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Reduction of silver amine complexes by carbon monoxide Nakamura, Shuzo 1962

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REDUCTION OF SILVER AMINE COMPLEXES BY CARBON MONOXIDE by SHUZO NAKAMURA Bo Sc. in Engineering, Kyoto University, 1960 A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF SCIENCE in the Department of CHEMISTRY We accept this thesis as conforming to the required standard THE UNIVERSITY OF BRITISH COLUMBIA August, 1962 In presenting this thesis in partial 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 is understood that copying or publication of this thesis for financial gain shall not be allowed without my written permission. The University of British Colombia, Vancouver 8, Canada. Department i i ABSTRACT The kinetics of the reduction of silver amine complexes in aqueous solution by carbon monoxide were investigated. For a number of amines including ethyl-, methyl-, diethyl-ethanol-, diethanolamine and some primary diamines, the rate law was found to be of the form; d(Ag(I)) d(CO) [AgL2+3[CO] " dt = ~ 2 ~~dt~ = k e x P [LH+j - 2k (L-Ag-OH ]tCOj (1) where L denotes the amine. These kinetics were interpreted in terms of the following mechanism. AgL 2 + + H20 ^ r = ± : L-Ag-OH + LH+ (Rapid) (2) k L-Ag-OH + CO L-Ag-COOH (Rate-determining) (3) L-Ag-COOH + Ag(I) >• Products (Rapid) (4) The rate constant of the rate-determining step (3) was found to be nearly independent of the nature of the amine molecule, L, coordinated to silver ion, using the basicity constants of the amines and dissociation constants of the corresponding silver amine complexes. The actual overall rate of the reaction varied with the nature of amine but this was attributable only to the different equilibrium concentrations of L-Ag-OH. The rate of this rate-determining bimolecular process was found to be surprisingly fast; k__o = 5x10 mole 1. sec. , i i i A H ^ 9 Kcal. mole" and A S — -15 e,u. The reduction of s i l v e r ion by CO i n a c i d i c or neutral media i s known to be very slow and t h i s can now be at t r i b u t e d to the base catalyzed nature of the reaction. S i l v e r complexes of primary diamines (ethylenediamine, 1,3-diaminopropane, etc.) were reduced more slowly; t h i s was att r i b u t e d to the s t a b i l i z a t i o n of mono-complexed s i l v e r (I) species by chelate formation. In the case of ammonia normal k i n e t i c s were observed at higher pH but at Lower pH the rate became second order i n (Ag(I))and inversely second order i n (NH^ +). This was a t t r i -buted to competition between decomposition of the intermediate complex and i t s further reaction with another Ag(I) species to give m e t a l l i c s i l v e r and carbon dioxide. Evidence for si m i l a r competition was found with two t e r t i a r y amines, i . e . , triethylamine and triethanolamine. ACKNOWLEDGEMENTS The author wishes to express h i s sincere g r a t i t u t e for the continuing advice, help and encouragement given by Dr. J . Halpern, who suggested and directed t h i s study, and to Dr. E. Peters and Mr. R. T. McAndrew of the Department of Mining and Metallurgy for information about t h e i r r e l a t e d work. He i s also g r a t e f u l to Dr. C. A. McDowell, Head of the Department of Chemistry, who enabled him to work i n t h i s Department. Support of t h i s work by the A l f r e d P. Sloan Foundation and the National Research Council of Canada i s also g r a t e f u l l y acknowledged. V TABLE OF CONTENTS Page I. INTRODUCTION 1 II. EXPERIMENTAL. . 6 MATERIAL. 6 ANALYSIS 7 PROCEDURE .7 Kinetic Measurements 7 Stoichiometry Measurements 11 III. RESULTS AND DISCUSSIONS 13 STOICHIOMETRY 13 KINETICS AND MECHANISM. 18 Ethylamine Complex. 19 Other "Standard" Systems 25 Triethylamine Complex 33 Triethanolamine Complex 40 Ammonia Complex . . . . . . . . 47 Diamine Complexes 56 GENERAL DISCUSSION 64 REFERENCES 71 APPENDIX I. Selected Thermodynamic Properties of Amines and Silver Amine Complexes 72 v i LIST OF TABLES Table No. Page l o Results of Stoichiometry Measurements (I) 1 3 2 . Results of Stoichiometry Measurements (II) . 1 4 3 o Results of Stoichiometry Measurements (III) . 1 8 4 . Rate of Reaction of Ethylamine and Related Amine Complexes of Silver . . . . . . 2 2 5. Summary of Kinetic and Related Thermodynamic Data for "Standard" Systems at 25°C. . . . . . . 3 0 6. Apparent Enthalpy and Entropy of Activation for "Standard" Systems 3 2 7. Rate of Reaction of Triethylamine Complex at 2 5 ° C . . . 3 4 8 . Rate of Reaction of Triethanolamine Complex . . . . . . 4 1 9 . Rate of Reaction of Ammonia Complex . . . . . . . . . . 4 8 1 0 . Rate of Reaction of Diamine Complexes . . . . . . . . . 57 11. Rates of Diamine Complexes and Their Stability 1 2 . Summary of Kinetic and Related Thermodynamic Data . . . 65 v i i LIST OF FIGURES Figure No. Page I . Gas bubbling glass apparatus. 8 I I . Typical tit r a t i o n curves for f i n a l reaction solutions . . . . . 16 I I I . Typical rate plots for ethylamine complex 20 I V . Dependence of rate on carbon monoxide concentration at 25°C. for ethylamine complex 21 V . Dependence of rate on ammonium ion concentration at 25 C. for ethylamine complex 24 V I . Typical rate plots for ethylamine-type complexes. . . 26 V I I . Arrhenius plots for ethylamine-type complexes . . . . 31 V I I I . Typical rate plots for triethylamine complex. . . . . 35 I X . Dependence of rate on free amine concentration at 25°C. for triethylamine complex . . 38 X . Typical rate plots for triethanolamine complex. . . . 42 X I . Arrhenius plots for triethanolamine and ethylene-diamine complexes 46 X I I . Typical rate plots for ammonia complex. 50 X I I I . Dependence of rate on ammonium ion concentration at 30°C, for ammonia complex 51 X I V . Arrhenius plot for ammonia system 54 X V . Typical rate plots for diamine complexes 58 X V I . Dependence of the rate on free amine concentration at 25°C. for 1,3-diaminopropane 60 I. INTRODUCTION Recently, molecular hydrogens although unreactive toward the majority of common inorganic oxidizing agents, was found to be oxidized under relatively mild conditions in aqueous 2+ solution by a few metal ions and complexes, notably Cu s Ag +, Hg2+, Hg^s, and MnQ^~„ Halpern and his coworkers (1) have studied these systems extensively and have elucidated the mechanism through which the relatively strong H-H bond, -1 having a dissociation energy of 103 Kcal. mole , i s activated, and the dependence of the reactivity on the electron configura-tion of the central metal ions. These studies have prompted similar studies on other inert reducing agents. The mechanism of these reactions and the nature of metal ions and complexes which activate those inert molecules are of great interest for the study of chemical reactivity, in general, and especially of the catalytic activity of transition metals and their compounds. Amongst other relatively inert reducing agents whose reac-tions have been investigated in this laboratory are carbon monoxide and formic acid. The reactions of the latter with - 94- 2+ %4-MnO^  s Hg , Hgj and T l were studied by Taylor and Halpern (2) and their kinetics and mechanisms were elucidated. Harkness and Halpern (3) examined the reaction of CO with those metal ions which are active toward molecular hydrogen and found that 2 only Hg and MnO^  showed measurable reactivity toward CO in homogeneous aqueous solution under moderate condition. They found Fe^ + 9 Tl"**" and Cr^O^- also to be inactive. At elevated temperature and pressure Bauch et a l . (4) observed that silver sulfate and cupric sulfate in aqueous solution also were reduced by CO.. They reported that the rate of the former reaction was second order in Ag(I) and was enhanced by buffering the solution with ammonium acetate. However, they did not study the dependence of the rate on pH. Following this work, and after commencement of the present study, Peters and McAndrew (5) studied the reaction of silver acetate in aqueous solution with CO in further detail under experimental conditions similar to those of Bauch et a l . (4). These related studies are summarized below. 2+ Hg (3)...This i s the only metal ion which was found to oxidize CO in aqueous solution under relatively mild condi-tions (atmospheric pressure and below 80 oC) in the absence 24-of complexing agents. For the reduction of Hg , i.e., 2Hg 2 + + CO + H20 HgJ + + C0 2 + 2H+ (1-D kinetic measurements in dilute HC10. solutution over the temperature range 26 to 54 C. yielded the pH-independent rate law (1-2) * -1 "k with AH = 14.6 Kcal. mole and AS = -13 e.u. This was interpreted in terms of the following mechanism. 0 4--Hg z + 0H 2 + CO *- -Hg-C-OH + H (slow) (1-3) -Hg-C-OH > Hg + C0 2 + H (fast) (1-4) Hg + Hg 2 + Hg 2 + (fast) (1-5) Support for the proposed intermediate complex i s provided by the P, isolation of a stable analogue, AcO-Hg-C-OCH^, formed by reaction of CO with mercuric acetate in methanol solution (6). Mn04~ (3)...The reduction of MnO^ " by CO (to Mn02 in acidic and neutral solutions and to MnO^  in basic solutions) was found to proceed readily over the temperature range 28 to 50°C. The complete rate law was found to be " ^dT = kCC0J[Mn04") (1-6) with AH = 13 Kcal. mole" 1 and AS = -17 e.u., both substantially constant over the pH range L to 13, which confirmed and extended the earlier kinetic measurements on this system by Just and Kauko (7). Harkness and Halpern (3) also found that this system shows a remarkable catalytic effect on the addition of Ag + and Hg 2 + (but not Cu 2 +, Fe 3*, Cd 2 +, or T l 3 + ) which they attributed to favorable reaction paths involving intermediate such as Ag-CO-OMnOg analogous to that postulated in the reaction of Hg 2 + with CO. 4 Ag2SO^ and CuSO^ (4)...Bauch et a l . studied the aqueous Ag2SO^ system over the temperature range 70 to 110°C. under CO pressure up to 50 atmosphere and found the reaction Ag + + %C0 + %H20 > Ag + %C02 + H + (1-7) proceeded according to the rate law given by (1-8). " ^ d f ^ = ^ S x l O 6 U g + ) 2 P c o e - 1 4 » 0 0 ° / R T (1-8) They also observed that the rate was increased by buffering the solution with ammonium acetate, the kinetics for the buffered silver sulfate system being given by (1-9). - = 6.02x10* U g + ) 2 P C Q e - 9 ' 3 0 0 / R T (1-9) They attributed this difference in rate to the favorable dependence of the equilibrium on increasing pH, and proposed the same mechanism for both buffered and unbuffered system, i.e., Ag + + CO 5 = = ^ Ag(C0) + (rapid equilibria) (1-10) Ag(C0) + + Ag + ~ = r Ag2(C0)"H" Ag 2(C0) 4 + + H20 >2Ag + C0 2 + 2H + (rate determining) (1-11) In support of this mechanism, they cited the existence of carbonyl complex of the " f i r s t subgroup" such as (Cu(Cl,Br)C0)^, Ag2(C0)S0^ and (AuCl'CO)^. It i s obvious from their results that there must be some pH-dependent process contributing to the overall rate but this was not elucidated. For the reduction of CuSO^ by CO they also observed second order dependence of the rate on [Cu ), the rate over the temperature range 160 to 190 C. being given by 3 d ICu dt 2.56X1013 ( C u ^ ) 2 * e " 3 3 ' 5 0 ° / R T (1-12) In this system the effect of buffering was not reported because of experimental d i f f i c u l t i e s . Recently, Peters and McAndrew (5) have extended this work on the reduction of silver salts in acidic solution. Both in acetate-buffered and perchlorate media the rate was found to be very slow, requiring the use of elevated temperature (>90°C.) and CO pressure (10 to 30 atm.). The results of this work w i l l be considered later. The present study i s concerned with the reduction of silver amine complexes by CO in basic media. In contrast to the behaviour in acid solutions the reaction under these conditions i s rapid and readily measureable at room temperature and atmospheric pressure. 6 II. EXPERIMENTAL MATERIALS Silver perchlorate was G. F. Smith Reagent grade and was unaffected by recrystallization. Perchloric acid was Baker and Adamson 60% Reagent grade. Ethylenediamine, Fisher c e r t i -fied reagent, was used without further purification. D i s t i l l a -tion of this product had no effect on the reaction rate. Matheson triethylamine, which contained a reducing impurity, was purified by passing through a molecular sieve column and then d i s t i l l e d under 120 mm. Hg nitrogen atmosphere. Matheson 33% aqueous solution of ethylamine, and diethylamine (b.p. 55-56°C); B.D.H. 25/30% methylamine aqueous solution, pure-ethanolamine, diethanolamine and triethanolamine, and K & K Laboratories' 1,3-diaminopropane were used without further purification. K & K Laboratories' 1,4-diaminobutane was redis-t i l l e d at 20 mm. Hg before use. Ordinary d i s t i l l e d water was used in the preparation of a l l solutions and gave rates identical with those obtained with water d i s t i l l e d from alkaline perman-ganate. Nitrogen gas was supplied by the Canadian Liquid Air Co. Carbon monoxide (CP. grade) and CO-^ gas mixtures were obtained from Matheson of Canada Ltd. The chromatographic analysis of a l l these gases revealed substantially no contamination by oxygen. In a l l experiments the amine perchlorate was prepared 7 by n e u t r a l i z i n g the amine with p e r c h l o r i c a c i d . The experimental solutions were prepared by d i l u t i n g aliquots of standardized stock solutions. ANALYSIS The normality of amines and aqueous amine solutions was determined by t i t r a t i o n with standard hydrochloric a c i d . S i l v e r ion concentration was determined by thiocyanate t i t r a t i o n i n a c i d i c solution with f e r r i c i n d i c a t o r . Carbon monoxide and nitrogen gas mixtures were analyzed with a Beckman GC-2 gas chromatograph using a molecular sieve column. PROCEDURES Ki n e t i c measurements Except for the ammonia system, rates of a l l the reactions were determined at atmospheric pressure, by bubbling the CO gas (or a CO-N^ mixture) through the solution i n the glass apparatus depicted i n Figure I. The gas was passed through a presaturator f i l l e d with aqueous solution of NaNOg and the amine to e s t a b l i s h the same p a r t i a l pressure of water and the amine as the reaction solution, and was then dispersed through a sintered glass plate into the reaction s o l u t i o n . The ef f l u e n t gas was l e d to a gas flame and was burned. The whole apparatus was immersed i n a constant temperature bath thermostated to ~ 0.03°C. I t was established that the flow rate of the gas d i d not a f f e c t the observed reaction rate; hence i t may be assumed that the solutions were saturated with the gas. F i g . I. Gas Bubbling Glass Apparatus A. Gas Inlet B. Presaturater Solution C. Sintered Glass Plate D. Reaction Mixture E . Gas Outlet F. Gas Outlet Stopper G. Sampling Tube H. Sample Outlet CO 9 After placing 250-500 ml. of the reaction mixture of the desired composition in the apparatus, the system was allowed to attain thermal equilibrium under nitrogen flow. The s t a b i l i t y of the reaction mixture was checked by sampling and analyzing the solution several times under the nitrogen flow and then the gas flow was switched to CO or to a CO-^ mixture. The solution was sampled periodically and the samples were analyzed as described previously. The time required for saturation of the solution with the gas was usually negligible; less than 30-60 sec. For reaction solutions in which the total solute concentrations were lower than 0.5-0.6 molar, the part i a l pressure of the gas was assumed to be atmospheric pressure minus the vapor pressure of pure water at the reaction tempera-ture. In the case of triethylamine, which required a very high amine concentration to obtain stable solutions, Lattey°s (8) data for the total vapor pressure of triethylamine-water mixtures were used. Variation of the CO partial pressure was achieved, when desired, by using analyzed CO -N2 mixtures. With ammonia, whose partial pressure i s very high and whose reaction rate was very low at atmospheric pressure, an autoclave was used. The apparatus used was a Parr Series 4500 autoclave with a glass-lined stainless steel reaction vessel, provided with a s t i r r e r , gas inlet tube, sampling tube f i t t e d with a stainless steel f i l t e r , pressure gauge and thermowell, surrounded by an electric heating mantle controlled by a rheostat. 10 Fine temperature control was achieved by use of an a u x i l i a r y e l e c t r i c heater, immersed i n the reaction mixture through the g l a s s - l i n e d thermowell, and controlled by a Thermistemp Tempera-ture Controller (Model 71) actuated by a thermistor immersed i n the solution. This arrangement gave temperature control of d= 0.3°C, A 1 s500 ml. reaction mixture was made up from stock solutions and placed i n the reaction v e s s e l . Nitrogen gas was run into the mixture through the sampling tube and the porous s t a i n l e s s s t e e l f i l t e r , under a g i t a t i o n by the s t i r r e r , for some f i v e minutes; then the vessel was sealed and brought to the desired temperature. The s t a b i l i t y of the solution was established by taking samples and analyzing them f o r s i l v e r ion over a one hour period. The i n t e r n a l pressure was then reduced to one atmosphere by opening the gas out l e t once, and the CO gas was introduced from a CO cylinder and maintained at a desired pressure. The p a r t i a l pressure of CO was calculated as the difference between the t o t a l pressure and the combined vapor pressure of the solution and r e s i d u a l nitrogen pressure. Samples were taken at appropriate i n t e r v a l s and analyzed for s i l v e r ion. After each sampling more CO gas was supplied to * The nichrome wire c o i l of the a u x i l i a r y heater was made to occupy the lower h a l f of the thermowell and a volume of solution s u f f i c i e n t to f i l l the glass l i n e r was used so as to avoid undesirable superheating of the thermowell. 11 the system to establish a constant pressure. The s t i r r e r was rotated, usually at 600 r.p.m., after establishing that the rate of reaction was independent of the s t i r r i n g rate. Stoichiometry measurements The amount of CO gas consumed was measured by means of a gas burette apparatus. A 125 ml. conical flask reaction vessel was placed in a small poly-ethylene water bath (diameter ca. 5 in.) which was connected to a thermostated water circulator. The reaction vessel flask was connected to a 50 c c . burette provided with a mercury balancing bottle. The reaction solution was stirred with a teflon-coated magnetic s t i r r e r . 100 ml. of reaction solution was placed in the reaction vessel; nitrogen gas was f i r s t passed through the solution to remove oxygen from the reaction system, and f i n a l l y after displacing the nitrogen by CO and f i l l i n g the reaction vessel with CO, the whole system was closed, the gas burette mercury level was balanced and the magnetic st i r r e r was turned on. A dibutylphthalate auxiliary manometer was used for fine adjustment of the mercury level. When the rate of gas uptake became sufficiently slow the burette was read, the reaction vessel was disconnected from the burette and the solution was analyzed. The volume of gas uptake was corrected for the solubility of CO, temperature and pressure. In several cases the reaction mixture was titrated for 12 amine and carbonate before and after the reaction. In those cases, an original reaction solution was flushed with carbon dioxide-free nitrogen gas before i t was placed into the reaction vessel. A 25 ml. aliquot was taken for analysis and 100 ml. was pipetted under into the reaction vessel. The original and the f i n a l solutions were analyzed for amine and carbonate by potentiometric tit r a t i o n with standard hydrochloric acid using a Beckman pH meter. Silver metal deposited in the reaction vessel was collected on a glass crucible and weighed. Its purity was determined by dissolving in n i t r i c acid and t i t r a t i n g with thiocyanate. The amount of carbonate in the f i n a l reaction mixture was determined more directly by a gravimetric method, as barium carbonate. Barium nitrate or barium hydroxide was used to precipitate the carbonate. For weakly basic amines barium hydroxide was used. In some cases ethylenediamine or propylenediamine was added to stabilize the remaining silver ion before barium hydroxide was added. 13 I I I . RESULTS AND DISCUSSIONS STOICHIOMETRY For a l l the amines which were used in the present study, when aqueous silver (I) amine solution was reacted with CO, a brown or black solid separated from the solution, which proved on analysis to be pure silver metal. In many cases the glass wall of the reaction vessel also was covered by a silver mirror. The amount of silver metal formed was found to be equal to the amount of silver ion reduced (Table 1). TABLE 1 RESULTS OF STOICHIOMETRY MEASUREMENTS (I) Amine Carbon- Silver I n i t i a l Per- Silver Monoxide Metal (Ag(I)] chlorate Amine Reduced Uptake Recovered Amine mole/1 mole/1 mole/1 mole/1 mole/1 mole/1 ethylamine 0.0299 0.000 0.300 0.0299 0.015 0.0295 1,3-diamino-propane 0.0297 0.000 0.300 0.0297 0.015 0.0295 dlethanol-amine 0.0296 0.000 0.300 0.0296 0.0129 0.0294 14 TABLE 2 Amine RESULTS OF STOICHIOMETRY MEASUREMENTS (II) • 1 • 2 I n i t i a l Amine Carbon- Hydrogen Carbon-(Ag(I)) Per- S i l v e r Monoxide Ion a t e 2 chlorate Amine Reduced Uptake Produced Produced mole/1 mole/1 mole/1 mole/1 mole/1 mole/1 mole/1 ethylamine 0.0300 0.000 0.0963 0.0300 0.015 0.0606 0.0151 methylamine 0.0299 0.000 0.0937 0.0299 0.015 0.0605 0.0150 d i e t h y l -amine ethanol-amine dlethanol-amine 0.0296 0.100 0.0933 0.0292 0.0146 0.0620 0.0145 0.0300 0.000 0.100 0.0178 0.00809 ***3 ***3 3 3 0.0296 0.000 0.0988 0.0250 0.0114 (0.0532)(0.0133) u 3 0.0297 0.000 0.300 0.0296 0.0129 ***" t r i e t h a n o l -amine 0.0298 0.000 0.300 ***-ethylene-diamine 0.0199 0.0091 0.0300 0.000 0.0648 0.0298 0.0150 0.0618 0.0148 1,3-diamine-propane 0.0297 0.000 0.300 0.0297 0.015 *** *** 1 In those experiments where the reaction was continued to comple-t i o n a slow further uptake of CO was observed even when a l l the s i l v e r had been reduced. This zero order uptake of CO i s presumably a t t r i b u t a b l e to reaction of CO and H2O on the s i l v e r metal surface to form formate (9) «Jj > 2 These are the r e s u l t s from p H - t i t r a t i o n of the f i n a l s olution. 3 In the case of ethanolamines, which are the lea s t basic of a l l these amines, pH t i t r a t i o n was hot r e a d i l y applicable. For diethanolamine, t h i s method yielded only the combined concentra-t i o n of amine and carbonate. These r e s u l t s i n the parentheses were calculated from t h i s assuming hydrogen ion produced: / carbonate = 4s1, hence subject to error. ' 15 Results of CO uptake measurements in Table 2 show that two gram-ions of silver ion were reduced for each mole of * CO . The pH tit r a t i o n of the f i n a l solution (results are summarized in Table 2 and typical t i t r a t i o n curves are given in Figure II) showed that two gram-ions of hydrogen ion and one-half gram ion of carbonate (identity of this product i s to be discussed later) were produced for each gram-ion of silver ion reduced by one-half mole of CO. The overall reaction can therefore be represented by 2AgL 2 + + CO + 2H20 s*2Ag + C0 3~ + 4LH + (3-1) under conditions where the amine (represented by L) i s present in excess. The results of direct carbonate determinations (gravimetri-cally as barium carbonate) on the same f i n a l reaction solutions are shown in Table 3. Only triethylamine, diethylamine and ammonia systems gave those yields of carbonate expected from the ti t r a t i o n results. In the other cases a portion of the * In the case of the three ethanolamines the CO uptake was about 10% lower than the theoretical value (Table 2). In these cases, some silver apparently also was reduced by the amines or by an impurity. These side reactions were most pronounced at the high pH of these CO uptake experiments in which; in order to increase the rate of reaction, no amine perchlorate was added. Such solutions deposited some metallic silver on standing even in the absence of CO, while solutions containing amine perchlorate (i.e. those used for the kinetic experiments) were stable. The amount of silver reduced by these side reactions seemed to be directly dependent on the amine concen-tration as indicated by the two experiments with diethanola-mine in Table 2. 5 io 15 20 25 30 0.1 U Standard HCl, ml. Pig. II. Typical Titration Curves of Pinal Reaction Mixtures —'•— See Table 2 for experimental conditions. 17 "carbonate", resulting from the oxidation of the CO apparently combines with the amine to form a substance which decomposes on acidification. Decomposition with release of carbonate also occurred on treatment with base. Thus, when the f i n a l reaction solution was l e f t standing with an excess of barium hydroxide, the amount of barium carbonate precipitated increased slowly with time. In the case of methylamine, boiling with barium hydroxide resulted in a 100% yield of barium carbonate. These results suggest that the product in question i s a carbamate or similar compound, which i s known to form by reaction of carbon dioxide and amines or. .ammonia under moderately basic condition, and which i s decomposed by acid or by strong base. However, attempts to isolate, and characterize this product were unsuccessful and some question as to i t s identity remains. The representation of the reaction products by equation (3-1) i s thus subject to qualification, in certain cases, although the reactant stoichiometry appears to apply in every case. 18 TABLE 3 RESULTS OF STOICHIOMETRY MEASUREMENTS (III) I n i t i a l Amine Per- Silver Carbonate Carbonate (Ag(l)) chlorate Amine Reduced Produced Yield Amine mole/1 mole/1 mole/1 mole/1 mole/1 CO^/^Ag Triethyl-amine 0.0374 0.1 0.9 0.0356 0.0179 99% Diethylamine 0.0314 0.100 0.386 0.0281 0.0136 97% Ammonia 0.0300 0.00 0.300 0.0172 0.0083 97% Methylamine 0.0500 0.00 0.300 0.0455 0.0069 31% Ethylamine 0.0400 0.00 0.300 0.0317 0.0053 33% Ethanolamine 0.0500 0.00 0.300 0.0301 0.0086 57% Dlethanol-amine 0.0500 0.00 0.300 0.0339 0.0094 55% Triethanol-amine 0.0304 0.010 0.100 0.0243 0.00871 72% Ethylene-diamine 0.0114 0.00 0.200 o.oiii 0.001 20% 1,3-Diamino-propane 0.0400 0.00 0.300 0.0345 0.0084 49% 1 CO uptake was 0.0105 mole l " 1 (86%). KINETICS AND MECHANISM Among the silver amine complexes which were examined in this study, a number, including the complexes of ethylamine, methylamine, diethylamine, ethanolamine and dlethanolamine exhibited very similar behaviour (designated as "standard") and w i l l be discussed f i r s t . The triethylamine-, triethanolamine-19 and certain diamine- complexes, exhibited some departures from this "standard" behaviour and w i l l be considered later. ETHYLAMINE COMPLEX This system9 typical of those exhibiting "standard" behaviour, w i l l be considered in some detail. The disappearance of Ag(I) at constant CO pressure obeyed f i r s t order kinetics in a l l experiments as shown by the typical f i r s t order plots of log [Ag(I)) vs. time in Figure III. This was verified by the fact that the same rate constant was ob-tained for two different i n i t i a l s ilver ion concentrations keeping the other conditions unchanged (Experiments l c and l g ) . The rate-law obeyed during the course of each experiment i s thus where (Ag(I)) i s the total concentration of a l l the Ag(I) species, t i s time in seconds and log i s common logarithm. When a l l the other conditions were kept constant at 25°C, k" exhibited f i r s t order dependence on the CO parti a l pressure as shown by the plot of k" vs. [CO) in Figure IV. The CO concentration in the reaction solution was calculated using the solubility data of Seidell (10), assuming that the solution i s saturated with CO and the solubility can be approximated by that in pure water. Equation (3-2) can then be rewritten, as a second order rate-law. dt (3-2) 20 • 0 1,000 2,000 3,000 4,000 5,000 Time, sec. Pig. I I I . Typical Rate Plots f o r Ethylamine Complex See Table 4 for experimental conditions. 2 1 i 2 4 6 8 10 (CO], 10~4 mole - 1 Pig.IY.Dependence of Rate on Carbon Monoxide Concentration at 25° for Ethylamine Complex; (LH+)=0.1 & (L)=0.2 mole 1 where - k9 (Ag(X)) (CO) 2.303 22 (3-3) Values of the second order rate constants, k°, measured under various conditions are summarized in Table 4. TABLE 4 RATES OF REACTION OF ETHYLAMINE AND RELATED AMINE COMPLEXES OF SILVER AND OF SOME OTHER SIMILAR SYSTEMS Amine Ethylamine Methylamine Initial CO Amine Pressure Per- Si Ug(I)l mm. Amine chlorate M OCt o MxlO"3 Hg. M M T°C. sec. 1 xl0~ z No. 10.0 730 0.200 0.0500 25 0.310 1.55 la 10.0 730 0.200 0.0750 25 0.206 1.55 lb 10.0 730 0.200 0.100 25 0.152 1.52 lc 10.0 730 0.200 0.200 25 0.077 1.54 Id 10.0 730 0.300 0.100 25 0.153 1.53 le 10.0 730 0.100 0.100 25 0.156 1.56 If 20.0 730 0.200 0.100 25 0.149 1.49 lg 10.0 562 0.200 0.100 25 0.156 1.56 lh 10.0 490 0.200 0.100 25 0.154 1.54 l i 10.0 395 0.200 0.100 25 0.152 1.52 l j 10.0 234 0.200 0.100 25 0.143 1.43 lk 10.0 742 0.200 0.100 20 0.103 1.03 11 10.0 723 0.200 0.100 30 0.223 2.23 lm 10.0 712 i 0.200 0.100 35 0.337 3.37 In 10.0 730 0.200 0.100 25 0.283 2.83 2a 10.0 730 0.200 0.200 25 0.152 3.04 2b 10.0 730 0.200 0.300 25 0.094 2.88 2c 10.0 730 0.100 0.200 25 0.152 3.04 2d 20.0 730 0.200 0.200 25 0.152 3.04 2e 10.0 562 0.200 0.100 25 0.283 2.83 2f 10.0 395 0.200 0.100 25 0.287 2.87 2g 10.0 234 0.200 0.100 25 0.286 2.86 2h 10.0 747 0.200 0.100 15 0.110 1.10 2i 10.0 742 0.200 0.100 20 0.174 1.74 2j 10.0 723 0.200 0.100 30 0.424 4.24 2k 23 TABLE 4 (Continued) Initial CO Amine k Pressure Per- exp (Ag(I)l mm. Amine chlorate / M see." Amine MxlQ~J Hg. M M T°C. sec."-1 xl0" z No. Diethylamine 10.0 234 0.200 0.100 25 0.83 8.3 3a 10.0 234 0.200 0.200 25 0.435 8.7 3b 10.0 234 0.200 0.300 25 0.271 8.1 3c 10.0 234 0.400 0.100 25 0.69 6.9 3d 10.0 234 0.100 0.100 25 0.87 8.7 3e 20.0 234 0.200 0.100 25 0.76 7.6 3f 10.0 730 0.200 0.200 25 0.425 8.5 3g 10.0 562 0.200 0.200 25 0.445 8.9 3h 10.0 395 0.200 0.200 25 0.435 8.7 3i 10.0 742 0.200 0.200 20 0.295 5.9 3j 10.0 728 0.200 0.200. 30 0.66 13.2 3k 10.0 712 0.200 0.200 35 0.91 18.3 31 Ethanolamine 10.0 736 0.200 0.0500 25 0.057 0.29 4a 10.0 736 0.200 0.0250 25 0.109 0.27 4b 10.0 736 0.100 0.0250 25 0.109 0.27 4c 10.0 728 0.200 0.0500 30 0.091 0.46 4d 10.0 718 0.200 0.0500 35 0.147 0.74 4e Diethanola- 10.0 736 0.200 0.0250 25 0.122 0.31 5a mine 10.0 736 0.200 0.0500 25 0.060 0.30 5b 10.0 736 0.100 0.0250 25 0.127 0.32 5c 10.0 728 0.200 0.0500 30 0.107 0.53 5d 10.0 718 0.200 0.0500 35 0.170 0.85 5e The rate is seen to be independent of the free amine concentration, but inversely dependent on the concentration of the conjugate acid of the amine (amine perchlorate, designated as LH+) as shown in Figure V. Thus the complete rate-law is <i(Afi(I)) , (Afi(I)1 foO) - dt 8 8 kexp (LET] <3"5> where (3-6) 24 for Ethylamine Complete j (Lj- 0.2 H, p - 730 mmHg Referring to (3-1) and (3-5), some retardation of the rate as the reaction proceeds due to increase of (LH +) i s expected. Usually the i n i t i a l concentration of amine perchlorate was made s u f f i c i e n t l y high so that t h i s e f f e c t was n e g l i g i b l e and good f i r s t order plots were obtained. However, i n some cases at low (LH +), (Ag(I)) vs. t plo t s showed the expected retarda-ti o n i n l a t e r stages of the reaction (Expt. 1-a i n Figure I I I ) . In these cases the i n i t i a l l i n e a r portion of the rate p l o t s was employed for the determination of the rate constant. OTHER "STANDARD" SYSTEMS Among those amines which were investigated, methylamine, diethylamine, ethanolamine and diethanolamine showed the same type of k i n e t i c s as ethylamine, i . e . , f i r s t order dependence on Ag(I) and CO, independence of free amine concentration and inverse f i r s t order dependence on ammonium ion concentration. The experimental r e s u l t s expressed i n terms of the second order rate constant, k' (3-3) and k e X p (3-5) are summarized i n Table 4. Most of the k i n e t i c measurements were done at 25°C. Typical rate p l o t s f o r these systems are shown i n Figure VI. In some cases the f i r s t order rate plots exhibited some downward concavity i n the l a t e r part of the reaction. This may be att r i b u t a b l e to some zero order reaction of s i l v e r (I) with a l i t t l e impurity i n the amines or amine i t s e l f , or i t may be due to some heterogeneous reaction on s i l v e r metal. 26 0 500 1,000 1,500 2,000 2,500 Time, sec. Pig. 71. Typical Rate Plots for Ethylamine-type Complexes See Table 4 for experimental conditions 27 In those c a s e s s l o p e o f the i n i t i a l linear portions of the rate plots were employed to determine the rate constants. Mechanism The inverse dependence of the rate on (LH+) may be understood by taking account of the following equilibria which prevail in the solution. AgL,+ J^U AgL+ + t; tAfr+KL? = (3-7) * UgL.2; ; i L + H20 LH + + OH"; j j ° H " ' ' = 1^ (3-8) Referring to the stability constants of silver amine complexes which are summarized in Appendix I, i t is seen that the silver ions in these solutions are present predominantly as the bis-complex, AgL 2 +, so that (Ag(I)) & (AgL 2 +) (3-9) The resulting rate expression obtained from equations (3-5, 7, 8 and 9) is - 4<Mm . ^ ^-l K ^ - l (AgL+) (OH') (CO) (3-10) This may be identified with the following reaction mechanism AgL2 + H20 L-Ag-OH + LH (Rapid equilibrium) (i) 0 k II L-Ag-OH + CO > L-Ag-C-OH (Rate-determining step)(ii) 0 II L-Ag-C-OH + Ag(I) >• Products (Rapid) ( i i i ) The apparent rate constant of disappearance of Ag(I), k^p* defined by equation (3-5) and the bimolecular rate constant, k, 28 of the process (ii) are thus related through = 2kK = 2 1 ^ ^ ^ (3-11) 4- -where is the association constant of AgL with OH , i.e., 4- - Kk fL-Ae-OH} LAg 4- OH L-Ag-OH; (AgL+} (0H~) = \ ( 3 " 1 2 ) The factor of 2 reflects the fact that «ach rate-determining reaction results in the reduction of two silver ions. Hence, i t is more appropriate to express the rate of the reaction in terms of the rate of consumption of CO, i.e., the rate of the rate-determining step. Thus, ~ i L d | 1 = " % d ( d t ( I > ^ = k (L-Ag-OH) (CO) (3-13,a) = k^fAgL4") (OH"] (CO) (3-13,b) • W d , <3-13'c> Processes (3-7), (3-8) and (3-12), which are involved in equilibrium (i) or the process (i) itself is presumably suffici-ently rapid that (i) can be regarded as a pre-equilibrium. The overall stoichiometry requires that the reaction intermediate containing a CO molecule, and one silver ion, L-Ag-COOH, reacts with another silver (I) species (the identity of which is to be discussed later). However, this step ( i i i ) appears to be fast, compared with ( i i ) , so that the kinetics are fir s t order in Ag(I). 0 This intermediate complex, L-Ag-C-OH, is analogous to the one 29 which was previously proposed by Harkness and Halpern (3) as an intermediate complex in the reaction of Hg 2 +, i.e., -Hg-C^ -OH. Support for the structure of the latter was provided by the observation by Halpern and Kettle (6) that when methanolic solution of mercuric acetate takes up CO under similar conditions, a stable methylformate derivative, AcO-Hg-H-OCH^, analogous to the proposed complex was formed, isolated and spectroscopically identified. They also reported that attempts to prepare analogous CO adduct of silver acetate were unsuccessful, but this can pre-sumably be attributed to the poor solubility of silver acetate in methanol and instability of the CO adduct toward decomposi-tion into metallic silver. In terms of this mechanism i t might be expected that the rate constants k (and also kKn) should be relatively insensitive to the nature of L and thus that the large dependence of k^^ on the nature of the amine should reflect largely the variation of and K^ . This is shown to be the case in Table 5 where i t is seen that notwithstanding a 30-fold variation in k^p -3 -2 -1 (which ranges from 2.8x10 to 8.6x10 sec. , for the five amines under consideration) the value of kK^ is substantially 5 -2 2 -1 constant, (1x10 mole 1. sec. ) for a l l the systems. 30 TABLE 5 SUMMARY OF KINETIC AND RELATED THERMODYNAMIC DATA FOR "STANDARD" SYSTEM AT 25°C. 1) 2) 3) % xlO* kKh lc _____ xlO 4 xlO" 5 exp xlO 4 _ 9 mole -mole mole. mole 1? Amine sec."*" l . - l l . " 1 I"? sec." C2H5NH2 1.55xl0"2 1.2 6.5 7.6 1.0 CH3NH2 2.85xl0"2 2.9 5.2 15 0.9 (C2H5)2NH 8.6 xlO" 2 5.0 9.1 46 1.0 HOC2H4NH2 2.8 xlO" 3 2.8 0.55 1.5 1.0 (HOC3H4)2NH 3.1 xlO" 3 16 0.10 1.6 1.0 1) and 2) Refer to Appendix I. 3) Calculated by use of equation (3-11). The temperature dependence of. the rate constants was determined for a l l these systems over the temperature range 15 to 35°C. In a l l cases good linear Arrhenius plots were obtained, which are given in Figure VII. The activation parameters are summarized in Table 6. 3 1 3 . 2 3 . 3 3.4 3 . 5 T" 1, 10~3 dee'1 Fig. VII. Arrhenius Plots for Ethylamine-type Complexes 32 TABLE 6 APPARENT ENTHALPY AND ENTROPY OF ACTIVATION FOR "STANDARD" SYSTEMS 1) * A H AS 3 exp exp exp Amine Kcal. mole ** e.u.mole C 2 H - N H 2 7.8 xlO" 3 14.3 -20.0 CH_NH2 1.43xl0"2 15.5 -14.9 (C2H-)2NH 4.3 xlO" 2 14.1 -17.5 HOC2H4NH2 1.4 xlO' 3 17.4 -13.1 (HOC2H4)2NH 1.6 xlO" 3 18.3 - 9.9 -1 1) %k = kKuKj K, is used for the calculation of the activa-' - exp h d^ T> tion parameters. The present result shows that the only silver (I) species that is active toward CO under the conditions investigated for these five amines is the hydrolyzed mono-amine complex, L-Ag-OH while other silver (I) species involving the bis-complex, AgL2+, and free silver ion or aquo complex, Ag+, make a negligible contribution. Although metallic silver was precipitated during the course of reaction, its heterogeneous catalytic effect, at least during the early stages, was small. 33 TRIETHYLAMINE COMPLEX The triethylamine complex exhibited somewhat different kinetic behaviour, from the "standard" systems described above, notably in that the rate of reaction was no longer independent of the free amine concentration but exhibited an inverse depen-dence on the latter. The rate was inversely proportional to ammonium concentration but the effect of silver concentration was somewhat more complicated. The experimental results are summarized in Table 7 and typical rate plots are shown in Figure VIII. This complex exhibited the fastest overall rate of a l l the amine complexes investigated in this study, so that a 32.3% CO - 67.7% N 2 mixture was used in most of the experiments to obtain a reaction rate convenient for measurement. A fairly high concentration of triethylamine (up to 0.8 mole'l )^ was employed to prevent the hydrolysis of silver and precipitation of silver oxide because the triethylamine-silver complex is much less stable than the complexes of primary and secondary amines, while the basicity of the amine is almost the same. Because of the high vapor pressures of the resulting solutions Lattey's (8) data on the vapor pressure of aqueous triethylamine solution were used to calculate the partial pressure of CO. 34 T A B L E 7 R A T E O F R E A C T I O N O F T R I E T H Y L A M I N E C O M P L E X A T 2 5 ° C . 1) Amine Vapor C O Initial Amine Per- Pres- Pres- K , k o v n (Ag(I)h ( L ) chlorate sure sure mole'l exp Expt. mole°l mole-l"1- mole-l" 1 mmHg mmHg sec"*- sec"*- No. O o O l O 0.809 0.600 79 220 1.02 0.61 6a 0.010 0.802 0.480 79 220 1.31 0.63 6b 0.010 0.794 0.400 78 220 1.71 0.68 6c 0.010 0.784 0.300 78 220 2.27 0.68 6d 0.010 0.205 0.600 47 230 2.32 1.39 6e 0.010 0.250 0.600 51 229 2.16 1.30 6f 0.010 0.263 0.600 53 228 2.12 1.27 6g 0.010 0.309 0.600 57 227 1.91 1.15 6h 0.010 0.342 0.600 60 226 1.74 1.04 6i 0.010 0.412 0.600 65 224 1.62 0.97 6j 0.010 0.508 0.600 70 223 1.45 0.87 6k 0.010 0.619 0.600 74 222 1.20 0.72 61 0.005 0.080 0.300 34 234 5.1 1.53 6m 0.005 0.180 0.300 45 231 3.4 1.02 6n 0.005 0.180 0.200 45 231 4.9 0.98 6o 1) Reference (8). Although triethylamine was very carefully purified as described in the experimental section, there was an appreciable amount of i n i t i a l reduction of silver (up to 25% of total silver (I) concentration) which may be attributed to some impurity in the amine, and which introduced some complications, * The reaction mixture was usually stable toward the i n i t i a l reduction under nitrogen flow especially at high ammonium con-centration. This i n i t i a l reduction occurred usually after C O gas was introduced into the reaction vessel, as can be, seen in the rate plots (Figure V I I I ) . The amount of the i n i t i a l reduc-tion of silver was directly dependent: on the total amine concen-tration (free amine and ammonium) in the reaction solution. However, i t seemed to have no significant effect on the rate of the subsequent reduction by C O . 3 5 36 particularly in the determination of the effect of i n i t i a l silver concentration on the reaction rate, since at low i n i t i a l (Ag(I)) a large portion of Ag(I) was lost by the. i n i t i a l reduction while high (Ag(I)J could not be employed because of the insta-b i l i t y of the complex toward hydrolysis of silver ion and preci-pitation of silver oxide. Two alternative interpretations of the inverse dependence of the rate on the free amine concentration were considered. The fir s t of these involved the possibility that not only the mono-complexed silver ion but also uncomplexed ion (or aquo ion) contributes to the overall reaction to an appreciable extent. This failed to yield a rate-law which fitted the observed dependence of the rate on the silver (I) and amine concentrations in detail, and, furthermore, required that the reactivity of free silver ion, which is present in very low concentration compared to the mono-complex, be much higher (some 300-fold) than that of the latter. This could not be readily reconciled with the observed insensitivity of the rate of various other mono-complexes to the nature of the amines, and with the measure-ments of Peters and McAndrew (5) on the reduction of uncomplexed Ag + ions by CO in perchlorate media. Consequently this inter-pretation was rejected. A more satisfactory account of the. kinetic behaviour of this system was obtained in terms of the following mechanism in which the competitive reaction of the intermediate, L-Ag-COOH, 37 are considered. kL L-Ag-OH + CO v x- L-Ag-COOH ( i i ' ) k - l + k2 L=Ag-COOH + LA.g > Products (iv) Here, processes k ^  and are competing for the intermediate complex, L-Ag-COOH, the back reaction of the fir s t step (ii) no longer being negligible. Assuming steady state approximation for L=Ag-COOH, the kinetics are found to be _ d(CO) m , dCAn(I)) _ . tAg(I)] (CO) dt * dt % Kexp (LH+) . k l W r i k2VAg(I))/(L) k l K h K b K d L (LH*") k.j+kj^  CAg(D)/(L) * 3 1 4> k K , k K K2 (Afi(I)j2(CO) 1 ,~ , . V h VVSi, (LH+) k -CL)+k.K (Ag(I)j ^3-15> •L ~ 1 I. d^ kexp " ^ i V W ^ * * 1 " t . l t L W y (Ag(I)) ( 3" 1 6> Thus, k is no longer a constant. Taking reciprocals k t p " " 2 k l V k 2 K > (Ag(I)) fr^CL^K (Ag(I))) (3-17) d l From the equation (3-25), we expect a linear relation between 1/k^p and ( L ) at a fixed total silver ion concentration, (Ag(I)) . The plots of l/kov„ vs. ( L ) in Figure IX based on the experimental results in Table.7 are in complete accord with this. From equation (3-17) the intercept and the slope of the linear plots are given by 3 8 0 . 0 0 0 .2 0.4 0.6 0.8 ( L ) , mole l " 1 Pig. IX. Dependence of Rate on Free Amine Concentration at 25° for Triethylamine Complex 39 Intercept - 2fc g v K (3-18) 1 n D d. Slope = A „ Hr 1 „ rL/ T0 (3-19) k2 K d 1 [ A 8 ( I ) ) Thus, the intercept should be independent of (Ag(I)) and the slope should be inversely proportional to (Ag(I)) s in accord with Figure IX, which also shows k. _ to be independent of (LH+}. From the intercept and slopes of the plots in Figure IX, and using equation (3-18) and (3-19), the following values were obtained for the rate constants. k-^ - 2.5xl05 mole"2«12.sec."1 (3-20) k 2 ^ k - l = 3 ' 4 x l ° 3 mole"1'! (3-21) The value of kjK^ in this case is about 2.5 times as large as that found for ethylamine and related complexes. It is obvious that when the second term in the denominator of the equation (3-16) is much larger than the fir s t term, (i.e., k^K^ (Ag(I))^ k_^ L ), the overall kinetics approach those for the ethylamine complex. For some reason, in the case of triethylamine, these two terms are comparable in their magni-tudes. Since for triethylamine is actually larger (70 times) than for ethylamine, the observed kinetics must reflect either an abnormally large value of k_^ or an abnormally small value of k£. 40 TRIETHANOLAMXNE COMPLEX The only other tertiary amine that was investigated in this study was triethanolamine. The silver complex of this amine exhibited almost the same kinetics as ethylamine, the rate in this case being almost independent of the free amine concentra-tion;, with only a slight dependence in the opposite direction to that for triethylamine (i.e., the rate increasing with amine concentration) and inversely proportional to the ammonium concen-tration. The experimental results are summarized in Table 8 and typical rate plots are given in Figure X. The dissociation constant of the triethanolamine complexes is so large (Kd = 0.046) that It is no longer valid to approximate [AgL2 ) by (Ag(I)) . This would give rise to a dependence of the rate on the free amine concentration even i f the reaction followed the same kinetics as for the ethylamine complex. The observed dependence of the rate on the amine concentration is in the direction expected from this (i.e., the rate increases with the amine concentration), but is much smaller than predicted. This suggests that there is also superimposed upon this an inverse dependence of the rate on the free amine concentration, similar to that found for triethylamine complex. This is not unexpected in * Actually, using the value for K^., given above, the calcula-tion shows that more than 30% of the silver is present in the form of AgL+ at (L) = 0.1 mole«lm^0 41 TABLE 8 RATE OF REACTION OF TRIETHANOLAMINE COMPLEX Amine CO Initial (Ag(I)1 mole*I"1 Amine Per-L chlorate mo 1 e • 1" ^mo 1 e • 1 °° 1 Pres-sure nnnHg T°C. k« mole°*l»l sec'l. kexp sec'l Expt. No. 0.010 0.100 0.0500 705 40 0.23 1.2x10' -2 7a 0.010 0.100 0.100 705 40 0.13 1.3x10' -2 7b 0.007 0.100 0.200 705 40 0.069 1.4x10' -2 7c 0.07 0.100 0.100 705 40 0.13 1.3x10' -2 7d 0.007 0.200 0.100 705 40 0.13 1.3x10" -2 7e 0.007 0.040* 0.0500 730 25 0.049 2.4x10" -3 7f 0.007 0.100* 0.0500 730 25 0.060 3.1x10' -3 7g 0.007 0.190* 0.0500 730 25 0.062 3.1x10' -3 7h 0.014 0.040* 0.0500 730 25 0.051 2.6x10" -3 7i 0.007 0.040* 0.0250 730 25 0.091 2.3x10' -3 7j 0.007 0.100 0.0500 728 30 0.106 5.3x10' -3 7k 0.007 0.100 0.0500 718 35 0.17 8.5x10' -3 71 * Free amine concentrations corrected for the dissociation of silver complex. view of the similar stability constants of the two complexes. The dissociation constants of both complexes are much larger than that of the ethylamine complex and in both cases the order of the first and second dissociation constants (K^^ ) is the reverse of that for ethylamine and other primary and 42 -1.8 Log(Ag(l)] 0 1,000 2,000 3,000 4,000 5,000 Time, sec. Fig. X. Typical Bate Plots for Triethanolamine Complex - i — See Table 8 for experimental conditions. 43 secondary amines (Appendix I). Applying the same mechanism to triethanolamine, as previously to triethylamine, and assuming (Ag(I) ) = UgL 2 +) + [AgL +) (3-22) (3-23) Kd + CD d l and neglecting other silver species, then the rate law is expected to be of the form , d ( A n ( I > ] . (Afi(I)_j (CO) * dt . * kex P f u r ) dlCOJ _ dt ^ ^ ""exp k2(AgL+) k l K h(AgL +) (OH") (CO) ^J^^ (3-24) 2 s f D k^(K d +(Lj)^fk 2K d (Ag(I)}(Kd+(L)) 1 1 1 (3-25) It was not possible to make sufficiently accurate kinetic measurements to test this equation in detail and determine a l l the rate constants involved. However, an attempt was made to estimate the rate constant k^ from the experimental rate constants. It is seen from equation (3-25) that when ( D is very small, the first term in the denominator becomes negligible compared to the second term and the kinetics approach the following form. = k l K h K b K d 1 f AeCLH^f)C°'1 K,f+(L) (3-26) d[CO dt 44 Among the experimental data Expt. 7-f is the one done at lowest ( L ) (=0.040 mole*!"*-). For the experimental condition of this experiment the ratio of the two terms in the denominator of (3-25) was calculated to be k - l ( Kdj + (IQ) 2 k 2 K^(Ag<I))(K d+(L]) " °' 0 8 ( 3 " 2 7 ) using the value of given in Appendix I (K^ = 0.046) and 1 1 assuming the same value of k.^/k^ as for triethylamine given by (3-21). This ratio seems to be sufficiently small to approxi-mate the kinetics at these experimental conditions of Expt. 7-f by (3-26). (Simple comparisons of the numerical values of (L)/(K, + ( L ) ) with these experimental results in Table 8 show that below ( L ) = 0.1, the results agree fairly well with the kinetics given by (3-26) but deviate from i t very rapidly as ( L ) increases. This is expected since the f i r s t term in the denominator of (3-25), which was neglected in (3-26), is second order in (L).) From (3-26) k ^ is given by (3-28). k-K. = kexp d l (3-28) Using the data of the Expt. 7-f, k^K^ was estimated to be 5 =2 2 -1 1x10 mole *" 1 sec . Although this is a very rough estimate, i t agrees with that of ethylamine (kK^ = 1.0x10"*), at least in order of magnitude. A further expected feature of the kinetics in the observed 45 curvature of Arrhenius plot (Figure XI) of the apparent rate con-stant, k j at (Lj = 0.100 mole-l" 1. As can be seen from the * exp' several experiments at 40°C. (Table 8) the kinetics at this temperature were closer to the "standard" ones than at 25°C. However, at (L) - 0.1 and at 25°C. the kinetics appear to be close to those for ethylamine and the two other ethanolamine complexes and the apparent energy of activation under this condi-tion and at this temperature is also very close to those for two * -1 other ethanolamines ( AE ~ 18 Kcal. mole ). This mechanism, comprising a sequence of ( i i 1 ) and (iv), appears to give a satisfactory account of the kinetic behaviour of the triethylamine and triethanolamine complexes, although only a semi-quantitative discussion was possible for the latter. As mentioned previously, the necessary condition for this mechanism to hold is that either k_^ is abnormally large or k 2 abnormally small compared to those for ethylamine. This behaviour may be related to the fact that the dissociation constants (especially K^) of the silver complex of tertiary amines are abnormally large (K^ is 70 times for triethylamine and 400 times for triethanolamine complex as large as that of ethylamine; Appendix I), that i s , their bis-complexes seem to be abnormally unstable toward the loss of the second coordinated group. A similar instability of the intermediate complex L-Ag-COOH toward decompo-sition could account for an abnormally large value of k_^. These observations provide some information about the 46 - 1 . 0 -1 .5 Log k exp - 2 . 0 - 2 . 5 3.1 Ethylene-,! amine ChH+] - 0.100 mole 1" (L) - 0.200 mole 1* Trie thanoland ne (LH+J - 0.050 (L ) - 0.100 mole 1 3 .2 3.3 T"1, 1 0 " 3 degT1 3.4 Pig. XI. Arrheniue Plots for Triethanolamine and Ethylenediamine Complexes 3.5 47 second step of the reaction, beyond that which could be deduced from the results of the ethylamine-type systems„ In particular, they indicate that the second silver species which reacts with the intermediate complex is also a mono-complexed species. This point is to be discussed again later. AMMONIA COMPLEX The ammonia complex exhibited the slowest overall reaction rate of a l l the amine complexes investigated in this study. Most of the rate measurements for this system were made in an autoclave at 30°C. and at high CO pressure (up to 20 atm.). The experimental results are summarized in Table 9 and some typical rate plots are given in Figure XII. The dependence of the rate on the ammonium ion concentration was complex. The log k' vs. log (LH+) plot in Figure XIII shows that the rate is inversely proportional to (LH ) at low ammonium ion concen-trations but that at higher ammonium ion concentrations the inverse order increases to two or higher. This implies that the reaction mechanism may be different in the two regions or the rate is controlled by different steps. The effects of silver ion concentration, CO pressure and free ammonia concentration were studied at both low (0.02 mole'l"''") and high (0.1 mole»l~*) NH^+ concentration and these results also are summarized in Table 9. In both NH^  region, the reaction rate was first order in CO and only slightly 48 TABLE 9 RATE OF REACTION OF AMMONIA COMPLEX Initial CO Ammonia Ammonium (Ag(I);) Pressure CU (LR+) mole°l atm. mole-1" mole•1~1 T°C. k* mole=^»l sec'^- sec exp 1 Expt, No. Effect of Ammonium 0.010 9.2 0.180 0.0200 30 4.6x10 -2 9.2x10 -4 8a 0.010 9.2 0.180 0.0236 30 3.3x10 -2 7.4x10 »4 8b 0.010 9.2 0.180 0.0300 30 2.8x10 -2 8.4x10 -4 8c 0.010 9.2 0.180 0.0500 30 1.65x10 -2 8.3x10 -4 8d 0.010 9.2 0.180 0.0700 30 7.6x10 -3 5.3x10 -4 8e 0.010 9.2 0.180 0.0850 30 8.3x10 -3 7.0x10 -4 8f 0.010 9.2 0.180 0.100 30 5.5x10 -3 5.5x10 -4 8g 0.010 9.2 0.180 0.125 30 5.0x10 -3 6.3x10 -4 8h 0.010 9.2 0.180 0.150 30 3.1x10 -3 4.7x10 -4 81 0.010 9.2 0.180 0.175 30 1.8x10 -3 3.1x10 -4 8j 0.010 9.2 0.180 0.200 30 1.2x10 -3 2.4x10 -4 8k Effect of CO Pressure 0.010 4.1 0.180 0.020 30 5.3x10 -2 10.5x10 -4 81 0.010 9.2 0.180 0.020 30 4.6x10 -2 9.2x10 -4 8a 0.010 14.1 0.180 0.020 30 4.9x10 -2 9.8x10 -4 8n 0.010 19.2 0.180 0.020 30 5.0x10 -2 10.0x10 -4 8o 0.010 4.1 0.180 0.100 30 5.7x10 -3 5.7x10 -4 81' 0.010 9.2 0.180 0.100 30 5.5x10 -3 5.5x10 -4 8g 0.010 14.1 0.180 o.ido 30 5.7x10 -3 5.7x10 -4 8n' 0.010 19.2 0.180 0.100 30 7.2x10 -3 7.2x10 -4 8o» Effect of Initial Silver Ion 0.0047 9.2 0.180 0.020 30 4.2x10' -2 8.5x10' -4 0.0047 9.2 0.180 0.020 30 4.6x10' -2 9.2x10" -4 0.0195 9.2 0.180 0.020 30 4.8x10' -2 9.6x10' -4 0.0048 9.2 0.180 0.100 30 2.6x10 -3 2.6x10" -4 0.0099 9.2 0.180 0.100 30 5.5x10' -3 5.5x10" -4 0.0196 9.2 0.180 0.100 30 9.8x10' -3 9.8x10" -4 8p 8a 8q 8p' 8g 8q» 49 TABLE 9 (Continued) Initial CO Ammonia Ammonium k° k (Ag(I)} Pressure (L j (LH+) mole"1-! e x p Expt mole•I"1 atm. mole-l 1 mole4!"1 T°C. sec"l sec" 1 No. Effect of Free Ammonia 0.010 9.2 0.080 0.020 30 3.9x10 -2 7.8x10 -4 8r 0.010 9.2 0.180 0.020 30 4.6x10 -2 9.2x10 >4 8a 0.010 9.2 0.280 0.020 30 3.1x10 -2 6.2x10 -4 8s 0.010 9.2 0.080 0.030 30 2.7x10 -2 8.2x10 =4 8r" 0.010 9.2 0.180 0.030 30 2.8x10 -2 8.4x10 -4 8c 0.010 9.2 0.280 0.030 30 2.2x10 -2 6.6x10 -4 8s" 0.010 9.2 0.080 0.100 30 0.010 9.2 0.180 0.100 30 0.010 9.2 0.280 0.100 30 Effect of Temperature 0.010 9.2 0.180 0.050 25 0.010 9.2 0.180 0.050 30 0.010 9.2 0.180 0.050 35 0.010 9.2 0.180 0.050 40 0.010 9.2 0.180 0.050 50 7.6xl0-3 7.6x10-4 8re 5.5x10-3 5.5x10"4 8g 5.1x10*3 5.1x10-4 8s» 1.06x10-2 5.3xl0"4 8t 1.4x10-2 7.2x10-4 8d 2.0xl0~2 10.2x10-3 8u 3.1x10*2 1.5x10-3 8v 6.6x10-2 3.3xl0"3 8w dependent on the free NH^  concentration. No trend, however, was discernible in the latter dependence. The dependence on (Ag(I)j was first order at low (NH^+) and second order at high (NH4+) (i.e., k' was proportional to (Ag(I)).) This suggests that at high (NH^  ) the rate is controlled in part by the second step in which the second silver (I) species takes part. The high inverse order dependence on [ NH^  J in this region further suggests that the second silver species also is hydrolyzed. 50 Time, sec. Pig. XII. Topical Rate Plots f o r Ammonia Complex See Table 9 for experimental conditions. -1.0 51 1.5 -l.o -0.5 Log [LH+] Pig. XIII. Dependence of Rate on Ammonium Ion Concentration at 30° for Ammonia Complex; [ L j = 0.200 mole l " " 1 52 The following mechanism is consistent with these observa-tions t k l . L-Ag-OH + CO ^  L-Ag-COOH ( i i ' ) -1 k 2 L-Ag-COOH + L-Ag-OH > 2Ag + C02 + 2L + H_0 (v) where two processes with rate constants k_^ and k2 are competing for intermediate complex, L-Ag-COOH. The kinetics, assuming steady state concentration of L-Ag-COOH, are thus of the follow-ing type. - "^T1 " " k ^ I P 1 " CAg(X)) (CO)/(LH+) = k1k2K^Kg41(Ag(I))2(C0) " (LHT)^k^^K^K^(Ag(I))/tLH*")} (3-29) At low (NH^ where k_x « k ^ I ^ K ^ (Ag(I))/[LH+), the kinetics approach those for ethylamine while at high (NH^4) where k_^» k2KjiKbKd^(Ag(I))/(LH ], the reaction is second order on (Ag(I)) and inverse second order on (LH+J, i.e., L o w ( L H + ) 8 d M m V h V ^ J A i g ^ f i i ( 3. 3 0) . ,TW+, d(00) 2 2 2 (A f i(I)) 2fCOj High (LH ) s - d t = k - 1 K ^ K ^ (LH+J- ( 3 _ 3 1> From the experiments at low (LH+) (0.020 mole l " 1 ) kexp - 2 k l W d l = 9 X 1 0 ' 4 S e C " X ( 3 _ 3 2 ) Using the values for and Kd^ at 30°C. in Appendix I, 53 (3-33) On the other hand at high (LH*") k e X p corresponds tos kexp - ^ 1 K ^ K ^ J A ^ I J ( 3 „ 3 4 ) The experimental rate constant at (LH ; = 0.100 mole"1 (Expt9s. 8p°, 8g, 8q') exhibits a good fir s t order dependence on (Ag(I)) and gives a constant value for k /(Ag(I)) (5.4, 5.6 -9 -1 -1 and 5.0x10 A mole •l-sec x, respectively), i.e., at (LH+) = 0.100 mole'l" 1 2 2 2 k 2k, k 9 KhKbKdT _9 , ( A i ^ r • "WT • 5 X 1 0 2 - l - - 1 - ! - ^ " 1 0-35) Substituting (LH+J = 0.100 mole l * * 1 and using (3-32) and the values for and again we get k2 Kh k i -1 2xl09 mole"2.!2 (3-36) Al l these results are for 30°C. This temperature was chosen for the most of kinetic studies instead of 25°C. because of difficulty in controlling the autoclave temperature at 25°C. To compare the rate constant with those of other systems, i t is necessary to correct i t to 25°C. Arrhenius plot at (LH+) • 0.050 mole 1 * given in Figure XIV also reflects the complex nature of the kinetics of this system. It was difficult to measure the temperature effect at lower (LH*) than this because 54 55 of the fast rates and poor rate plots. The Arrhenius plot i n Figure XIV is linear at high temperature but is concave upwards at low temperatures. This indicates that the linear portion represents the temperature dependence of the apparent rate constant of the form given by (3-34) and not (3-32). The appa-rent energy of activation calculated for linear portion of the Arrhenius plot is 13.7 Kcal. mole \ which is very close to the values for k (^kH^K^R^ ) for other systems (14-18 Kcal. ~1 ^ mole , Table 6). It thus seems reasonable to calculate the rate constant given by (3-32) for 25°C. using a temperature co-* »i efficient corresponding to AH «= 14-18 Kcal. mole . . Because the temperature interval involved is small (5°C), the error involved in the extrapolation can hardly be very large. Thuss from (3-32) exp - 210^1^^ - 6xl0~3 sec" 1 (at 25°C.) (3-37) Again, using the values for K and K, for 25 C. in Appendix I. b °i Vh " - 1 « * x l ° 5 mole'2»12-sec"1 (3-38) This is very close to the values previously found for ethylamine and related systems (1.0x10^). + -1 At (NH^ ) exceeding 0.1 mole'l , the dependence of the rate on (NH^+) seems to exceed inverse second order (Figure XIII). However, in this region the actual rate of the reaction was extremely slow so that the measurements are considered unreliable. The mechanism involved here is substantially the same as that derived previously for the two tertiary amines«, triethylamine and triethanolamine9 apart from a difference in the nature of the Ag(I) species participating in the second step of the reaction (L-Ag-OH and LAg , respectively). No explanation for this difference is available. Recently, Peters and McAndrew (5) reported the kinetics of the reduction of silver in perchlorate media at 70°C. to be In this case the second silver (I) species reacting with the intermediate complex is an unhydrolyzed ion analogous to LAg in the tertiary amine cases. The slowness of the second step in the case of the ammonia complex which leads to the departure from the simple ethylamine-type kinetics may be due to the much lower basicity of ammonia (i.e., to a smaller L-Ag-OH concentration). However, the failure of diethanolamine, which is even less basic, to exhibit the same type of kinetics as ammonia, throws some doubts on this. DIAMINE COMPLEXES In addition to the above monoamines, three primary diamines, ethylenediamine, 1,3-diaminopropane and 1,4-diaminobutane were examined. The experimental results are summarized in Table 10 (3-39) which was interpreted by the following mechanism Ag + + CO + H20 -=_± AgCOOH + H + (Rapid equilibrium) AgCOOH + Ag + > 2Ag + C0o + H + (Rate determining) 57 TABLE 10 RATE OF REACTION OF DIAMINE COMPLEXES Initial CO Amine 8 (Ag(I)) Pres- Amine Per- k k sure (L) chlorate mole"1.! exp Amine mole.I"1 mmHg moleX 1mole ol- 1T°C. sec™1 sec".1 No. Ethylene-diamine 1,3-Diamino-propane 1 s4= Diamine-butane 0.013 705 0.013 705 0.013 705 0.012 705 0.020 705 0.020 508 0.020 239 0.010 741 0.010 742 0.010 730 0.010 728 0.010 712 0.010 699 0.010 683 0.010 664 0.010 736 0.010 736 0.010 736 0.005 736 0.010 736 0.010 736 0.010 736 0.010 736 0.010 736 0.010 736 0.010 736 0.010 736 0.010 736 0.010 736 0.010 736 0.010 736 0.300 0.100 0.300 0.300 0.100 0.100 0.200 0.200 0.200 0.100 0.200 0.100 0.200 0.100 0.200 0.100 0.200 0.100 0.200 0.100 0.200 0.100 0.200 0.100 0.200 0.100 0.200 0.100 0.200 0.100 0.183 0.100 0.183 0.050 0.183 0.025 0.192 0.100 0.0356 0.025 0.0643 0.025 0.0840 0.025 0.183 0.025 0.282 0.025 0.190 0.0177 0.190 0.0250 0.190 0.0354 0.090 0.0250 0.050 0.0250 0.010 0.0250 0.005 0.0250 40 0.34 40 0.112 40 0.34 40 0.17 40 0.34 40 0.34 40 0.34 15 0.132 20 0.193 25 0.216 30 0.220 35 0.228 40 0.354 45 0.480 50 0.668 25 7.1xl0-2 25 1.6X10"1 25 2.8x!0<=1 25 7.1x10-2 25 1.7x10-1 25 2.2X10"1 25 2.5X10"1 25 2.8x10"1 25 3.0x10-1 25 4.3x10-1 25 3.2x10-1 25 2.3x10-1 25 3.3x10-1 25 3.4x10-1 25 2.6X10"1 25 2.4X10"1 0.034 9a 0.034 9b 0.034 9c 0.034 9d 0.034 9e 0.034 9f 0.034 9g 0.0132 9h 0.0193 9i 0.0216 9j 0.0220 9k 0.0228 91 0.0354 9m 0.0480 9n 0.0668 9o 7.1x!0"3 10a 8.2x10-3 10b 7.1xl0"3 10c 7.1x10-3 10d 4.2x10-3 I0e 5.5x10-3 I0f 6.1x10-3 log 7.1x10-3 lOh 7.6x10-3 10i 7.6xl0-3 11a 8.1x10-3 lib 8.3x10-3 He 8.3x10-3 Hd 8.4x10-3 l i e 6.4x10-3 l l f 6.0xl0°3 l i g and some typical rate plots are given in Figure XV. The ethylene-diamine complex was first investigated at 40°C. and found to 58 0 1,000 2,000 3,000 Time, sec. Pig. XV. Typical Rate Plots for Diamine Coplexes ••• See Table 10 for experimental conditions. 59 exhibit kinetics similar to those of the ethylamine complex as can be seen from Table 10. However, the study of the temperature dependence of the rate revealed more complicated behaviour at lower temperature and the results failed to yield a linear Arrhenius plot (Figure X V I ) . Attempts to elucidate the kinetics at 25°C. and obtain directly the rate constant at this tempera-ture for comparison with the other amine complexes were unsuc-cessful. Thus the rate of this system at this temperature exhibited a very complicated dependence on amine, amine perchlor-ate and silver ion concentration which was not altogether reproducible. At higher temperature, however, (above 35°C.) the system appeared to be well behaved. Extrapolation of the linear portion of the Arrhenius plot in this region yielded a value of 1.07x10 * sec for at 25°C. Using the values for and of ethylenediamine in Appendix I , we get kK. = kexp = 0.30xl05 mole' 2.l 2.sec" 1 (3-40) 2 K b K d ] L This is only about one-third of the value for ethylamine and related amine complexes. The silver complex of 1,3-diaminopropane has an abnormally large first stability constant (KQ-J_ 88 8.9x10"*) as can be seen in Appendix I , and there has not been reported any data on its second stability constant. This is due to stabilization of 61 the mono-complex by chelation, with the result that there i s l i t t l e tendency to add a second amine molecule to form the bis-complex. This great s t a b i l i t y of the monochelate complex i s reflected in the extremely large f i r s t s t a b i l i t y constants the second s t a b i l i t y constant, presumably being very small s so that the approximation9 [Ag(I)) = CAgL^J i s obviously invalid in this case. Therefore, equation (3-41) (as in the case of triethanolaminej, which also has a small K^£» i.e., large Kd^) must be used for the concentration of AgL + . Assuming (AgL +) + (AgL 2 +) = (Ag(I)) (3-22) [AgL^) (AgU*" )(L) = K12 = K where K i s the second s t a b i l i t y constant of the silver complex. Then i f ethylamine-type behaviour applies also in this case, the overall kinetics w i l l be as follows: COl _ k J jA^j I lL _ fgxp_ CAR(I)) (COj dt * dt 2 TLH+I Then = " V b (1+KCL))(LH+) ( 3 4 1 ) kexp 8 5 2 k K h K b 1+K(L) (3-42) Taking the reciprocal of k^^ >4r = 2 E^(lT7 +«) 62 The experimental free amine concentration was first calculated assuming and on this basis the experimental results for different (L Jwere plotted as 1/k vs. 1/(L). The resulting plot, which was not quite linear but slightly concave upwards9 was used to obtain a rough estimate of K, which in turn was used to improve the free amine concentration. This procedure was repeated to self-consistency and yielded a good linear plot of l/k e Xp vs. l/(L) shown in Figure XVI. The con-stants derived from this, using the value of in Appendix I are kK^ = 5.1xl02 mole" 2-l 2osec" 1 (3-44) K - 26 mole"1.! (3-45) This value of kK^ is only 1/200 of the "normal" value for ethyl-amine, etc. (kK^ = 1.0x10^). The much lower reactivity of this complex presumably reflects blocking of the reaction site by chelation, i.e. H 0 N — A g + — N H 2 HoC - , C H n H 2 The rate constant of ethylenediamine is also smaller than that of ethylamine, etc., but the reduction factor is only 1/3 in this case. Presumably this also is attributable to chelation but in this case the chelation tendency is much smaller than 4. 1,3-diaminopropane, because of the preference of Ag for linear coordination which, for steric reason, is more readily realized with 1,3-diaminopropane than with ethylenediamine. This is 63 reflected also in the corresponding stability constants of the complexes. Thus the fi r s t stability constant of ethylenediamine silver complex, while larger than that of ethylamine, is much smaller than that of 1,3-diaminopropane. These data are summa-rized in Table 1 1 . TABLE 1 1 RATE OF DIAMINE COMPLEXES AND THEIR STABILITY CONSTANTS kKjXlO5 _2 mole Log K Q 1 Log K 1 2 Log K Q 2 1*. 3.37 3.93 7.30 1 4.62 2.92 7.54 0.3 5.77 1.42 1) 7 . 1 9 1 ) 0.005 sec"*1 CH3CH2NH2 H2NCH2CH2NH2 H NCH CH CH NH 2 2 2 2 2 1 ) Results from the present study. Other stability constants are from Appendix I. It is seen that K^2 and kK^ both of which should reflect (inversely) the chelating tendency of the mono-complex indeed follow closely parallel trends. 1,4-diaminobutane which also has an abnormally large KQ^ value, similar to that for 1,3-diaminopropane, was expected to show the same type of behaviour as the latter. However, the rate in this case was almost independent of the amine concentra-tion and the kinetics were similar to those for ethylamine, although the actual overall rate of reaction (k e x p) w * 8 almost the same as for 1,3-diaminopropane (cf. Table 10). cance of this behaviour is not understood. 64 The signifi-GENERAL DISCUSSION A common feature of the systems examined in the course of this study is that in every case CO apparently reacts with a species of the composition L-Ag-OH. The i n i t i a l reaction in each case can be represented as L-Ag-OH + CO k > L-Ag-COOH ( i i ) The rate constant, k, of this process could not be measured directly, but the data yielded values of kK^. In some cases, the back reaction of (ii) was sufficiently fast to compete with the subsequent reaction of the intermediate, L-Ag-COOH, with another silver ion. , 1 L-Ag-OH + CO „ L-Ag-COOH ( i i 1 ) k - l A l l the values of kK^ (or k^K^) for various amine complexes investigated in this study are summarized in Table 12, together with the values of K s^ Kd^ and kK^K^K^. The value of the latter is identical to %kfiXp in the case of ethylamine-type "standard" amine complexes. For a l l the mono-dentate amines i t is seen that * kK^ is substantially independent of the nature of the amine. * Triethylamine is the only case where the deviation of the va-lue of kK^ from the "standard" value of 1x10^ appears to l i e out-side experimental error. For this amine the values of and both of which were used for the determination of kKjj were less precise than those for the other amines; only one significant f i -gure was available from the literatures. Furthermore, in this sys-tem i t was necessary to use high (LH+} in most of the experiments so that ionic strength effects which have not been taken into account may be important. 65 TABLE 12 SUMMARY OF KINETIC AND RELATED THERMODYNAMIC DATA Amine xlO 2 sec" 1 xlO 4 mole'l" 1 xlO 4 mole'l" 1 xlO 8 xlO" 5 mole2°1"2 mole°°2.l2. sec" 1 NH3 0.03 1.2 0.18 0.22 1.4 CH3NH2 1.43 2.9 5.2 15 0.9 C2H5NH2 0.78 1.2 6.5 7.6 1.0 H0CoH.NHo 2 4 2 0.14 2.8 0.55 1.5 1.0 (C2H5)2NH 4.3 5.0 9.1 46 1.0 (HOC2H4)2NH 0.16 16 0.10 1.6 1.0 (C 2H 5) 3N 120 80 5.9 470 2.5 (HOC2H4)3N 0.36 460 0.0079 3.6 1 H2NCH2CH2NH2 0.5 12 1.5 18 0.30 H2N(CH2)3NH2 0.4 380 4.4 1700 0.0025 The constancy of kK^ over a 2,000 fold variation of K^K^ (and hence of kK^K^K^) is striking. The constant kK^ may be identified with the rate constant of the alternative and kinetically equivalent representation of ( i i ) , i.e., with the rate constant of the termolecular process ( i i " ) , , kKh LAg + CO + OH" =^ L-Ag-COOH ( i i " ) It seems likely that K^  also is insensitive to the nature of 66 L, and indeed probably does not differ greatly from the hydrolysis + 2 ~1 constant of the free Ag ion whose value is about 2x10 mole 1 (11). Using this value for K^ , k is estimated to be about o - l - l 5x10* mole "l.sec . This participation of hydroxide ion in the reaction (base catalysis) accounts for the low reactivity of Ag + toward CO in acidic media. The insensitiyeness of the reactivity of L-Ag-OU to the nature of L, suggests that the amine molecule in L-Ag-OH is acting only to solubilize AgOH and prevent precipita-tion of silver oxide. The enhancement of reactivity in these amine-buffered systems would appear to be due mainly to the high pH, rather than to specific complexing effects. On the other hand, the rate of the back reaction of step (ii) and the rate and nature of the second step of the reaction do appear to vary with the nature of amine, L. This is shown particularly by a comparison of the ammonia and triethylamine complexes. * The apparent enthalpy and entropy of activation, A H exp and A for k^^ « kK^ K^ K^ ^ for "standard" amine complexes listed in Table 6 correspond to A H e x p " A H * + A H h + A H b + A H d i (3-46) A s e x p - ^S* + A S h + AS b + A S d i (3-47) where AH and AS correspond to the enthalpy and entropy of activation of the bimolecular process (ii ) and the other terms with subscripts h, b and d^ correspond to enthalpy and entropy 67 changes of the following equilibrium processes. h : LAg+ + OH" v LAgOH - A H ^ - A $ h (3-12') b s L + H20 v LH + + OH" - AH^ - A S b (3-8") d± : AgL 2 + LAg+ + L - A H ^ - A S ^ (3-7B) Few thermodynamic data relating to these processes are available. Only in the case of ethylamine, has i t been possible to obtain fairly reliable values of A H ^ , A H d , A S ^ and A S d 9 using 1 1 available data (11, 12). The values of A H ^ and A H d ^ thus obtained for ethylamine are 0.7 and 6.4 Kcal. mole"1 and A S ^ and A S ^ are -12.2 and 3.6 e.u. respectively. The values of A H^  and A S ^ as well as that of K h can be approximated to that for the free Ag + ion as pointed out previously, which are esti-mated to be about -2 Kcal. mole"1 and 4 e.u., respectively. Hence, the enthalpy and entropy of activation of the bimolecular process (ii) are estimated to be A H * ~ 9Kcal. mole"1 and if n AS <~ -15 e.u. (the bimolecular rate constant being k ~ 5x10 mole"1 -I"sec" 1 at 25°C). Although reliable values for AIL^, A S d _ , etc. for other amines are not available, i t is expected 1 •k & that A H and A S for these systems will not differ greatly from those for ethylamine. Peters and McAndrew (5) have recently reported the following kinetics for the reaction of aqueous silver acetate with CO in acetate-buffered acidic media at 90°C. and high CO pressure. 68 + k K (Afi+)(A^OAc)(CO ) 3 c [Kf-) , (3-48) and interpreted these in terms of the following mechanism, AgOAc + CO =^ ->AgCOOAc (slow) (vi) AgCOOAc 4- Ag(I) > Products (rapid) (vii) Ag + 4- CO 4- H20 K c AgCOOH 4- R"*~ (rapid) (viii) AgCOOH 4- Ag + — k 2 > 2Ag 4- C0 2 4- H + (slow) (ix) AgCOOH 4- AgOAc 2Ag 4- C0 2 4- HOAc (slow) (x) In this case there appears to be a contribution to the reaction, not only from AgOH but also from AgOAc. The process (viii) which corresponds to ( i i ) , ( i i 8 ) or (ii") in the case of amine complexes, is faster than the subsequent steps and thus corres-ponds to a pre-equilibrium. Hence, i t is not kinetically dis-tinguishable whether this is a base-catalyzed process as in the case of amine complexes or rather an acid-inhibited process; in other wordsi whether the silver species which is directly reactive toward CO molecule is a hydrolyzed species, AgOH, or an unhydrolyzed ion Ag . In the case of amine-buffered solution L-Ag-OH was seen to be the only reactive species. It is of great interest, therefore, to see i f the contribution of unhydrolyzed Ag + is detectable in the acetate-buffered solution. It can easily be seen that rate constants and equilibrium constants of these two processes, base-catalyzed (ii") and acid-inhibited (viii) are related as follows. 69 k, K k = FT" (3-49) - l Vw Suppose the base-catalyzed process (ii") is the only process contributing to equilibrium ( v i i i ) , then the kinetics observed by Peters and McAndrew requires that ko (Ag"*1 2l J «1 (3-50) k - l and hence k l K f a k ^ 8 ^ «i k ^ (3-51) From the data of Peters and McAndrew the value of k^Knk2(Ag+J/k=^ + 6 —2 —1 (= k^^fAg J/K^) can be estimated to be about 10 mole »l«sec at 90°C. On the other hand the insensitiveness of the kK^ (rate constant of the base-catalyzed process (ii") ) for amine complex to the nature of L over a 1,000-fold variation in its basicity suggests that the reactivity of uncomplexed Ag + for the same type of process should also be close to that of amine-complexed * species. Using the value of kK^ at 25 C. and the previously determined temperature coefficient, the value of kK^ at 90°C. 5 - 2 2 - 1 can be estimated to be 9x10 mole »1 "sec . These results are not in accord with (3-51). This suggests that in the case of * The value of and A are not reliable so that kK^ is to be used for comparison. The value for A H* + ethyl-amine has been estimated to be 7.2 Kcal. mole"1. This tempera-ture coefficient being small and fairly reliable, the resulting kKh for 90°C. is also fairly reliable. 70 acetate-buffered solution there may indeed be a reaction path involving direct reaction between unhydrolyzed Ag + and CO represented by ( v i i i ) . However, the rate constant of this process is not obtainable from these data. This possibility that the unhydrolyzed Ag + ion, may also be reactive toward CO is not altogether unexpected, since, as mentioned earlier, i t has already been found that in acidic 2+ solutions the reduction of Hg by CO proceeds through reaction with the unhydrolyzed ion (3). That such a path (i.e., OH -independent) is not observed in the case of the silver amine complex may simply be due to the high pH of the solutions, resulting in enhancement of the OH -dependent path. Another interesting result of Peters and McAndrew's work is the evidence suggesting that AgOAc is also reactive toward CO, presumably through an intermediate complex analogous to L-Ag-COOH, i.e., Ag-COOAc. In this reaction the first step (vi) is rate-determining with a rate constant at 90°C. esti-mated to be -2 -1 -1 k^ = 3.6x10 mole 'l-sec . 3 -1 -1 This compares with a value of 2.5x10 mole ^l' s e c estimated for the corresponding rate constant of L-Ag-OH toward CO, i.e., L-Ag-OH appears to be about 10"* times as reactive toward CO as AgOAc. The activation energy of the process (vi) has been determined to be 15 Kcal. mole"1 while that for the process (ii) -1 is 9 Kcal. mole 71 REFERENCES 1. Halpern, J., Advances in Catalysis. XI, 301 (1959) 2. Halpern, J . and Taylor, S. M., Disc. Faraday Soc, 29, 174 (1960). 3. Harkness, A. C. and Halpern, J., J. Am. Chem. Soc, 83, 1258 (1961). 4. Bauch, G., Pawlek, F. and Plieth, K., Z. Erzbergbau und Metallhutenwessen, XI, 11 (1958). 5. McAndrew, R. T. and Peters, E., X l l l t h International Congress of Pure and Applied Chemistry, Montreal, Canada, August, 1961 and unpublished results. 6. Halpern, J. and Kettle, S. F. A., Chem. Ind., 668 (1961). 7. Just, G. and Kauko, Y., Z. phisik. Chem., 82, 71 (1913). 8. Lattey,' J. Am. Chem. Soc, 1959, 29 (1907). 9. Von Georg-Maria Schwab, et al., Z. anorg. Chem.,, 252, 205 (1944). 10. Seidel, "Solubilities of Inorganic and Metal Organic Compounds" 3rd Edition, Vol. 1, p. 217 (1940). 11. "Stability Constants of Complex Salts", special publication of the Chemical Society. 12. "Selected Values of Chemical Thermodynamic Properties", U.S. Bureau of Standards circulation. 72 APPENDIX I SELECTED THERMODYNAMIC PROPERTIES OF AMINES AND SILVER-AMINE COMPLEXES *2 *4 *1 Stability Constants*3 d l Amine pKa LogK-^ LogK12 LogK2 xlO 4 xlO 4 NH3 9.25 3.31 3.92 7.23 0.174 1.2 *6 (9.09) (3.24) (3.81) (7.05) (0.182) (1.5) CH-NH- 10.72 3.15 3.53 6.68 5.2 2.9 C2H5NH2 10.81 3.37 3.93 7.30 6.5 1.2 (C2H5)2NH 10.96 3.06 3.30 6.36 9.1 5.0 (C 2H 5) 3N 10.77 2.6 2.1 4.76 5.9 80 HOC2H4NH2 9.74 3.13 3.55 6.68 0.55 2.8 (HOC2H4)2NH 9.00 2.69 2.79 5.48 0.10 16 (HOC2H4)3N 7.90 2.30 1.34 3.64 7.9xl0"3 460 H2NCH2CH2NH2 10.18 4.62 2.92 7.54 1.5 12 H2N(CH2)3NH2 10.64 5.77 1.42*5 7.19*5 4.4 380 H2N(CH2)4NH2 10.82 5.9 6.6 *1 (L) (H*"] / (LH+) - Ka *2 (LH+nOH-VCL) = Kb -*3 (AgL+)/(Ag+)(L) = K Q 1 > (AgL2+J/(Ag+L)CL3 - K 1 2, K Q 2 - K^-K^ *4 (AgL^(L)/(AgL 2tl - K d i These data are from "Stability Constants of Complex Salts", special publication of the Chemical Society, corrected, where necessary, to 25°C. using the known temperature coefficient for Ag(NH2C2H5)2"*" (d log K m/dt - d log K 1 9/dt - -0.016 73 * 5 Estimated from the result of the present work (cf, p. 62). * 6 Values in parentheses are for 30°C, 

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