A STUDY OF THE ELECTRON EXCHAN® REACTION BETWEEN TIN(II) AND TIN(I?) IN AQUEOUS SULFURIC ACID SOLUTIONS By Gilbert Gordon A THESIS Subaitted to the School for Advanced Graduate Studies of Michigan State University of Agriculture and Applied Science in partial fulfillment of the requirements for the degree of DOCTOR OF PHILOSOPHY Department of Chemistry 1959 i ProQuest Number: 10008634 All rights reserved INFORMATION TO ALL USERS The quality of this reproduction is dependent upon the quality of the copy submitted. In the unlikely event that the author did not send a complete manuscript and there are missing pages, these will be noted. Also, if material had to be removed, a note will indicate the deletion. uest ProQuest 10008634 Published by ProQuest LLC (2016). Copyright of the Dissertation is held by the Author. All rights reserved. This work is protected against unauthorized copying under Title 17, United States Code Microform Edition © ProQuest LLC. ProQuest LLC. 789 East Eisenhower Parkway P.O. Box 1346 Ann Arbor, Ml 48106- 1346 (3- I 5 & > n u/iPjw ACKNOWLEDGMENT The author is pleased to acknowledge gratefully the invaluable theoretical counsel and many helpful suggestions of Professor C. H. Brubaker, under whose guidance this research was conducted. iii A STUDY OF THE ELECTRON EXCHANGE REACTION BETWEEN TIN(II) AND TIN(IV) IN AQUEOUS SULFURIC ACID SOLUTION By Gilbert Gordon AN ABSTRACT Submitted to the School for Advanced Graduate Studies of Michigan State University of Agriculture and Applied Science in partial fulfillment of the requirements for the degree of DOCTOR OF PHILOSOPHY Department of Chemistry Year Approved 1959 14- > iv ABSTRACT The kinetics of the exchange reaction between tin(II) and tin(IV) in aqueous sulfuric acid were studied in the temperature range of 25 °C. to 50 °C. Experiments were designed to determine the effect of variation in concentration of tin(II), tin(IV), sulfate, hydrogen, and chloride ions. It was found that the reaction is first order with respect to tin(ll) and tin(IV), and that the effect of increased hydrogen ion is to decrease the rate of exchange, and the effect of increased sulfate ion is to increase the rate of exchange. to be SnOH+ and Sn^O^)^, The predominant tin species appear A combination of the following exchange pro­ cesses *SnO(SO^)| -h SnOHf SnOCSO^)' + *SnOH + or *Sn(OH)t + 3 SnSO. + S0,r 4 4 Sn(OH)^ 3 *SnSO, 4 S0,= 4 and *SnO(S0^2 + SnSO^ SnO^O^ + *SnS04 lead to R/ab = 0.674(H+)"2 + 0.0725(S04=)(H+)“1 The addition of chloride ion to the reaction mixture in 3»00 M sulfuric acid increased the rate by a factor of 100. be described by the following exchange process v The reaction can *SnCl£ + SnCl£ ^ SnCl^ + *SnCl^ which leads to R =4.94[Cl3[Sn(Ilj]fsn(IV)] The chloride ion concentration probably appears in the rate expression because SnCl ** is most likely the major tin(II) species. 3 Spectrophotometric examination of aqueous solutions of tin(II) in perchloric and sulfuric acids have been interpreted in terms of the hydrolysis of Sn ^ . A hydrolysis constant equal to 24.5 has been eval­ uated. The absorption spectra of tin(IV) solutions and tin(II)- tin(IV) solutions in 3.00 M sulfuric acid and tin(II)-tin(IV) solutions in sulfuric acid containing 1.00 M chloride ion (3.99 M H'O are given. The spectra have been interpreted in terms of an interaction dimer containing one atom of tin(II) and one atom of tin(IV). vi VITA Gilbert Gordon Candidate for the degree of Doctor of Philosophy- Major Fields Inorganic Chemistry Minor Fieldst Physical Chemistry, Mathematics Biographical: Born 11 November, 1933 in Chicago, Illinois, Undergraduate Studies at Bradley University Peoria, Illinois, 1951-55. Graduate Studies at Michigan State University, 1955-59. Experiences Graduate Assistant in Chemistry at Michigan State University, 1955-58; DuPont Teaching Assistant, 1958-59; National Science Foundation Fellow, Summer 1959. Member of the American Chemical Society, and The Society of Signa Xi, vii TABLE OF CONTENTS Page INTRODUCTION................................................ 1 HISTORICAL.................................................. 2 THEORETICAL................................................. 7 EXPERIMENTAL................................................ 12 RESULTS..................................................... 30 A, B. Spectrophotometric Observations........................ Kinetic Observations...................... 30 46 DISCUSSION................... 65 SIMMART..................................................... 74 LITERATURE CITED............................................. 76 APPENDICES.................................................. 79 viii LIST OF TABLES TABLE I. II* III* Page 11 Standard Deviation as a Function of (1-f)............... Tin(II) Preparation. ••••••• Tin(Il) Concentration and Absorption Spectra as a Function of Time for a Typical Exchange Experiment. 17 24 IV. The "Observed" Absorbancy Index and Other Derived Quantities for 2.00 x 10*3 Sn(ll) in Perchloric Acid at 230 np........ 7........................ 38 V. The "Observed" Absorbancy Index and Various Other Derived Quantities for 2.00 x 10”^ M Sn(II) in Perchloric Acid at 230 mp ..................... VI. VII. Dependence of Exchange Rate on the Concentration of Tin(II) and Tin(IV)..................................... 39 51 Dependence of Exchange Rate on the Concentration of Hydrogen Ion. ............................. 52 VIII.• Dependence of Exchange Rate on the Concentration of Sulfate Ion. ....................... IX. Dependence of Exchange Rate on the Concentration of Hydrogen and Sulfate Ions...................... 54 X. Dependence of Exchange Rate on Temperature............... 57 XI. Catalysis of Exchange Rate by Chloride Ion................ 59 XII. XIII. XIV. Catalysis of Exchange Rate by Chloride Ion with Variation in Total Tin Concentration................. ••••• 62 Dependence of Exchange Rate on the Concentration of Tin(II) and Tin(IV)..................................... Dependence of Exchange Rate on Temperature...... .. ix 64 64 LIST OF FIGURES FIGURE Page 1. Apparatus for Preparation and Storage of Aqueous Sn(ll)..• 15 2. Glass Adaptor for Voluaetric Flask.......... . 22 3. Ultra-Violet Spectra of Tin(II) and Tin(IV) Sulfate in 3.00 M Sulfuric Acid.................................. 31 A. Absorbancy Index of Tin(II) in Perchloric Acid as a Function of Wave Length.......... .................. . 33 5* Absorbancy Index of Tin(II) in Sulfuric Acid as a Function of Wave Length...................... ......... 34 6. Absorbancy Index of Tin(II) as a Function of Hydrogen Ion Concentration. ...................... . 35 7. Ultra-Violet Absorbancy of SnCSO/^ • H20 in 3.00 M Sulfuric Acid as a Function of Time........ ............ 41 8* Absorbancy as a Function of Tin Product, (jC(II)C(IV)} 44 9* Excess Absorbancy as a Function of Tin Composition........ 45 10. Excess Absorbancy as a Function of Tin Composition in 1.00 M Chloride. ............................. •••••• 11. Typical Rate Data. .................................... 47 48 12. Typical Rate Data for Precipitation from Two Different Media................................................ 49 13. Log(l-f) as a Function of Time at 49.7 °C........... . 56 14* Exchange Rate as a Function of Chloride Ion Concentration. 60 15. Log(l-f) as a Function of Time at 49*7 °C. in 1.00 M Chloride. ............ *...................... 63 x INTRODUCTION The availability of isotopic tracers has made possible the study of exchange reactions between different oxidation states of the same element in which the initial and final states are identical. The exchange reaction Sn(ll) Sn*(lV) ^ Sn*(ll) + Sn(lV) was studied in aqueous sulfuric acid in the temperature range of 25 to 50 degrees Centigrade. The primary interest is in connection with reactions which may proceed either by direct transfer of several elec­ trons or alternatively by the stepwise exchange of one electron at a time. The exchange reaction was studied kinetically by the use of radioactive tracer in an attempt to elucidate a mechanism for the elec­ tron transfer process. The catalytic effect of chloride ion was also studied such that a comparison of the reactions in sulfuric acid and sulfuric acid containing chloride could be made. The ultraviolet absorption spectra of tin(II) in perchloric and sulfuric acids and of tin(II) - tin(IV) mixtures in sulfuric and hydrochloric acids were also studied in an attempt to predict quantitatively the behavior of the aqueous tin species. * By common usage, these asterisks indicate radioactive tracers. 2 HISTORICAL AND THEORETICAL CONSIDERATIONS Historical Tin metal has been known since the time of the ancients when it was known as diabolus metallorum-the devil of the metals (l). The physical properties of tin metal have been elucidated in recent years to facilitate its use in electroplating and alloying* Few studies have been made of the properties of tin(II) and tin(IV) in aqueous media other than hydrochloric acid solutions. The ions, tin(ll) and tin(lV) hydrolyze readily and even in the presence of chloride ion, hydrous tin oxide is known to precipitate. The well-known nature of SnO is closely associated with the tend­ ency of tin(II) to undergo extensive hydrolysis in aqueous solutions. Denham and King (2) have studied the Sn0-S0^-H20 system at 25 °C. Several solid phases, including SnSO^* SnO and SnSO^ were isolated but the data do not indicate the formation of any aqueous tin(II) sulfate complexes. Studies of the conductance (3, 4) of aqueous sulfuric acid solu­ tions containing tin(II) sulfate suggest that two types of complexes are formed, although these studies in very dilute solutions ~ 2 x 10"’’3 M Sn(ll) , in the absence of excess acid, suggest complete dissociation. The hydrolysis of tin(II) can be described by several equations (5, 6, 7) 3 1) Snff+ 2) 2 Sn* + H20 ^ SnOH+ + 2 H20 ^ H+ Sn2(OH)2 -+- 2 H + Dimer and polymer formation (Equation 2) is always favored in solutions of low acidities as in the case of Fe(III) (8). the However, studies of Gorman (9) show that the variations of the equi­ librium quotient with ionic strength are consistent with Equation (l)• At 25 °G, the constant is given as 0.02. It appears that when tin(lV) sulfate dihydrate is dissolved in dilute sulfuric acid, the species, S n S O ^ , are formed (10). In more concentrated solutions of sulfuric acid, higher completing occurs. Brubaker (11, 12) has studied solutions of tin(IV) in sulfuric acid spectrophotometrically and these data indicate that the following reactions take place: 3) SnSO^ -t S0U= ^ 4) Sn(S04)2 + H2S04 ^ Sn(S04)2 H2Sn(S04)3 kx k2 The ratio of Sn(S0^)2/SnS0^+ is ^ 100 in 3.00 M sulfuric acid. The absorbancy index for SnS04 , obtained by comparison of data in perchloric and sulfuric acids and in view of the findings (10) that i-l SnS04 is the predominant species in dilute sulfuric acid, is approx­ imately 2500. The complex acid suggested in Equation (4) is present in very low concentrations in 3*00 M sulfuric acid. In an attempt to clarify reactions in which several electrons are transferred,Brubaker and Court (13) have studied the oxidation of 4 tin(II) with cerium(IV) in sulfuric acid, with most of their data being collected at 0 °C. The following mechanism is consistent with the second-order reaction which occurs at moderate sulfate ion con­ centrations s 5) SnSO^ ■+■ Ge(SO^)^ — tin intermediate 6) tin intermediate ■+• Ce(SO^)^ SnSO^- + 2 Ce(S0^)2 Equally probable products for reaction (5) could be SnSO^ Ce(SO^)2 + S0^= or a complex, SnCe(SO^)x -+■ . The experimental evidence is compatible with the possibility of the "intermediate" ex­ istence of Sn(III), but the data do not exclude the other possibilities shown above. Boyle, etal., (14) reactionbetween interest have shown that Y -rays can catalyse the Fe(III) and Sn(II) in sulfuric acid. thatthedata can sient intermediates, ofSn(I) be interpretedby It isof theformation and Sn(III).These some of tran­ unstable intermed­ iates, however, are then destroyed by reaction with other Y -induced radicals. Weiss (15) has shown, that in the reduction of ferric chloride by tin(II) salts, the tin(II) chloride enters as the complex anion, SnCl£ , in a two stage reaction of a simple electron transfer pro­ cess. 7) Fe(lII) + SnCl.= 4 8) Fe(III) + SnCl” 4 ^ Fe(II) + SnCl“ 4 Fe(ll) + SnCl, 4 5 This is consistent with the observation of Duke and Pinkerton (17) that a minimum of three chlorides are necessary for the activated com­ plex to yield any appreciable reaction rate* The reaction rate is greatly depressed when large amounts of tin(II) are present, at con­ stant chloride ion concentration, due to complexing and the removal of halide ions* Davidson and co-workers (18) have studied the inter­ action between tin(II) and tin(IV) in hydrochloric acid solutions* A mixed solution of tin(II) and tin(lV) absorbs light of longer wave lengths than either solution alone* The observed absorbancy is pro­ portional to the product of the tin concentrations (Equation 9). Ai £\, Sn(Il), Sn(Ivjjj ~ k± A[sn(Ilj)£sn(lV)j The absorption spectra of tin(II) solutions did not obey Beer’s law and this was probably due to the presence of 1-5$ of tin(IV) in the solution studied* The result of equation (9) implies that the species responsible for the interaction absorption is a complex of tin(II) tin(IV) and not a complex containing a single atom of tin(III)* The absorption spectra of the interaction complex are characteristic of the long wave length tails of electron transfer bands (19)* Kinetic (18) and photochemical (20) studies on the above system, suggest that the optical interaction complex is unsymmetrical and that exchange takes place through this complex. The activation energy of 11 kcal* suggests that this is the minimum amount of energy necessary to activate the interaction complex into a configuration in which 6 exchange can occur. Meyer and Kahn (21) have studied the exchange reaction SnCl2 + Sn*Cl2 + SnCl^ in absolute ethyl alcohol. Sn*Cl^ The reaction is first order with respect to tin(II) and tin(lV) chloride concentra­ tion. In accord with their kinetic data, they have postulated these reactions as contributing to the mechanism 10) Sn*Cl4 ;=± Sn*Cl| 11) S n C ^ + Cl" ■c==^ SnCl^ SnCl3 ■+ Cl" 12) SnCl3+ + 13) O ^ n C L ^ M L j * = * Sn*Cl2 + (rapid) (rapid) CL^SnCl^n^eig SnCl^ (rapid) (dete^ lng) If the exchange occurs as suggested* the SnCL, could approach the SnCl^ in such a way as to form an activated complex consisting of two tetrahedral units sharing two chlorine atoms along thus a common edge, necessitating steric orientation before exchange could occur. 7 Theoretical The rate of exchange can be studied by independent variation of the concentrations of each of the reactants by use of the logarithmic form of the first order exchange lav/ (22, 23) ii.) m |j-xA Since 0's — }ds/dz| C/ z, and the error occurring is mainly due to radioassay technique, not statistical counting error, d In s d In z - i m i 11 It is possibleto obtain a family of curves, one for each a/b ratio, by plotting against s. By choosing times or In (l-f) values corresponding to the minimum in this curve, the error in the rate constant, R, can be minimized. Similarly, in the case in which the activity increases in a component that was originally inactive, one obtains 24) d In (a) d — exp(s) — (y/yCoqi)) 1 s and only one curve is observed for all mole fractions. Use of optimum degree of exchange and calculation of the standard deviation for each point , depending on the number of samples taken and the minimum In(l-f) value, the standard deviation in R (Table I) can be calculated. TABLE I STANDARD DEVIATION AS A FUNCTION OF (l-f) Lowest (l-f) value Number of samples r (in %) 0.350 0.550 0.450 0.350 0.350 10 10 9 9 9 8 5.26 6.48 0.450 4.12 5.60 6.56 6.67 EXPERIMENTAL Materials The isotope used in the kinetic studies was the 112 day ^gSn"*"^, which decays by K-electron capture to an excited state of This decays to the ground state with a half-life of 1.73 hours by emission of a 0.392 K.e.v. gamma ray (29)* All samples were mounted and Y-counted in new one-dram screw-cap vials, in a well-type Nal(Tl) scintillation counter. The radioactive tin was a sample of neutron activated tin metal from the Oak Ridge reactor (30). acid solution of tin(IV). It was obtained as a hydrochloric Besides Sn^ ^ activity, this contained the 2.7 year beta active Sb - ^ isotope. The sample was diluted to 10 ml. with 6 M hydrochloric acid, 4.64 mg. Sb+3 and 2,5 gpu of oxalic acid were added. The solution was heated to boiling and hydrogen sulfide was bubbled into the solution until Sb2S-j precipitated. The super­ natant liquid was removed by centrifugation and the above procedure was repeated twice. Concentrated nitric acid was added to the supernatant liquid to decompose the oxalic acid present. This solution was then heated with concentrated sulfuric acid for 250 hours and diluted with water to give a solution of active tin(IV) for the exchange experiments. An aluminum absorption curve was made by mounting a sample two centimeters below the 1.1 inch mylar (1.5 mg./sq. cm.) window of a Tracerlab G-M tube. The observed break in this absorption curve at 13 110 £ 2 mg./cau corresponds to an electron energy of 0.37-0.01 M.e.v. (31) in exact agreement with that expected from a k-conversion elec­ tron from 0.31 M.e.v. gamma ray from metastable indium. A lead absorption curve was made on the same sample and a lead half thickness of 2.77-0.04 ©n./sq.cm. corresponding to a gamma energy of 0.3910*01 M.e.v. (31) was observed. The half-life was determined by comparison to an aqueous sample of Go^® in the scintillation counter. The half-life was observed to be 113 1 2 days which is in good agreement with the value of 112 days reported in the literature. Throughout this experimental work, the sulfuric and hydrochloric acids used were DuPont A. R. grade and the 70% perchloric acid was Kallinckrodt A. R. grade. Aqueous solutions of these acids were standardized with carbonate-free sodium hydroxide which had been standardized against potassium acid phthalate, using phenolphthalein as the indicator. Stannous sulfate solutions were prepared indirectly, using the method of Noyes and Toabe (32). Pure cupric oxide was dissolved by boiling it with an excess of dilute sulfuric acid. To this copper sulfate solution, a threefold excess of 20 mesh granular tin was added. It was found that in sulfuric acid concentrations of 3 M to 5 M the displacement reaction 25) Sn(0) d- Cu(II) ^ Cu(0) + Sn(ll) 14 occurred. If the mixture was kept at 100 °C for two hours, the re­ action proceeded to completion without complication. Lower acid con­ centrations lead to rapid hydrolysis, whereas at sulfuric acid concen­ trations above 5 M> rapid hydrogen evolution occurs. These situa­ tions made regulation of the tin(II) concentration rather difficult. Solutions of tin(II) are readily oxidized by air, therefore, all prep­ arations were carried out in an atmosphere of specially purified ni­ trogen. A drawing of the apparatus used to prepare, purify and store tin(II) solutions is given in Figure 1. Mallinckrodt, pre-purified nitrogen was passed over copper wire in a tube furnace (A) at 450 °C, and then passed over activated copper (B) which had been deposited on kieselguhr (33)* According to Meyer (33)> this activated copper is capable of reducing the oxygen content to less than 10“® moles/liter. The purified nitrogen was bubbled through two pyrogallol towers(D, E) to act as purifiers if the electricity should fail, or to act as indicators if the activated copper were to become exhausted. The pyrogallol solutions were prepared by the dissolution of 30 ©a. of Merck N. F. pyrogallic acid in 300 ml. of water and made alkaline by the addition of potassium hydroxide pellets. A third tower (F) con­ tained 300 ml. of 3 M sulfuric acid which was used to maintain con­ stant water vapor pressure in the gas phase. This was necessary to minimize volume changes during storage of tin(II) and tin(II) - tin(lV) mixtures. Previous to preparation of tin(II) solutions, the apparatus was completely assembled, evacuated to a pressure less than 20 mm. of mercury and filled with purified nitrogen. This operation was re- 15 16 peated twice* Calculated volumes of copper sulfate and sulfuric acid were added to reaction vessel (G) via the condenser (H) under a posi­ tive pressure of nitrogen* The system was closed to the atmosphere and evacuated twice as above* Nitrogen was bubbled through the solu­ tion for 10-20 hours, after which a threefold excess of Baker C. P. Analyzed 20 mesh tin was added through condenser (H)* The system was re-evacuated and nitrogen was bubbled through the solution, until one atmosphere pressure was obtained. The reaction mixture was heated for two hours (twice the time necessary for the blue color of cupric ion to disappear) and cooled under nitrogen atmosphere. The freshly pre­ pared tin(II) solution was diluted to the proper volume with distilled water contained in flask (I)* The water in flask (I) was prepared by cooling freshly boiled distilled water under an atmosphere of the specially prepurified nitrogen. After dilution, the tin(XI) solution remained in contact with tin metal for ten hours to reduce any tin(IV) formed. The tin solution was transferred to the storage flask (K) by reducing the pressure in the storage flask, closing stopcock (L) and increasing the nitrogen pressure. A fritted glass filter (M);midway between the reaction vessel and the storage vessel, facilitated the removal of excess tin and copper metals. After the transfer operation, flask (K) was usually at a pressure less than one atmosphere. Con­ tinuous nitrogen flow equalized the pressure throughout the system and also served to mix completely the tin(II) solutions. The nitrogen was vented through stopcock (N) to flasks in the constant temperature bath* 17 Nitrogen was bubbled through buret (0) for ten minutes prior to filling the buret vdth tin(ll) solution# The buret was rinsed twice with tin(II) solution, which had been forced into it by use of excess nitrogen pressure# Measured volumes of solution were removed from the buret under an atmosphere of nitrogen# For tin(II) and hydrogen ion concentrations see Table II# TABLE II SN(II) PREPARATION Code ml* 0,138 M Cu ml. Cone. HgSO^ Acid 2SD 75 ml. 50 ml. 2.214 0.0257 0.0015 2SF 160 ml. 60 ml. 2.406 0.0250 0.0013 2SG 150 ml. 100 ml. 3.448 0.0325 0.0024 2SH 250 ml. 75 ml. 2.307 0.0418 0.0021 2SI 300 ml. 75 ml. 2.980 0.0643 0.0007 2PA 300 ml.* 200 ml.f 1.710 0.0182 0.0000 2SJ 175 ml. 75 ml. 2.036 0.0298 0.0006 2SK 250 ml. 100 ml. 3.227 0.0364 0.0030 2SL 300 ml. 100 ml. 4.254 0.0484 0.0034 * 0.05 M Cu(C10^)2 Final Concentration M (H^) M Sn(II) M Sn(lV) ? 70% HCIO^ Three methods were used for the preparation of tin(IV) sulfate in aqueous sulfuric acid solutions. An evaluation of these methods 18 appears in the discussion section of the report. (I) Stannic hydroxide was prepared from Mallinckrodt stannic chloride pentahydrate by precipitating the tin with concentrated ammonia solution. The stannic hydroxide thus obtained was repeatedly washed and digested with distilled water until no chloride ions could be detected in the washings. thirty days. These purifications required twenty to The precipitate was stored under water until needed. The precipitate was dissolved in 3 M sulfuric acid and evaporated on the steam bath until Sn(S0^)2 * 2 (12, 34) began to crystallize. Complete crystallization occurred on cooling. The stannic sulfate dihydrate was dissolved in aqueous sulfuric acid of appropriate con­ centrations. Duplicate determinations were made for tin and acid con­ centrations. Solutions of stannic sulfate prepared by this method were unstable, with respect to formation of white tin(IV) oxide, and were undesireable for use in prolonged kinetic studies. (II) Hydrogen trisulfatostannate(XV) monohydrate was prepared by the reaction of Baker C. P. Analyzed 20 mesh tin in concentrated sul­ furic acid at 190 °C (35). Subsequent fuming of this solution almost to dryness served to coagulate any colloidal sulfides that appeared and these were removed by filtration through fritted glass filter after dissolution of the tin sulfate in hot sulfuric acid. Fuming almost to dryness, filtering and re-dissolving were repeated twice. The result­ ing salt was hydroscopic and had the composition corresponding to hydrogen trisulfatostannate(lV) monohydrate. 19 Anal* 0,47. Calcd. for HgSnCSO^^ • HgOs Found: Sn, 27.8; H, by titration, Sn, 27.8; H, by titration, 0.48* Solutions prepared from this compound by heating it with dilute acid were less desireable for kinetic studies, due to hydrolysis, than those prepared by method (III). (Ill) Stable solutions of stannic sulfate were prepared by dis­ solution of 0*2 moles of hydrogen trisulfatostannate(IV) monohydrate in one liter of boiling sulfuric acid. These solutions were kept at the boiling point (with addition of concentrated sulfuric acid to main­ tain constant volume) for 200-300 hours. After this prolonged heating, the solutions were diluted to 10-12 M acid and analyzed. Further dilu­ tion resulted in solutions which were stable with respect to hydrol­ ysis for more than 60 days. All attempts to prepare stable solutions of tin(IV) in perchloric acid by methods similar to those used for the preparation of tin(IV) sulfate were unsuccessful. Dissolution of freshly precipitated meta- stannic acid was also unsuccessful. Even dissolved crystalline hydro­ gen trisulfatostannate(IV) monohydrate was unstable in perchloric acid. Mallinckrodt A. E. grade oxalic acid was recrystallized from dis­ tilled water, dried in vacuo over concentrated sulfuric acid, and ground to a fine powder. Mallinckrodt analytical reagent grade lithium sulfate monohydrate was recrystallized from distilled water as the monohydrate. The puri­ fied salt was dried in vacuo for 24 hours and then in an oven at 180 °C. 20 for 48 hours in order to prepare the white anhydrous form. Lithium perchlorate was prepared from Mallinckrodt analytical reagent grade lithium carbonate by the addition of perchloric acid in a slight excess. The solution was evaporated on a steam bath. After cooling, the salt was separated by filtration and recrystallized from distilled water seven times. The salt was dried over anhydrous lithium perchlorate which had been dried in vacuo over concentrated sulfuric acid* In this manner, lithium perchlorate, trihydrate was obtained. A sample of the trihydrate was dried in vacuo over concentrated sul­ furic acid, and then heated at 200 °C.for several days. The loss of weight corresponded to that calculated for three moles of water. The tin(Il) concentration was obtained by titration in an atmos­ phere of carbon dioxide with 0.07 N iodine solution. Prior to each series of titrations, the iodine solution was standardized with arsenous oxide in a bicarbonate buffered medium using freshly prepared starch solution as the indicator (36). The stability of a properly prepared and preserved iodine solution is illustrated by the experi­ ence of Washburn (37) who found that the normality of an iodine solu­ tion that had been in frequent use during an interval of two months changed less than 0.002$. The sensitivity of starch as an indicator corresponds to an iodine concentration of 2 x 10“ ^ u. Thus, the con­ centration of the most dilute tin(II) solutions could be known better than five parts per thousand. The total tin concentrations were eval­ uated by a modification of Farnsworth’s method (38) in which the absorption of a dispersion of tin(II) toluene 3-4 dithiolate was meas- 21 ured at 536 mp. Matheson Coleman and Bell reagent grade thioglycollic acid was used to reduce the tin species to tin(II), Santomerse SX* was the dispersing agent, and Eastern Chemical Company toluene 3-4 dithiol was used to form the red complex with tin* A standard tin solution containing 393 ^ tin/ml. was prepared to be used as a reference. Aliquots of the standard solution, when treated as indicated, were found to obey Beer*s law. Dithiol reagent was freshly prepared before each determination by dissolving approximately 0.03 @ns. of dithiol in 10 ml, of 2% NaOH. The addition of two drops of thioglycollic acid increased the stability of the dithiol reagent from several hours to several days. Previous workers (38, 39) have noted poor stability for the tin dithiol complexes. The author has been able to observe stabilities (no spectral changes) of greater than 36 hours by maintaining the fol­ lowing concentrations in the final solution: Sulfuric acid concentra­ tion equal to 0.21 ±0,05 M and tin concentrations from 0.8 to 8.0 T/ml. A known volume of the appropriate solutions, three reference solutions (100-500 X of stock solution), the calculated amount of sul­ furic acid and 3 drops of thioglycollic acid were thoroughly mixed and diluted to 20 ml. Five drops of Santomerse SX were added, and one half milliliter of dithiol reagent were added, and the solutions were mixed, very carefully. The solutions were diluted with water to 25 ml. and the absorbancy determined at 536 mp. Distilled water, treated as above was used as the reference solution. * Trademark, Frontsanto Chemical Company 22 Kinetic Studies The exchange studies were carried out in solutions with the molar ionic strength of 4*9# except for the series used to study very high hydrogen and sulfate ions, where it was necessary to vary the ionic strength also. The actual kinetic studies were carried out in 100 ml. volumetric flasks. All glassware used was carefully cleaned, dried, heated, blown out with purified nitrogen, and stoppered before use. For all samples, except those in which tin concentrations were varied, the total tin concentration was 2.5# x 10 M. Preliminary studies indicated that the half lives would be rather long (300-600 hours), thus it was necessary to keep a stream of nitrogen flowing into the reaction vessels which were stored in a constant temperature bath. A 14/35 T male joint was fitted with a 6 mm. glass inlet and an outlet tube (see Figure 2). The vertical tubing served as the gas inlet, and the horizontal tubing served as an outlet to the next inlet tube. Four such adaptors were placed in series. Four series of adaptors came from one common manifold. By maintaining a positive nitrogen pressure above the liquids continuously, it was possible to main­ tain 10"2 M tin(II) solutions for peri­ ods up to 100 hours without detectable Figure 2 oxidation. 23 Previous te each run, the glass-stoppered 100 ml* volumetric flasks were placed into the hath and flushed with nitrogen for 6-6 hours before the addition of any solutions* The solutions used, and the volumes of each were calculated using equations 63-65 in Appendix B. The volumes of sulfuric acid, the volumes of minor acids, the vol­ umes of lithium salts, and the water were added in that order. The flasks were replaced into the constant temperature bath and were cooled and flushed with nitrogen. The calculated volume of tin(lV) was added, and the flasks were flushed with purified nitrogen for 45-60 minutes. The calculated volumes of tin(II) solutions were added under a constant flow of nitrogen. The flasks were stoppered, inverted 10 times and were replaced into the constant temperature bath and flushed with ni­ trogen. At this point, the flasks contained approximately 96 ml. The reaction flasks were allowed to come to thermal equilibrium for twelve hours. After attaining equilibrium (see Table III), three ml. of Sn^^ were added and the flasks were inverted 10 times to insure thorough mixing. The flasks were placed in the bath again and nitrogen flow was begun immediately. Zero time was taken as the time when one-half of the activity had drained into the flask. The overall mixing pro­ cedure took less than 20 seconds. As soon as possible, after mixing, the tin(ll) concentration was determined iodometrically and the absorption spectrum was measured (see Table III). Data from kinetic runs were recorded only -when the absorption spectra and tin(II) concentrations remained constant for two half lives. 24 TABLE III TIN (II) CONCENTRATION AND ABSORPTION SPECTRA AS A FUNCTION OF TIME FOR A TYPICAL EXCHANGE EXPERIMENT t(hrs#) M(Sn(II}) A250 wp a270 mp 0.5 0.0126 1.89 0.189 50.0 0.0126 1.89 0.192 100.0 0.0127 1.87 0.191 300.0 0.0126 1.87 0.189 600.0 0.0127 1.88 0.193 900*0 0.0127 1.89 0.191 Separation Procedure (I) Oxalic acid was used to precipitate stannous oxalate. The separation was effected by adding 2 ml. of reaction solution to a small beaker containing 1 gpu of powdered oxalic acid and 2 ml. of saturated oxalic acid solution. To insure complete precipitation, even when tin(II) concentrations were lowered to 0.005 M, 1 ml. of 0.100 M tin(Il) in 2.00 M hydrochloric acid was added. The solutions were stirred rapidly with a magnetic stirrer for two minutes and fil­ tered through a stainless steel filter with a removeable chimney and holding 22 mm. Whatman No. 42 paper# The rate of filtration was con­ trolled by an aspirator connected to a micro bell jar with a ground i< flange and opening for the steel funnel. The filtrate was collected in 25 a 10 ml, volumetric flask. The beaker and precipitate was washed with two 2 ml. samples of saturated oxalic acid solution. The filtrate and washing were diluted to 10 ml. with water and mixed thoroughly. Four ml* samples were withdrawn and were mounted for X -counting in new 1dram screw-cap vials. The samples were counted 20-24 hours after sepa­ ration to insure that equilibrium was established with the In^^m present. Periodically, duplicate samples were taken to determine the sampling error. The precipitation of stannous oxalate was satisfactory for samples containing less than 0.6 M chloride ion. Above 0.6 M chloride ion, the precipitation was kinetically hindered by the chloride and a second method of separation was utilized. (II) Cesium hexachlorostannate(lV) was precipitated in the high chloride samples. The exchange reaction in the sample was quenched by the addition of 2 ml. of reaction solution to 1 ml. of 0.087 M cesium chloride in 11 M hydrochloric acid. seconds and filtered The solution was stirred for 60 through the same apparatus as in method (I). filtrate was collected directly in a 1-dram screw-cap vial. The The beaker and precipitate were washed with 1 ml. of concentrated hydrochloric acid. Duplicate samples gave results which agreed within expected counting statistics. The cesium hexachlorostannate samples were not counted until 20-24 T1Qwj hours after separation; the In activity at this time indicates the amount of Sn 113 present in the filtrate. 26 It is possible to obtain equilibrium activities, i.e. activities when t = oo, by taking samples after ten half lives. With very long half lives, or as in this case where possible oxidation becomes a factor, a calculated value for the activity must be used. This method has been used successfully by several authors (18, 40) and was also used, in part, in this study. Since the decrease of activity in an initially active component was observed, (1-f) has been rewritten 26) 5(11)1© Where X is the activity at any time, t, n(II) is the average mole fraction of tin(ll) in the sample, and X© is the initial activity of the radioactive component. The (l-f) values, Appendix A, pp. 80-102, have been calculated using Equation (26). To facilitate future calculations, the c/s(IV) values listed in Appendix A have been calculated using 27) and the c/s(IV) when t = oo(c) are 28) thus c.M l V ) '2 < K ce cm H _ ro Ll I t3r CD o co m < ( 32 were made at 240 - 260 mp, whereas at higher chloride ion concentra­ tions, absorbancies were measured around 300 mju In contrast to the failure of Beer's law, observed for dilute sol­ utions of tin(II) in 3*00 to 10,00 M chloride by deMaine and deMaine (42), all of the tin(II) solutions followed Beer's law to±2%, It was necessary to correct the tin(II) solutions for the tin(IV) present and the absorption spectra of tin(II) solutions containing more than 2^ tin(IV) were found to disobey Beer's law, unless they were so corrected. Careful preparation of tin(ll) resulted in less than 2% tin(lV) and little or no correction was necessary. The absorp­ tion spectra of these solutions were found to follow Beer's law very well from 230 - 260 mpu The log of the "observed" molar absorbancy index, where the absorpancy index is defined as 30 a(obs) ” A (obs)/C^Sn^II0 ; has been plotted as a function of wave length between 220 mp and 260 mp for perchloric and sulfuric acids. These absorption spectra, for sol­ utions of tin(II) in perchloric acid, are shown in Figure 4. Figure 5 illustrates some typical absorption spectra for tin(II) solutions in sulfuric acid. Qualitatively, the curves in Figures 4 and 5 appear to be quite similar. The hydrogen ion concentration was calculated for the sulfuric acid solutions, using the data of Smith (41), and Figure 6 shows a(obs) at 230^240 and 250 mp as a function of the calculated hydrogen ion concentration. 33 H CIO, 1.28 3.84 6.41 7.80 8.96 450- 40C- 10.0 350 300 250 200 150 100 >i ., 9 0 AS A FUNCTION H2 S04 SOLUTIONS AT 250 m/A H2 S04 SOLUTIONS AT 240 m/t (D) H2S04 SOLUTIONS AT 2 3 0 36 At 240 and 250 mji, the absorption spectra of tin(ll) are virtually identical in both sulfuric and perchloric acids, suggesting the follow­ ing type of equilibriums Sa~ + H2Q 35) If SnOH V ^ H* % is the absorbancy index for SnOH+, and a2 is the absorbancy || index for Sn , we can assume that a # = if we also assume that SnOH and Sn a ^ and that aQ = a-^. Now, are the only "colored" species observed at 240 mji, we can write M *(ob,) = ^ i f s « D H ^ + a2 |Sn'^|'c( Sn)| 37) »(ob»l - »„ = (ll***? + *2l?»*J - S ° ( s J / ( f C < s J then a «o ” *0 a2 “ al and using K-^ from equation (35) and the definition of and a0, we can write ,81 Hob.1 - *0 _ a A / O O + a2 - .^/(H*) £a2 - a^JTK/H-) + lj H Equation (38) can be inverted and rewritten as 39) 81«° !a_ = 1 + k/(h*) a(obs) “ ao The value aQ = 179 is obtained by extrapolation of curve B in Figure 6 to zero hydrogen ion concentration, A plot of 37 S od “ ^//^(obs) " ^ as a function of hydrogen ion yields a good straight line if 320. 24.5± 0.05. — is determined from the slope and is These data are summarized in Table (IV). If the species present can be adequately described by Equation (35) t then the spectra of tin(II) in perchloric acid at 230 mp should also follow this description. k ) where a'denotes * (obs> - Thus, let /Ki/HfJ + 1 the absorbancy index at230 mp. Since has been evaluated, it can be substituted in Equation (40) and by subsequent rearrangement, we obtain Equation (41). W) 4 f * ' ( o b s ) ] [ i t- Ki / « ^ A plot of /p'(obs)^lj- + ^ as a faction of with a slope, a^ := 300 and an intercept, a^ = Jk^/H +J is 250. linear These spectro- photcmetric data are summarized in Table (V). The proposed mathematical model which quantitatively describes the equilibrium between SnOH and Sn can also be used qualitatively to explain the general shape of the tin(Il) spectra. Since the maximum in the absorption spectra appears to be in the vicinity of 240 mp and the absorbancy index for Sn0H+ increases toward shorter wave length, a minimum could occur when these two curves were added together, and this minimum would be expected to shift toward shorter wave lengths as the hydrogen ion increased, because the absorption due to Sn* would be 38 TABLE IV THE "OBSERVED" ABSORBANCY INDEX AND VARIOUS OTHER DERIVED QUANTITIES FOR 2.00 x 10“3 M TIN(II) IN ACID SOLUTION AT 240 mji “ H (H*) aoba (l/H +) a oo “ ao aobs 1.00 185 1.000 23.5 1.40 187 0.714 16.6 1.60 190 0.625 14.1 2.00 192 0.500 1C.8 3*00 199 0.333 7.05 4.00 207 0.250 5.C5 5.00 217 0.200 3.74 6.00 227 0.167 2.94 7.00 238 0.142 2.39 8.00 251 0.125 1.96 9.00 265 0.111 1.64 10.00 281 0.100 1.38 11.00 298 0.091 1.18 ao 39 TABLE V THE "OBSERVED" ABSORBANCY INDEX AND VARIOUS OTHER DERIVED QUANTITIES FOR 2.00 x 10"3 H TIN(II) IN PERCHLORIC ACID AT 230 mf K ( H+) a(obs) [a '(obs)J^1 + Kl/Hj/ 1.30 285 19.9 5860 2.62 269 10.9 2940 3.30 262 8.4 2200 3.99 255 7.1 1810 4.61 250 6.3 1575 5.27 246 5.65 1390 5.90 243 5.15 1250 6.52 241 4.75 1145 7.IS 239 4.41 1051 7.76 240 4.16 1000 8.27 242 3.96 959 8.82 245 3.78 926 9.30 250 3.63 910 9.76 259 3.51 910 40 increasing and that due to SnOH+ decreasing. The spectra of tin(II) in sulfuric acid at 230 suggest that some highly "colored" tin(II) species is formed in dilute sulfuric acid, and that similar species to those observed in perchloric acid exist in more concentrated sulfuric acid. This suggests possible ion- pair formation between the partially hydrolized species, S n O H ^ , and SOjT or H50j~ • No suitable mathematical model has been found that can describe these data* Curves B and G in Figure 3 are representative of the absorption spectra observed for solutions of Sn(lV) in 3»00 K sulfuric aeid. The absorbancy index for SnCSOj^ * 2 H^O as a function of sulfuric acid concentrations have been reported (11), and those observed in curve B are in agreement with these values reported in the literature. It is also reported that these solutions are somewhat unstable, becoming cloudy in less than a week, and finally, hydrous white tin(lV) oxide is precipitated. Preliminary exchange studies in sulfuric acid indicated that the half lives would be on the order of hundreds of hours. Thus, it was essential to obtain tin(IV) species whose absorption spectra would remain constant for this period of time. It was assumed that a change in the absorption would be an indication that the species were changing, and, therefore, the rates of exchange observed would be meaningless. Figure 7 shows how the absorption spectra of these tin solutions change with time. No Tyndall beam was observed for the first 60 hours. By 75 hours, a Tyndall beam was visible and at 100 hours, a visible FIGURE 7 o 00 o ro ID CO < TIME o IN 8 3.00 M H2S04 AS 3 O X S*{S04)2.2H 20 O ro — OF OF CO FUNCTION ABSORBANCY o: ULTRA-VIOLET u < 42 precipitate was observed. Very small absorbancy indices were observed for solutions of tin (TV) that had been boiled in sulfuric acid for prolonged periods of time /curve C, Figure 5/ • The absorption spectra did not change with time, i.e., the absorbancy index was constant for 800 hours. Beth tin(IV) solutions were observed to obey Beer's law although the curves did not intercept the abscissa at zero, which is indicative of more than one tin species being present. The absorption spectra of mixed solutions of tin(II) and tin(IV) were dependent on the history of the tin(IV) species present. The mixed solutions, in all cases, absorbed more energy than the sum of the absorptions of the individual components. All of the following spectra were studied in 3*00 K sulfuric acid, except for the studies which were 1.00 M in chloride ion, and these solutions were 3*99 M in hydrogen ion (the H + concentration of 3»00 M sulfuric acid). The interaction spectra for two different tin preparations will be treated separately. Dissolution of Sn(S0^)2 * 2 HgO in 3*00 M sulfuric acid provides solutions in which the tin(IV) spectra change with time (see Figure 7), but the spectra are repreduceable, at any given time after dissolution. Thus, the absorbancies were determined fifteen minutes after the solu­ tions were mixed. The absorbancy indices for the tin(II) and tin(IV) were the same order of magnitude (424, and 630 @ 240 mji) and suggest that these solutions will support the model proposed by Davidson (18) for the tin(II) - tin(IV) - chloride system. Equation (9) defines 43 this relationship and Figure 8 displays some of the data that support Equation (9). These data indicate that the interaction absorption is that of a dimeric species, formed by 42) Sn(ll) -j-Sn(IV) Sn(XI)Sn(lV) and not the formation of an intermediate species, 43) Sn(Il) + Sn(IV) ^ 2 Sn(lII) If Equation (43) were correct, a plot of (C(II) • C(jyJ)i against A (obe) would be linear, but this is not the case (Figure 8). the system can be described by A& — Mathematically, kJc(Il) - C(DjJ£c(lV) - C(T)Jj^ where C(I1) and C(XV) refer to the molar concentration of tin(ll) and C(B) is the molar concentration of the interaction dimer. However, the use of Equation (9) implies that the concentration of interaction species must be very low (although it may have a rather large absorb­ ancy indea^because the cross terms seem t© be negligible. Several solutions, containing tin in both oxidation states, were prepared using tin(IV) that had been boiled in sulfuric acid for 250 hours. Since the absorbancy indices of these tin(IV) solutions were very low, it was possible to use the method of continuous variation to study the complex formation. A graph of A. A as a function of the tin(Il) - tin(TV) ratio is exhibited in Figure 9 • A maximum occurs in this graph at a ratio of 1:1 which supports the formation of an interaction dimer. Variation of sulfuric acid concentration affects the absorbancy indices of the individual species, b u t ^ A remained u 1.40 1.30 1.20 ABSORBANCY 1.00 0.90 0.80 0.70 0.60 0.50 0.40 0.30 0:20 0.10 5 10 15 I08 FIGURE 8. 20 [C (H ) • COffj] ( A ) ABSORBANCY AS A FUNCTION OF ( c ( I ) C(DT))fe AT 2 4 0 m/* ( B ) ABSORBANCY AS A FUNCTION OF COE) CUSH AT 2 4 0 mp (C ) ABSORBANCY A S A FUNCTION OF C (E) CUE) AT 260 45 0.070 0.060 EXCESS ABSORBANCY ^ A 0050 0.040 0.030 0.020 0.010 0.0001 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 FRACTION OF Sn (Iff) FIGURE 9. EXCESS ABSORBANCY AS A FUNCTION OF TIN COMPOSITION. 2 .5 < I0 '3 M TOTAL TIN IN 3.00M Y^SO^ 46 constant, even though the sulfuric acid concentration was varied from 2.5 to 4.5 M. Evidence for complex formation between tin(II) and tin(IV) in hydrochloric acid solutions has been given by Davidson (18), and in a series of papers by de Maine and de Maine (42). Analysis of these spectra indicate that the observed enhancement is due to interaction between tin(II) and tin(IV) species. Figure 10, curve B exhibits a graph of A A against the tin(II) - tin(IV) ratio in 1 M chloride. In conformity with Davidson's results, a maximum is observed at a ratio of one to one, indicating an interaction dimer. Kinetic Observations Variation of the concentration of each reactant is necessary to determine the order of that component with respect to the overall exchange rate. Figures 11, 12 against t for several runs. shows some typical graphs of leg(l-f) Curves A and B, Figure from the oxalate precipitation of tin(II). 11 include data Curves B and C, Figure 12 give a comparison of the half lives obtained by the two different pre­ cipitation methods. Curve B is typical of the oxalate precipitate in high chloride, curve A is from the cesium precipitation of tin(IV) for the same run, and curve C is typical of other runs in which tin(lV) was precipitated. Values of the exchange rate R were computed from the half time obtained from graphs, such as those in Figures 11 and 12. It has been shown by Prestwood and Wahl (46) that incomplete separation or partial 47 0.070 0.060 EXCESS ABSORBANCY A a 0.050 0.040 0 .0 3 0 - 0 .020- 0.010 02 0.1 0.3 0.4 0.5 06 0.7 0.8 0.9 FRACTION OF Sn ( H ) FIGURE 10. EXCESS ABSORBANCY AS A FUNCTION OF TIN COMPOSITION. 2 .5 x |0- 3 M TOTAL TIN IN, 1.00 M HCl , WITH 3.99 M H* AND 0.795 M S04=. 48 0 .9 0.8 0.7 0.6 0 .5 0.4 0.3 0.2 100 300 200 400 TIME, HOURS FIGURE II. TYPICAL ( A ) 3 .00M H2S04 , RATE DATA (B )0 .0 7 7 5 MCI" IN 3-00 M H2 S04 500 k9 0.8 0.7 0.6 r -/) 0.5 0.4 0.3 0.2 5 10 15 20 T IM E , FIGURE 12. TYPICAL RATE DATA FOR TWO DIFFERENT MEDIA. HOURS PRECIPITATION (A ) Cs^SnCI 0 FROM 1.00 M Cl" ( B) (C ) Cs2SnCI6 SnC2 04 FROM 0 6 0 M Cl" FROM 0.60 M Cl” 25 FROM 50 catalysis of an exchange reaction by the separation method does not affect the value of R determined in this way, provided these affects are reproducible in each run. In these experiments, the apparent zero time exchange for the cesium precipitation of tin(IV), i.e., the value of 100 l/ljjj at t = 0 , averaged about k5%» The agreement of duplicate runs (curves B and C, Figure 12) by two different methods and the statistical agreement with the predicted exchange law is a good indication that the method is justified. A log-log plot of the rate of exchange, from independent variation of tin(II) and tin(IV) concentrations, against tin concentration gave the expected linear relationship. When the tin(IV) concentration was maintained constant, and tin(XI) varied, it was found that t 0.04. — 0.98^ Constant values of tin(II) and variation of tin(IV) concen­ tration resulted in ^2 - 0.99^- 0.01*;. The constant k or R/ab can be evaluated for 3.00 M sulfuric acid (3.00 M hydrogen ion and 0.99 M sulfate ion)from the above data. Table VI summarizes these data. A value of 0.0621i 0.0017 liters/mole-hour was obtained for R/ab. In order to evaluate the effect of hydrogen ion on the rate of exchange, a series of experiments were made in which the ionic strength was maintained at 4.98 and the sulfate ion concentration was kept at 0.99 M. The equations used to satisfy these conditions are listed in Appendix B. A summary of these data appear in Table VII, where C(Sn) is the total M tin concentration. The variation of R can be shown in a log-log plot of the rate constant against hydrogen ion^ tion, where a slope of -1.5 ±0.1 was observed. concentra­ An increase in the ^Calculated hydrogen ion concentration as evaluated on page 28 51 TABLE VI DEPENDENCE OF EXCHANGE HATE ON CONCENTRATION OF TIN(II) AND TIN (IV) (H^) = /Sn(ll2/ M 3.99 H; (SO^~) = 0 . 9 9 M? f /Sn(IVj/ M =4.98 1. mole”^ hrs.”^ k(obs) k(I) k(II) 0.0077*7 0.0129 0.062 0.061 0.062 0.0097 0.0126 0.061 0.061 0.062 0.0104 0.0129 0.064 0.061 0.062 0.0129 0.0129 0.062 0.061 0.062 0.0156 0.0129 0.060 0.061 0.062 0.0129 0.0066^ 0.064 0.061 0.062 0.0129 0.00^ 0.062 0.061 0.062 0.0129 0.0106 0.062 0.061 0.062 0.0129 0.0129 0.064 0.061 0.062 0.0129 0.0205 0.063 0.061 0.062 52 TABLE VII DEPENDENCE OF EXCHANGE RATE ON THE CONCENTRATION OF HYDROGEN ION J1 = 4.98; (SO^=) = 0.99 H; C(Sn) = 0.0258 M C*+3 M M M M 3.30 2.31 O .345 O .345 3.60 2.61 0.200 0.200 3.99 3.00 — 4.40 3.21 4.80 5.00 A 1010!? M -1 1. moles”1 hrs. k(obs) k(I) k(II) — 0.082 0.083 0.084 — 0.072 0.073 0.072 — — 0.060 0.062 0.061 0.21 — 0.57 0.054 0.054 0.051 3.43 0.3 8 — 0.18 0.047 0.047 0.045 3.53 0.48 __ 0.045 0.044 0.042 TABLE VIII DEPENDENCE OF EXCHANGE RATE ON THE CONCENTRATION OF SULFATE ION p - 4 .98; (H+) =0.99; C(Sn) ~ 0.0258 M /so4-] [ksaj /hcio0 SagioO M M M 0.70 2.12 1.17 0.80 2.42 0.77 0.99 3.00 — — 1.20 3.215 — — 1.30 3.32 — 1.40 3.42 — M 4 W 1. moles”1 hrs •-1 M k(obs) k(I) k(II) 0.29 — 0.057 0.057 0.056 0.19 — 0.058 0.059 0.057 — 0.060 0.062 0.061 0.425 0.065 0.065 0.064 — 0.630 0.066 0.066 0.065 — °.83o 0.066 0.067 0.067 53 hydrogen ion concentration decreased the rate of exchange. ations in hydrogen ion were limited by two factors r (a) the The vari­ tin (IV) spectra were stable as a function of time only above 3*00 M hydrogen ionj (b) an ionic strength of 4.98 limited the maximum hydrogen ion at 5.00 M. Data appear later in which the ionic strength, jx, was allowed to vary and hydrogen ion concentrations up to 6.67 M were obtained. The rate law can now be written: 44) B = k/sn(II)7/sn(IV)] The effect of variation in sulfate ion is shown in Table VIII. These data indicate that the effect of sulfate ion is very small, but would correspond to a value of 0.25 £ 0.04 in the mathematical expres­ sion for the rate law, Due to the insolubility of tin(IV) in perchloric acid, it was impossible t® diminish the sulfate ion concentration below 0.70 M at constant ionic strength and constant hydrogen ion concentration. Constant ionic strength limited the maximum sulfate ion concentration t# 1,40 M. In •rder to study the effect of higher hydrogen and sulfate ions than were allowed withal = 4.98* a series of measurements were made in which the ionic strength was allowed to vary from run to run. are tabulated in Table IX. These results The value, k(obs), is the apparent rate constant calculated from the observed half life (see Equation (16)) K) x: a- n X 4> sD sulfate, hydrogen, and chloride ions. It was found that the reaction is first order with respect to tin(ll) and tin(IV), and that the effect of increased hydrogen ion is to decrease the rate of exchange, and the effect of increased sulfate is to increase the rate of exchange. be SnOH*" and Si^SO^^. The predominant tin species appear to A combination of the following exchange pro­ cesses *Sn0(S0^)j •+SnOH+ Sn0(S0^)| + ^SnOH^ or *Sn(0H)^ + SnSQ^ + S0£ Sn(OH)^ + *SnS0^ +■ S0^ and *Sn0(S04)“ + SnSO^ ^ Sn0(S04)| + *SnS04 lead to R/ab — 0.674(H )*2 -+- 0.0725(S0^=) (H^)”1 The addition of chloride ion to the reaction mixture in 3.00 M sulfuric acid increased the rate by a factor of 100. be described by the following exchange process The reaction can 75 *SnClg -h SnCl£ SnCl= -h *SnCl = which leads to n = K.%[crj /si>(IljJ/sn(IVj) The chloride ion concentration probably appears in the rate expression because SnCl^” is most likely the major tin(ll) species. Spectrophotometric examination of aqueous solutions of tin(II) in perchloric and sulfuric acids have been interpreted in terms of the hydrolysis of Sn-^. uated. A hydrolysis constant equal to 24.5 has been eval­ The absorption spectra of tin(IV) solutions and tin(II) - tin(IV) solutions in 3*00 K sulfuric acid and tin(II)-tin(IV) solutions in sulfuric acid containing 1.00 M chloride ion (3.99 M H*) are given. The spectra have been interpreted in terms of an interaction dimer containing one atom of tin(II) and one atom of tin (IV). 76 LITERATURE CITED 1* Mantel, C. L., Tin, New York, Reinhold Publishing Corp., 1949* 2. Denham, H. G., and W. E. King, J. Chem. Soc., 1251 (1935). 3. Miyamoto, S., Bull. Chem. Soc. Japan, 2, 56 (1932), C. A. 26 2916"5. 4. O'Connor, E. A., Nature, 122, 151 (1937). 5. Tobias, R. S., Acta Chem. Scand., 12, 198 (1958). 6. Tobias, R. S., J. Chem. Education, 22, 592 (1958). 7. Prytz, E., Z. Anorg. u. Allgem. Chem., 174. 355 (1928). 8. Hedstrom, B. 0. A., Arkiv. Kemi, £, 457 (1953), C. A. 4£ 119381. 9. Gorman, M., J. Am. Chem. Soc., 61, 3342 (1939). 10. Brubaker, C. H., ibid.. 72, 5265 (1955). 11. Brubaker, C. H., J. Phys. Chem., 61, 696 (1957). 12. Brubaker, C. H., J. Am. Chem. Soc., 76. 4269 (1954). 13. Brubaker, C. H., and A. J. Court, ibid.. 78, 5530 (1956). 14. Boyle, J. W., S. Weiner, and C. J. Hochhandel, J. Phys. Chem., 62, 892 (1959). 15. Weiss, J., J. Chem. Soc., 309 (1944). 16. Rabinowitch, E., and W. H. Stockmayer, J. Am. Chem. Soc., 64, 335 (1942). 17. Duke, F. R. and R. C. Pinkerton, ibid.. 73. 3045 (1951). 18. Browne, C. I., R. P. Craig, and N. Davidson, ibid.. 22, 1946 (1951). 19. Rabinowitch, E., Revs. Mod em Phys. 14. 112 (1942). 20. Craig, R. P., and N. Davidson, J. Am. Chem. Soc., 22, 1951 (1951). 21. Meyer, E. G., and M. Kahn, ibid., 22, ^950 (1951). 22. McKay, H. A., Nature 11£, 997 (1938). 77 23. Friedlander, G., and J. W. Kennedy, IntroductiontoRadiochemistry, New York, John Wiley and Sons, Inc. (1955). 24. Frost, A. A ., and R. G. Pearson, Kinetics andMechanisms,NewYork, John Wiley and Sons, Inc. (1953). 25. Job,^Ann.Chim., 10 %.> H3> (1928). 26. Vosburgh, W. C., and G. R. Cooper, J. Am. Chem. Soc., 63 . 437 (1941). 27. Katzin, L., and E. Gebert, ibid.. 22,5455 (1950). 28. Davidson,N., and J. H. Sullivan, ibid., 21, 739 (1949). 29. Seaborg, G. T., Revs. Modern Phys., 20, 585 (1948). 30. ’’Isotopes, Radioactive and Stable”. (Oak Ridge National Laboratory, Oak Ridge, Tenn., Tenn., July 1952). 31. Glendenin, L. E.,Nucleonics, 2, 26 (1948). 32. Noyes, A. A., and K. Toabe, J. Am. Chem. Soc., 1537 (1917). 33. Meyer, F. R., and G. Range, Angew. Chem., j>2, 637 (1939). 34. Ditte, A., Compt.rend., 104. 172 (1887). 35* Mathers, F. C. and H. S. Rothrock, Ind. Eng. Chem., 2£, 831 (1931). 36. Vogel, A. I., Quantitative Inorganic Analysis, New York, Longmans,Green and Co. (1953) p. 341. 37. Washburn, E. W., J. Am.Chem. Soc., J!0, 41 (1908). 38. Farnsworth, M., and J. Pekola, Anal, Chem., 26, 735 (1954). 39. Court, A. J., Ph.D. Dissertation, Michigan State University, (1956). 40. Sincius, J. A., Private Communication. 41. Smith, H. M., Ph.D. Dissertation, University of Chicago (1949) and Record Chem. Progress, 12, 81 (1951). 42. de Maine, M. M., and P.A. D. de Maine, J. Inorg. & Nuclear Chem., in press ; and Private Communication. 43. Vanderzee, C. E., and D. E. Rhodes, J. Am. Chem. Soc., 74, 3552 (1952). 78 44. Prytz, M., A. Anorg. Chem., 219. 89 (1934). 45. Smith, L., ibid., 176, 155 (1928). 46. Prestwood, R. and A. Wahl, J. Am. Chem. Soc., Jl, 3137 (1949). 79 APPENDICES 80 APPENDIX A ORIGINAL KINETIC DATA 81 KINETIC STUDIES IN 3.00 M SULFURIC ACID, - A.98 TABLE XV STUDY OF THE EFFECT OF VARIATIONS IN TIN(II) CONCENTRATIONS ON THE RATE OF EXCHANGE c/s(rv)* 0.0077*1* Sn(II) 0.0129 M Sn(IV) 3.99 H H+ 0.99 H SOr tl~ 540 hrs. 0.0104 M Sn(lI) 0.0129 M Sn(lV) 3.99 M H+ 0.99 M sor ti =481 hrs. 2 - 0.0129 M Sn(Il) 0.0129 U Sn(lV) 3.99 M H+ 0.99 M SOf t i =434 hrs. 5 * c d-f)+ t(hrs) 30.72 30.34 29.44 28.28 26.68 22.92 19.72 31.35 0.980 0.968 0.939 0.902 0.851 0.731 0.629 — 13.82 25.00 47.33 78.62 126.91 238.11 359.11 oo(c) 30.29 29.82 28.78 27.46 25.79 21.87 20.83 18.29 31.42 0.964 0.949 0.916 0.874 0.821 0.696 0.664 0.582 — - 13.85 25.03 48.05 78.65 126.95 238.11 271.37 359.21 00(c) 33.71 32.41 31.42 29.99 28.08 23.41 22.21 19.11 34.12 0.988 0.950 0.921 0.879 0.823 0.686 0.651 0.560 13.95 25.83 48.08 78.70 126.95 238.17 271.42 359.17 00(c) — — (Xoo-X) in equation 27, page 26 Value calculated on basis of mole fraction, equation 26, page 26 Calculated 82 TABLE XV cont. 0.0156 M Sn(II) 0.0129 M Sn(IV) 3.99 M H+ 0.99 M S0.= tjL = 409 hrs. 0.0970 M Sn(ll) 0.0126 M Sn(IV) 3.99 U H+ 0.99 M S O T tj = 519 hrs. c/s(IV) (l-f) t(hrs) 41.17 40.38 39.49 37.77 35.34 30.18 29.04 25.64 41.97 0.981 0.962 0.941 0.900 0.842 0.719 0.692 0.611 —— 13.97 25.13 48.13 78.80 127.00 238.22 271.47 360.00 oo(c) 20.84 19.32 18.67 17.97 16.97 15.72 14.99 13.30 11.12 24.18 0.862 0.799 0.772 0.743 0.702 0.650 0.620 0.550 0.460 — - 26.00 58.00 102.00 153.00 191.00 240.00 290.00 360.00 500.00 oo(c) TABLE XVI STUDY OF THE EFFECT OF VARIATIONS IN TIN (IV) CONCENTRATIONS ON THE RATE OF EXCHANGE 0.0129 M Sn(II) 0.00663 M Sn(lV) 3.99 M 0.99 HSO,H ~ 555 hrs. c/s(IV) (1-f) t(hrs) 42.52 41.71 41.15 40.25 39.13 38.06 36.95 32.31 25.83 42.91 0.991 0.972 0.959 0.938 0.912 0.887 0.861 0.753 0.602 ,B"■ 5.75 17.75 30.25 51.43 72.37 93.95 117.90 212.32 401.21 oo(c) 83 TABLE XVI cont. 0.0129 M Sn(II) 0.0086 M Sn(IV) 3.99 M 0.99 M SO." t^ = 518 hrs. 0.0129 M Sn(II) 0.0106 M Sn(IV) 3.99 MH + 0.99 M S0 = t^ = 480 hrs. 0.0129 M Sn(lI) 0.0129 M Sn(lY) 3.99 HH^ 0.99 M S04t |-420 hrs. 0.0129 M Sn(II) 0.0205 M Sn(IV) 3.99 H** 0.99 11 SOr =330 hrs. c/s(IV) (l-f) t(hrs) 38.90 38.24 37.81 36.53 35.47 34.58 33.10 31.07 22.62 38.94 0.999 0.982 0.971 0.938 0.911 0.888 0.850 0.798 0.581 — - 5.73 17.73 20.23 51.45 72.38 93.95 117.90 212.13 401.21 co(c) 33.99 33.49 32.82 31.71 30.96 29.82 28.86 25.01 35.63 0.954 0.940 0.921 0.890 0.869 0.837 0.810 0.702 -- 5.70 17.70 30.20 51.50 72.40 93.95 117.90 212.13 oo(c) 25.86 24.68 21.93 20.16 18.91 18.58 15.88 14.73 11.64 26.74 0.967 0.923 0.820 0.754 0.707 0.695 0.594 0.551 0.435 —— 26.00 58.00 102.00 153.00 191.00 240.00 290.00 360.00 500.00 oo(c) 24.78 24.11 23.58 22.48 21.58 20.85 19.42 15.81 25.06 0.989 0.962 0.941 0.897 0.861 0.832 0.775 0.631 5.62 17.62 30.12 51.31 72.30 93.88 119.83 212.13 oo(c) — 84 TABLE XVII STUDY OF THE EFFECT OF HYDROGEN ION ON THE RATE OF EXCHANGE 0.0128 M Sn(II) 0.0130 M Sn(IV) 3.30 M H* 0.99 M SOr =328 hrs. 0.0126 M Sn(IX) 0.0132 M Sn(IV) 3.60 M H^ 0.99 M sor ti =371 hrsT 2 0.0124 M Sn(II) 0.0134 M Sn(IV) 3.99 M 0.99 M SO4= 430 hrs. 0.0124 M Sn(II) 0.0134 M Sn(IV) 4.40 M 0.99 m sor t^ = 498 hrs. c/s(IV) (l-f) t(hrs) 52.61 45.00 40.42 34.28 28.80 25.37 21.08 57.60 0.910 0.780 0.699 0.594 0.499 0.439 0.365 — 50.00 100.00 170.00 245.00 325.00 400.00 505.50 oo(c) 62.58 56.21 50.77 43.49 36.52 32.83 26.76 70.21 0.892 0.800 0.723 0.619 0.520 0.468 0.381 50.00 100.00 170.00 245.00 325.00 400.00 505.00 00(c) 58.33 53.85 46.95 35.82 31.02 25.58 21.93 17.14 63.96 0.912 0.842 0.734 0.560 0.485 0.400 0.343 0.268 57.25 52.17 47.99 38.93 31.63 28.76 24.77 65.21 0.878 0.800 0.736 0.597 0.485 0.441 0.380 — - — - — - 50.00 100.00 175.00 350.00 450.00 550.00 650.00 800.00 00(c) 50.00 100.00 175.00 350.00 450.00 550.00 650.00 00(c) 85 TABLE XVII cont. 0.0125 M Sn(II) 0.0133 M Sn(IV) 4.80 M H+ o.99 m sor t i =575 hrs. 2 0.0124 M Sn(II) 0.0134 M Sn(IV) 5.00 M H+ 0.99 M SO^11 -603 hrs. 2 c/s(IV) (l-f) t(hrs) 57.68 54.44 49.32 40.77 35.65 31.69 29.10 23.98 64.81 0.890 0.840 0.761 0.629 0.550 0.484 0.449 0.370 — 50.00 100.00 175.00 350.00 450.00 550.00 650.00 800.00 00(c) 57.28 54.50 49.18 39.02 36.83 32.26 29.06 24.21 68.21 0.840 0.799 0.721 0.572 0.540 0.473 0.426 0.355 — 50.00 100.00 175.00 350.00 450.00 550.00 650.00 800.00 00(c) TABLE XVIXX STUDY OF THE EFFECT OF SULFATE ION ON THE RATE OF EXCHANGE 0.0120 H Sn(II) 0.0132 M Sn(XV) 3.-99 K H + 0.70 K SO^=466 hrs. c/s(IV) (l-f) t(hrs) 57.41 53.48 48.37 43.50 38.36 33.56 29.69 64.48 0.890 0.830 0.750 0.675 0.596 0.521 0.462 50.00 100.00 170.00 245.00 325.00 400.00 505.00 00(c) 86 TABLE XVIII cont. 0.0128M Sn(ll) 0.0130 M Sn(lV) 3.99 M H+ 0.00 M S0~ ti — 456 hrs. 0.0127 M Sn(Il) 0.0131 M Sn(IV) 3.99 MH + 0.99 m sor ti— 447 hrs. 2 0.0127 M Sn(II) 0.0131 M Sn(IV) 3.99 KH + 1.20 M S0^~ ti —414 hrs. 2 0.0128 M Sn(H) 0.0130 M Sn(TV) 3.99 MH + 1.30 M SO.f t i — 410 hrs. 2 c/s(IV) (l-f) t(hrs) 64.21 57.99 52.63 47.4-2 41.72 35.38 30.80 68.26 0.941 0.849 0.773 0.694 50.00 100.00 170.00 245.00 325.00 400.00 505.00 00(c) 63.50 57.42 52.11 46.41 39.00 36.37 31.49 68.42 0.929 0.838 0.761 0.678 0.571 0.532 0.460 — - 50.00 100.00 170.00 245.00 325.00 400.00 505.00 00(c) 58.66 51.93 47.13 34.53 29.33 24.80 16.53 66.66 0.880 0.779 0.707 0.518 0.440 0.372 0.319 0.248 — 50.00 100.00 175.00 350.00 450.00 550.00 650.00 800.00 00(c) 56.62 52.01 46.42 34.07 29.12 23.99 20.48 15.28 65.01 0.871 0.800 0.714 0.524 0.448 0.369 0.315 0.235 — 50.00 100.00 175.00 350.00 450.00 550.00 650.00 800.00 00(c) 21.26 0.611 0.518 0.451 - — 87 TABLE XVIII cont. 0.0126 M Sn(II) 0.0132 M Sn(IV) 3.99 M H+ 1.40 h sor ti - 405 hrs. 2 c/s(IV) (l-f) t(hrs) 58.43 57.17 48.67 29.55 30.68 26.56 21.58 16.27 66.40 0.880 0.861 0.733 0.445 0.462 0.400 0.325 0.245 50.00 100.00 175.00 350.00 450.00 550.00 650.00 800.00 00(c) - — TABLE XIX STUDY OF THE EFFECT OF SULFATE AND HYDROGEN IONS ON THE RATE OF EXCHANGE 0.0129 M Sn(H) 0.0129 M Sn(lV) 4.95 M H^ 1.21 M SO4ju - 6.16 ti -578 hrs. 2 0.0129 M Sn(ll) 0.0129 M Sn(IV) 5.08 M H + 1.24 M SO£ )i- 6.32 ti = 625 hrs. 2 c/s(IV) (l-f) t(hrs) 26.70 24.28 23.50 22.42 21.42 19.75 18.37 16.84 14.04 27.05 0.962 0.875 0.847 0.808 0.772 0.712 0.662 0.607 0.506 — — 26.00 58.00 102.00 153.00 191.00 240.00 290.00 360.00 500.00 00(c) 58.93 56.51 55.70 51.53 46.61 43.65 41.32 36.02 29.68 25.39 61.86 0.955 0.917 0.900 0.830 0.737 0.706 0.667 0.584 0.480 0.410 25.00 74.CO 122.00 175.00 245.00 300.00 365.00 460.00 655.00 792.00 oo(c) — — — 88 TABLE XIX cont. 0.0127 M Sn(II) 0.0131 M Sn(XV) 5.99 M H f 1.43 M S0.= m = 7.39 t|= 630 hrs. 0.0129 M Sn(II) 0.0129 M Sn(IV) 5.74 M H+ 1.40 M so r u =7.71 ti =690 hrs. z c/s(IV) (l-f) t (hrs) 26.50 25.35 24.36 22.92 21.21 19.97 18.32 18.10 15.54 13.00 26.74 0.991 0.948 0.911 0.857 0.793 0.747 0.685 0.677 0.581 0.516 — 26.00 58.00 102.00 153.00 191.00 240.00 290.00 360.00 500.00 600.00 00(c) 63.44 59.92. 57.31 53.62 51.39 47.65 42.71 39.02 33.35 0.929 0.866 0.841 0.786 0.753 0.699 0.625 0.584 0.489 0.425 23.00 71.00 118.00 172.00 242.00 297.00 362.00 457.00 652.00 789.00 00(c) 68.15 0.0129 H Sn(Il) 0.0129 M Sn(lV) 5.42 M 1.73 M SOr ju =8.00 4 ti =605 hrs. 5 63.14 60.99 56.77 53.66 48.75 45.23 41.28 38.77 34.77 26.57 69.92 — 0.904 0.873 0.811 0.769 0.698 0.648 0.592 0.555 0.497 0.380 — 25.00 74.00 121.00 175.00 245.00 300.00 365.00 460.00 655.00 792.00 00(c) TABLES XIX cont. 0.0129 M Sn(ll) 0*0129 M Sn(IV) 6.40 M H^“ 1.52 m sor~ u- 8.10 - 810 hrs. 0.0129 M Sn(II) 0.0129 M Sn(lT) 6.67 M H^~ 1.53 MSO^ n -8.20 ti'865 hrs. 2 0.0129 M Sn(II) 0.0129 M Sn(rV) 5.57 M H* 1.94 M S04" — 8.77 -600 hrs. c/s(IV) (l-f) t(hrs) 62.58 60.89 57.60 56.43 51.62 48.84 45.17 42.36 37.32 31.85 27.31 67.77 0.924 0.897 0.850 0.815 0.761 0.770 0.665 0.624 0.550 0.470 0.403 25.00 74.00 121.00 175.00 295.00 300.00 365.00 460.00 655.00 792.00 963.00 00(c) 62.67 59.09 56.83 53.63 51.93 49.15 46.12 42.89 37.47 33.49 29.12 70.50 0.889 0.842 0.806 0.762 0.735 0.697 0.640 0.608 0.532 0.475 0.413 25.00 74.00 121.00 175.00 245.00 300.00 365.00 460.00 655.00 792.00 963.00 00(c) 61.04 58.08 54.63 51.54 47.51 44.42 40.87 36.86 29.56 26.20 65.99 0.925 0.880 0.828 0.781 0.720 0.072 0.619 0.558 0.448 0.397 23.00 71.00 118.00 172.00 242.00 297.00 362.00 457.00 652.00 789.00 00(c) ----- — — 90 TABLE XIX cont. 0.0120 M Sn(XI) 0.0129 M Sn(lV) 5.70 MH+ 2.15 M SO," -9.51 -595 hrs. c/s(IV) (l-f) t(hrs) 63.16 59.76 55.15 50.75 48.91 46.47 41.15 38.20 30.19 26.11 65.12 0.970 0.918 0.846 0.779 0.751 0.697 0.631 0.586 0.403 0.401 23.00 71.00 118.00 172.00 242.00 297.00 362.00 457.00 652.00 789.00 co(c) — - TABLE XX STUDY OF THE EFFECT OF TEMPERATURE ON THE RATE OF EXCHANGE c/s(rv) (1-f) t(hrs) 0.0129 M Sn(II) 0.0129 M Sn(IV) 3.99 MH + 0.99 M SO," 37.8 °C ti =127 hrs. 5 43.79 40.40 35.98 30.93 26.31 22,95 17.09 12.91 0.823 0.760 0.676 0.581 0.495 0.431 0.321 0.243 25.00 40.00 67.00 90.00 118.00 150.00 195.00 245.00 oo(c) 0.0129 M Sn(II) 0.0129 M Sn(lV) 3.99 M H+ 0.99 M SO," 37.8°C k =127 hrs. 43.92 40.44 35.55 32.24 26.27 23.04 18.29 12.05 0.809 0.762 0.670 0.608 0.495 0.434 0.348 0.228 25.00 40.00 67.00 90.00 118.00 150.00 195.00 245.00 oo(c) 91 TABLE XX cont. 0.0120 M Sn(II) 0.0132 M Sn(IV) 3.99 HH + 0.99 M 30.“ 49.5 °C 4 t^ = 19 hrs. 0.0127 M Sn(II) 0.0131 K Sn(IV) 3.99 M H* 0.99 M SO/,49.5 °C ^1=20 hrs. A c/s(IV) (1-f) t(hrs) 46.00 35.52 27.13 22.02 17.76 16.30 15.31 13.82 11.42 9.55 61.13 0.746 0.580 0.441 0.360 0.290 0.268 0.251 0.226 0.187 0.156 — - 14.00 24.00 37.00 50.10 75.00 100.00 125.00 170.00 245.00 325.00 00(c) 44.02 35.50 26.89 22.13 17.62 16.81 15.90 14.23 12.12 10.32 61.64 0.714 0.576 0.436 0.358 0.286 0.272 0.258 0.231 0.196 0.168 14.00 24.00 37.00 50.00 75.00 100.00 125.00 170.00 245.00 325.00 00(c) 92 KINETIC STUDIES WITH 3.99 M HYDROGEN ION, CATALYZED BY CHLORIDE ION, - 4.98 TABLE XXI STUDY OF THE EFFECT OF VARIATIONS IN CHLORIDE ION ON THE RATE OF EXCHANGE c/s(IV)* 0.0129 M Sn(ll) 0.0129 M Sn(IV) 3.99 M H+ 0.015 M Cl" ti z =425 hrs. 0.0129 M Sn(II) 0.0129 M Sn(IV) 3.99 M H^ 0.030 M Cl" ti