May, 1963
RATEO F ELECTROX TRAKSFER BETWEEK
I R O N ( I 1 ) AXD I R O K ( I I 1 ) I O K S
1025
THE EFFECT OF SEVERAL OXY-ACIDS ON THE RATE OF ELECTRON TRANSFER BETWEEK IRON(I1) AND IRON(II1) IOKS I N PERCHLORIC ACID BY JOHN C. S H E P P A RAND D ~ ~LARRYC. BROTTS'~ Department of Chemistry, Sun Diego State College, San Dirgo, Calzfornza Receaved October 6 , 196W The specific rate constants for the iron(I1)-iron( 111) sulfate, acid phosphate, and oxalate paths for electron transfer between these ions in perchloric acid have been determined. Assuming that these reactions are between iron(II1) complexes and iron(I1) ions and the heats of formation of the iron(II1) complexes of these anions are negligible, the specific rate constants a t 0' for these paths are: sulfate ion, 100 f 20 J4-l see.-'; acid phosphate ion, 520 f 50 M-l sec.-l; and oxalate ion, 250 f 20 A4-l set.-'. The formal activation energies for these paths are 13.5 f 2, 15 f 3, and 21 f 2 kcal./mole, respectively.
The effect of several anions on the rate of electron transfer between iron(J1) and iron(II1) ion has been extensively studied2-6 with the purpose of elucidating the mechanism of the oxidation-reduction process. Since the energies and entropies of activation as well as the specific rate constants for most paths of these reactions are similar, suggestions6-8 have been made that a common path exists. The purpose of this research was the investigation of the influence of several large oxygenated anions on the rate of electron transfer between iron(I1) and iron(II1) ions. Sulfate and acid phosphate ions were chosen for this study because these ions are large, symmetrical, and about the same size. The results may be compared to the effects observed for other anions. Oxalate ion catalysis is of interest because this ion has a different geometry relative to phosphate and sulfate ions. Experimental Iron Tracer.-Iron-59, obtained from the Oak Ridge National Laboratory, was used in these experiments. The tracer contained less than 0.2% cobalt-BO at the time it was received and was further purified by the following method. The tracer, along with a few milligrams of cobalt(I1) chloride, was added to a few ml. of concentrated hydrochloric acid and was absorbed on an IRA-400 anion-exchange resin column, previously washed with concentrated hydrochloric acid. After the iron(II1) and cobalt(I1) were absorbed ort the column, the cobalt(I1) ion was eluted with a 4 M hydrochloric acid and discarded. The iron-59 tracer was then eluted with 1 144 hydrochloric acid and the eluate was evaporated to dryness with concentrated perchloric acid to remove any chloride ion present. The residue containing the tracer was then taken up in 6 M perchloric acid and small aliquots of the resulting solution were used in the exchange experiments. Radiochemical purity of the tracer was established by the fact that it decayed over a period of several months with a 45 day half-life. &o, comparison of the gamma spectrum of this tracer with an iron-59 standard indicated that its radiochemical purity was greater than 98%. Exc hange experiments under conditions used by Silverman and DodEionZaas well as Hudis and Wahl*b gave essentially identical results and this was taken as additional proof of radiochemical purity of the tracer. Reagents.-All reagents used in this study were reagent grade and were used without further purification. Exchange Procedure.--The reaction vessel used in these experiments was a pear-shaped, three-necked flask. The flask, along (1) (a) Now with General Electric Company, Hanford Laboratories. Richland. Washington; (b) National Suience Foundation undergraduate research participant, 1960 nnd 1961. ( I ? ) ( a ) J. Silverman and R. W. D i d s m , J . P h y s . , Chem. 66, 846 (1952); ( b ) .J. Hudis and .4. C. Fl'ahl. J . A m . Chem. Soc., 76, 4153 (1953). (3) G. d . Laurence, Trans. Faraday Sor., 63, 1326 11957). (4) R. A . H?rne, J . Phys. Chem.. 64, 1512 (1960). (5) 13. Bunn and F. 5. Dainton, T ~ W L FSw. a d a y Soc., 66, 1267 (1959). (6) R , W. Dodson and N. Drtvidson, J . Phys. Chem., 66, 866 (1952). (7) W. L. Reynnlds and R . IT-. L u m r y , J . Chem. Phys., 23, 2460 (1955). (8) D. R. Stranks, "Modern Co-nrdination Chemistry," J. Lewis and R . G. W'ilkins. editarg, lnteraicience Puhliebersi Inc,, New York, Na Y , , 1960, p, 154.
with its contents, was maintained at the desired temperature by immersing it in a refrigerated, constant temperature bath. It is estimated that the temperature of the reaction solution varied not more than 0.1" during these experiments. To 100 ml. of the solution containing the desired concentration of the anion of interest as well as 0.53 M perchloric acid, enough iron(I1) and iron( 111)perchlorate were added t o make the total M . The iron concentration about 1.0 X 1 0 - ~to 2.5 X exchange reaction was initiated by injerting a small aliquot of the iron(II1) tracer into the reaction mixture and mixing was accomplished by the use of a small electric motor-driven glass stirrer. During the course of the reaction about eight 5-ml. aliquots were removed at appropriate time intervals by means of an automatic pipet and injected into the quenching solution. Two of the aliquots were infinite time samples. The quenching solution consisted of 5 ml. saturated sodium acetate, 5 ml. 0.1% bipyridyl, and 1 ml. of 0.1 M lanthanum nitrate. One ml. of 15 M ammonium hydroxide was then added to precipitate lanthanum hydroxide which coprecipitates m y iron( 111)hydroxide present. The quenched reaction solutions were allowed to stand a few minutes to permit the formation of lanthanum hydroxide and then centrifuged. Five ml. of the supernatant solution, containing only the iron(I1) bipyridyl complex, was transferred to a counting tube and counted in a well-type scintillation counter until the counting error was about one per cent
Results Under the conditions used in these experiments, this reaction obeyed the exponential exchange law. Assuming that the reaction is first order with respect to the iroii(I1) and iron(II1) ion concentrations as found by the other investigators,2 the specific rate constant, ' k ' , for some experimental condition is related to the half-time ti/, of the exchange by
The data found in Fig. 1, 2 , and 3 were obtained using this equation. Analysis of the exchange curves having half-times greater than 0.76 minute indicates thdt the error in the specific rate constants was less than 5%. Those runs with half-times of less than 0.75 minute had errors of not more than 10%. It is interesting to note that the 10 and 20' oxalate ion data are in agreement with those of Horiie4 but the 0' data are not. In an effort to resolve this discrepancy the following were done: Two different sources of oxalate ion, oxalic acid, and ammonium oxalate were used. The reaction was initiated in a different way by having the tracer present in the iron(II1) solution before mixing with the iron(I1) solution. Both investigators ran exchange experiments. Within the experimental error of 5% the results of the experiments described above agreed with Horiie's4 data a t 10 and 20' but not at 0'. The experimental errors reported by Home4for the data a t these temperatures are large, especially at
JOHN c. S H E P P 4 R D A N D LARRYc. BROWK
1026
0.
440-
400360-
320-
7'01. 6'7
/
0
'k.
k.
[H~SO x lo3 ~ F. Fig. 1.-The effect of sulfuric acid on the rate of electron transfer between iron(I1) and iron(II1) ions in 0.81 formal perchloric acid and a t an ionic strength of 1.0 ionic strength adjusted with sodium perchlorate.
Fig. 3.-The effect of oxalic acid on the rate of electron transfer between iron(I1) and iron(II1) ions in 0.53 formal perchloric acid.
90
3.8
I
I
80
I
1
15 i 3 kcallmole
E~,=
3.4
70
GO
EaCts 21 t 2 kcallmole
-
_y
-g
3.0
'k'. 50 40
2.6
-
f
30 2.2
I
I
.3
3.4
-
I
3.5 1lT x
3.6
3.8
Id.
Fig. 4.-An Arrhenius plot of log k against 1 / T , yielding the formal activation energies for oxalic, phosphoric, and sulfuric acid paths of the iron(I1)-iron(II1) exchange reaction.
Fig. 2.-The effect of phosphoric acid on the rate of electron transfer between iron(I1) and iron(II1) ions in 0.53 formal perchloric acid.
0'. When these errors are considered in the iiiterpretation of the data, widely different results and coilclusions can be obtained.
Discussion who have studied the iiiflueiice Most of anions on the rate of electron-transfer between iron(11) and iron(II1) ions have assumed that complexes of iron(1II) bring the anion into the transition state. Accordingly, entropies and energies were calculated for these paths, aiid proposalso-8 concerning the nature of the electron-transfer process itself or the geometry of
RATEOF ELECTROS TRANSFER BETWEEN IROX(II)AND IROS(III) Ioss
May, 1963
the transition state were made. Some proposals have been made 011 the basis of even less kinetic data.g As Taubelo has pointed out, it is not possible from rate law studies to establish either the geometry or the mode of formation of the transition state for electron-transfer reactions involving substitution labile reactants and products. Iron(I1) and iron(II1) ions fall into this category, consequently there is an indeterminacy associated with the interpretation of the catalytic effect of anions on the rate of electron transfer between iron(1I) and iron(II1) ions. 'Therefore, it is not possible a t t,he present state of the art to specify that iron(II1) complexes react with uncomplexed iron(I1) ions or iron(I1) complexes react with uncomplexed iron(II1) ions. It can only be argued that certain paths are more probable on the basis of relative abundances of the various complexes in solution,2b favorable considerations of coulombic repulsion, and by analogy. With respect to arguments of analogy it is of interest to note the observation" made on the chloride ion catalysis of the chromiuni(I1)-iron(II1) reaction. It is chromium(1J) that brings the anion into the activated stat'e at -50'. This suggests that the iroii(I1)-iron(II1) reaction procedes by a similar mechanism. But, as pointed out by these investigators, this does not mean that the iron(111) complex cannot contribute significantly at, higher temperatures. The same argument holds for the iroii(11)-iron(II1) reaction. With the understanding that other anion catalyzed paths are available and possible, the rate data were analyzed assuniing that the iron(II1) complexes, which predominate in these reaction solutions, bring the anion into the activated state. At learjt under these conditions the results can be compared on the same basis used by earlier inves1,igators.2--5 The second assumption is that the heats of formation of the sulfate, oxalate, and acid phosphate complexes are negligible, and is made because of lack of data concerning the beats of formation of the sulfate, oxalate, and acid phosphate complexes of iron(II1). There is, however, evidence to the effect that the heat of formation of the iron(II1) sulfate complex may be less than one kcal./niole. Three values reportedl2-I4for the first iron(II1) sulfate complex were essentially the same for a ten degree temperature range. The situations with respect to heats of formation of the oxalate and acid phosphate complexes of iron(II1) are less firm, however Home4gives evidence in support of a small value for the former. At low anion and constant hydrogen ion concentrations the data for the iron(I1)-iron(II1) reaction can be represented by the equation
,1i, = kl
+ klK(a[anion] + K7[anion]
1
which is the form used by previous investigators. 'k', and K 3 are, respectively, the observed rate constant in the absence of the coniplexiiig anion, the rate constant for the anion catalyzed path, and the forma/cot kl,
(9) K. H. Lieser a n d H Schroeder. J . Inorg. Suclear Chem.. 14, 98 (1960). (IO) H. Taube, Aduan. Inorg. Chem Radaochem., 3, 1 (1959). (11) M. Ardon, J. Levitan, a n d H. Taube, J A m Chem. Soc., 84, 872 (1962). (12) R. A. Whitaker a n d N. Davidson, abad., 76, 3081 (1953). (13) B. N. Mattoo. Z. p h y s z k . Chem. (Frankfurt), 19, 156 (1959). (14) K. W. Sykes, J . C h e n . Soc., Spec. Publ. No. 1, 64 (1954).
1027
tion constant for the iron(II1) complex. At even lower anion concentrations data of this type can be fit to the equation ,
kI
=
ko
+ klKg[anion]
(3) The data shown in Fig. 1, 2, and 3 yielded straight lines when 'k' was plotted against the anion concentrations; however deviations were observed a t higher anion conceiitrations. From the slopes, klK3, of these lines and by use of the appropriate KB the values for k1 were obtained. Table I includes the values for lcl as well as the K3's used to obtain them. Also included in this table are the respective energies and entropies of activation. Arrhenius plots for the data found in Table I are shown in Fig. 4. The ionization constants for the acids which were used t o obtain the free anion concentration are listed in Table 11. Fukushima and Reynolds15 recently reported a rate constant for the iron(I1)-iron(It1) sulfate path of 680 h1-l sec.-l a t 25' and an ionic strength of 0.25. When differences of ionic strength are considered, this value is in good agreement with 760 1W-l set.-' calculated from the data in Table I. TABLEI RATEDATAFOR THE SULFATE, 24CID PI
