C. F. DECK,P. KING,JR., 672 R. J. CAMPION,
ANI)
A. C. WAHL
Inorganic Chemistry THE DEPARTMENT OF CHEMISTRY, \vASHINGTON USIVERSITY, ST. LOUIS,MISSOURI
CONTRIBUTION FROM
Kinetics of Electron Exchange between Hexacyanoferrate(I1) and -(III) Ionsla BY RAYMOND J. CAMPION,’b CHARLES F. DECK, PERRY KIKG,
J R . , ~ AKD ~
ARTHUR C. LVAHLld
Received A’overnbev 19, 1966 The rate of isotopic exchange between Fe(CN)s4- and Fe(CN)e3- has been measured in aqueous solutions containing a number of different electrolytes. The rate was found to be first order in the concentration of each reactant and to depend markedly on the nature and concentration of the cations present. Very little dependence was found on the anion concentration or on the ionic strength at constant cation concentration. T h e following values of the rate constant (4l-l sec-I) a t 0.1O and the activation energy (kcal/mole), respectively. were deterniined for alkaline solutions containing 0.01 M of the 1260 and 0.3; J C B H S ) ~ S250 + , and 0.6; (n-CdH;)?N-, 41 and 5.2; (n-C4Hg)rN+, 23 and 5.0; indicated cation: (CH3)4NC, (n-C6Hll)4S+,16 and 7.8; (C6Hj)4hs+,28 and 6.3; Co(CbHs)a+, 1050 and -1.0; K C , 230 and 6.0. T h e large and specific effects of different cations is attributed to the participation of cations in the reaction by rcduction of the Coulombic repulsion between the reactants in an activated complex. There is no evidence, however, that the cations participate in the actual transfer of an electron. Extrapolation to zero cation concentration gave values of 6.0 4l-Isec-I and 9.1 kcal/mole for the rate constant at 0.1’ and the activation energy, respectively, for very dilute solutions. From these values the free energy and entropy of activation were calculated to be 15.6 kcal/mole and -24 cal deg-’ mole-’, respectively. These values are compared with values calculated from a model based on the Marcus theory of electron transfer.
Introduction centration of the cations present, even those which are normally considered kinetically inert such as potasThe kinetics of isotopic exchange between the hexasium ion and the tetraalkylammonium ions. Difficulty cyanoferrate(I1) and -(III) ions is of considerable interwas experienced with the presence in reaction solutions est because exchange must occur by electron transfer, of trace impurities which affected the exchange rate and the complex ions being inert with respect to ligand exmade reproducible results difficult to achieve. In order change.2 Franck-Condon restrictions on the rate of to obtain reproducible results, small amounts of ethylelectron transfer should be minimal since the ions have enediaminetetraacetic acid were added to many resimilar structure^;^ on the other hand, Coulombic efaction solutions ; this addition resulted in slightly lower fects on the rate should be prominent because of the rates. large ionic charges. For these reasons, the experimental While this article mas being prepared for publication, results are of interest in themselves and also for coma study of the exchange reaction by a nuclear magnetic parison with theoretical models of electron-transfer resonance method was The nmr experireactions. ments were carried out a t much higher electrolyte conThe earliest investigations of the rate of isotopic centration than were the isotopic exchange experiexchange between ferro- and ferricyanide ions indicated ments, so direct comparison of the results is not posthat the rate was “unmeasurably large”; we found, sible; however, within large uncertainties of long exhowever, that the rate though large mas m e a s ~ r a b l e . ~ , ~ trapolations the results appear to be consistent. This article describes the results of our studies of the kinetics of the isotopic exchange reaction in aqueous Experimental Section solutions. Radioactivity.-The tracer used was either FeS5or Fe69. For Most of the solutions studied were alkaline because radioactivity measurcmcnts, samples containing Fe65 were early in the investigation i t was found that hydrogen mounted as zinc ferricyanide for counting, and the intensity of the 5.9-kev X-rays emitted was measured with a side-window, ion, even a t small concentrations, catalyzes the exflow-type, proportional counter. The counter window was made change, and we were unable to keep the pH constant of Cellophane or Scotch Tape and w-as coated with a thin layer of near neutrality without the use of buffers. Buffers rvere graphite (Aquadag) to make it conducting. The counter was to be avoided because, as reported in this article, the operated on a 907,argon-lOyo methane gas mixture (P-10 gas). rate is markedly dependent on the natwc and conFor radioactivity measurements of F P , solutions were placed (1) (a) Supported by t h e National Science Foundation under G i a n l s G196, 6-553, G-9493, G-24219, and GP-5939X and by the Washington University Computing Facilities through NSF Grant G-22298. A bstracted in part from thefollowing P h . D . theses: C. F. Deck (1956), P. King, J r . (iYSO), and R. J. Campion (1963), Washington University, St. Louis, Mo. 1’1.esented in p a r t a t American Chemical Society Xational Meetings in St. Louis, Mo., March 1961 and in Atlantic City, h-,J., Sept 1952. ( b ) David Sarnoff Fellow, 1959-1961. (c) Phillips Petroleum Fellow, 1957-1958; General Atomics Fellow, 1958-1959. ( d j Author to whom inquiries should be addressed. (2) A. G. MacDiarmid and S . F. Hall, J . A m . C h c m .Soc., 7 6 , 422 (1954). (3) T h e Fe-C-S distances in Fe(CK)s4- and Ire(CNj6’- have not been measured accurately. However, t h e Fe-Fe (Fe-C-S-Fej distances in Prussian blue, KFelll[Fell(CNjs], and Berlin green, [FeIII(CN)sl, have both been determined t o be 6.1 A [J. F. Keggin and F.D . Lliles, Natzwe, 137, 677 (1936)l. (4) A . C. Wahl and C. F.Deck, J . A m Cizenz. SOC.,7 6 , 4054 (1934). (5) A. C. Wahl. Z . Elektrokhenz.. 64. 90 f l 9 6 0 ) .
in test tubes, and the intensity of the y rays emitted was nieasured with a well-type sodium iodide scintillation counter. Relativc specific activity values were calculated by dividing the measured counting rate for an aliquot of ferricyanide solutio11 by the absorbance at 4200 4 , determined for a 1-cm path length of solution. Tagged ferrocyanide ion was prepared from iron(II1) chloride containing Fe56 or FeSg. First the iron(II1) chloride was purified by two extractions from 6-8 Af HC1 into isopropyl ether. Then iron(II1) hydroxide was precipitated from ammonia solu(6) A. Loewenstein, M . Shporer, and G . Navon, J . A m . Chenz. SOC.,85, 2855 (1963). (7) M. Shporer. G . R o n , A . 1.oewenstein. and G. Navon, I x o v e . C h e m . , 4 , 361 (1965).
Vol. 6 , N o . 4, April 1967
ELECTRON EXCHANGE BETWEEN HEXACYANOFERRATE(II) AND -(III) IONS 673
tion, and the precipitate was suspended in a concentrated potasprepared from potassium bromide and silver oxide by the procedure described for the tetraalkylammonium hydroxide solutions. sium cyanide solution. The suspension was heated in a steam bath for approximately 1 hr until the solution was clear and Tetraphenylarsonium chloride was obtained from a variety of faintly yellow in color. Tagged potassium ferrocyanide was sources and was purified by recrystallization twice from sodium precipitated by addition of ethanol and purified by recrystallizachloride solution and twice from water. Ethylenediaminetetration from water-ethanol solution. acetic acid was obtained from Versenes Inc. Water was redisTagged tetraphenylarsonium ferricyanide was prepared from tilled from alkaline permanganate solution. All other reagents were of analytical grade. aqueous solutions of Fe*(CN)e4-. A solution was treated with a Reactant Solutions.-For the early exchange experiments in solution of iodine in carbon tetrachloride, and Fe*(CNha- was potassium hydroxide solutions without added ethylenediamineextracted into a chloroform solution of tetraphenylarsonium tetraacetic acid, KdFe*(CN)e and K3Fe(CN)6 were dissolved in chloride. The solution was washed twice with water, and then appropriate solutions. For most other experiments, [ (C&)4diethyl ether was added to precipitate [ (C6H5)4As]sFe*(CN)6. The compound was purified by repeated precipitation from were dissolved in sepaAs]3Fe(cN)~and [ (CsHs~4As]3Fe*(CNj~ rate basic solutions containing EDTA“$ Fe*(CN)eS- was then chloroform and was then dissolved in water. The solution was filtered through a Millipore-VC filter and was allowed to evaporate reduced by hydrogen a t a platinum plate; and each solution was diluted to volume, filtered through Millipore-VC filters on stainto dryness in the air stream coming through the filter. The radiochemical purity of each preparation of a tagged less-steel funnels, and collected in quartz or polyethylene containers. compound was checked by isotopic exchange between Fe*(CiX)e4- and Fe(CT\’)a3-. After essentially complete exchange had According to our interpretation of the data (see Discussion) occurred the specific activities of the two reactants were comthe small concentrations of (CsHs)aAs+ (usually 0.0006 M ) introduced with reactant ions had little effect on the rate, increasing pared, or the specific activity of ferricyanide ion was compared with the average specific activity determined after oxidation of the rate constant by only -1 M-’ sec-l a t 0.1”. We checked FefC”le4- to Fe(CN)aa-. For all preparations the specific acthis conclusion by removing (C,&)&+ from two sets of reactant solutions by passing them through an Amberlite IR 120 cationtivity values compared were the same within an experimental exchange column in the hydrogen form and into basic solution, uncertainty of a few per cent. R e a g e n t ~ . - K ~ F e ( C N )and ~ K ~ F ~ ( C N ) C ~were H~O analytical 0.01 M KOH or 0.003 M [ ( ~ - C ~ H O ) ~ N ] OThe H . measured rates agreed with those predicted. grade reagents recrystallized twice from water. [ (CsH5)rAs]J?e(CN)s was prepared from &Fe(CN)e and [ ( C ~ H ~ ) ~ A S ] C ~Separation of Reactants.-The reactants were separated by by a procedure similar to the one described for the tagged comextracting Fe(CN)e3- into a solution of tetraphenylarsonium pound. KaCo(CN)e was prepared by the method of Bigelows chloride in chloroform12 in the presence of Co(CN)e*- and Ruand recrystallized four times from water. K~Ru(CN)G was pur(CN’Js4-, ions which minimized the “zero-time’’ e x c h a r ~ g e . ~ , ~ ~ chased from the K and K Laboratory or was prepared by the In most experiments 8 ml of aqueous reaction solution was method of DeFord and Davids0n.O The prepared material was M in mixed first with 2 ml of aqueous quench solution 6 X purified by recrystallization after treating a solution with iodine &Co(CN)s and 6 X M in K&u(CN)s and then with 5 ml of in carbon tetrachloride to oxidize Fe(CN)e4- impurity to Fechloroform solution 0.1 M in [ ( C ~ H ~ ) ~ A S ]Mixing C ~ . was continued for 5 sec, and then the phases were separated by centrifu(CiY)e3-, which does not crystallize with K ~ R u ( C N ) B * ~ H ~ O . Solutions of te traalkylammonium hydroxides were prepared gation for 30 sec. The Fe(CN)eS- was back extracted from the by mixing solutions of the recrystallized bromide salts (Eastchloroform solution into an aqueous 0.75 M KNOs solution for man Organic Chemicals) with excess silver oxide (Mallinckrodt specific-activity measurements. “Zero-time” exchange varied purified). After reaction was complete the basic solution was somewhat with the composition of the reaction solution but was filtered, passed slowly through a Dowex 1-X4 anion-exchange reproducible for a given set of conditions. column in the hydroxide form, and filtered through a MilliporeApparatus.-A diagram of the apparatus used for most experiVC filter. ments is shown in Figure 1; 4 ml of one reactant solution was Solutions of [ ( CG&’JIAS] OH were prepared by mixing soluplaced in reservoir R-1 and 4 ml of the other was placed in resertions of [ (CeHs)&] c1 and gMnO4 to precipitate [ (CeH5)&]voir R-2, 2 ml of aqueous quench solution was placed in vessel MnO4, which was filtered, washed thoroughly with water, and Q-1,and 5 ml of chloroform solution was placed in vessel Q-2. treated with 30y0 hydrogen peroxide. After removal of manThe reactant solutions were mixed by applying a pressure of 15 psi ganese dioxide by centrifugation, the solution of [ (CBHQ)~AS]OH of nitrogen through tube S-1 to force the solutions together in was passed through either a Millipore-VC filter or a column of stopcock M and through it into reaction vessel RV where the Amberlite IRA-400 (OH-) resin. stirrer was rotating a t -2000 rpm. The reaction was quenched A solution of [Co(CbH5)2]OH was prepared from a benzene by applying the pressure of nitrogen through tube S-2 forcing the solution of cobaltocene supplied by Araaphoe Chemicals, Inc. aqueous quench solution and then the chloroform solution into The cobaltocene was oxidized by aqueous hydrogen peroxide, and the reaction mixture. cobalticinium picrate was precipitated from the aqueous phase. The application of pressure was controlled by solenoid valves. The method of WilkinsonIo was used for purification and converThe time of exchange was taken to be the time interval between sion to the chloride. A solution of the [Co(CaHs)n]Clwas treated activation of the valves and was measured by an electric timer with silver oxide, and the resulting solution was passed through a with a precision of ctO.01 sec. Any difference between this time Dowex 1-X4 (OH-) anion-exchange column and then through a interval and the “true reaction time” was constant for a given Millipore-VC filter. set of conditions so affected only the intercept, not the slope, of an Solutions of potassium hydroxide were prepared in several exchange curve. ways. In the early work, before ethylenediaminetetraacetic The apparatus used in the early work differed from the one acid was added to reaction solutions, Baker and Adamson described by keeping the reaction solution under pressure in a “Special Low Carbonate” potassium hydroxide was used. Later, reservoir hetween two stopcocks for the desired time interval when ethylenediaminetetraacetic acid was added, no difference and then delivering the reaction solution into a centrifuge tube was observed in exchange rates determined using solutions precontaining the two-phase quenching mixture, which was being pared from Baker and Adamson “Reagent Special” potassium stirred. This apparatus worked well for experiments in which hydroxide and filtered through Millipore-VC filters and those
( 8 ) J. H. Bigelow, Inovg. Syn., 2, 225 (1946). (9) D. D. DeFord and A. W. Davidson, J . A m . Chem. SOL.,78, 1469 (1951). (10) G . Wilkinson, ibid., 74, 6148 (1952).
(11) The symbol EDTAI- is used for the ethylenediaminetetraacetate ion. (12) L. Eimer and R . W. Dodson, Brookhaven National Laboratory Report No. BNL 93 (S-s), 1950, p 69 (unpublished). (13) K . Wolfsberg, M.S. Thesis, Washington University, S t . Louis, Mo. (unpublished).
674 R.J. CAMPION, C. F. DECK,P. KING,JR.,
AND
A. C. WAHL
Inorganic Chemistry
n
I
Figure l.-Drawing of the apparatus used for measurement of exchange rates. The operation of the apparatus is described in the text. reaction times were short; during long reaction times leaks somctimes occurred. Interpretation of Data.-The fraction exchange was calculated by dividing the Fc(CI\T)e3- specific activity by its value after a long exchange time when isotopic equilibrium had been established or by the average specific activity of the reaction mixture determined after oxidation of Fe(Cx)s4-. The precision of the fraction-exchange determinations was estimated to be ~ 2 7 ~ . The half-time for exchange, til2, was taken from an exchange curve, a semilogarithmic plot of 1 - fraction exchange ~ s the . reaction time. A typical exchange curve is shown in Figure 2. The rate of exchange was calculated from the half-time through the familiar relationshipI4
0,2
I
0
=
k[Fe(Cx)64-] [ F e ( C N ) P l
(3)
in which the concentrations are in units of gram-formula weights per liter, applies under a variety of conditions, and we assume that it applies under all conditions that we have studied. Of course the value of k depends on particular conditions, as is illustrated in Figure 3, and this effect is discussed in later sections. (14) See, for example, 0. E. Meyers and R. J. Prestwood, “Radioactivity Applied t o Chemistry,” -4. C. Wahl and N. A. Bonner, Ed., John Wiley and Sons, Inc., New York, N. Y., p 7.
I
I
1
I
50
100
150
200
250
TIME, sec
The uncertainty in measured values of k was estimated to be 5%. The uncertainty in concentration values was titimated to be
