3388
GEOFFREY DAVIES AND KAY0. WATKINS
The Kinetics of Some Oxidation-Reduction Reactions Involving Cobalt(111) in Aqueous Perchloric Acid1 by Geoffrey Davies2 and Kay 0. Watkins3 Chemistry Department, Brookhaven ATational Laboratory, Upton, New York 11978
(Received January
14,1970)
The stoichiometries and kinetics of the reactions of cobalt (111) with hydrogen peroxide, nitrous acid, hydroquinone, bromide, iodide, and thiocyanate have been investigated in aqueous perchloric acid solution over wide concentration ranges. Stoichiometries were determined using standard analytical procedures and kinetic data were obtained from conventional and stopped-flow spectrophotometry. The reactions were all found to be first order in each reactant and intermediate complex formation could not be detected either kinetically or by spectrophotometric means. The second-order rate constants generally obey the relationship leobsd = a b/(H+) , where a and b are interpreted as referring to the reactions of CO,,~+ and COOH,~~+, respectively, with the predominant reductant species. The mechanisms of these reactions are discussed and a comparison is drawn between the rates of oxidation and complexation reactions of cobalt(II1) and the rates of complexation reactions of iron(II1). It is concluded that the rate-determining steps in the reactions between CoOH,,*+ and HzOz, “ 0 2 , Br-, and SCN- involve the replacement of water molecules in the inner sphere of the oxidant. The faster reactions with hydroquinone and iodide evidently proceed by a different mechanism in which the OH- ligand may be involved as a bridging group between the reactants.
+
Introduction The instability of perchloric acid solutions of cobalt(111) has limited the number of investigations involving C0aq3+and CoOHaq2+ions. Some systems which have been studied previously are the reactions of cobalt(II1) with water14-* chloride ionslg malic acid,1° iron(II),” and cerium(III).l* These studies have led to conflicting conclusions13 concerning the nature of the oxidizing species. We have investigated a series of oxidations by Coaq3+ions in perchloric acid solution in an attempt to obtain additional information concerning the extent of hydrolysis of C O , ~ ~ + , ~ the J O Jexistence ~ of polymeric specie^,^-^^^^ and the mechanisms of reactions of cobalt (111) under weakly complexing conditions. I n this paper we report on the stoichiometry and kinetics of the oxidation of hydrogen peroxide, hydroquinone, nitrous acid, and bromide, iodide, and thiocyanate ions by cobalt (111).
Experimental Section Materials. Stock solutions approximately 7 M in sodium perchlorate were prepared by neutralizing sodium carbonate with perchloric acid and were standardized gravimetrically. The stock solutions of cobalt(II1) were prepared as described p r e v i o u ~ l yusing ,~ cobalt(I1) perchlorate hexahydrate as starting material. Spectrophotometric measurements were used to determine the cobalt (111) and cobalt(I1) concentrations (€606 35.3 and E609 4.84, respectively). The conce,ntration of cobalt(II1) was also checked by adding an aliquot to an excess of standardized iron(I1) solution and titrating the excess iron(I1) with standard chromium(VI) solution, using diphenylamine as indicator. The Journal of Physical Chemistry, Vol. 74, No. 18, 1970
All other reagents were of analytical grade and triply distilled water was used throughout. Hydrogen peroxide solutions (prepared by dilut,ion of 30% wt/vol hydrogen peroxide) were standardized either by titration with cerium(1V) using ferroin as indicahor, or by estimation as the titanium(1V) complex (eqI5 731).14 Sodium nitrite solutions were standardized either by addition of excess standard manganese(VI1) , followed by addition of excess iodide and titration of the resulting iodine with standard thiosulfate solution, or spectrophotometrically as nitrous acid 13.7, e337 23.3, e347 36.7, ~ 3 50.3, 6 ~ €372 52.0, and €386 30.6). Bromine was determined gravimetrically as AgBr and iodide as AgI. Hydroquinone stock solutions were always (1) Research performed under the auspices of the U. S. Atomic Energy Commission. (2) Address inquiries to University Chemical Laboratory, University of Kent, Canterbury, Kent, England. (3) On sabbatical leave from Adams State College, Alamosa, Colorado. (4) C. E. H. Bawn and A. G. White, J . Chem. Sac., 331 (1951). (5) D. W. Weiser, Ph.D. Thesis, University of Chicago, 1956. (6) J. H. Baxendale and C. F. Wells, Trans. Faraday Soc., 53, 800 (1957). (7) H . Taube, J . Gen. PhysioZ., 29, 49 (1966). (8) M , Anbar and I. Pecht, J . Amer. Chem. SOC., 89, 2563 (1967). (9) T. J. Conocohioli, G. H. Nancollas, and S . Sutin, Inorg. Chem., 5, 1 (1966). (10) J. Hill and A . McAuley, J . Chem. Sac. A , 1169, 2405 (1968). (11) L. E. Bennett and J. C. Sheppard, J . P h y s . C‘hem., 66, 1275 (1962). (12) L. H . Sutcliffe and J. R. Weber, Trans. Faraday Soc., 52, 1225 (1956). (13) C. F. Wells, Discuss. Faraday Soc., 46, 197 (1968), and refer-
ences therein. (14) R. Bailey and D. F. Bolts, AnaZ. Chem., 31, 117 (1959).
KINETICSOF OXIDATION-REDUCTION REACTIONS INVOLVING COBALT(III) freshly prepared in darkened flasks using deoxygenated water and were standardized by titration with cerium(IV) immediately before use. Thiocyanate was determined by the Volhard procedure. Stoichiometry Measurements. Most stoichiometry measurements were made by adding a solution of cobalt(II1) to an excess of reductant. In some cases the concentration of the reductant remaining after reaction was deternlined, while in others the concentration of the oxidation product was estimated spectrophotometrically. The stoichiometry of the reaction with hydrogen peroxide was determined by estimating the remaining reductant, as described above. The concentration ranges used were (CoIII) = 0.62-6.7X M, (H202) = 2.322 X M , and (H+) = 0.15-1.50 M . The reaction with nitrous acid was studied by titrating the remaining oxidant or estimating the remaining nitrous acid spectrophotometrically, as described above. The concentration ranges used were (CoIII) = 9.93 M ,("02) = 1.06-11.1 X X 10-5 to 1.19 X 10-3 M ,and (H+) = 0.15-3.10 M . Bromine formed in the oxidation of excess bromide was determined by adding the product solution to an excess of iodide and titrating the iodine formed with standard thiosulfate solution (starch end point) or by spectrophotometric measurements as Bra-, 4.28 X lo2 at (Br-) = 1.84 X 10-1 M). The concentration ranges used were (CoIII) = 1.7-7.1 X 10-3 M ,(Br-) = 1.00-1.84 X lo-' M, and (H+) = 0.30-1.50 M . Iodine formed in the reaction with iodide was determined as described above. The concentration ranges used were (CoIII) = 0.35-10.4 X 10-3 M , (I-) = 0.2-0.4 M, and (H+) = 0.30-1.80 M . The product p-benzoquinone formed in the oxidation of p-hydroquinone, H2Q, was determined spectrophotometrically ( ~ 2 41.82 ~ X lo4, €260 2.11 X lo4, €260 0.88 X lO4)l5 in the range (Co'II) = 5.44-43.5 X M , (H2Q) = 2.44 X M, and (H+) = 0.12-0.96 M . The stoichiometry in the presence of an excess of thiocyanate was determined in a three-chamber mixing apparatus as follows. Various measured volumes of a standardized cobalt(II1) solution in 3 M perchloric acid were added to 10 ml of 1.114 M sodium thiocyanate solution, and after vigorous stirring the mixture was rapidly added to 10 ml of 5 M sodium iodide solution.16 After dilution to about 200 ml, the iodine formed was titrated as described above. The concentration ranges used were (CoIII) = 2.466.82 X lop3 M , (SCN-) = 0.28-0.56 M , (I-) = 1.02.5 M, and (H+) = 0.6-1.5 M . Variation of the thiocyanate and iodide concentrations had no observable effect on the stoichiometric results. The product of oxidation was identified as follows. Cobalt(II1) solution (25 ml) (2.51 X lop2 M in 3 M HClO4) was added to 10 ml of 1.111l.I sodium thiocyanate solution in a separatory funnel. Diethyl ether (25 ml) was then added and the mixture was shaken for 5 min. After
3389
standing a further 10 min the yellow solid which had formed at the interface was filtered off. The combined products from repeated experiments were washed several times with water followed by ether and were then dried overnight in vacuo over calcium chloride." The yellow-orange solid was then analyzed for carbon, nitrogen, and sulfur, with the following results. l8 Anal. Calcd for (CNS),: C, 1; Tu', 1.17; S, 2.67. Found: C, 1;N, 1.26;S, 2.71. Kinetic Measurements. Most of the kinetic measurements were performed using the stopped-flow apparatus described previ0us1y.l~ The Cary Model 14 spectrophotometer was used for some of the slower runs. In all cases the reductant was present in sufficient excess to ensure pseudo-first-order conditions. Each rate constant is the average of at least six independent measurements. The reactions with hydrogen peroxide, nitrous acid, and thiocyanate were followed by measuring the rate of disappearance of cobalt(II1) in the wavelength range 240-605 nm. The choice of wavelength depended on the initial cobalt (111) concentration employed. In the reaction with bromide the appearance of tribromide ion was monitored in the wavelength range 250-325 nm. At lower bromide ion concentrations the disappearance of cobalt(II1) was followed at 240-250 nm. The reaction with iodide was followed by measuring the rate of appearance of triiodide ion at 288-330 nm. The rate of appearance of p-benzoquinone was measured at 247-250 nm in the reaction with hydroquinone.
Results The stoichiometric measurements established that the following reactions occurred under the experimental conditions used.20
(15) G. Davies and K. Kustin, Trans. Faraday Soc., 65, 1630 (1969). (16) C. E. Vandersee and A. S. Quist, Inorg. Chem., 5 , 1238 (1966). (17) Control experiments showed that no solid material appeared in the, absence of cobalt(II1) under the experimental conditions. (18) Only the relative composition is presented since the solid product was contaminated with sodium thiocyanate and water. The carbon, nitrogen, and sulfur accounted for about 70% of the material. (19) G. Dulz and N. Sutin, Inorg. Chem., 2, 917 (1963). (20) The upper limit of uncertainty is 2% of the given stoichiometry in all cases. The Journal of Physical Chemistry, Vol. 7 4 , N o . 18,1970
3390
GEOFFREY DAVIES AND KAY0. WATKINS
The ultraviolet and visible spectra of the immediate products of the reaction of cobalt(II1) with hydroquinone, bromide, and iodide ions were identical with those of p-benzoquinone (A, 247 nm) , l 5 tribromide ion (A, 265 nm), and triiodide ion (maxima at 288 and 348 nm), respectively. The kinetic dataz1indicate that each reaction is accurately first order in each The observed second order rate con~ ~ wavestants are independent of (CoIII), ( C O I I ) , and length of measurement. 2E The latter observation suggests that competing photochemical reactions play only a minor role and that the slow formation of appreciable concentrations of cobalt(II1)-reductant complexes, such as those found in the reactions with chlorideg and malic acid,lO does not occur during the course of the reactions studied in this work. The rate law for each reaction is thus
1 d(ColI1) d(product) n dt dt
_ - _ _ _ . -
=
kobsd(Co"I))(B)
(7)
where
In these equations B is the reductant, n is the appropriate stoichiometric factor, and (H+) is assumed equal to (HC10.J. The parameters a and b may be obtained from the intercept and slope, respectively, of plots of kobsd us. l/(HClOS. Examples of plots of eq 8 for reactions a t 25.0" are given in Figure 1, which illustrates the point that for some reactions the intercepts of plots of kobsd us. l/(H+) are very uncertain (e.g., for H20zand Br-), while for others the intercept is much better established (e.g., for SCN- and H2Q). Thus, only upper limits could be obtained for a in the reactions with hydrogen peroxide and bromide. The following mechanism is consistent with the observed rate law
(13) in which steps 10 and 11 are rate-determining and B is a radical species. The observed rate law (eq 7) is predicted by this mechanism if Kh/(H+)
