5530
CARLH. BRURAKER, JR., [CONTRIBUTIONFROM
THE
A N D ANITA J E A N N E C O U R T
\'ol. TS
KEDZIECHEMICAL LABORATORY, MICHIGAN STATE UNIVERSITY AGRICULTURE AND APPLIEDSCIENCE]
OF
A Study of the Kinetics of the Reaction between Tin(I1) and Cerium(1V) in Aqueous Sulfuric Acid Solutions BY CARLH. BRUBAKER, JR.,
AND
ANITAJEANNE COURT'
RECEIVED MAY10, 1956 The kinetics of the reaction between tin(I1) and cerium(1V) in aqueous sulfuric acid was studied in the temperature range from 0 to 25'. Three types of reaction were observed corresponding to (1) low sulfate ion and tin(1V) concentration, (2) high sulfate ion and tin(1V) concentrations and ( 3 ) the presence of colloidal tin(1V) in the solutions. Only the first !ype seems subject to interpretation and appears to represent the stepwise oxidation of tin(I1) by the tris-sulfatocerate( IV) ion.
We have selected the reaction between tin(I1) copper was displaced from sulfuric acid solution by metallic and cerium(1V) in sulfuric acid solutions for study tin; aliquots of these solutions were analyzed by the dithiol method'* using Santomerse 30m as surfactant. Lithium sulfor several reasons. Our primary interest has fate was recrystallized, dried until the anhydrous form was been in connection with oxidations which may pro- obtained and weighed directly. Tin(1V) sulfate was prepared ceed either by direct transfer of several electrons by reaction of the metal with concentrated sulfuric acid a t from reducing to oxidizing agent or alternatively 190" and was recrystallized several times from the concenacid; tin(1V) oxide was obtained by ignition of the by a stepwise loss of one electron a t a time. We trated sulfate for 12 hours a t 500". Reagent grade sulfuric and perhave also been intrigued by the possibility of find- chloric acids were standardized volumetrically. ing evidence for the existence of tin(II1) as an interIn general, particular pains were taken to exclude air from mediate during oxidations of tin(I1) or reductions all solutions which were t o be used in the rate studies, although no different results were obtained when the reacof tin (IV) .?, tions were allowed to take place in solutions which were satuIn hydrochloric acid the reaction between tin(I1) rated with air (just before mixing the reactants). The soluand cerium(1V) is fast,3 but other reactions in this tions were prepared with air-free distilled water and were medium, such as the tin(I1)-uranium(V1) reac- stored under nitrogen. The spectrophotometer cells in which the reactions were carried out, were capped and made tion4 and the reduction of iron(II1) by tin(II)?r5-' reasonably air-tight with stopcock lubricant around the caps. have been found slow enough that the rates could Spectrophotometric Determination of the Rate of Reacbe measured. These reactions appear to go step- tion.-At 315 mp, absorbancies of tin(I1) and (IV) in sulfuric wise and in the second case the existence of a tin- acid are negligible; extinction coefficientsof cerium(1V) and are e4 = 5.58 X l o 3 and = 2.7 X 102,13respectively, (111) intermediate has indeed been postulated , 2 (111) in sulfuric acid concentrations from 1 t o 3 M . Thus the although the over-all mechanism appears to be a rate of reaction can be followed by absorbancy measurematter of interpretation.8.9 ments upon the solution, by making suitable correction for
Experimental Equipment and Reagents.-Reaction rates were determined by using a Beckman Model DU spectrophotometer. The solutions were kept in 1 cm. spectrophotometer cells throughout the course of the reaction and the temperature of the cell housing was controlled to A0.2' over the temperature range 0-25'. The source of each reagent and method of determination of its concentration was as follows: ammonium hexanitratocerate(IV), obtainable from the G. F. Smith Chemical Company, was recrystallized, converted to the hydroxide then dissolved in concentrated sulfuric acid according to the method of Smith and Flylo and the resulting solution was diluted with distilled water and standardized with iron(I1) in the presence of phosphoric acid using barium diphenylamine sulfonate as an indicator. Solutions of cerium(II1) were obtained by reduction of cerium(1V) solutions with sulfur dioxide and after oxidation with sodium bismuthate, cerium concentrations were determined as before. Tin( 11) solutions mere prepared by the method o f Noyes and Toabe" in that (1) Abstracted from a thesi4 submitted b y Anita Jeanne Court t o the School of Advanced Graduate Studies, Michigan S t a t e University, in partial fulfillment of the requirements for the Doctor of Philosophy Degree. (2) J . Weiss, .I. ( . / w ~ ?,So,. , , :WR ( l 9 t 4 ) . (3) A. E. Remick, THISJ O U R N A L 69, , 94 (1947). (1) R . L . h f w r e , thzil., 77, 1.504 (1955). ( 5 ) A . A . S o y e s , %. p h y s i k . Chenz., 1 6 , 546 (1695). (0) W. P.Timofeem, C;. E. Mucliin and W.G . Gurewitsch, ibid., 116, 161 (192.51. ( 7 ) 1'. 12. Duke and 12. C . Pinkerton, THIS JOURNAL,1 3 , 3045 (1931) ( 5 ) J . Gorin, ibid., 58, 1787 (1936). ( 0 ) R . A Robinson a n d N. II. Law, Ti,ai?s, Faraday SOC.,31, 8!W ( I 935). (10, ( > , 1' Smith arid Tt- 11 I:ly, A i r ~ l ( ' h e n , , 21, 1233 (1949). (11) A . A . S o y e i and ii. 'l'oabe, THISJOURNAI,, 61, 33-12 (1939).
cerium( 111) present. Solutions of tin(1I) and cerium(1V) were mixed in the spectrophotometric cells. Generally, 1 ml. of tin(I1) was forced into 2 ml. of cerium(1V); zero time was taken as the time mixing began; absorbancies were recorded for about 5 min. and then a t t m = 50 min. Whenever cerium(1V) > 2 tin(I1) a t t = 0, concentrations of both reactants were obtainable from the corrected optical densities. Because initial tin( 11)concentrations were only determined from cerium(1V) absorbancies, a t t m , and determinations of total tin, then if cerium(1V) < 2 tin(I1) a t t = 0 , a t t = 5 min. more standard cerium(1V) was added such that, a t t m , cerium(1V) > 2 tin(I1) = 0. In this way tin(I1) concentrations were calculable from the absorbancy data. The numerical values of k which are presented in this paper were determined graphically. The number of readings taken per experiment ranged from 15 to 25 with an average of 20 for all experiments reported. Sulfate Ion, Bisulfate Ion, Hydrogen Ion and Ionic Strength Dependence.-Since the reaction rate was most easily measured a t O", a study of the effect of the concentration of these ions and the ionic strength was performed a t this temperature. Raman spectral data (leading to activitycoefficient ratios) are available for sulfuric acid solutions at 25".14 For the purposes of this study, concentrations of thc various ions were estimated in the following manner. Young, Klotz and Singleterry's values of the thermodynamic dissociation constants for bisulfate ion were used and it was assumed that (1) the molal activity coefficient ratio, YH +. ~ s o ~ - / Y-, ~ sisoessentially ~ constant over the temperature range 0-25", (2) the correspondence (as obtained by these approximations a t 0 ' ) of the molal dissociation constant of W. Dupraw, private communication; rf A I . I'arnsworth Pekola, Anal. Chem., 26, 736 (1'353). A. I. Medalia and B. J. Byrne, ibid., 23, 463 (1951). H. M. Smith, P h . D . Dissertation, University of Chicago, 19-19 (15) I
