactured from stainless steel (type 316) by Pressure-Products-Industries Inc., Hatboro,1 Pennsylvania andxdesigned :for: working:1.pressures :.up. to, 7200,-psi.,. .A schematici.diagram, of the..reactor system is given in Figure 2 and a section drawing of the reactor is shown in Figure 3- Pressure-tight closure of the vessel was achieved by using a totally enclosed stainless steel 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 in a vertical position on a shaker mechanism which when activated reciprocated horizontally at 275 o s c i l  lations per minute with a l - l / 2 - i n stroke. Gas inlet- and outlet lines of flexible l / l 6 - i n o.d. stainless steel (type 3^7) capillary tubing, and a stainless steel clad l / l 6 - i n o.d. iron- constantan thermocouple were connected through the bottom of the reactor by means of Ermeto stainless steel (type 17-4 PH) sleeve f i t t i n g s . Also a short l/8-in o.d. li q u i d sampling line 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 in the gas inlet line. The transducer consisted essentially of strain-gauge windings - 13 - A E V 3 A potentiometer PT pressure transducer B recorder TH thermocouple C temp controller TP thermister probe D reactor V! gas inlet valve E CO cylinder V 2 gas outlet valve SW selector switch v 3 l i q u i d sampling valve Figure 2 . Schematic Diagram of Reactor System - Ik - 1/8" Ermeto Set screw (8 at k^>°) Cover, 316-S.S. Thrust washer Cover nut, A-5136 steel Gasket, 316 S.S. + Teflon + asbestos --Gasket, Teflon Thermowell Reactor body, 316 S.S. 1/8" Ermeto l/l 6 " Ermeto (k at 9O0) Figure 3- Stainless Steel Reactor (full-scale) - 15 - connected in 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 el e c t r i c a l output proportional to the applied pressure. The bridge cir c u i t was excited by a battery of dry cells supplying up to five volts at 10 ma, and the output was measured with a high precision 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 excitation voltage of 5.O volts dc by the manufacturer and had a nominal 20 mv full-scale output. From the manufacturer's specifications i t i s estimated that the pressure measurements had an accuracy of better than 2$. II-3 Temperature Control and Measurement The temperature of the experimental solution was regulated to ± 0.3,oC using a Yellow Springs Instrument Co. Thermistemp Temperature Controller (Model 71) with a stainless steel clad thermister probe (No. k-06) mounted in a thermo- well in the reactor l i d . The heating unit was a 600-watt external band heater connected to a variable transformer through the controller which was shunted by a four-ohm nichrome resister when in the "off" position. Regulation was made by adjustment of the transformer to give a heat input sufficient to main tain the temperature at a value slightly below the operating point when the regulator was in the "off" position, and slightly above the operating point when the regulator was in the" "on" position. The top of the reactor was wrapped with removable asbestos lagging to reduce heat losses. Independent temperature measurements were made with an iron-constantan thermocouple sheathed in stainless steel inserted through the bottom of the reactor. The thermocouple was calibrated at the melting point of indium (156.4°C), the steam point and at room temperature by comparison with a - 16 - precision mercury-in-glass thermometer graduated in tenths of a degree. II-k Materials A l l chemicals were of reagent grade and used without further p u r i f i  cation. Carbon monoxide (99-5$ min) was obtained from the Matheson Co. Analysis of several gas samples indicated that N 2, 0 2 and H 2 impurities were less than 0.2$, 0.1$ and 0.002$ respectively. -Experimental solutions were prepared by mixing'and diluting aliquots of standard stock solutions. AgC104 solutions were f i l t e r e d after preparation and thus any trace amounts of C l - or C103~ impurity were removed as insoluble AgCl. NaOAc and HOAc solutions were prepared with known weights of reagent and were not standardized further. II-5 Chemical Analysis Silver solutions were analyzed by t i t r a t i o n in dilute n i t r i c acid with standard ammonium thiocyanate using fe r r i c nitrate 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 pH- meter (Model G). Gas samples were generally analyzed in a Beckman GC-1 chromatograph using molecular sieve and s i l i c a gel 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. When temperature control had been achieved the shaking mechanism was 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 li q u i d and gaseous phase was attained within two to four minutes after agitation had been resumed. Rates for reactions involving the reduction of AgC104 in HOAc-NaOAc buffered solutions were generally measured by a pressure-drop method in which the decrease in pressure of the closed reactor system was recorded. The majority of the rate data was obtained by measuring the slope of each pressure- time record after CO saturation was complete (two to four minutes). These slope measurements were converted to rates expressed as M s e c _ i * using CO and CO2 solubility data and the estimated volumes of gas and l i q u i d in 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 is given in Appendix A. To augment meager published data (34) the solubility of CO in 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 solubility of H 2 in water measured under similar condi tions. Also included are a few measurements on the solubility of C0 2 in water and HOAc-NaOAc solutions at 90°C and 2.6 atm. * A l l concentrations are expressed in terms of l i t r e s measured at room temperature (20-25°C) - 18 - III RESULTS AND DISCUSSION I I I - l Rate of Carbon Monoxide Reduction of Silver(I) Early experiments demonstrated that at low pH the reduction of AgC104 solutions was very slow even at temperatures of 175°C and CO pressures of 20 atm. When the solutions were buffered with NaOAc and HOAc, however, the rates were readily 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 typical pressure-time records is shown in Figure k. The fast i n i t i a l decrease in pressure corresponds to saturating the solution in CO and the slower pressure decrease thereafter corresponds to chemical reaction exclusively. The marked difference in the rate of pressure drop in the two regions of the curve is clear evidence that after i n i t i a l CO saturation the measured rate of CO consumption is effectively independent of mass transfer of CO between the gas and li 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 risen to ± 20$ or higher for faster or slower rates. The reproducibility of the rates, however, was generally better than ± 5$ a s shown in Table II. Results included in the table also indicate that the rate i s independent of the surface area of stainless steel and precipitated silver in contact with the solution. The reaction therefore is not hetero geneous in nature but must occur homogeneously in the l i q u i d phase. No effect on the rate was detected when the C0 2 concentration present at the time the i n i t i a l slopes were measured was increased a factor of five. The presence of small amounts of a i r (equivalent to about one atmosphere) was also without - 19 - TABLE I Reduction Rate of Silver(I) by CO Under Various Experimental Conditions Temp °C [Ag(I)] M [NaOAc] M [Acid] M pco atm R x 10sa M s - i Method 13 175 .03 • 05c 20 0.36e [Ag(I)] 90 .05 -- .002c 53 0.58f PH 90 1.0 .05c 53 12.0f pH 60 .10 • 13 • 77d 5^ O.95S Pressure 60 .10 • 13 .05^  5-3 5.81S Pressure 80 .11 • 19 .06d 5A I8.58 Pressure 90 .ok .66 .10d 12.4 4l.4§ Pressure 90 .2k .Ok • 77d 27.1 83.0S Pressure 110 .11 • 21 .28d 4.6 40.86 Pressure 110 .10 .18 .05d •^5 93-7g Pressure a - R = -d[C0]/dt = -0.5d[Ag(I)]/dt = -0.5d[H+]/dt b - [Ag(I)] = silver analysis of periodic l i q u i d samples pH = pH determination of periodic liquid samples Pressure = pressure-drop method c - Present as HC10 4 d - Present as HOAc e - Average rate between 2 and 15 hours f - Average rate after 15 minutes g - I n i t i a l rate after CO saturation complete (2-4 minutes) - 20 - Expt [HOAc] • I n i t i a l Rate No M 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 15 Time (min) Figure k. Typical Pressure-Time Records (0.235 M AgC104; 0.090 M NaOAc; 11.7 atm CO; 90°C) - 21 - TABLE II Reproducibility of Rate Measurements Expt Wo Temp °C •i-.. I n i t i a l Concentrations R x 106 M s- 1 R'x 10 s M a-i s- 1 Average Deviation i [AgC104] M [WaOAc] M [HOAc] M pC0 atm 139 90 .100 .090 .180 26.6 21.4 .804 6.2 14 0 90 .100 .090 .180 26.7 20.2 • 755 -0.3 l 4 l 90 .100 .090 .180 26.3 18.7 .711 -6.1 161 90 .115 .045 .766 27.7 14.4 .520 1.8 l 6 4 a 90 .115 .045 .766 27.5 14.2 .516 1.0 I65 b 90 .115 .045 .766 25.6 .12.7 .496 -2.9 265 90 .144 .195 • 195 5.16 19.4 3.75 -3.9 264G 90 .144 .195 • 195 5.10 20.9 4.09 4.9 266d 90 .144 .195 .195 5.04 20.3 4.02 3.1 267e 90 .144 .195 • 195 4.78 I8.5 3.87 -0.8 268^ 90 .144 .195 .195 5.04 ' 18.9 3.75 -3.9 198 110 .115 .195 .023 4.49 124. 27.6 9.1 200 110 .115 • 195 .023 4.53 111. 24.6 -2.8 201 110 .115 .195 .023 4.49 106. 23.7 -6.3 a - Contained 0.1 g Ag precipitated during previous experiment plus 2.0 g silver sponge obtained from the Consolidated Mining and Smelting Co. Ltd. b - Reactor thoroughly leached with HN03 prior to charging to remove a l l traces of previously precipitated Ag. c - Air evacuated from solution in reactor prior to heat-up. d - Added C0 2 to 5.7 psi prior to CO addition. e - Contained 3-3 g fine 316 S.S'. f i l i n g s . f - Contained 1.0 g Ag precipitated during previous experiments plus J.O g silver sponge obtained from the Consolidated Mining and Smelting Co. Ltd. - 22 - effect on the rate. I l l - 2 Chemistry and Stoichiometry of the Reaction When dissolved silver undergoes reduction by carbon monoxide in acid solution the reaction products are silver metal and carbon dioxide according to the reaction: The Ag(I)-CO-C02 stoichiometry of reaction IV was examined at several tempera tures by analyzing l i q u i d and gas samples taken after about ^>0-6ofo reaction had occurred. Solution composition and experimental conditions that gave relatively fast reduction rates were selected. The concentrations of CO and C0 2 were estimated from their p a r t i a l pressures using solubility coefficients given in Appendix B, and assuming ideal gas laws. The results of these measurements, as summarized in Table III, agree within 10$ with the stoichi ometry represented by reaction IV*. No H 2 was detected in any gas sample. the presence of aldehydes and ketones using 2,4-dinitrophenylhydrozine (35), and for compounds containing a CH3C0 group using the Iodoform test (36,37)- Liquid samples were analyzed qualitatively in a Beckman GC-2 chromatograph containing a Carbowax 1000 column using a thermal conductivity bridge detector, and also in a Perkin-Elmer Vapor Fractometer (Model 15^C) containing a dinonyl phthalate column using a flame ionization detector**. No organic compounds * The metallic silv e r precipitate from a typical experiment was confirmed by X-ray diffraction. Mr. D.J. Rose performed the analysis and his assist ance i s gratefully acknowledged. (IV) Several f i l t e r e d l i q u i d samples of reaction products were tested for ** Thanks are extended to Professors C.A. McDowell and J. Halpern of the U.B.C. Chemistry Department for allowing these analyses to be performed in their laboratories. TABLE I I I Stoichiometry of Acetate-Buffered CQ - S i l v e r P e r c h l o r a t e Reaction Expt No Temp °C I n i t i a l Concentrations [Ag(I)] Consumed M [CO] Consumed M [co2] Produced •M A [ A g ( I ) ] A[C0] A[C0 2] A [CO] [HOAc] M [NaOAc] M [Ag(I)] M c^o atm 228 90 0.222 0.115 5.44 O.O69 0.037 O.O38 1-9 1.0 229 110 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 for saturation of the aqueous solutions. The presence of previously precipitated silver, metal had no effect on the CO absorption in Ag(I)-free solutions. In the absence of CO no reduction of Ag(I) was detected in HOAc-NaOAc solutions over a period of several hours at 90°C. These observations are consistent with the absence of side reactions affecting the CO-reduction of AgC104 in NaOAc-HOAc buffered solutions and also with the stoichiometry represented in-reaction-IV. Ill-3 Effect of Carbon Monoxide Pressure Figure 5* shows that the reduction rate is directly proportional to CO pa r t i a l pressure up to at least 30 atm. The CO parti a l pressures were estimated by subtracting the vapour pressure of H20 and the parti a l pressure of a i r from the i n i t i a l saturation pressure. Small corrections were also made to take account of the reaction which occurred during CO saturation. The deviations from Henry's law for the solubility of CO in H20 at 30 atm have been.estimated to be only a few per cent (see Appendix B) and since the pos sible errors in the i n i t i a l rate measurements were between ± 10 and 20$, the reduction rate is proportional to the concentration of CO in solution, within experimental error. Figure 5 also indicates that the rate is favoured by a low HOAc/NaOAc ratio or by a high NaOAc concentration, or both. I l l - k Effect of Acetic Acid In view of the large variation in rate with pH.(Table I) and HOAc/ NaOAc ratio (Figure 5) a series of experiments was conducted over a wide * Data for rates shown in Figures and- Tables throughout the text are tabulated in 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 AgC104. The amount of free acetate i n i t i a l l y present was therefore identical in each experiment and [H+] was directly proportional to [HOAc] through the relation: [H+] = K ± .{HOAc] ( 1 ) LOAc-J v ' where Kj_ is the ionization constant of HOAc. The results of this series of experiments, as depicted in Figure 6, indicate that the rate (reduced to unit pressure*) is inversely proportional to [HOAc]. The non-zero intercept of R' vs [HOAc] - 1 on the R'- axis indicates that there is also an acid-independent contribution to the overall rate. Therefore at a constant degree of si l v e r - acetate complexing (i.e. constant amounts of added AgC104 and NaOAc) the apparent rate law i s : R = -d[C0]/dt (2) = I ' Pco + S'PC0/[H0Ac] (3) R' = R/ P c o (k) = I"! + S* /[HOAc] (5) = I' + D' (6) where R' = total rate of CO consumption at unit pressure (M atm-i sec-i) I' = acid-independent contribution to the total rate at unit pressure and constant complexing (M atm -i sec-i) * Since i t has been demonstrated in Section III - 3 that the rate is directly proportional to the pa r t i a l pressure of CO, rates referred to in the remainder of the text have generally been reduced to unit pressure and represented as R' (i.e. • R1 = -d[CO]/Pco = R / P c O M atm-i sec-i). dt / - 27 - Figure 6. Dependence of Rate on [HOAc] and [HOAc]-1 (0.115 M AgC104; O.I95 M NaOAcj 90°C) - 28 - D' = acid-dependent contribution to the total rate at unit pressure and constant complexing (M atm - 1 sec - 1) S' = acid-proportionality constant for the acid-dependent contribution to the total rate at unit pressure and constant complexing (M2 atm - 1 s e c - 1 ) . In terms of this nomenclature I 1 and S1 are respectively the intercept and slope of R' vs [HOAc] - 1 plots at constant Jcomplexing (e.g. Figures 6, 7 a n d 8). 1 Since the experimental solutions were buffered, an inverse dependence of R*on [HOAc] at constant degree of complexing corresponds to an inverse dependence on [H +] or a direct dependence on [0H-] for the acid-dependent con tribution to the rate. This effect is similar to that reported for the CO- reduction of aqueous basic silver amine solutions at room temperature and atmospheric pressure (14). In these amine solutions no pH-independent con tribution to the reduction rate was observed. Ill-5 Effect of Acetate Complexing The effect on the rate of varying the amount of NaOAc added at constant AgC104 addition is shown in Figure 7 while the effect of varying the amount of AgC104 added at constant NaOAc i s shown in Figure 8. The intercepts (I') and slopes (S') increase with increasing amounts of NaOAc and AgC104 indicating that the overall rate is favoured by complexing between silver and acetate. Data of R1 vs [HOAc]"1 at 90°C were obtained for fifteen different degrees of silver-acetate complexing as summarized in Table IV. I l l - 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]-1 at Various NaOAc Levels (0.117 M AgC104 initially; 12 atm CO; 90°C) - jo - 0 2 k 6 8 10 12 [HOAc]"1 (M-i) Figure 8. Dependence of Rate on [HOAc]"1 at Various AgC104 Levels (0 .090 M NaOAc; 12 atm CO; 90°C) - 31 - TABLE IV Summary of Intercepts and'Slopes from R' vs [HOAc] - 1 Plots at Various Degrees of Acetate Complexing * Series Code Wo of Expts [Ag(I)] M [WaOAc] M [Ag(I)][WaOAc] M2 I'x 10s M,a-i s- 1 S'x 106 M2 a-i s - i 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 . 1 7 8 . O O 7 8 .43 .081 G 4 .042 .215 .0090 .43 • 077 H 3 .230 .041 .0094 .50 .133 I 3 .114 . O 8 7 .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 * Concentrations of Ag(I) and WaOAc are the averages of the corrected values of a l l experiments in a particular R' vs [HOAc] - 1 series as tabulated in Appendix D. - 32 - total [Ag(I)] at various NaOAc levels produces a group of non-linear curves passing through the origin, as shown in Figure 9- When I' is plotted against [NaOAc] at various Ag(I) levels a similar set of curves is obtained, as shown in Figure 10. Combining the two sets of curves by plotting I' vs [Ag(I)][NaOAc], as in Figure 11, produces a single curve which passes through the origin and approaches an asymptote with increasing [Ag(I)][NaOAc]. This evidence suggests that the rate-determining step i n the acid-independent reaction involves a CO molecule and a complex between Ag + and OAc". The total concentration of silver complexed with acetate can be approximated by an average value for 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 contribution to the rate law at unit pressure might then be expressed as: I' = kx' [AgOAc] (7) = ki'KaUg+HOAc"] (8) where K a is the average s t a b i l i t y constant of AgOAc. Substituting for [Ag +] in terms of total [Ag(I)], K a and free [OAc-] gives**: T' - v [Ag(I)][QAc-] , . or upon rearrangement: [Ag(i)3 = 1 + 1_ ( 1 0 ) * Based on available s t a b i l i t y constant data at room temperature (38) the ratio of AgOAc to Ag(0Ac)g is about 5:1 at 0.2 M NaOAc. ** [AgOAc] = K a[Ag +][0Ac"] = [Ag(I)] - [Ag +] .-.[Ag+] = [Ag(I)j/(l + K a[0Ac"]) - 33 - 1.3 1.2 l . l 1.0 • 9 .8 • 7 .6 • 5 .4 • 3 .2 .1 [NaOAc] (M) / / - •/ © - 0.7 O- 0.2 0.08 Q- 0.0k _ / 1 : •/'• O / _ / . / / o - / / / A i / ) -i / / 1/ // A$ •\ y/ is i 1 .05 .1 .15 [Ag(I)] (M) • 25 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 .3 .h .5 [Ag(I)] [NaOAc] (M.2) Figure 11. Dependence of Acid-Independent Reaction on [Ag(I)][NaOAc] (90°C) - 36 - If the proposed form of the acid-independent contribution to the rate law i s correct then a plot of [Ag(I)]/l' vs [OAc -]- 1 according to equation 10 should be linear with a slope of l/k!'K a and an intercept of l / k ^ . Figure 12 depicts such a plot where [0Ac~] has been approximated by [NaOAc], which is a good estimate at high [NaOAc]. From Figure 12, ki' has a value of about 3 x 10-s atm-i s e c - 1 while K a has a value of about 2 M - 1. If corrections are made for the amount of acetate effectively removed from solution through complexing with silver, assuming a value of 2 M-i for Ka, the intercept, and hence ki' , remains essentially unchanged while the slope becomes fl a t t e r (i.e. the value for K a increases). Using an iterative procedure of this type values of K a between about two and five are obtained from the slope of [Ag(I)]/l' vs [ 0 A c ~ ] - i plots. E.M.F. measurements using cells of the type: Ag AgOAc NaOAc KN0 3 Saturated AgN0 3 NaN0 3 Ag have been made at room temperature (39) at ionic strength to about two. Using this reported data, assuming complete dissociation of AgN0 3 and representing the concentration of Ag(I) present as silver-acetate complexes by [AgOAc], values for the ratio [AgOAc]/[Ag+][OAc-] = K a were found to l i e between 2.8 and 3.4 M - 1 at ionic strengths between 0.1 and 0.9 increasing to ^.0 M - 1 at inf i n i t e dilution. To evaluate the temperature coefficient of Ka, further 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 cells of the type: Ag AgC10 4 NaOAc NaC10 4 8 .M AgC10 4 NaC10 4 Ag Complete details of these measurements and a summary of the results are given in Appendix C. Values of [AgOAc]/[Ag+][OAc"] = K a assuming complete dissocia tion of AgC10 4 were found, within experimental error, to be effectively Figure 12. Dependence of Acid-Independent Reaction on [Ag(I)] and [NaOAc]; p l o t t e d according t o equation 10, assuming [0Ac~] = [NaOAc] (90°C) - 38 - independent of temperature to 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 for K a of 3-7 M"1 to calculate the free acetate con centration [OAc-]* at 90°C, a plot of Ag(I)/l'vs [OAc-]"1, as in Figure 13, has an intercept equivalent to ki' = 2.5 x 10-s atm-i sec-i and an average slope equivalent to K a = 3-7 M-1. The slight curvature of the plot in Figure 13' might be due to a small systematic error in the estimation of [ 0 A c ~ ] , or to higher order silver-acetate complexes (e.g. A g ( 0 A c " ) 2 ) affecting the rate at high'acetate levels. A plot of I 1 / [Ag +] vs [OAc-], as i n Figure Ik, exhibits positive deviations from linearity with increasing [OAc"] and thus supports the latter explanation for the curvature in Figure 13. The average slope of Figure Ik is equivalent to k^ K a and yields a value for k^ of 2.5 x 1 0 - 5 atm - 1 sec"i assuming K a = 3•! M - 1. The dependence of the acid-independent reaction on silver-acetate complexing i s also illustrated in Figure 15 where a plot of I' vs [AgOAc] is linear, passing through the origin with a slope equivalent to kx' = 2.5 * 0.6 x 1 0 - 5 atm - 1 sec~i. The acid-independent contribution to the overall rate therefore has the form: I' = ki'' [AgOAc] (7) or I = ki' P c o [AgOAc] (11) = ki[C0][AgOAc"] (12) = k 1 K a [ C 0 ] [ A g + ] [ 0 A c _ ] (13) where I = the acid-independent contribution to the overall rate of CO con sumption (M sec - 1) k i = the apparent second-order specific rate constant for the acid- independent reaction (M _ 1 sec - 1) * For [OAc -], [AgOAc] and [Ag+] values using K a = 3.7 M _ 1 see Table V, Section III-7. - 39 - 50 o .5 io 15 20 25 30 35 o^ 45 50 [OAc-] (M _ 1) Figure 13. Dependence of Acid-Independent Reaction on [Ag(I)] and [OAc -]; plotted according to equation 10, assuming K a = 3.7 M _ 1 (90°C) - ko - [OAc"] (M) Figure lk. 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 solubility of CO in water of 6.9 x 10 - 4 M/atm at 90°C (see Appendix B) the experimental bimolecular rate constant kj. in equations 12 and 13 has a value of 0.04 ± .01 M-i sec" 1 at 90°C. Il l - 7 Acid-Dependent Reaction The variation of the acid-dependent reaction rate with i n i t i a l [Ag(I)] and [NaOAc], as summarized in Table IV (Section III-6), is depicted i n Figures 16 and 17. The slope of the S' vs[Ag(I)] curves at constant [NaOAc] (Figure 16) increases with increasing [Ag(I)] and thus indicates that the acid-dependent reaction has a higher than first-order 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 in that the slopes decrease with increasing [NaOAc], This indicates that S' has approximately the same depend ence as I' on [OAc -], which has been shown to be first-order. It has also been demonstrated from Rr' vs [HOAc] - 1 plots at constant degree of complexing that the acid-dependent reaction is inversely proportional to [H+] or [HOAc]/ [0Ac~], The slopes, S' , of R1 vs [H0Ac]-i plots therefore have a first-order [0Ac~] factor incorporated in them from the acid dependence. A simple rate law, consistent with the shape of plots in Figures 16 and 17, to describe the acid-dependent reaction might involve a second-order dependence on uncomplexed silver and an inverse dependence on [H+] or [HOAc]/ [OAc -]. In this case the slopes of R' vs [HOAc] - 1 plots at constant degree of complexing would be given by: S' = kg1 [Ag +] 2[0Ac-] (14) If equation 14 accurately represented the variation of S' then S'/[Ag +] 2[OAc -] should be constant for a l l degrees of complexing. The experimental data, as - 43 - Figure 16. Dependence of Acid-Dependent Reaction on [Ag(I)] at Various NaOAc Levels (90°C) - 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 not the case, but r a t h e r S'/ [Ag +][OAc -] increases w i t h i n c r e a s i n g [OAc~] or decreasing [ A g + ] , This v a r i a t i o n i s taken as evidence f o r at l e a s t a two-term r a t e law t o describe the acid-dependent r e a c t i o n , e.g.: S* = k 2' [Ag+]2[0Ac-] + k 3"[Ag +]2[0Ac-]2 (15) or S 1 = k 2' [Ag +] 2[OAc-] + k 4"[Ag +][OAc-] (16) I f e i t h e r equation 15 or 16 represents the c o n t r i b u t i o n t o the rate f o r the acid-dependent term then p l o t s of S r"/[Ag +] 2[OAc -] vs [OAc~] or [ A g + ] _ 1 should be l i n e a r w i t h i n t e r c e p t s of k2'- and slopes of ~k3" or k 4" r e s p e c t i v e l y . Figure 18 i n d i c a t e s t h a t S'/[Ag +] 2[OAc _] vs [OAc -] i s l i n e a r as p r e d i c t e d by equation 15 w i t h a slope of 3 . 2 ± 0 . 8 x 10~3 M"2 atm~i s e c - i and an i n t e r c e p t of 7 ± 7 x 10 - 5 M"1 a t m - 1 s e c " 1 . Figure 19 shows t h a t a p l o t of S'/[Ag +] 2[OAc -] vs [Ag"1"]"1 i s not l i n e a r , except perhaps a t low [ A g + ] - 1 v a l u e s , i n d i c a t i n g t h a t equation 16 i s inadequate t o describe the experimental observations over the e n t i r e concentra t i o n range i n v e s t i g a t e d . I n c r e a s i n g p o s i t i v e d e v i a t i o n s a t higher [ A g + ] _ 1 v a l u e s , which a l s o correspond t o higher [OAc"],•might be evidence f o r g r e a t e r a c t i v i t y of Ag(0Ac) 2 r e l a t i v e t o AgOAc*. Subsequent measurements presented i n Sections I I I - 8 and I I I - 9 , however, favour equation 15 t o describe the a c i d - dependent c o n t r i b u t i o n t o the r a t e law. Thus the r a t e law f o r the a c i d - dependent r e a c t i o n can be expressed as: D' " k = , [ A 8 + ] 2 l i i i T + vibrio*,-! ( I T ) • k = ' [ A g + ] 2 l i s T + [ A B + ] ™ { i ^ T <18> * Based on a v a i l a b l e s t a b i l i t y constant data a t room temperature ( 3 8 ) , the r a t i o of AgOAc t o Ag(0Ac) 2 i s about 5 : 1 at 0 . 2 M NaOAc. TABLE V Summary of the Dependence of S'on [Ag +]* at 90°C *-* Series Code S'x 10s M2 a - i s- 1 [Ag(I ) ] M [NaOAc] . M [Ag+] M [OAc-] M [Ag*]" 1 M"1 S'x 104 [AgOAc] M [Ag +] 2l0Ac-] M-i a ' l s - i B .026 .096 .041 .010 .086 .031 11.6 1.1 H .133 .230 .041 .018 .212 .023 4.7 1.3 D .066 • 115 .043 .011 .103 .031 9.7 2.1 J .384 .225 .082 .034 .191 .048 5.3 2.2 E .074 .090 .080 .017 .063 .073 15.9 2.6 A .017 .048 .041 .OO56 .042 .035 23.6 2.7 I .178 .114 .087 .022 .092 .065 10.9 3-3 L .320 .110 .190 .039 .071 .151 14.1 4.3 K .322 .110 .173 .036 .073 .136 13.7 4.4 C .058 .045 .089 .010 .035 .078 28.6 6.1 F .081 .044 .178 .016 .027 .161 36.4 7.0 G .077 .042 .215 .017 .024 • 197 4l.2 7-0 N .275 .056 .550 .036 .019 .513 51.3 14.8 0 .340 .056 .780 .041 .015 • 739 66.7 20.4 M .249 .040 .670 .028 .012 .642 84.4 27.4 Values of [Ag ] and [AgOAc] were estimated using Ka = 3-7 M" See Appendix D for Series Code references - 4 7 - Figure 18. Dependence of Acid-Dependent Reaction on Acetate Complexing; plotted according to equation 15 (90°C) Figure 19. Dependence of Acid-Dependent Reaction on Acetate Complexingj plotted according to equation 16 (90°C) - k9 - The overall rate law including both the acid-independent and the acid-dependent reactions then becomes: R' = k l ' [AgOAc ] + k2' [ A g + ] 2 | ^ l | . + k 3' [ Ag +][AibAc"]-[|^ (19) R = kx[CO] [AgOAc"] + katCOltAg+FJI^j + k 3 [C0] [Ag+] [AiOAc"]-[|A|l| (20) From Figure 18, k2' = 7 ± 7 x 10-5 M"1 atm"1 sec" 1 and k3' = k 3"/K a = 9.0* 2.1 x IO"4 M _ 1 atm - 1 s e c - 1 . Using a value of 6.9 x IO - 4 M/atm for the solubility of CO in water (see Appendix B) yields values for k 2 and k 3 in equation 20 of 0.1 ± 0.1 M-2 s e c - 1 and 1.3 - 0.3 M - 2 s e c - 1 respectively, at 90°C. From Section III-5 k x has a value of 0.0k ± .01 M _ 1 s e c - 1 . I l l - 8 Acetate-Independent Reaction In equation 19 the k2' -term is acetate-independent since [HOAc]/ [OAc-] is simply a measure of [H +]. TO verify the form of the acetate- independent contribution to the rate law several series of experiments were conducted at various low WaOAc levels using constant [Ag(I)] and HOAc/WaOAc ratios. Extrapolation of each series to zero [NaOAc] yields a measure of the reduction rate in the absence of acetate complexing. Several plots of this type are depicted in Figure 20. Although the total [Ag(I)] and HOAc/NaOAc ratio were held constant during each series, [Ag +] and [HOAc]/[OAc-] (or [H +]) varied with each experi ment due to differing degrees of silver-acetate complexing at the various acetate levels used. To minimize these effects plots of R'[HOAc]/[AgOAc] vs [OAc-] calculated using Ka = 3'7 M-i were extrapolated to zero [OAc-] as shown in Figure 21. Rearrangement of equation 19 according to equation 21, - 50 - Figure 20. Dependence of Rate on [NaOAc] at Constant [Ag(I)] and [HOAc]/[NaOAc] (90°C) - 51 - 0 .02 .Ok .06 .08 .10 .12 [OAc-] (M) Figure 21. Dependence of R1 [HOAc]/[AgOAc] on [ 0 A c ~ ] at Constant [Ag(I)] and [HOAc)/[NaOAc]; plotted according to equation 21 (90°C) - 52 - indicates that the intercepts of plots of this latter type are equivalent to k2' [Ag +]/K a while the slopes are equivalent to average values of ki' [HOAc]/ [OAc-] + k 3'[Ag +]. Table VI summarizes the intercepts, ^R' [HOAc]/[AgOAc]) Q , obtained from ten R"[HOAc]/[AgOAc] vs [OAc-] plots including those shown in Figure 21. Since the intercepts are equivalent to k2' [Ag +]/K a a plot of (R1, [HOAc ]/[AgOAc ] ^ 0 vs [Ag +] should be linear passing through the origin with a slope of k2'/Ka. Such a plot, as depicted in Figure 22, indicates that the data, although exhibiting 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), is consistent with a linear relationship having a slope of 8 ± 4 x IO - 5 atm - 1 sec - 1. Thus the acetate-independent contribution to the overall rate law can be expressed as: V = y t A g ^ O A c - ] = k , K [Ag+f ° [HOAc] [H+J K ' • or R - 1 : [CO][Ag +F[OAc-] = y K [CO][Ag +] 2 f25) R° ~ K 2 THOACT K 2 K L — [ I T — ( O ; where is the ionization constant of acetic acid. From the slope of Figure 22, using a value of 3.7 M _ 1 for Ka, k2' = 30 i 15 x 10-5 M - 1 atm - 1 s e c - 1 , and using a value of 6.9 x 10 - 4 M/atm HOTOLQQ, k 2 = 0.4 t 0.2 M-2, sec-i at 90°C. In Section IJI-7, k 2' and k 2 were estimated to be 7 * 7 x 10-5 M _ 1 atm - 1 s e c - 1 and 0.1 ± 0.1 M - 2 sec- 1 respectively. I l l - 9 Reduction of Unbuffered Silver Perchlorate A more direct estimation of the acetate-independent contribution to the overall rate can be made from rate measurements in unbuffered AgC104 solutions. The perchlorate ion is among the weakest complexing ligands for metal ions in aqueous solution (40). Thus in a solution of AgC104 in 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] / Ri[HOAc] \ \ [AgOAc]/o M a-i s - i x 10s 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 for Series Code references - ^ - ,[Ag(I)] (M) Figure 2 2 . Dependence of Acetate-Independent Reaction on [Ag(I)] ( 9 0 ° C ) - 55 - absence of other complexing agents. the silver can be considered to exist exclusively as the aquo-complex. Unfortunately the rate of CO consumption in such a system is too slow to be easily measured using the pressure-drop method. By measuring the pH of periodic li 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 + H20 >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. The stoichiometry between the consumption of Ag + and the production of R"1* as shown in Table VII is 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 = k 2'-Ki[Ag+]2/[H +] (22) The duration of each experiment reported in the present section was equivalent to about 15$ reduction and hence [Ag +] as well as the CO pressure remained essentially constant during each experiment. Equation 22 is thus equivalent to: d[H+]/dt = v2d[C0]/dt = k 2"/[H +] (25) The effect of CO pressure was not investigated and is assumed to be f i r s t - order. Since [Ag+] remained constant during each experiment i t is 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 additional saturated potassium chloride and f i l t e r paper salt bridge. - 56 - TABLE VII Stoichiometry of Unbuffered CO - S i l v e r 'Perchlorate Reaction 1 at"9'0°C and 55 atm CO Expt No [AgC104] (M) [H +] (M) A[Ag +] A[H +j I n i t i a l Final Consumed I n i t i a l Final Produced 262 .0530 .0458 .0072 7 x 10"6 .0081 .0081 0.9 260 '.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- 4 .204 .204 1.0 - 57 - necessary to assume a second-order [Ag +] dependence in unbuffered AgClC>4 solution. The rate constant k2'" in equation 23 i s thus equivalent to: k2"' = 2k 2'Ki[Ag+] nP C 0 (2k) Integration of equation 23 from zero to t yields: [H +] 2 - [H +]| = 2k2"*t (25) But rapid i n i t i a l reaction quickly makes [ H + ] 0 negligibly small with respect to [H +]. Taking logarithms of equation 25 gives: log [H+] = 0.5 log t + log k2"' (26) pH = -0.5 log t - log k2"' (27) Plots of pH vs log t at constant [Ag +] and Pco should be linear with slopes of -0.5- Figure 23 shows that this prediction is essentially correct with slight deviations from linearity being explained by the decrease in [Ag +] and the simultaneous formation of small amounts of C0 2 during each experiment. Values of k2'" can be obtained from Figure 23 using the differential logarithmic form of equation 23: k 2 - = M f *(log[H+]) t d(log t) v ' A summary of k2"' values obtained in this manner, assuming d(log[H +])/d(log t) = 0.5, is given in Table VIII. The apparent dependence of k2'" on [Ag +] can also be found from Figure 23 by plotting pH values, obtained at equal times for 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 I I I I I I I I I I I I 2.0 .2 .4 .6 .8 3.0 .2 .4 .6 .8 4.0 .2 .4 log tjme Figure 23. Reduction Rate of AgC104 in Unbuffered Solution; plotted according to equation 27 (53 atm; 90°C) TABLE VIII Summary of Experimental Rate Constants in Unbuffered Silver Perchlorate Solutions at 90°C and 53 atm CO [Ag +] M log [Ag +] t = 1000 sec t = 10,000 sec Mean k 2" a-i s - i pH [H +] M k2"'* M2 s" 1 k 2" ** a- 1 s - i pH [H+] M k 2 m * M2 s- 1 k 2" ** a-i s - i x 103 x 109 x 103 x 109 x 109 0.051 2.71 2.65 2.24 2.6 x 10-9 9-5 2.14 7.25 2.6 x IO" 9 9-5 9-5 0.097 2.99 2.40 3.98 7.9 x 10-9 7-9 1.88 13.2 8.7 x IO" 9 8.7 8.3 0.195 1.24 2.11 7.76 3.0 x 10-8 7-5 1.55 28.2 4.0 x IO" 8 9-9 8.7 O.96 I.98 1.33 46.8 1.1 x 10"6 11.3 O.79 162. 1.3 x IO"6 13.3 12.3 * > - - [H +] 2 d(log[H +]) n„ o Q K 2 - • ) ° L J' according to equation 28 t d(log t) k 2" = 0.5 k 2"'/[Ag +] 2 P c o according to equation 34 - 60 - k2'" = constant x' [ H + ] 2 (28) = constant x-[Ag+]n (24) .". pH = -0„5n log [Ag +] + constant (29) Two plots of this type prepared using the data summarized in Table VIII are presented in Figure 2k, and indicate that the rate of CO-reduction of unbuffered A g C 1 0 4 solution is second-order in [Ag +], In constructing Figure 2k no account was taken of the possible effect on ks'" of ionic strength which varied from about 0.05 to 1.0 throughout the experimental series (but remained essentially constant during each individual experiment). Rate measurements of the CO-reduction of unbuffered AgC10 4 solutions, assuming a first-order dependence on Pco, a r e consistent with a one-term experimental rate law of the form: . R Q = -d[C0]/dt = 0.5 d[H+]/dt (30) = k 2 " [ A g + ] 2 PC0/[H+] (31) or R0'. = k 2"[Ag + ] 2/[H +] (32) Equation 32 . i s of the same form as equation 22, thus supporting the experi mental rate law developed to represent the rate of CO-reduction of acetate- buffered AgC10 4 solution. From a comparison of equations 32, 22 and 2k i t is evident that: k 2" = k2''K± (33) k 2" = 0.5 k 2"'/[Ag + P P C O (34) , Average values of k 2" calculated using equation 34 are included in Table VIII and l i e in the range 8.3 - 12.3 x 1 0 - 9 atm-i sec-i. The 40$ variation in k 2" may be due largely to ionic strength effects particularly with the solution O.96 M in AgC10 4. The average of the other three values i s 8.8 x 1 0 - 9 atm-i sec-i and may be compared with a k 2 ' o f 3 . O i l . 5 x 1 0 - 4 M _ 1 atm - 1 sec-i Figure 2k. Dependence of Rate on [Ag +] in Unbuffered Solution (53 stm; 90°C) - 62 - estimated in Section III-8 for the acetate-independent reaction in buffered solutions. Thus from equation 33 Ki should have a value of about 3 x IO - 5 M to make the rate measurements in the buffered and unbuffered systems consist ent . The above result is in agreement with accepted values for Kj_ taking into account the effect of ionic strength on the ionization of weak acids in aqueous solution. For example the ionization constant of HOAc in NaCl solution at 25°C varies from 1.75 x IO - 5 to a maximum of 3-32 x 10 -5 between an ionic strength of zero and 0.5 (kl). The effect of temperature causes the value of Ki for HOAc in water to pass through a maximum of I.76 x 10 - 5 at about 25°C and f a l l to I.55 x 10-5 at 60°C (kl). Extrapolation of this data to 90°C yields a value of about 1.2 x 10 - 5. Thus i t appears from a comparison of the rates of CO-reduction of AgC104 in buffered and unbuffered solutions that the ionic strength rather than the temperature has a greater effect on the ioniza tion of acetic acid in HOAc-NaOAc solutions at 90°C. The results of the rate measurements made in buffered and unbuffered solutions are in good agreement considering particularly the widely different methods used to obtain the measurements and the approximately hundred-fold difference in overall rate. Ill-10 "Best Value" Rate Parameters at 90°C The CO-reduction of AgC104 in acetate-buffered solution is consist ent with a rate law represented by equations 19 and 20: R' = kx' [AgOAc ] + k2' [ A g + ] 2 ^ A g l j + y [ A g+] [ A^oAc"]|P^l| ( 1 9 ) R = k1[C0][Ai0Ac"] + k 2[C0][Ag + ] a | 0 A c 4 + k 3 [CO ] [Ag+] [AioAc"]-^!! (20) LHOAc J IHOAc J - 63 - Values for the rate constants at 90°C in equation 19, found by graphical analysis in Sections III-6, III-7 and III-8 are summarized in Table IX* Also included are the results of least square regression analyses on equation 19 using various sets of experimental data. The regression coefficients were evaluated on an IBM 1620 d i g i t a l computer using a slightly modified S3-4 s t a t i s t i c a l analysis program available from the program library of the U.B.C. Computing Centre. Unfortunately the program considers the experimental rate data to have equal possible errors on an absolute rather than a relative 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 in Table IX, a set of "best value" parameters at 90°C for equations 19 and 20 have been selected and pre- sented in Table X. The r e l i a b i l i t y of the acid-independent parameter is 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 is estimated to be ± 50$. As a further verification of the proposed rate law, equation 19 was integrated numerically using the "best value" rate constants from Table X and compared with the pressure records of several extended reduction experiments. Figure 25 shows the results for five experiments, two of which represent about 50$ reduction of AgC104. Experiments of longer duration were not made because of the pos s i b i l i t y of minute gas leaks producing erroneous results for extended measurements of slow rates. A sample calculation for the numerical integration is given in Appendix F. The calculated pressure-time curves are in good agree ment with the experimental curves and thus equation 19, developed from measure ments of i n i t i a l rates, is adequate to describe the CO-reduction of AgC104 in acetate-buffered solutions for at least 50$ reaction. * The rate constants in equation 20 can be obtained by dividing the corres ponding constants in eqution 19 by aQQ = 6.9 x 10 -4 M/atm. - 6k - TABLE IX Summary of Rate Parameters for Equation 1 9 at 90°C No of Expts Rate Parameters Reference ki'x 105 a" 1 s- 1 k 2'x 104 M-i a-i s" 1 k 3'x 10* M-i a- 1 s - i Varied [HOAc] at constant [NaOAc] and [AgC104] ". Series A-0 67 67 2.5 ± 0.6 0.7 ± 0.7 9.0 ± 2.1 Figures Ik, 15 Figure 18 67 3.0 0.9 5-7 Regression Analysis Varied [NaOAc] at constant [AgC104] and [HOAc]/[NaOAc] Series P-Y 51 — 3-0 ± 1.5 -- Figure 22 31 2.3 2-5 3-0 Regression Analysis A l l Experiments 119 3-0 l.k 5-8 Regression Analysis Equation 19 R' = k x' [AgOAc] + k2' [Ag +] 2 + kg' [Ag +] [AgOAc ]-[§|f^ * To obtain rate parameters in terms of [CO] divide values in Table IX by a = 6.9 x IO- 4 M/atm. 65 TABLE X "Best Value" Rate Parameters at 9Q°C Equation 19 Equation 20 Reli a b i l i t y = 2.7 X 10-5 a- i s - i = 0.04 M-i s - 1 ± 25$ k2' = 2.1 x 10"4 M-1 a- 1 s " 1 k 2 = 0.3 M-2 S" 1 ± 50$ k3' = 6.2 x 10"4 M-i a- i s - i k 3 = 0.9 M-2 s - i ± 50$ Equation 19 R' = kx' [AgOAc] + k2' [Ag + ] 2 j | A ^ | + kg' [Ag +] [AgOAc ]|^£l-| Equation 20 R = kitCO][AgOAc] + k 2 [ C 0 ] [ A g + j 2 [ 0 A c ~ ] + k 3 [CO ] [Ag +] [AgOAc ]• [HOAc J OAc-] .HOAc ] o> 0\ Time (min) Figure 25. Comparison of Experimental and Calculated Pressure Records ( I n i t i a l Conditions: 0;115 M AgC104; O.I95M.NaOAc; 5.4 atm CO; 90°C) - 67 - III-11 Proposed:Mechanism The kinetics of the CO-reduction of acetate-buffered AgC104 solutions are described by the experimental rate law: R = k x[CO] [AgOAc"] + k 2 [ C O ] [ A g + ] 2 j P i ^ + katCO] [Ag+][AiOA7 ] l |£ |^ (20) In terms of [H+] equation 20 may be re-written as: R = kJCOHAiOAc-] + k 2 K i [ C ° ^ f ] g + ^3K±[C0]^llf^l (35) A mechanism which is" consistent with equation 35 c a n be represented by the following scheme: -. Ag + OAc- v, AgOAc (rapid equilibrium) (a) AgOAc + CO k a >Ag-JJ-0Ac (b) 0 Ag-C-OAc + Ag+ + H20 f a s t > 2Ag + C02 + HOAc + H + (c) Ag + + CO + HaO - J&s Ag-C"-0H + H + (rapid equilibrium) (d) S-H O (V) Ag-^-OH + Ag + k b > 2Ag + C02 + H + (e) Ag-C-OH + AgOAc —££->2Ag + C02 + HOAc (f) The rate law derived from this sequence corresponds to: R = = k a[CO][AgT0A7] + k ^ M ^ p + W00" Ag^ [AgOAc] ^ which is identical with equation 35 i f k a = k x, kbKc = k 2Ki, k cK c = k 3Ki and i f [H20] is incorporated in Kc. Support for the nature of the proposed intermediates is drawn from studies on the CO-reduction of Hg + + in dilute HCIO4 (15), the reduction of Ag(I) amines in basic solution (Ik), and the Hg + + and A g + catalyzed reductions of Mn04~ in both acid and basic solution (15)• Analogous intermediates were proposed in these cases with the evidence for (HgJj-0H) + being particularly - 68 - strong because of the existence of the stable methyl formate derivative AcO-Hg J!-OCH3 formed when CO reacts with methanolic solutions of Hg(OAc)2 (30). The oxidation of CO in aqueous solution is apparently f a c i l i t a t e d by the presence of an oxygen-donating base (e.g. OH-, OAc", H20, Mn04"). The reduction of silver acetate by a pH-independent mechanism is similar to the reduction of Hg + + and Mn04~ (see mechanism III, Section 1-4), while the pH-dependent contribution is similar to the reduction of silver amines in basic solution (see mechanism II, Section 1-4). In the latter case the formation of the LAgJ!-OH complex (where L denotes an amine ligand) is rate-determining as evidenced by a first-order dependence on [Ag(I)] and [OH-]. Apparently at [NH 4 +] greater than 0.02 M (i.e. lower pH) with NH3 as the amine ligand the decomposition of the complex to form f i n a l products is retarded to" the extent that the rate is determined by competition between this step and a back-reaction to form i n i t i a l reactants. In the reduction of Ag(I) in acid solution the rate of complex decomposition to products relative to back- reaction 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 is determined exclusively by attack of another Ag(I) species (e,.g. Ag + or AgOAc) on the complex to form f i n a l products. A hydrolyzed silver(I) species (e.g. AgOH) is not required to explain the observed kinetics in acid solution. The association constant of Ag + and OH- is given as about 10 4 at room temperature in dilute solution (42). Thus while LAgOH i s the predominant silver species in basic amine solutions (assuming for AgL + is of the same order as for Ag +) AgOH is present only in trace amounts in acid solution. In the proposed mechanism for reduction of Ag(I) in acid solution reaction V(d) is indistinguishable from 0 Ag + + CO + 0H"^=^ Ag-C-OH - 69 - preceded by the water dissociation equilibrium : H20 ^ =£- H + + 0H- .TnepH effect arises in an equilibrium preceding the rate-controlling step and therefore i t s exact nature cannot be determined from the present kinetic study. The mechanism proposed to describe the CO-reduction of AgC104 is probably the simplest which is consistent with the present kinetic study in 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 to assess their v a l i d i t y on an experimental basis. For example Ag(0Ac) 2 might attack the intermediate complex formed in V(d) by a reaction par a l l e l to V(e) and V(f), viz: 0 Ag-C'-OH + Ag(OAc)2 ^d > 2Ag + C0 2 + HOAc + OAc" (V)(g) Also AgOAc might form an intermediate by an equilibrium analogous to V(d), viz: , 0 AgOAc + CO + H20 ^ K c ^  (AcOAg-C'-OH) " + H + (V)(h) This intermediate could be attacked through processes parallel to reactions V(e), V(f) and V(g), viz: ( AcOAgil-OH)" + A g + — k e _ > 2 A g + C0 2 + HOAc (V){i) 0 . ( AcOAg-fi-OH)" + AgOAc —^£_> 2Ag + C0 2 + HOAc + OAc" (V)(j) 0 (Ac0Ag-6-0H)" + Ag(OAc)2__kJL_> 2Ag + C0 2 + HOAc +20Ac" (V)(k) If the association constants of AgOAc and Ag(OAc)2 from Ag + and OAc- are represented by Kx and K 2 respectively and [H20] is incorporated in each of K c and Kc' , the acid-dependent reduction rate, considering a l l contributions from - 70 - reaction V(d) to V(k), can be represented by: D = [C0][Ag + ]2 (k bK c + (k cK cKi + k eK c'Ki) [0Ac~] + (k dK cK 2 + k f K c ' K ^ ^ O A c " ] 2 [H+] * . " + kgKc'KiKatOAc-]3) (37) The sum of the constants of like terms in 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 in [OAc-] to describe the kinetics of the pH-dependent CO-reduction of acetate-buffered AgC104 solutions. D = 2 C j O A c - f " 1 (38) [H +] n = i By a suitable choice of parameters equation J>Q can be fi t t e d to almost any set of rate measurements, be they good or bad, with a high degree of precision. Graphical analyses of the measurements made in the present investigation are consistent with a rate expression to describe the pH-dependent reaction that involves only the i n i t i a l two terms in equation 38 and reactions V(d), V(e) and V(f) give rise to a rate law of this form. Similar considerations to those presented above can be applied to the acid-independent reduction process and again i t i s concluded that the rate law arising from reaction V(b) i s of the simplest form consistent with experi mental observation. Many ions with a d 1 0 structure, including Ag +, Cu +, Au + and Hg + +, tend to form linearly co-ordinated complexes'* and thus i t is unlikely that Ag(0Ac) 2 is capable of reacting without prior dissociation to form-an inter mediate complex similar to those described by reactions V(d) and V(h). For * It has been suggested (43) that this effect is due to hybridization of d and s orbitals which remove charge from the region between the metal ion and i t s ligands and thus favours linear co-ordination for those &10 ions with sufficiently low d-s separations. - 71 - similar reasons i t is doubtful that the Ag(0Ac)3 complex proposed to help explain some equilibria studies in aqueous solutions (hk) plays a role in the reduction mechanism, i f in fact the'complex exists. In the interpretation of another study on equilibria in aqueous silver acetate solutions an Ag 2OAc + complex has been postulated (^ 5). The existence of such a complex reflects an a f f i n i t y of AgOAc for another silver ion. Because of the higher basicity of water compared to HOAc, which i s analogous to AgOAc, this second silver ion should strongly prefer to remain in the simple hydrated form, and therefore the existence of significant concentrations of Ag 2OAc + i s doubted. The amount of silver, present in the various acetate complexes has been approximated by an average [AgOAc] calculated from a mean association constant (K a). If the reactivity of AgOAc and Ag(OAc)2 were significantly different deviations from the proposed rate law would be expected particularly at high [OAc -]. . Reference to Figure 18,- Section III-7, in which S'/ [Ag +] 2[OAc~] is plotted against [OAc -], shows that no deviation in the acid-dependent rate expression is apparent at [OAc~] up to about 0.8 M, where AgOAc and Ag(0Ac) 2 should be present in approximately equal concentrations, based on published room temperature complexity constants.(38). It i s therefore concluded that AgOAc and Ag(0Ac) 2 are approximately equally reactive in the pH-dependent reduction process. Reference to Figure lk, Section III-6, in which l'/[Ag +] is plotted against [OAc -], gives some evidence that Ag(0Ac) 2 may be more reactive than AgOAc toward direct attack by CO. Silver carbonyl intermediates have been proposed to explain the experimentally observed rate of CO-reduction of dilute Ag2S04 solutions buffered with O.65 M NH40Ac- (13). At the concentrations used the majority of the silver was complexed with ammonia. The effect of pH was not investigated and thus a first-order dependence on [CO] and an apparent second-order depend ence on [Ag(I)] was taken as evidence for the rate-determining step involving - 72 - the reaction of H20 with Ag 2C0 + + formed in a pre-equilibrium^(see mechanism I, Section 1-4). More recent studies (14) have shown that at [NH 4 +] greater than about 0.02 M the reduction of Ag(NH3)0H + approaches second-order in both 0 [Ag(I)] and [OH-]. In this latter case a mechanism involving NH3Ag-Cl-0H as an intermediate was proposed. The existence of a silver carbonyl complex is s t i l l feasible particularly since analogous stable Cu(I) carbonyl complexes are known (25). Thus the observed kinetics for the CO-reduction of unbuffered AgC104 might be explained by a mechanism similar to: Ag + + CO ^ ---^ AgC0+ AgC0+ + OH" > products A rate law developed from this scheme using the steady-state approximation for [AgC0+] successfully described the experimental observations in unbuffered solution. When acetate contributions were considered, however, a satisfactory simple expression could not be developed. It was therefore concluded that such a mechanism i s not responsible for the CO-reduction of Ag(I) in acid solu tion . Silver carbonyl complexes formed in pre-equilibria might take part in the formation of Ag-C^ -OH in V(d), e.g.: Ag + + CO ;==^ AgC0+ AgC0+ + H 20^=^ AgJ-OH + H + This implies that a CO molecule f i r s t co-ordinates with an Ag + ion before 0 reacting further with an H20 molecule to form the Ag-£-0H intermediate. Kinetically, such a process i s indistinguishable from the direct insertion of CO between Ag + and a co-ordinated H20 molecule. Silver hydride complexes, which are postulated to be active inter mediates in H 2-reduction processes in aqueous solution (46,47), apparently do - 73 - not influence the kinetics of corresponding CO-reduction processes. The mechanism by which each gas reacts with Ag(I) i s specific; H 2 is activated by dissociation, while CO is' oxidized by the transfer of an oxygen atom from a donor-base (e.g. H20, OAc-). In both cases basic ligands increase the metal ion reactivity through stabilization of protons released in the reduction processes. It i s unlikely that nucleation or growth of silver crystals influence the reduction kinetics. Nucleation generally involves a high-order dependence on metal ion concentration, for example, the disproportionation of Cu(I) is tenth-order in [Cu(I)] (48), and growth rates for most metals are also fast as evidenced by low overvoltages required for electrodeposition (49). Trace amounts of precipitated silver from previous experiments were usually present in the reactor and served to minimize possible nucleation effects. Ill-12 Effect of Temperature The effect of temperature on the rate of CO-reduction of AgC104 in acetate-buffered solution was determined at 60, 80, 90 and 110°C by measuring the rate of CO consumption at a constant degree of acetate complexing at a number of HOAc concentrations. The results of these measurements are shown as R' vs [H0Ac]-i plots in Figure 26 and summarized in Table XI. In Section III-4 the overall reaction was shown to be made up of an acid-independent and an-acid-dependent component, i.e.: R' = I' + D' (6) = I' + S/[H0Ac] (5) Analysis of the intercepts of R' vs [HOAc] - 1 plots at 90°C indicated that: I' = ki''[AgOAc] (7) - 74 - 30 0 - 5 1° 15 20 25 30 [H0Ac]-i (M-i) Figure 26. 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 ° K _ I [Ag(I)] M [NaOAc] M [AgOAc] M I ' X 106 M a" 1 s- 1 S' X 106 M2 a - 1 s - 1 ki'x 106 a-i s - i log kx' log j s 1 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 I n i t i a l Conditions: 60°C _ 0.100 M AgC104; 0.135 M NaOAc; 5 atm CO 80, 90, 110PC _ 0.115 M AgC104; O.I95 M NaOAc; 5 atm CO - 76 - while analysis of the slopes indicated that: S' = k2' [Ag+]2[OAc-] + kg' [Ag+][AgOAc][OAc"] (15) The effect of acetate complexing on the reduction rate was studied only at 90°C and thus i t was not possible to calculate the temperature dependence of k 2' and k3' individually. An average activation energy for the acid-dependent reaction was estimated, however, from the slope of a log S' vs l/T plot at a constant degree of complexing. -Such a plot is given in Figure 27 for 80, 90 and 110°C and yields an average activation energy of 17 ± 3 kcal/mole. Also included in Figure 27 is a plot of log I'/[AgOAc] vs l/T at 60, 80, 90 and 110°C which gives the temperature dependence of ki' and yields an activation energy of 15 ± 2 kcal/mole for the acid-independent reaction. The experimental activation energies incorporate the enthalpies of a l l equilibria 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 in 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 dissolution and AgOAc formation are small (see Appendices B and C) as is the heat of formation of HOAc (e.g. ..-0.1 kcal/mole at 25°C (hh)). No information i s available for the heat of formation of the intermediate com plex. If OH- rather than H20 is involved in the formation, the experimental activation energy for the acid-dependent reaction includes the heat of dissociation of H20 which is about Ik kcal/mole i n dilute aqueous solution Combining the experimental activation energies with the "best value" rate parameters at 90°C l i s t e d in Table X (Section III-10), assuming equal activation energies for each of the acid-dependent reactions, gives the f o l  lowing Arrhenius expressions for the temperature dependence of the rate Figure 2 7 . Arrhenius Plots for Acid-Independent and Acid-Dependent Reactions - 78 - parameters over the range investigated: k i = 1 0 7 . 6 ± 1 . 3 expt-(15 ± 2) 103/RT] M-I sec-i (39) k 2 = IO 9- 7 ± 2-° exp[-(17 ± 3) 103/RT] M -2 sec-i (kO) k 3 = 10 1 0- 2 * 2-° exp[-(17 ± 3) 103/RT] M" 2 sec-i (4l) The experimental activation entropies corresponding to the frequency factors in the above expressions are: ASf = -26 - 6 e.u., A S 2 * = -16 * 9 e.u. and AS 3* = -l4 - 9 e.u. (based on a standard state of one mole per l i t r e ) . These experimental activation entropies also incorporate contribu tions from equilibria preceding the rate-determining step. In the acid- independent reaction these contributions may be small and the true activation entropy may correspond to that found from equation 39- This'value is abnormally low for a bimolecular reaction (50) and may reflect steric hindrance for the insertion of a CO molecule into the silver-oxygen bond. Alternatively the reactive silver species may not be undissociated AgOAc molecules but rather one present in very low concentration (e.g. ion pairs). A further explanation may involve a large solvent ordering effect in the formation of the activated complex. Such an effect is not generally expected for reaction between undissociated molecules. - 79 - IV CONCLUSION The CO-reduction of silver perchlorate in acid solution is described by the overall reaction: 2Ag+ + CO + H20 > 2Ag + C0 2 + 2H + In acetate-buffered solution the reaction proceeds homogeneously in the liquid phase by two paral l e l routes, one of which is pH-independent and the other, pH-dependent. The pH-dependent route is favoured by increased pH and is made up of both an acetate-independent and acetate-dependent contribution. The observed kinetics are consistent with the following mechanism: Ag + + 0Ac-^=^ AgOAc (rapid equilibrium) AgOAc + CO — K A > AgJi-OAc (slow) 0 Ag-C-OAc + Ag + + H20 >2Ag + C0 2 + HOAc + H + (fast) 0 Ag + + CO + H20 v. ^ Ag-d-OH + H + (rapid equilibrium) 0 Ag-C-OH + Ag + — 2 A g + C0 2 + H + (slow) AgJLoH + AgOAc — ki^_>2Ag + C0 2 + HOAc (slow) In the pH-dependent reaction the reactivity of silver-acetate com plexes is about a factor of three greater than the reactivity of simple hydrated silver ions. This enhanced reactivity is attributed to stabilization by the basic acetate anion of the proton released in the reduction process. Buffering silver perchlorate solutions with sodium acetate and acetic acid increases the reduction rate by (i) increasing the pH, ( i i ) i n  creasing the reactivity of silver ions in the pH-dependent reaction through complexing and ( i i i ) providing an alternate pH-independent route for reduction. The effect of increased pH i s much greater than the specific effects of - 8 0 - silver-acetate complexing. The CO-reduction of silver(I) in acid solution is consistent with the formation of intermediate complexes by the insertion of a CO molecule between a silver 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 total pressure vs time record (e.g. Figure k} Section III-l) can be converted to the rate of CO consumption in fundamental units (e.g. M sec _i) from a knowledge of (i) the gas-liquid ratio in the reactor, ( i i ) the solubility of CO and C0 2 under the experimental conditions and ( i i i ) the stoichiometric relationship between the consumption of CO and the produc tion of C0 2. The mathematical expression used for this conversion is given by equation A-1: -d[CO] = -dPT f(F + aC0)(F + QCCOgA r A _ - , N dt dt ^ ( aco 2 - aco) J [ > where -d[CO]/dt = rate of CO consumption (M s e c ~ i ) * dPip/dt = slope of a pressure-time record (atm sec- 1) aC0 = solubility of CO (M/atm)* a C 0 2 = solubility of C0 2 (M/atm)* and F = a gas-liquid volume factor (M atm) = (V gAi)(1000/RT) where 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 in terms of l i t r e s of solution measured at room temperature (20-25°C). - 82 - The derivation of equation A-l i s given below. A particular gas (e.g. CO or C0 2) present in the reactor w i l l be distributed between the gas and li q u i d phase in fixed proportion dependent on the ratio of gas and liq u i d volumes and on the solubility of the gas in the liq u i d phase. Assuming that the ideal gas law applies with sufficient accuracy at the temperature and pressure of interest, the concentration of a gas in the gas phase is given by: [gas]g = ^ £ (moles per V i mis of solution) = P x V g 1 0 0 0 (moles per l i t r e of solution) V l RT = F x P where P = parti a l pressure of gas (atm). The concentration of a gas in the li q u i d phase is given by Henry's law: [gas]j_ = a x P (moles per l i t r e of solution) Thus the total concentration of a particular gas present in the reactor is [gas] = P(F + a) From the stoichiometry of the Ag(I)-CO reaction in acid solution, viz : 2Ag(I) + CO + H20 > 2Ag + C0 2 + 2H4" (A-2) the quantity of CO consumed equals the quantity of C0 2 produced. Thus the change in concentration of each gas is given by: X = A[gas] = A P C 0 (F +aC0) = AP C Q (F + «C0 2) - 83 - The observed pressure change during a reduction experiment is the difference between the decrease in CO pressure and the increase in C0 2 pressure, i.e. : APr = AP C 0 - AP C o 2 X _ X (F + QT'CO) (F + a C0 2) y( ( a C0 2 - a CO) (F + ©C!0)(F + OC C0 2) or A [CO] = A P T / ( F + <* Co)(F + <* C0 2) (Qco2 " Qco) Thus the rate of CO consumption is given by: -d[Cpi = -dPr /(F + a C 0 ) ( F + oc C0 2)\ ( A _ 1 } d t . dt V ( a C0 2 " cc CO) The values for the solubility coefficients used in equation A-l were obtained from measurements described in Appendix B and are summarized in Table A-I. TABLE A-I Solubility of CO and C0 2 in Water at 60, 80, 90 and 110°C Temp °C CC c o x 104 M/atm a ™ x 104 M/atm 60 7.0 163-5 80 6.8 122.5 90 6.9 IO5.O 110 7-3 84.0 - 8k - The gas volume in each experiment was estimated by subtraction from the total 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 variation in 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 in 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 in the reactor during each determination. Reactor volumes determined in this way were estimated to be accurate to ± 0„3 mis in a total 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 AgC104, NaOAc, HOAc and 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 + H20 + 2NaOAc 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 AgC104; 0.090 M NaOAc; 0.090 M HOAc; 12.8 atm CO; 90°C I n i t i a l Slope (dP T/dt) = -5.76 x 10 -3 atm sec" 1 after 2-l/k min Gas Solubility at 90°C a C Q = 6.9 x 10"4 M/atm a C 0 2 = 105.0 x IO" 4 M/atm Volume of solution added (Vl) = 101.5 mis Total volume of reactor at 90°C = 120.0 mis Density of H20 = O.9653 g/ml at 90°C; O.9982 g/ml at 20°C -d dt - 8 5 - VK = 120.0 - 101.5 x 0.9982/0.9655 = 15.1 mis F = x , 1 0 0 0 , = i+9.9 x 10-4 M/atm 101.5 82.05 x 363 [CO] = = 7 6 x 1 0-3 / (49.9 + 6.9) IO' 4 (49-9 + 105-0) !Q-4\ i t J % 1 ^ (105.0 - 6.9) 10-4 y = 5I.8 x 10-6 M s e c - i Using this 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] = -0.5 A [NaOAc] = - A [ C 0 ] = 0.5 A [HOAc] = 2-1/4 x 60 x 5I.8 x 10"6 = 0.007 M .The concentrations at the point the rate was measured are therefore estimated to be [Ag(I)] = 0.235 - 0.014 = 0.221 M [NaOAc] = 0.090 - 0.014 = O.O76 M [HOAc] = 0.090 + 0.014 = 0.104 M P c o = 12.8 - 0.007/(49.9 + 6.9) 10-4 = 11.6 atm - 86 - APPENDIX B SOLUBILITY OF CARBON MONOXIDE, CARBON DIOXIDE AND HYDROGEN IN WATER Data for the solubility of carbon monoxide in water are available (52) at atmospheric pressure and temperatures to 100°C while carbon dioxide solubility 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 solubility of carbon monoxide in water were extended to 63 atm and 220°C using the previously described reactor system (Section II-1). The solubility of hydrogen in water was also measured at about 25 atm and temperatures to 225°C. A value for the solubility of C0 2 in water and the salt effect of sodium acetate - acetic acid mixtures were determined at ^0°C and 2.6 atm. Experimental The gas outlet line inside the reactor was bent so that liquid samples could be drawn through the l/l6-±n o.d. capillary tubing and collected over mercury in the ^O-ml water-jacketed burette shown in Figure B-l. S u f f i c i  ent 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 boiling the solution and steam flushing at atmospheric pressure. After 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 line was flushed with a few mis of solution and a 10 to 30-ml liq u i d sample collected in the burette while the shaking mechanism was stopped. The burette valves were closed and the excess gas flushed from solu tion by rapidly raising and lowering the mercury level. After equilibrium had - 87 - from reactor Y l / l 6 " o.d. capillary tubing Water manometer- K 5 S Levelling bottle -J—3-way microvalve S 3 - Thermometer 50-ml burette -Water jacket — Mercury Figure B-l. Measuring Burette System for Gas Solubility Determinations - 88 - been established in the burette the volumes of li q u i d and gas were measured at the temperature of the water jacket and at atmospheric pressure as deter mined by balancing the two legs of a water manometer. Several samples could be drawn from a single charge. Corrections were made for the residual amount of gas remaining in the burette solution and for the vapour pressure of water both in the burette and the reactor. The pressure in 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 purification. Hydrogen was of commercial grade (99-8$ min) supplied by the Canadian Liquid Air Co. and was used without further purification. Degassed d i s t i l l e d water was used in a l l determinations. Results (a) Solubility of CO in Water Data obtained for the solubility of CO in water at 25 atm from room temperature to 220°C are summarized in Table B-I and shown in Figure B-2. Published data available to 100°C at atmospheric pressure (52) when extrapo lated 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 df0 at 100°C. The solubility coefficient of CO in water passes through a minimum between 50° and 100°C, a characteristic exhibited by other permanent gases (5*0- The effect on the CO solubility of pressure to 65 atm at 4l.5°C and to 40 atm at l40°C is summarized in Table B-II. Deviations from Henry's law - 89 - TABLE B-I Solubility* of CO i n Water at 25 Atmospheres Temp °C PCO atm s mls/g S 2 5 mls/g a x 10* M/atm 24 .4 25.2 .567 .561 10.00 2k. k 25.2 .568 .564 10.05 2k.6 25.5 •565 • 554 9.87 2k. 9 25.6 .560 .548 9.76 25.4 25A • 53.2 .524 9.34 26.1 25.6 • 559 .545 9.71 26.2 25A .562 • 552 9.84 26.4 24.9 .528 • 530 9.43 27.1 24.7 .522 .529 9.43 41.5 21.4 • 390 .455 8.10 41.5 27.2 .489 .450 8.02 76.5 25.9 • 387 • 383 6.82 100 24.8 .403 .406 7.24 100 25.O • 394 • 394 7.01 120 25.O .431 .431 7.67 120 25.1 .425 .424 7-55 iko 24.6 A75 .484 8.61 iko 25.4 .480 .472 8.40 i4o 27.0 .522 .484 8.64 iko 27.5 .516 .470 8.36 160 24.5 • 5^ 5 • 555 9-89 160 24.8 .545 .548 9.76 180 24.1 .624 .647 11.5 180 25.1 .638 .635 11.3 200 23.4 .718 .768 13.7 200 24.2 • 731 • 757 13.5 220 25.5 .940 .920 16.4 220 25.9 .928 .898 16.0 220 26.2 .909 .869 15.5 220 26.5 .981 .924 16.5 220 26.6 • 907' .852 15.2 220 27.1 • 983 .905 16.1 * S = mis gas (measured at S.T.P.) per gram H20 S 2 5 = mis gas (S.T.P.) per gram H20 corrected to 25 atm assuming Henry's law- a = moles gas per l i t r e H20 (measured at 20°C) per atmosphere of gas = S 2 5 x I.78I x IO"3 (M/atm) - 90 - Figure B-2- Solubility of CO and H2 in Water from 25 to 225°C - 91 - ' TABLE B-II  Effect of Pressure on Solubility* of CO in Water Temp PCO s a x 104 °C atm mls/g M/atm 41.5 11A .209 8.17 11.7 .216 8.26 'i4.o .251 7-97 21.4 .390 8.10 27,2 .489 8.02 34.1 .601 7.84 41.4 .725 7.82 48 :o .830 7.70 56.9 • 973 7.61 64.3 1.079 7.48 64.3 1.096 7-59 i4o 8.6 .177 9.15 13.6 .274 9.00 19.0 .367 8.59 20.7 .400 8.61 24.6 .475 8.61 25.4 .480 8.40 27.O .522 8.64 27.5 .516 8.36 33-5 .625 8.31 34.3' .635 8.26 39-5 .727 8.21 * S = mis gas (measured at S.T.P.) per gram H20 CC = moles gas per l i t r e H20 (measured at 20°C) per atmosphere of gas = S 2 5 x I.781 x 10-3 (M/atm) - 92 - were observed above about JO atm. Sample Calculations The solubility (S) in 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 in solution in the burette was made on the basis of Henry's law by subtracting the parti a l pressure of the gas in the burette from the experimental reactor gas pressure. At higher temperatures the correction was made by adding the volume of gas remaining in the burette solution, as estimated from room temperature solubility data, to the measured gas volume. The two methods give identical.results and were used in their respective temperature regions merely to f a c i l i t a t e computation of the data. The reproducibility of duplicate determinations was better than the estimated maximum possible error of 3 to ^  for both CO and H 2 measurements. (i) At room temperature: P = atmospheric pressure (mm Hg) P w = vapour pressure of water (mm Hg) at T V = gas volume (mis) at T and P D V w = water volume (mis) at T d = density of water at T S = gas solubility mis (S.T.P.) per g water at the experimental temperature (T a) and pressure (P a) reduced by the part i a l pressure of gas in the burette. S (B-1) where T temperature of water jacket (°C) - 93 - Example; Solubility of Carbon Monoxide in Water at 24.9 ± 0.3°C (T a) V g = 15.1 ± 0.1 mis d = O.9977 g/ml V w = 24.4 ± 0.1 mis P = 762.O ± 0 . 5 mm T = 22.2 * 0.1°C P w = 20.1 ± 0.2 mm Measured CO pressure (P a) = 26.6 ± O.5 atm Corrected CO pressure = 26.6 - 762.0 - 20.1 = 25.6 atm 760 S = 15.1 \ (762.0- 20.l\ f__273__ 24.4 x 0.9977/ I 76O j I22.2 + 273y = O.56O mis CO (S.T.P.) per g water at 25.6 atm CO Possible error = 'AV2 AVW A(P - PV) + — — — + — — — — + AT 0.1 Vw + J4- + P - Pi w T + 275 AP« + — + AT a Ta + 273 0.7 0.1 15.1 24.4 741.9 295.2 + + 0-3 26.6 279.9 100 = 3-1$ ( i i ) At elevated temperature: v g + s x.a .Vw (T + 273)/273\ /P - P^  d«V, w 760 273 \T + 273 (B-2) where T, P, PW, Vg, V w and d have the same meaning as previously and S i = gas solubility mis (S.T.P.) per g water at T and 1 atm S = gas solubility mis (S.T.P.) per g water at experimental temperature (T a) and pressure (Pa) Example: Solubility of carbon monoxide in water at 200.0 ± 0.3°C (T_) V g = 12.5 ± 0.1 mis V w = I5.8 ± 0.1 mis S x = 0.0214 mis (S.T.P.)/g Measured CO pressure (P a) = 23.4 ± 0.7 atm d = O.997O g/ml P = 752.8 t 0.5 mm Lw 24.0 ± 0.2 mm - 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 ) 2^98.2) = O.718 mis (S.T.P.) per g water at 23.4 atm CO r> . M + A Vw + A ( P - Pw) + AT AP a + ATa \ Possible error = -rr^- —— —r — + •———• V V g V w P - P w T + 273 P a T a +" 273 J r = ( o-1 + + Q-7 + Q-i + _PVL + 00 \ 1 0 0  \12.5 15.8 728.8 298.2 23.4 473 J (b) Solubility of H 2 in Water Data obtained for the solubility of H 2 in water at 25 atm from room temperature to 225°C are summarized in Table B-III and shown in Figure B-2. These results are in excellent agreement with other published values (55,56, 57). Figure B-2 indicates that CO and H 2 are equally soluble in water at about 50°C. At higher temperatures H 2 is the more soluble while at lower temperatures the reverse is true. The solubility of H 2 in water obeys Henry's law to about 100 atm (55). (c) Solubility of C0 2 in Water and Acetate Solutions The solubility of C0 2 in water and acetate solutions was determined at 90°C from the observed pressure drop due to C0 2 saturation of a measured quantity of solution in the reactor. The results of these measurements as summarized in Table B-IV indicate that 1:1 solution mixtures of NaOAc and HOAc up to 2 M have l i t t l e effect on the solubility of C0 2. The value for the C0 2 solubility obtained in this way agrees with a value interpolated from published data (58) at 25 atm, i f a 10$ deviation from Henry's law is assumed. Other published data at 15°C (59) indicate that such a deviation is reasonable. TABLE B-III  S o l u b i l i t y * of H g i n Water Temp PC0 S S 2 5 a x 104 °C atm mls/g mls/g M/atm 28.1 24.9 .434 • 435 7.75 • 28.1 25.2 .438 .434 7-73 28.2 25.4 .447 .440 7.84 50 24.7 .414 .420 7.47 50 24.6 .414 .421 7A9 50 24.6 .402 .409 7.28 100 24.3 .440 .452 8.05 100 24.4 .445 .456 8.12 100 24.8 .444 .448 7-97 125 25.3 • 535 .529 9.4i 125 25.9 .547 .528 9.40 125 26.4 • 551 .521 9.28 125 25.8 .542 .526 9-58 150 25.7 .626 .609 10.8 150 25.8 .634 .616 11.0 150 24.8 .615 .619 11.0 175 24.6 .742 • 753 13.4 175 25.1 .743 .739 13.2 200 25.8 .900 .871 15.5 200 25.5 .908 .889 15.8 225 25.3 1.101 1.089 19.4 225 24.9 1.075 1.081 19.3 * S = mis/gas (measured a t S T.P.) per gram H 20 S 2 5 = mis gas (S.T.P.) per gram H2Q corrected to 25 atm assuming Henry's law a = moles gas per l i t r e H20 (measured at 20°C) per atmosphere of gas = S 2 5 x I.78I x 10-3 (M/atm) - 96 - TABLE B-IV Solu b i l i t y * of C02 in Acetate Solutions at 90°C [NaOAc] [HOAc] PC0 2 s Si a x 102 M M atm mls/g mls/g M/atm -- — 2.61 .605 .232 1.02 — — 2.66 .627 .236 1.05 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 Si = mis gas (S.T.P.) per gram solution corrected to one atmosphere assuming Henry1s law 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 coefficient for C02 at 90°C together with other data for the solubility of C02 in water at one atmosphere and tempera tures to 60°C as included in a recent review (53) are summarized in Table B-V and are shown in Figure B-3-- 98 - TABLE B-V Solubility of C02 in Water at One Atmosphere Temp °C C02 Solubility g/100 g * M/atm 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 — 1.05 * C02 solubility in grams C02 per 100 g H20 from reference 53 ** a, = moles gas per l i t r e H20 (measured at 20°C) per atmosphere of gas = 0.227(g C02/l00 g H20) - 99 - Figure B-3. Solubility of C0 2 in Water at One Atmosphere - 100 APPENDIX C SILVER-ACETATE COMPLEXING FROM E.M.F. MEASUREMENTS Equilibria i n silver acetate solutions have been extensively studied at room temperature by solubility measurements (1+5,60,61,62,63,64,65) and by E.M.F. measurements (39>^>66»). The effect of temperatures to 90°C on sil 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 cells of the type: Ag AgC104 NaOAc NaC104 8 M AgC104 NaC104 Ag Experimental A l l c e l l solutions contained 0.0104 M AgC104 and appropriate amounts of either NaOAc or NaC104 i n equal concentrations to maintain constant ionic strength in both sides of each c e l l . Silver f o i l electrodes, approximately 4 x 4-cm with a 5- c m lug, were conditioned by electrolyzing approximately M AgC104 solution acidified with HC104. 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 silver deposit. Current densities of about 0.05 amps/cm2 gave suitable deposits within one to two minutes. The electrolysis conditions were not c r i t i c a l . The electrodes were then formed into cylindrical shapes, washed and stored in acidified d i s t i l l e d water u n t i l required. The experimental c e l l , depicted in 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 NaC104 Salt Bridge Figure C - l . Experimental C e l l for E.M.F. Measurements (approximately half-scale) - 102 - pyrex tubing and had a total volume of about 75 mis. A capillary tube for 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 salt-bridge from each compartment and solution contact was made in the plugs. The c e l l was about 80$ immersed in a. stirred, 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 Controller (Model 71) with a stainless steel-clad thermister probe (No. k-06) . E.M.F. measurements were made with a high precision Leeds and Northrup potentiometer (No. 7552). The general procedure consisted of pipetting sufficient 8 M N a C 1 0 4 (previously prepared with degassed d i s t i l l e d water and stored in a stoppered flask) to bring the solution level in the salt bridge tubing to within about two cm of the bottom of each solution compartment. After the f i l t e r paper plugs were inserted approximately 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 in the glycol bath at room temperature and the solution compartments were flushed with a slow flow of helium for about five minutes. After flushing the gas outlet lines were clamped and a slight positive pressure of helium maintained over the solu tions . At room temperature the E.M.F. became constant within 2 to k hours and showed l i t t l e deviation for periods up to 12 hours. After the room temperature value had been determined the glycol bath was heated and the E.M.F. measured at successively higher temperatures, sufficient time being allowed to attain equilibrium as evidenced by constant cell-voltage read ings. The E.M.F. of cells 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 in li q u i d junction potentials is negligible. 2. Activity coefficients of Ag + are equal in a l l solutions of equal ionic strength. 3. AgC104, WaC104 and WaOAc are completely dissociated. 4. The average silver-acetate complex is represented by the formula AgOAc, On this 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 0.0104 (c-l) -nF 1 0 g [Ag+] k where E = c e l l E.M.F. (volts) R = universal gas constant (I.987 cal mole _i deg-i) F = Faraday constant (23.06 kcal/equivalent) T = absolute temperature (°K) n = number of volt equivalents (= unity) [Ag +] = silver ion concentration in the acetate solution compartment (M) 0.0104 = silver ion concentration in the perchlorate solution compartment (M) The concentration ratio (K a) for the formation of AgOAc is given by: K, a LAg+JlOAc"] The results of measurements made on six cells are summarized in Table C-I. K a estimated using equations C-l and C-2 is essentially independent of ionic strength between 0.1 and 0.9 and temperatures to 90°C and has an average value of 3-7 * 0.7 M"1. - 10k - TABLE C-I Effect of Temperature on Silver-Acetate Complexing Using 0 .0104 M AgC104 C e l l [NaOAc] Temp EMF [Ag +] No M °C mv M M-l 1 .10 23 9 . 0 ± 1 .0 • 0073 4 . 4 2 . 2 0 24 1 2 . 0 ± 1 .0 .OO65 3-1 . 2 0 25 1 2 . 5 * 1 .0 .0064 3-2 . 2 0 50 1 5 . 0 ± 1 .0 .0061 3 . 6 . 2 0 90 1 7 . 0 ± 1.5 .0060 3-7 3 .50 25 2 7 . 0 ± 1 .0 .0036 3 . 8 . 5 0 35 2 9 . 0 ± 1.5 • 0035 4 . 0 •50 5 0 3 1 . 5 * 1-5 .0034 4 . 2 4 . 5 0 23 2 4 . 0 ± 1 .0 .0041 3 .1 5 . 5 0 24 2 3 . 5 ± 1 .0 .0042 3 . 0 • 50 25 2 6 . 0 ± 1 .0 .0038 3 . 5 . 5 0 90 33-5 * 1.5 • 0037 3-7 6 • 90 24 3 5 . 0 ± 1 .0 .0027 3 . 2 . 9 0 35 3 9 . 0 ± 1 .0 .0024 3 . 7 . 9 0 50 4 5 . 0 ± 1.5 .0021 4 . 4 • 90 90 5 0 . 0 ± 2 . 0 .0021 4 . 4 K = 0.0104 - [Ag+]  a [Ag+]([NaOAc]-[Ag+]) - 105 - APPENDIX D SUMMARY OF SELECTED EXPERIMENTAL DATA FOR THE REDUCTION OF SILVER(I) SOLUTIONS BY CARBON MONOXIDE I. EFFECT OF CO PRESSURE AT 90°C (Figure 5) (a) I n i t i a l Conditions: 0.115 MAgC104; O.I95 M NaOAc; O.O78 M HOAc Expt Pco R* x 106 R'**x IO 6 No atm M s - i M a - i s-l 234 1.67 7.2 4.31 233 1.69 5.0 2.96 231 3.10 13.6 4.39 232 4.17 • 20.0 4.79 2l4 5-17 27.1 5.24 230 5.22 23.7 4". 35 207 5.24 23.7 4.34 236 9.28 42.0 4.57 235 12.8 43.3 3.38 237 13.3 54.4 4.09 238 29.8 132. 4.43 (b) I n i t i a l Conditions: 0.115 M AgC104; 0.045 M NaOAc; O.766 M HOAc Expt No PCO atm R* x 106 M. s - i R***x 106 M a-i s-i 162 11.4 5.8 .509 163 19.4 10.0 .515 165? 25.6 12.7 .496 l 6 4 b 27.5 14.2 .516 161 27.7 14.4 .520 a - Reactor thoroughly cleaned with HNO3 prior to charging. b - Contained O.lg Ag precipitated in previous experiment plus 2.0 g silver sponge obtained from Consolidated Mining and Smelting Co. Ltd. * R = -d[C0]/dt ** R' = R/P C 0 - 106 - II. EFFECT OF ACETIC ACID AT 90° C (Figure 6) I n i t i a l Conditions: 0 .115 M AgC104; O . I95 M NaOAc Expt [HOAc] [HOAc] - 1 PCO R' x 1 0 6 No M M"1 atm M a - 1 s."1 209 . 5 8 5 1 . 7 5.17 1 . 5 2 212 . 5 8 5 1-7 5.24 1.84 208 .160 6 . 3 5 . 3 1 3 . 6 1 210 . 1 6 0 6 . 3 5 . 3 1 3 . 8 9 238 .0978 1 0 . 2 2 9 . 8 4 . 4 3 237 .O898 1 1 . 1 1 3 . 3 4 . 0 9 235 .0886 1 1 . 3 1 2 . 8 3 . 3 9 236 . 0 8 6 8 1 1 . 5 9 . 2 8 4 . 5 3 230 .0832 1 2 . 0 5 . 2 2 4 . 5 4 207 . 0 8 3 0 1 2 . 0 5.24 4 . 3 4 214 .0828 1 2 . 1 5 . 1 7 5.24 232 .0818 1 2 . 2 4 . 1 7 4 - 7 9 231 .0805 12.4 3 . 1 0 4 . 3 9 233 .0786 1 2 . 7 1 . 6 9 2 . 9 6 234 .0786 1 2 . 7 1 . 6 7 4 . 3 1 213 .0624 1 6 . 0 5 . 2 4 6 . 8 8 206 .0460 2 1 . 7 5 . 1 7 7-73 211 .0334 2 9 . 9 5.17 1 0 . 1 III. EFFECT OF SILVER-ACETATE COMPLEXING AT 90°C (Figures 6 , 7 , 8 , - TablelV) A. I n i t i a l Conditions: O.O5O M AgC104j 0.045 M NaOAc Expt Ug(I)] [NaOAc] [HOAc]"1 PCO R' x 1 0 p No M M M-i atm M a - 1 s _ 1 A-180 .048 .043 4 . 4 2 8 . 0 .152 A-181 . 0 4 2 .037 1 8 . 9 2 7 . 9 • 559 A-191 .046 .041 1 0 . 6 2 8 . 1 .418 A-192 . 0 4 6 .041 5 A 2 8 . 1 .304 A-193 . 0 4 9 .043 1 . 2 2 7 . 8 .106 Average: . 0 4 8 .041 1 1) I' = 0 . 1 5 X 1 0 - 6 M atm - 1 sec - 1 S' = 0 .017 x 1 0 - 6 M2 atm - 1 sec" 1 - 107 - B. I n i t i a l Conditions: 0.100 M AgC104; 0.045 M NaOAc Expt [Ag(I)] [NaOAc] [HOAc]"1 PC0 R' x 106 No M M M-i atm M a" 1 .'s"1 B-lk2 • 097 .0k2 4.39 27.2 .420 -B-I4-9 .096 .01+1 4.36 27.6 .368 B-152 .O96 .041 1.30 27.6 .368 B-155 .095 .040 1.30 27.7 •.347 Average: .096 .04l I' = 0.31 X 10 " 6 M atm-1- sec - 1 S' = 0.026 x IO - 6 ;M2 atm' " 1 s e c - 1 C. I n i t i a l Conditions: 0.047 M AgC104; 0.090 M NaOAc Expt No [Ag(I)] M [NaOAc] M [HOAc] _ 1 pco atm R' x 10 s M a" 1 s _ 1 C-308 C-307 C-306 .046 .045 .044 .090 .089 .087 .64 4.41 10.7 12.4 12.2 12.1 .230 .504 .853 Average: .045 .O89 I' = 0.22 x IO" 6 Matm"1 secT 1 S' = O.O58 x IO" 6 M2 atm"1 sec" 1 D. I n i t i a l Conditions: 0.117 M AgC104; 0.045 M NaOAc Expt No [Ag(I)3 M [NaOAc] M [HOAc] _ 1 M-1 PC0 atm R1 x 10 s M a - 1 s - 1 D-299 D-298 D-297 .116 .115 .114 .044 .043 .042 .63 4.41 10.7' 12.1 12.1 12.4 • 552 .619 1.00" Average: .115 .043 I' = 0.32 x IO" 6 M atm-i sec"! S' = 0.066 x 10-6 M 2atm- 1 s e c - 1 - 108 - E. I n i t i a l Conditions: 0.100 M AgC104; 0.090 M NaOAc Expt [Ag(I)] [NaOAc] [H0Ac]-i PCO R'x 106 No M M M-i atm M a-i s" 1 E-159 .094 .084 5.38 26.6 .804 E-140 .094 .084 5.38 26.7 • 755 E - l 4 l .089 .079 5.24 26.5 .711 E-151 .095 .O85 2.20 27.4 .566 Average: .090 .080 I'' = 0.40 x IO - 6 Matm"1 sec" 1 S' = 0.074 x IO" 6 M2 atm"1 sec" 1 F. I n i t i a l Conditions: 0.047 M AgC104; 0.180 M NaOAc Expt [Ag(I)] [NaOAc] [HOAc]"1 Pco R'x 106 No M M M"1 atm M a- 1 s- 1 F-505 .045 .179 • 63 12.2 .467 F-504 .044 • 177 4.39 12.3 • 799 F-505 .043 .176 10.6 12.2 1.27 Average: .044 .178 ] :' = 0.43 x IO" 6 M atm"1 sec" 1 S' = 0.081 x 10-6 M2 atm"1 sec _ i G. I n i t i a l Conditions: O.O5O M AgC104; 0.225 M NaOAc Expt [Ag(I)] [NaOAc] [HOAc]-1 PC0 R'x 106 No M M M"1 atm M a - i s - i G-155 .033 .208 16.1 27.4 1.73 G-182 .o4i .215 4.3 28.2 .701 G-I89 .044 .219 2.1 27.7 • 5^ 9 G-190 .043 .218 1.6 28.4 .602 Average: .042 .215 I' = 0.43 x IO" 6 M atm - 1 sec-i S' = 0.077 x 10"6 M2 atm"1 sec" 1 - 109 - H. I n i t i a l Conditions: 0.235 M AgC104; 0.045 M NaOAc Expt [Ag(I)] [NaOAc] [HOAc]-1 PCO R'x 106 No M M M-i atm M a- 1 s" 1 H-293 .233 .0U3 • 63 11.6 •56 H-292 .231 .041 4.37 11.8 1.21 H-291 .229 .039 10.4 11.5 1.83 Average: .230 .041 I = O.5O x IO" 6 M atm"1 sec" 1 S' = 0.133 x IO"6 M2 atm"1 -1 sec I. I n i t i a l Conditions: 0.117 M AgC104; O.O9OM NaOAc Expt [Ag(I)] [NaOAc] [HOAc]"1 pC0 R'x 1 0 S No M M M"1 atm M a" 1 s- 1 1 - 2 9 6 . 1 1 5 . 0 8 8 • 63 1 1 . 8 . 5 6 1 - 2 9 5 . 1 1 4 . 0 8 7 4 . 3 7 1 2 . 0 1 . 1 1 1 - 2 9 4 . 1 1 0 .085 1 0 . 2 1 2 . 2 2 . 3 2 Average: . 1 1 4 .O87 I = 0 . 4 4 x 1 0 - 6 M atm"1 sec"i S = O . I78 X 1 0 - 6 M2 a t m - i sec" 1 J. I n i t i a l Conditions: 0.235 M AgC104; 0.090 M NaOAc Expt [Ag(I)] [NaOAc] [HOAc] _ 1 PC0 R'x 1 0 6 No M M M"1 atm M a - 1 s" 1 J - 2 9 0 . 2 3 1 .087 . 6 2 1 1 . 6 1 . 0 2 J - 2 8 3 . 2 3 0 .O85 1 . 2 6 1 1 . 9 1 . 3 4 J - 2 8 9 .227 .082 4 . 3 0 1 1 . 9 2 . 4 3 J - 2 8 5 . 2 2 1 .O76 9 - 6 0 1 1 . 6 4 . 4 7 Average: . 2 2 5 .082 I = O .78 x 1 0" s M atm"1 s e c - 1 S = 0 . 3 8 4 x 1 0 " 6 M2 atm" •1 s e c - 1 - 110 - K. I n i t i a l Conditions: 0.117 M AgC104; 0.180 M NaOAc Expt [Ag(I)] [NaOAc] [HOAc]"1 PC0 R' x 10s No M M M-I atm M a - 1 s" 1 K-302 .114 • 177 .63 12.1 .92 K-301 .111 .174 4.32 12.0 2.06 K-300 .105 .168 9.76 11-7 3.87 Average: .110 • 173 I' = 0.70 x IO" 6 M.atm-1 s e c - 1 S1 = 0.322 x 10-6 M2 atm"1 s e c - 1 L. I n i t i a l Conditions: 0.115 M AgC104; O.I95 M NaOAc Expt [Ag(I)] [NaOAc] [HOAc]"1 pC0 R'x 106 No M M M-i atm M a - 1 s - i L-209 .113 • 193 1.7 5.17 1.52 L-212, .113 .193 1-7 5.24 1.84 L-208 .111 .191 6.3 5-31 3.61 L-210 .111 .191 6.3 5-31 3.89 L-238 .096 • 175 10.2 29.8 4.43 L-237 .103 .183 11.1 13.3 4.09 L-235 .104 .184 11.3 12.8 3.39 L-236 .106 .186 11.5 9.28 4.53 L-230 .110 .190 12.0 5.22 4.54 L-207 .110 .190 12.0 5.24 4.34 L-214 .110 .190 12.1 5.17 5.24 L-232 . i l l .191 12.2 4.17 M 9 L-231 .112 .192 12.4 3.10 4.39 L-233 .114 .194 12.7 I.69 2.96 L-234 .114 .194 12.7 1.67 4.31 L-213 .108 .188 16.0 5.24 6.88 L-206 .108 .188 21.7 5-17 7-73 L-211 .105 .185 29.9 5.17 10.1 Average: .110 .190 ] :' = 0.90 x 10"6 M atm"1 s e c - 1 S' - 0.32 x 10-6 MS atm"1 s e c - 1 - I l l - M. I n i t i a l Conditions: 0.04-7 M AgC104; 0.676 M NaOAc Expt [Ag(I)] [NaOAc] [HOAc]"1 pC0 R'x 1 0 S No M M atm M a" 1 s - i M-511 .043 .672 • 63 12.1 1.04 M-310 • 039 .670 4.32 12.4 1.94 M-309 .036 .665 9.87 12.4 3.34 Average: .040 .670 I ' = O.87 x 10" 6 M atm - 1 sec" 1 S 1 = 0.249 x IO"6 M2 atm" 1 sec" 1 N. I n i t i a l Conditions: O.O58 M AgC104; O.556 M NaOAc Expt [Ag(I)] [NaOAc] [HOAc]"1 PCO R"x 106 No M M M"1 atm M,a-i,s-i N-254 .057 •555 •51 5-51 1.00 N-255 .O56 • 55^ 2.57 5.44 I . 5 8 N-256 • 055 • 553 8.78 5-57 5.24 Average: .O56 • 55^ I = O.85 x 10 " 6 M atm - 1 sec" 1 S = 0.275 x 10~ 6 M2 atm - 1 s e c - 1 0. I n i t i a l Conditions: O.O58 M AgCIO 4 ; O.778 M NaOAc Expt [Ag(I)] [NaOAc] [HOAc]-1 PC0 R'x 1 0 S No M M M-i atm M a" 1 s- 1 0-250 .057 • 777 .52 5.24 1.24 0-252 • 057 • 777 1.03 5-57 I . 5 8 0-251 .O56 .776 2.57 5.51 1.95 0-253 .O56 .776 5.10 5.51 2.92 0-249 .055 • 775 8.76 5.37 3 . 9 2 Average: .O56 .776 I = 1.10 x 10" 6 M atm"1 sec-i S = 0.540 X 1 0 _ s M2 atm" 1 sec" 1 - 112 - IV. REACTION RATES USED IN EXTRAPOLATION TO ZERO ACETATE AT 90°C (Figures 20, 21, 22; Table VI).  P. I n i t i a l Conditions: 0.050 M AgC104; [HOAc]/[NaOAc] = 1 R , [HOAci 1 0 6 M a - 1 s" 1 Expt No [Ag(I)] M [NaOAc] M [HOAc] M rC0 atm [AgOAT] M [OAc"] M R'xlOs M a - 1 s" 1 •184 -182 -I85 -183 -186 P-181 .OH-3 .04l .043 .04l .044 .042 .262 .215 .172 .126 .083 • 037 .278 .235 .188 .144 .097 .053 28.2 28.2 28.2 27.9 28.0 27.9 .0201 .0173 .0156 .0123 .0094 .0046 .245 .198 .156 .114 .074 .032 • 777 .702 • 759 .671 .650 .476 10.8 9-53 9.15 7.85 6.71 5.54 p.[HOAc] LAgOAc J = 4.6 x IO" 6 M atm - 1 sec~i Q. I n i t i a l Conditions: 0.100 M AgC104; [HOAc]/[NaOAc] = 1 Expt No [Ag(I)] M [NaOAc] M [HOAc] M PC0 atm [OAc"] .M R' xlO 6 M a-i s - i R.[HOAc] xlOe [AgOAc] M [AgOAc] M a-i s - i Q-137 Q-138 Q-136 .082 .O85 .O85 .207 .165 .120 .243 .195 .150 26.4 26.4 26.5 .0322 .0285 .0225 • 175 .137 .098 1.73 1.53 I.58 13.1 10.5 10.5 ( R , T A W T ) 0 = 6.8 x IO- M atm- s e c - R. I n i t i a l Conditions: 0.123 M AgC104; [HOAc]/[NaOAc] = 1 [HOAc] 1 Q 6 n [AgOAcJ M a- i s - i Expt No [Ag(I)] M [NaOAc] M [HOAc] M pco atm [AgOAc] M [OAc-] M R' xlO© M a-i s - i R-287 R-288 .115 .118 .O87 .017 .102 .029 12.7 12.8 .0223 .0051 .065 .012 2.48 1-75 11.4 10.0 ( ^..[HOAc] \ K [AgOAc J J o = 9.6 x IO- 6 M atm - 1 s e c - 1 - 113 - S. I n i t i a l Conditions: 0.235 M AgC104; [HOAc]/[NaOAc] = 1 Expt No [Ag(I)] M [NaOAc] M [HOAc] M PC0 atm [OAc-] M R'xlO6 M a- i s'"1 ,[HOAc] 6 [AgOAc ] M LAgOTSff 1 0 M a~i s _ : L S-285 S-286 .225 .226 .076 .013 .104 .032 11.6 11-9 .0313 .OO59 .045 .007 4.47 2.59 » 14.9 i4.o ( ( R ' S ) 0 = 15.7 x l O - e M atm- s e c - . *„ T. Initial'Conditions: 0.100 M AgC104; [HOAc]/[NaOAc] = 5 Expt No [Ag(I)] M . [NaOAc] M [HOAc] M •Pco atm [OAc"] M R'xlO6 M a - 1 s - i , iHOAc] [AgOAc] M R LAgOAcJ X l 0 6 M a - 1 s-1 T-150 T-154- T-151 T-lV? T-l42 .094 .095 .095 .096 .097 .129 .108 .O85 .041 .042 .681 .569 .455 .229 .228 27.5 27.5 27.4 27.6 27.2 .0259 .0228 .0187 .0099 .0102 .103 .O85 .066 .031 .032 .611 •575 • 555 .368 .419 16.0 14.4 13.5 - 8.5 9-4 ^ R' j ^  = 6.0 x IO"6 M atm"1 sec" 1 U. I n i t i a l Conditions: 0.064 M AgC104j [HOAc]/[NaOAc] = 8.7 •or [HOAc ] ^ c M a- i s-i- Expt ' No [Ag(I)] M [NaOAc] M [HOAc] M PC0 atm [AgOAc] M [OAc-] M R' xlO© M a - 1 s- 1 U-276 U-277 .065 .065 .090 .065 .785 •595 4.83 4.90 .0139 .0078 .076 .038 • 570 .55^  52.2 18.0 ( „,.[H0Ac] N\ K [AgOAc J J q = 4.4 x IO" 6 M atm - 1 s e c - 1 -114 - V. I n i t i a l Conditions: 0 . 1 3 0 M A g C 1 0 4 ; [HOAc]/[NaOAc] = 8.7 Expt No [Ag(I)] M [NaOAc] M [HOAc] • M PC0 atm [OAc-] M R1 xlO 6 M a" 1 s" 1 [HOAc] ^ [AgOAc] M R [AgOAcJ M a"i s " i V-281 V-280 .126 .127 .146 .022 1.314 .221 12.1 12.3 .0362 .0066 .110 .015 1.19 .649 43.3 21.8 ( R' rlP^ ] i 1 - 18.0 x 10-6 M atm"1 sec" 1 \ [AgOAc J J 0 W. I n i t i a l Conditions: 0.163 M AgC104; [HOAc]/[NaOAc]= 8.7 Expt No Ug(I)] M [NaOAc] M [HOAc] M Pco atm [AgOAc] M [OAc-] M R' xlO 6 M a-i s - i M a-i s - i W-278 W-279 .161 .162 • 073 .024 .655 .219 4.94 4.68 .0246 .OO87 .049 .015 1.13 •.908 30.1 22.9 / R, [HOAc] \ _ nq.u x 10-6 M a t T n - i R e n - i ^ [AgOAcV Q X. I n i t i a l Conditions: 0.235 M AgC104; [HOAc]/[NaOAc]= 8.7 Expt No [Ag(I)] M [NaOAc] M [HOAc] M Pco atm [OAc-] M R"xl06 M a-i s - i p,-[HOAc] , r l A B [AgOAc] M R LAgOA-c]XlCP M a-i s - i X-283 X-282 X-284 .230 .232 .232 .085 .042 • 053 • 793 • 399 .319 11.9 11.8 12.1 .0355 .0184 .0146 .050 .024 .018 1.34 .903 .890 29.9 19.6 19.4 ( E ' & ) 0 12.8 x 10~6 M atm-1 sec - 1 - 115 - Y. I n i t i a l Conditions: 0.270 M AgC104; [HOAc]/[NaOAc] = 8.7 •p R' xlO 6 *'[HOAc] X L 0 6 Expt [Ag(I)] [NaOAc] [HOAc] r co [AgOAc] [OAc-] [AgOAc] No M M M atm M M M.a"1 .s-i M s - i a-i Y-270 .267 .092 .829 5.05 .0419 .050 2.27 45.1 Y-274 .268 .046 .415 4.88 .0218 .024 1.55 29.5 Y-273 .269 .018 .167 4.92 .0087 .009 1.18 22.8 Y-272 .268 .0083 .084 4.97 .0041 .004 .969 19.8 ( R, [HOAc] -\ -= 18.4 X 10 " 6 M atm - 1 sec-.1 V [AgOAc J / Q V. P H CHANGES DURING REDUCTION OF UNBUFFERED AgC104 SOLUTIONS AT 90°C AND 53 ATM CO (Figures 23 and 24; Tables VII:and-VIII)  Experiment No. 262 Time sec [Ag+] M pH [H +] x 103 M log t 0 .0530 ~ 5.2 ^0.007 180 3.15 0.7 2.26 390 2.92 1.2 2.59 78O 2.71 2.0 2.89 1,800 2.55 2.8 3.26 3,600 2.32 4.7 3.56 6,300 2.27 5-4 3.80 9,000 2.14 7.2 3.96 12,780 .0458 2.09 8.1 4.11 The pH is defined in terms of hydrogen ion concentration rather than activity since the pH-meter was calibrated against mixtures of standard HC104 and AgC104 solutions. - 116 - Experiment No. 260 Time sec [Ag +] M PH [H+] x 103 M log t 0 .104 ~ 5.0 ~ 0.01 120 3.05 0.9 2.08 240 2.77 1-7 2.38 480 2.61 2-5 2.68 990 2.41 3-9 3.00 1,920 2.26 5-5 3.28 4,740 2.05 8.9 3.68 10,000 .090 1.87 13.5 4.00 Experiment No. 26l Time sec [Ag +] M PH [H +] x 103 M log t 0 .211 ~ 4.5 ^ 0.03 180 2.62 2.4 2.26 300 2.51 3-1 2.48 525 2.30 5.0 2.72 1,020 2.14 7.3 3.01 1,810 1.96 11.0 3.26 3,600 1.80 15.9 3.56 6,310 1.67 22.4 3.80 9,000 1.63 23.4 3.96 11,700 .179 1.47 33-9 4.07 Experiment No. 265 Time sec [Ag+] M pH [H+] x 103 M log t 0 1.06 ^4.0 —' 0.1 180 1.67 21 2.26 360 1-57 27 2.56 660 1.43 37 2.82 1,140 1.50 50 5.06 2,040 1.10 79 3.31 3,840 0.95 112 5.58 5,640 0.90 126 3.75 8,340 0.83 " 148 5.92 11,040 •85 0.69 204 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: 0.144 M AgC104; 0.195 M NaOAc; O.I95 M HOAc Expt [Ag(I)] [NaOAc] [HOAc] PCO R' x 106 No M M M atm M a-i s - i 265 .142 .192 .197 • 5.16 3-75 264a .141 .191 .198 5.10 4.09 266b ' .141 .191 .197 5.04 4.02 267c .142 .192 • 197 4.78 3.87 268d .142 .192 •197 5.04 3-75 a - Air evacuated from solution in reactor prior to heat-up b - Added C0 2 to 5.7 psi prior to CO addition c - Contained 3-3 g fine 316 S.S. fi l i n g s d - Contained 1.0 g Ag precipitated during previous experiments plus ).0 g Ag sponge (b) I n i t i a l Conditions: 0.111 M HOAc Expt [Ag(I)] [NaOAc] [HOAc] PC0 R'x 106 No M M M atm M a-i s - i 239 .068 .220 • 113 5.22 2.22 • 240 .068 • 330 .114 5.25 2.97 24l .070 .044 .112 5.45 .785 242 .069 .109 • 113 5.17 1-57 243 .133 .043 .113 5.11 1-57 244 .132 .108 .115 5.15 3.16 245 .129 .216 .118 5.11 6.34 246 .278 .104 .118 5.10 6.53 247 .044 .110 .112 5.27 • 955 248 .069 .110 .113 5-30 1.43 (c) I n i t i a l Conditions: 0.115 M AgC104; 0.045 M NaOAc; O.766 M HOAc Expt [Ag(l)] [NaOAc] [HOAc] Pco R'x 106 No M M M atm M a-i s - i 162 .114 .044 .767 11.4 .509 163 .110 .040 .771 19.4 .515 I 6 5 .111 .041 .770 25.6 .496 164 .110 .040 • 771 27.5 .516 161 .110 .040 • 771 27.7 .520 - 118 - VII. EFFECT OF TEMPERATURE. (Figure 26; Table XI) (a) I n i t i a l Conditions: 60°C, 0.100 M AgC104; 0.135 M NaOAc Expt [Ag(I)] [NaOAc] [H0Ac]-i Pco R'x 10s Wo M M M-I atm M a-i s - i 223 .100 ' .135 1.3 5.58 .177 222 .099 .134 11.0 5.51 .468 221 .099 . .134 21.5 5.30 .720 Average: .099 .134 I' = 0.14 x 10-6 M atm-i sec-i s' = 0.029 x 10-6 M2 atm _ 1 s e c - 1 (b) I n i t i a l Conditions : 80°C; 0.115 M AgC104; O.I95 M NaOAc Expt [Ag(I)] [NaOAc] [HOAc]-1 Pco R'x 106 No M M M-i atm M a-i s - i 220 • 113 • 193 1-7 5.4l 0.93 219 .112 .192 6.3 5.47 1.79 218 .111 .192 12.4 5.48 2.64 217 .110 .191 17.2 5.38 3.43 216 .110 .190 25.2 5.45 3-95 215 .107 .187 32.6 5.4.1 6.76 Average: .110 .191 I' = 0.64 x 10 -6 M atm-i.sec-i S' = O.I65 x IO - 6 M2 atm"1 sec" 1 (c) I n i t i a l Conditions: 110°C; 0.115 M AgC104; O.I95 M NaOAc Expt [Ag(I)] [NaOAc] [H0Ac]-i Pco R' x 106 No M M M-I atm M.a--L S" 1 205 .112 .192 1-7 4.56 5 .01 199 .109 .189 6.2 4.56 9 • 72 204 .105 .186 11.6 4.56 15 .6 203 .105 .183 20.0 4.49 20 • 9 202 .105 .183 20.0 4.49 20 • 9 197 .102 .182 19.7 4.65 21 .6 201 .101 .182 27.6 4.49 23 • 7 200 .100 .181 27.2 4.53 24 .6 198 .099 .180 26.1 4.49 27 .6 Average: .105 . I 8 5 I' = 5.2 X 10 -6 M atm"1 s e c - 1 S' = 1.04 x 10 -6 M2 atm-1 s e c - 1 - 119 - APPENDIX E THERMODYNAMICS OF THE OXIDATION OF CO, H2, HCOOH AND HCOO" IN AQUEOUS SOLUTION AT 25°C (Figure 1) The thermodynamics of the reactions considered in Figure 1 (Section 1-2) are summarized in terms of the free energy data (10) in 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 for gaseous species. The Inter national or Stockholm Convention ( 9 ) is used for the sign of electrode poten t i a l s . TABLE E-I Standard Free Energy at 25°C (kcal/mole) C0(g) -32.808 C0 2(g) -94.260 C03=(aq) -126.22 HC03"(aq) -140.31 HCOO-(aq) -80.0 HCOOH(aq) -85.I H+(aq) 0.0 H 2(g) 0.0 HaO(l) -56.690 AG° = -nFE° = -2.30 RT log K (E-l) E = E° - 2-30 RT l o g Q / E _ 2 ) nF where AG° = standard Gibbs free energy for the electrode reaction (kcal/mole) E° = standard electrode potential (volt/mole) E = electrode potential (volt/mole) n = number of volt equivalents - 120 - F = F a r a d a y c o n s t a n t (23.06 k c a l / e q u i v a l e n t ) R = u n i v e r s a l gas c o n s t a n t (I.987 c a l m o l e - 1 deg- 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 c o n s t a n t Q = a c t i v i t y q u o t i e n t • = 0.059 ( v o l t s / e q u i v a l e n t ) a t 25°C -F (1) C 0 2 ( g ) + 2H +(aq) + 2e = C0(g) + H 20(1) A G 0 = 4.762 k c a l / m o l e ; E° = 0.103 v o l t s E = -0.103 - 0.059pH - 0.030 log ( a c o / a c o 2 ) (2) HC0 3-(aq) .+ 3H +(aq) + 2e = 00(g) + 2H 20(1) A G 0 = -5.88 k c a l / m o l e ; E° = 0.128 v o l t s E = 0.128 - 0.089PH - 0.030 l o g ( a c o / a H C o 3 - ) (3) C0 3=(aq) + 4 H +(aq) + 2e = 00(g) + 2H 20(1) 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 l o g ( 0 0 0 / 8 ^ 0 3 = ) (4) 2H +(aq) + 2e = H 2 ( g ) AG° =0; E° = 0 E, = -0.059pH - 0.030 l o g ( a H 2 ) (5) C 0 2 ( g ) + 2H +(aq) = 2e = HCOOH(aq) AG° =9.16 k c a l / m o l e ; E° = -O.I99 v o l t s E: = -0.199 - 0.059PH - 0.030 l o g ( a H C 0 0 H / a c 0 2 (6) C 0 2 ( g ) + H + ( a q ) + 2e = HC00"(aq) AG" = 14.26 k c a l / m o l e ; E° = -O.3O9 v o l t s E = -0.309 - 0.030PH - 0.030 l o g ( a H C O o - / a c o 2 - 121 - HC0 3 _(aq) + 2H+(aq) + 2e = HCOO-(aq.) + H20(1) AG" =3.62 kcal/mole; E° = -O.O79 volts E = _o„079 - 0.059PH - 0.030 log(a H C 0 0-/a H C 0 3-) C0 3 = (aq.) + 3H+(aq.) + 2e = HC00_(aq) + H20(1) AG° = -10.47 kcal/mole; E° = 0.227 volts E = 0.227 - 0.089PH-0.030 log(aHCOO-/aco3") HCOOH(aq) = H+(aq.) + HC00_(aq) AG° =5.1 kcal/mole; log K = -3.7 pH = 3.7 + log (aHC00"/aHC00H) C0 2(g) + H20(1) = H+(aq) +, HC03"(aq) AG° = 10.64 kcal/mole; log K = -7.8 pH = 7-8 + log(a H C 0 3-/a C 0 2) HC03-(ao_) = H+(aq) + C0 3 = (aq) AG° = 14.09 kcal/mole; log K = -10.3 pH = 10.3 + log^ogz/ancog-) - 1 2 2 APPENDIX F NUMERICAL INTEGRATION OF EXPERIMENTAL RATE LAW (Figure 2 5 , Section 1 1 1 - 1 0 ) The rate of CO-reduction of silver perchlorate in sodium acetate - acetic acid buffered solution is given by equation 1 9 : R' = k l ( [AgOAc ] + k2' [ A g + P - [ | | l j + kg' [Ag+] [AgOAc-]{ j jg jgj . ( 1 9 ) A method for numerically integrating equation 1 9 is outlined i n the present section. From equation 1 9 , -d[CO]_ = p c o R. (M sec-i) (F-l) dt From Appendix A* : -d[CO] = dPr /(F + a C 0 ) ( F + a c o )\ 1 (M sec-i) (A-1) dt d t \ ( «C02 -aC0) Let rate factor (R.F.) = ( F + a C O ) ( F +aC02) ( M/ a t m ) . -dPT _ PCO R I. ( v p, "dt - o . R i FCO - ^CO \ R«. (atm sec-1) (F-3) R.F. i where PQQ = i n i t i a l CO pressure and APco =• decrease in CO pressure in time t From Appendix A, the amount of CO consumed in time t is given by: X - APco (F +«co) = APT x r- f- ( m) ( f J+) * See Appendix A for definition of terms. - 123 - A P C 0 = AP T B'*' (atm) (F-5) ^ "CO Substitution in F-3 gives: -dPrp / P c o " APT R.F./(F + aCo) dt V R.F. R' (atm sec - i ) (F-6) R.F. and F + QQ can be calculated by the method outlined in Appendix A. PQQ and APIJ can be obtained from an experimental pressure vs time record. R' can be evaluated in the following manner, using equation 19 and the stoichiometry of the overall reduction reaction represented by equation F-7: 2Ag(I) + CO + H20 + 2NaOAc > 2Ag + C0 2 + 2H0Ac + 2Na+ (F-7) Since X represents the amount of CO consumed in time t, then, i f [ ]^ denotes i n i t i a l concentration, [AgClOJi - 2X = [AgOAc ] + [Ag +] (F-8) [NaOAc]± - 2X = [AgOAc] + [OAc-] (F-9) [HOAcJi + 2X = [HOAc] (F-10) where X is given by F-k. [AgOAc] can also be expressed in terms of the association constant, Ka, according to equation F - l l : [AgOAc] = Ka[Ag+][OAc-] (F-ll) Substitution- for [Ag+] and [OAc-] from F-8 and F-9 gives: [AgOAc] = b - (b 2 - c)2" (F-12) where b = 0.5 ([AgClO^i + [NaOAc ] ± + l/K a - kX) c = ( [AgC104 ]i [NaOAc ]i - 2X ([AgC104]i + [NaOAc ]± )••+ kX) - 124 - Thus R1 can be evaluated in terms of F - 8 , F -9 , F-10 and F-12 using an experi 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 to t: -dPT dt~ ^ 0 - A ^ R . F . / ( F + ° c 0 ) j r, ( a t m s e c _ i } x l k m l ^0 - A P T ^ F . / ( F +aC0)y, ( p s i min-1)(F.15) 60 = f(APT) (psi min-i) (F-l4) ^t TAPT dt = t = -/ (f (APT)) _ 1 dP T (min) (F-15) o ~ o The time ( i n min) required for a decrease in total pressure (in psi) can be calculated from the area under a (f (ABj)) ""h/s APrpj plot-. The integral in 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 psi intervals. 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: 0 . 1 1 5 M AgC104; O.I95 M NaOAc; O . O 7 8 M H O A C ; 7 9 psi CO; 90°C Gas Solubility at 90°C OQ0 = 6 . 9 x 10-4 M/atm aco2 = IO5.O x 10-4 M/atm Volume of solution added (V]_) = 9 0 . 0 mis Total volume of reactor at 90°C = 120.0 mis Density of H 2 0 = O . 9 6 5 5 g/ml at 90°C; O . 9 9 8 2 g/ml at 20°C Vg = 120.0 - 9 0 . 0 x O . 9 9 8 2/O . 9 6 5 5 = 2 6 . 9 mis - 125 - F = !g x 1 0 0 0 = 26.9 1000 Vi RT 90.0 X 82.05 x 363 = 100.3 x 10"4 M/atm R.F. = (F + a c o ) ( F + a C 0 2 ) ( a c o 2 - a co) = (100.3 + 6.9)10~4 (100.5 + 105.0)10-4 (105.0 - 6.9)10-4 = 0.0224 M/atm Values for (f(APT)) _ 1 and calculated time computed from the above data at 1-psi intervals up to about 50$ reaction using the "best value" rate constants* for equation 19 from Table X, are given in Table F-I. Also included are the time values taken from the experimental records for experi ments 207 and 230. * kx' = k 3' = 2.7 x 10-5 atm-i sec-i; k2' 6.2 x 10-4 M_1 sec-i = 2.1 x 10"4 M-i sec-i; - 126 - TABLE F-I COMPARISON OF EXPERIMENTAL AND CALCULATED PRESSURE RECORDS (Figure 25, Section III-10) APT* (f (APr r j ) ) _ 1 min p s i - i , Time (min) Calculated Experimental No 207 No 230 0 0.67 0 0 0 1 0.75 0.7 1.0 0.9 2 O.85 1-5 1.8 1.9 3 O.95 2.4 2.8 2.7 4 1.07 3.4 3-7 3-8 5 1.21 4.6 --6 1.37 5.9 5-6 6.0 7 1.55 7.3 -- --8 1.76 9.0 8.7 10.3 9 2.00 10.9 --10 2.28 13.0 12. C 13.7 11 2.60 15.4 15.8 12 2.97 18.2 19.0 13 3 .41 21.4 --14 3.92 25.1 27.0 15 4.53 29.3 30.0 16 5.26 34.2 34.5 17 6.12 39.9 37-2 18 7.16 46.5 45.7 19 8.43 54.3 50.2 20 9.98 65.5 57-8 A total pressure decrease of 20 psi i s equivalent to about 50$ reaction. - 1 2 7 - VI REFERENCES 1. Forward, F.A., Trans. Can. Inst. Min. and Met., 5_6, 363 (1953). 2. Schaufelberger, F.A., J. of Metals, 8, 695 (1956). 3. Newton, J., "Extractive Metallurgy", John Wiley and Sons, Inc., New York, 1959, P 351- 4. Commings, E.W., Ind. Eng. Chem., 3_9_, 948 (194-7). 5. Lewis, R.M., Cookston, J.W., Coffer, L.W. and Stephens, F.M. Jr., J. of Metals, 10, 4l9 (I958). 6. Schaufelberger, F.A.. (to Chemical Construction Company) U.S. Patent 2,734,821 (Feb. l4, 1956). 7. VanHare, G. and Montgomery, R. (to Chemical Construction Company) U.S. Patent 2,733,990 (Feb. 7, 1956). 8. Begunova, T.G., Tsvet. Metal.., 3_0, 14 (1957). 9. Ives, J.G. and Janz, G.J., "Electrode Potentials", Academic Press, New York, I96I, p 26. 10. Latimer, W.M., "The Oxidation States of the Elements and Their. Potentials in Aqueous Solution", 2nd ed., Prentice Hall, Inc., Englewood C l i f f s , N.J., 1952. 11. Dobrokhtov, G.N., Zhurn. Priklad. Khim., 32, 78O (1959). 12. Fechner, J., Doctoral Dissertation, Technischen Universitat Berlin, i960. 13. Bauch, G. von, Pawlek, F. and Plieth, K., Z. Erzbergbau u. Metallhuttenw. , 11, 1 (1958). 14. Nakamura, S. and Halpern, J., J. Am. Chem. Soc, 83, 4102 (196l). l4a Nakamura, S. and Halpern, J., personal communication. 15. Harkness, A.C. and Halpern, J., J. Am. Chem. Soc, 83, 1258 (196l). 16. Weisser, H.B., "Inorganic Colloid Chemistry", vol I, John Wiley and Son, Inc., New York, 1955, p 4l. 17. Just, G. and Kauko, Y., Z. physik. Chem., 82, 71 (1913); Chem. Abs., J_, 3881.(1915). 18. Peters, E., unpublished observations. 19. Peters, E., personal communication. 20. Orchin, M. and Wender, I., in "Catalysis", vol V, edited by Emmett, P.H., Reinhold Publishing Corp., New York, 1957, p 55. - 128 - 21. Hirsch, E., unpublished observations. 22. Coulson, C.A., Quat. Rev., 1, l44 (1947). 23. Orgel, L.E., "An Introduction to 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. Soc, London, 1959, p 93. 25. Sidgwick, N.V., "The Chemical Elements and Their Compounds", vol 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., in "Catalysis", vol IV, edited by Emmett, P.H., Reinhold Publishing Corp., New York, I956, p 29. 28. Wender, I., Sternberg, Hr.W. and Orchin, M., in "Catalysis", vol 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 Chemistry, Special Publication No. 13, Chem. Soc, London, 1959, P 35- 30. Halpern, J. and Kettle, S.F.A.,•Chem. and Ind., 668 (1961). 51. Bjerrum, J., "Metal Ammine Formation in Aqueous Solution", P. Haase and Son, Copenhagen, 194l, p 134. 52. McDuffie, H.F., Compere, E.L., Stone, H.H., Woo, L.F. and Secoy, C.H., J. Phys. Chem., 62, 1030 (1958). 55. Vogel, A.T., "Quantitative Inorganic Analysis", 2nd ed., Longmans Green and Co., Toronto, 1951, p 256. 54. Seidell, A. "Solubilities of Inorganic and Metal Organic Compounds", 4th ed., vol I, edited by Linke,W.F., D. Van Nostrand Co., Inc., New York, 1958. 35- Shriner, R.L., Fuson, R.C. and Curtin, O.Y., "The Systematic Identification of Organic Compounds", 4th ed., John Wiley and Son, Inc., New York, 1956, p 29. 36. Brewster, R.Q., "Organic Chemistry", Prentice Hall Inc., Englewood C l i f f s , N.J., 1949, p 112. 37- Fieser, L.F., "Experiments in Organic Chemistry", 3rd ed., D.C. Heath and Co., Boston, 1957, p 86. \, 58. "Stability Constants - Part I, Organic Ligands", compiled by Bjerrum, | j . , Schwarzenbuch, G. and Sille'n, L.G., Special Publication No. 6, Cliem. Soc, London, 1957, p 3. 39. MacDougall, F.H. and Petersen, S., J. Phys. Chem., 5_1, 1346 (1947). - 129 - 4 0 . Monk, C.B., "Electrolytic Dissociation", Academic Press, New York, I 9 6 I , P 1 6 7 . 41. Harned, H.S. and Owen, B.B., "The Physical Chemistry of Electrolytic Solutions", Jrd ed., Reinhold Publishing Corp., New York, 1 9 5 8 , chap 1 5 . 42. "Stability Constants - Part II, Inorganic Ligands", compiled by Bjerrum, J., Schwarzenbuch, G. and Sille'n, L.G., Special Publication No. 7 , Chem. Soc, London, 1 9 5 8 , p 1 6 . 4 3 . Orgel, L.E., "An Introduction to Transition-Metal Chemistry", John Wiley and Sons, Inc., New York, i 9 6 0 , p 6 6 . 4 4 . Leden, I., Svensk Kern. Tidskr., 5_8, 129 ( 1 9 4 6 ) ; Chem. Abs. 4 0 , 6 0 2 1 ( 1 9 4 6 ) . 4 5 . .MacDougall, F.H., J. Phys. Chem., 46, 737 ( 1 9 4 2 ) . 46. Webster, A.H. and Halpern, J., J. Phys. Chem., 6 l , 1239 ( 1 9 5 7 ) . 4 7 . Webster, A.H. and Halpern, J., J. Phys. Chem., 6 1 , 1245 ( 1 9 5 7 ) - 48. Courtney, W.G., J. Phys. Chem., 6 0 , l 4 6 l ( 1 9 5 6 ) . 4 9 . Glasstone, S., "An Introduction to Electrochemistry", D. Van Nostrand Co., Inc., New York, , 1942, p 4 6 1 . 5 0 . Moelwyn-Hughes, E.A., "The Kinetics of Reactions in Solution", 2 n d ed., Oxford University Press, London, 1 9 4 7 , chap 3 - 5 1 . "Handbook of Chemistry and Physics", 3 6 t h ed., Chemical Rubber Publishing Co., Cleveland, Ohio, 1 9 5 5 , p I 9 6 5 . 5 2 . Winkler, L.W., Ber., 2 4 , 3 6 0 9 ( 1 9 O I ) ; summarized in reference 3 4 , p 4 5 5 . 5 5 - Dodds, W.S., Stutzman, L.F. and Sollami, B.J., J. Chem.. Eng. Data, 1 , 92 ( 1 9 5 6 ) . 5 4 . Himmelblau, D.M., J. Chem.. Eng. Data, 5_, 10 ( i 9 6 0 ) . 5 5 . Wiebe, R. and Gaddy, V.L., J. Am. Chem. Soc, 5_6, 76 ( 1 9 3 4 ) . 5 6 . Pray, H.A., Schweickert, C.E. and Minnich, B.H., Ind. Eng. Chem., 4 4 , 1146 ( 1 9 5 2 ) . 5 7 . Stephen,. E.F., Hatfield, N.S., Peoples, R.S. and Pray, H.A., "The Solubility of Gases in Water and Aqueous Uranyl Salt Solutions at Elevated Temperatures and Pressures", Battelle Memorial Institute, Columbus, Ohio, BMI-IO67, 1 9 5 6 . 5 8 . Wiebe, R. and Gaddy, V.L., J. Am. Chem. Soc, 6 1 , 315 ( 1 9 3 9 ) . 5 9 - Haehnel, 0 . , Centre. Min. Geol., 2_5_, 25 ( 1 9 3 0 ) ; summarized in reference 3 4 , p 4 6 1 . - 130 - 60. MacDougall, F.H., J. Am. Chem. Soc, 5_2, 1390 (1930). 61. MacDougall, F.H. and Rehner, J. Jr., J. Am. Chem. Soc, 5_6, 368 (1934). 62. MacDougall, F.H. and Bartsch, C.E., J. Phys. Chem., 40, 64-9 (I936). 63. MacDougall, F.H. and Larson, W.D., J. Phys. Chem., 4l, 4l7 (1937). 64. -MacDougall, F.H. and Allen, M., J. Phys. Chem., 46, 730 (1942). 65. MacDougall, F.H. and Allen, M., J. Phys. Chem., 4£, 245 (19^ 5). 66. MacDougall, F.H. and Topel, L.E., J. Phys. Chem., 5_6, 1090 (1952). - 131 - VII NOMENCLATURE USED IN RATE EXPRESSIONS Rate Functions R = -d[CO]/dt = rate of CO consumption (M sec-i) R' = R/Pco = rate of CO consumption at unit pressure (M atm - l sec-i) I = acid-independent rate (M sec-i) I' = acid-independent rate at unit pressure (M atm-i sec-i) D = acid-dependent rate (M sec-i) D' = acid-dependent rate at unit pressure (M atm _i sec - 1) S = acid-proportionality constant for acid-dependent rate (M2 sec-i) S' = acid-proportionality constant at unit pressure (M2 atm' - i sec-i) Ro = acetate-independent rate (M sec -i) Ho' = acetate-independent rate at unit pressure (M atm - 1 sec Rate Constants k 2, k 3 = experimental rate constants in terms of [CO] ki', k2', k 3 ! = experimental rate constants in terms of PQQ k2", k2'" = rate constants for reaction in unbuffered solution (Section III - 9 ) k3", k 4" = rate constants for acid-dependent reaction (equations 15 and 16) ka, kb, k c, k^, k e, kf, kg = rate constants for proposed mechanism (reaction V) Cn(n=l,2,3,4) = [OAc_]-power series coefficients in eqn 38 for acid-dependent react. Equilibrium Constants = ionization constant of acetic acid (M) K a = association constant for average silver-acetate complex (AgOAc) from Ag+ and OAc- (M~i) Ki, K 2 = association constants for AgOAc and Ag(OAc)2 from Ag + and OAc- (eqn 37) K c = formation constant, incorporating [H 2 0 ] , for intermediate complex from Ag +, CO and H 2 0 in proposed mechanism (reaction V(d)) Kc' = formation constant, incorporating [H 2 0 ] , for intermediate complex from AgOAc, CO and H 2 0 in proposed mechanism (reaction V(h)) a C 0 = CO solubility coefficient (M/atm) a C 0 2 = C 0 2 solubility coefficient (M/atm) 

Cite

Citation Scheme:

    

Usage Statistics

Country Views Downloads
United States 13 0
China 6 25
Canada 2 0
France 1 0
City Views Downloads
Unknown 5 4
Ashburn 4 0
Hangzhou 2 0
Matawan 2 0
Beijing 2 0
Shanghai 1 0
Charlotte 1 0
Paris 1 0
Parlin 1 0
Guangzhou 1 0
Mountain View 1 0
Redwood City 1 0

{[{ mDataHeader[type] }]} {[{ month[type] }]} {[{ tData[type] }]}
Download Stats

Share

Embed

Customize your widget with the following options, then copy and paste the code below into the HTML of your page to embed this item in your website.
                        
                            <div id="ubcOpenCollectionsWidgetDisplay">
                            <script id="ubcOpenCollectionsWidget"
                            src="{[{embed.src}]}"
                            data-item="{[{embed.item}]}"
                            data-collection="{[{embed.collection}]}"
                            data-metadata="{[{embed.showMetadata}]}"
                            data-width="{[{embed.width}]}"
                            async >
                            </script>
                            </div>
                        
                    
IIIF logo Our image viewer uses the IIIF 2.0 standard. To load this item in other compatible viewers, use this url:
http://iiif.library.ubc.ca/presentation/dsp.831.1-0105706/manifest

Comment

Related Items