R. A. HORNE
1512
Vol. 64
THE KINETICS OF THE OXALATE CATALYSIS OF THE IRON(I1)-IRON(II1) ELECTRON-EXCHANGE REACTION I N AQUEOUS SOLUTION' BY R. A. H O R N E ~ Chemistry Dthpartrnents, Brookhaven National Laboratory and Columbia University, Z'pton, N . Y., and New York City Received April 9.8, 1060
The kinetics of the electron-exchange reaction between the two oxidation states of iron is first order in both iron(I1) and iron( 111')and oxalate-catalyzed in aqueous perchloric-oxalic acid media of 0.55 ionic strength. The specific reaction rate constants for the reaction path involving ferrous ion and FeC204+are 700, 1100 and 2140 sec.-lf-" a t 0, 10 and 20°, respectively, corresponding to an activation energy of 9.2 kcal. mole-' and an entrop of activation of -14.1 oal. mole-' deg.-l. The specific reaction rate constants for the path involving ferrous ion and Fe(dO,)s- are in the range 2500-4500 sec.-l f-l. The exchange iii also catalyzed by silver foil and by acetate, succinate and phenolate anions.
Introduction Since the classical work of Hevesy and Zechmeist,er3on the electron-exchange between lead(I1) and lead(1V) a great deal of effort had been devoted to studies of oxidation-reduction processes in aqueous solution, especially to those involving electron-exchange between two valence states of the same element.4 I n particular, the kinetics of the iron(I1)-iron(II1) electron-exchange and the catalysis of this reaction by complexing anions have received considerable attention. Attempts t o interpret the kinetics and establish the mechanism of the iron(I1)-iron(II1) exchange have tended to fall into two principal categoriesanion bridging theoriess-l0 and water bridging t h e o r i e ~ ~ l -If~ the ~ electron is transferred across an anion bridge, one might reasonably expect that the activation energy of the exchange process should change as the complexing anion is changed. In the case #of the exchange involving Cr++ and (x&)5(2rA+--, Ogard and Taube18 have observed such a correlat,ion with the size of the complexing anion, the activation energies being 13.4, 11.1 and 8.3 kcal. mole-' for fluoride, chloride and bromide ions, respectively. However, if the exchange involves a water bridge a marked heavy water. isotope effect, even for the anion catalyzed
processes,17 and little dependence on the nature of the complexing anion is anticipated. In the case of the iron(I1)-iron(II1) exchange both anticipations have been v i n d i ~ a t e d . ~The ~ ~ ~activation energies for the exchanges involving Fe++, FeC104++, FeOH++, FeF++, FeCl++, FeSCN++, FeF2+, FeC12+ and Fe(SCN)2+ are 9.9, 9.5, 7.4, 9.1,8.8, 7.9,9.5, 9.7 and 8.6 kcal./mole respectively (see ref. 5,20, 5, 12, 5,21,12,5 and 21,respectively) -a range of values not outside the range of experimental error. The exchange involving FeNa++ appears to be an exception.22 But, it can be argued, fluoride, chloride, bromide, and even thiocyanate ion are similar. Any differences in the activation energies of their catalyses might be less than experimental error and hence escape detection. For this reason t,he catalysis of the iron(I1)-iron(II1) electron exchange reaction by a larger, non-halide, more structurally complex complexing anion, such as oxalate, is of interest. Experimental Procedures
and R. W. Uodsun, TIIISJOURNAL, 56, 848 (1952). 1 I?. W. Ihdson, J . A m . Chem. Soc., 78, 911 (1956). 7 , Tlrls J n z X N A L , 56, 803 (1952). ky Discus-ion of "Electron Transfer in Solution and rncli % c . Phys. Cliein.. Paris, M a y 8, 1951, ONR, London, '1ec:i. X e it. OSIiL-73-Zl (AUK.20. 1951). pp. 1-2. (9) 11. 'I'aiihe, ct at.. J . C h c m . S o r . . 75, 4118 (1053); 76, 2103. 4053 (19.3.); 7 7 , 4431 (19 ins, ib;d.. 80, 1091 (1958). (10) D. L. T?:i11 and E. I l l ) I T , L. Reynolds nn8l R. JV. L u n i t y . J . Chem. Phys., 23, 2160
Results and Discussion The experimental results for the oxalate catalysis of the iron(I1)-iron(II1) electron-exchange reaction in aqueous solution are shown in Table I and Fig. 1. In aqueous oxalate media, as in p e r c h l ~ r a t e , ~ * ~ ~ fluoride, l 2 ~ h l o r i d ebromide16 ,~ and thiocyanate18f21
The iron(I1) and iron( 111) perchlorate stock solutions, the tracer Fe56 solution, and the solutions of perchloric acid, sodium perchlorate and buffered 2,2'-dipyridyl were prepared from the same quality materials and purified and analyzed in the same manner as described previously by Silverman and Dadson.& The oxalic acid solutions (Baker and Adamson, A.C.S., reagent) were analyzed by (a) titration with standard sodium hydroxide and (b) titration with standard potassium permanganate.28 I n the present experi(1) Reszarch porforrnerl under the auspices of the U. S. Atomic ments the increases in the hydrogen ion concentration and in Energy Commission. the ionic strength due to the added oxalic acid were negligi( 2 ) J o w p h Iiaye and Co., Inc., 49 Hampshire St., Cambridge ble. Jlnssachii jetts. The details of the experimental procedure have been de( 3 ) G. 1-Ievesy :tnd L. Zechmeister, Be?., 53, 410 (1920). scribed previously by Silverman and Dodson.6 This pro( 4 ) (a) S o t r e Ejanie Symposium on Electron Transfer, THISJOCRcedure involves removing aliquots from the reaction mixture N ~ L 5C, , 80lff (19,jZ); (ir) C. B. Amphlett, Quart. Rev., London Chem. a t regular time intervals; quenching the reaction by adding Snc., 8 , 219 (1954): i c ) F. Basolo and R. G. Pearson, "Mechanisms of Inorganic Reactions," .John Wiley and Sons, Inc., New York, Ii. Y., each aliquot to a buffered 2,2'-dipyridyl solution, thus removing the iron(I1) as a complex; and precipitating, mounting l%8. eh. i : (11) Unirerpity of Toronto Symposium on Charge Transfer and counting the iron(II1) as the hydroxide. J C i , e m . , 37, 120E (1959). I
and A . C . 'Xdil, J . A m . Chem. Soc., 75, 4153 (1953). J . Cham. Phpa.. 19, 10fX (1951). (1.1) R. €'iatzrn:in and .I.Frank. Z. P h ! / s i k . 138, 411 (1964). (15) R. \Ii)odson . and S . I)avi&on, in the discussion of ref. 7. ( 1 6 ) R. A . Horne. Ph.D. Thesis. Columbia Cniversity, 1955. ( i i ) R . A. Iionie. I'ajier presented at tile 135th Nati. Meeting of the Ani. Chenl. Soc., 13oston, 1959. (18) A. E. Ogard and H. Taube. J . A m . Chem. Soc., SO. 1084 (1958).
(10) N. Slitin and R. W. Dodson, Paper presented at the 136th Natl llleccing of the Am. Chern. Soc., Atlantic Clty, 1959. (20) R. A. Horne, Nature, 181, 410 (1958). (21) G . S. Laurence, Trans. Faraday S o c . , 53, 1326 (1957). (22) D.Bunn, F. S. Dainton and S. Duckworth, Trans. Faraday SOC. S5, 1267 (1959). (23) L. F. Hamilton and S. C. Simpson, "Talbot's Quantitative Chemical Analysis," The Maemillan Co., New York, N. Y.,1946,pp. 129f. 160f.
Oct., 1960
OXALATE CATALYSIS OF
THE
FE(II)-FE(III) ELECTRON-EXCHANGE REACTIOX
1513
TABLE I Terms for polynuclear species, when significant, are readily added to expression 7. Equation 7 is IRON(II)-IRON(III) ELECTRON-EXCHANGE REACTION
'THE
AQUEOUS PERCHLORIC-OXALIC ACIDMEDIA Ionic strength = 0.55, ( H + ) = 0.548 asHClO4, (Fc(I1)) = 1.02 x 10-41, and (Fe(II1)) = 0.650 X lO-4f IN
Added oxalic acid
x
Ilalf-life min.
105, f
Specific reaction rate constant, k, sec.-l f - 1
ti,a,
0.000 2.00 4.00 6.00 8.00 10.0 100 9000
Temp. = 0.02 f 0.01' .... 1.31 f O.OS(av.) 17.7 3.89 f 0.83 8.21 8.46 11.16 5.48 12.6 1 2 . 1 4.55 15.2 11 . 6 3.70 20.3 1 9 . 7 (0.88) ( 79) (0.5) (137)
0.000 2.00 4.00 6.00 8.00 10.0
Temp. = 9.88 f 0.05" .... 3.44 6.50 10.6 4.85 14.2 3.55 19.5 2.67 26.0 2.16 32.0
0.000 2.00 6.00 8.00
Temp. = 20.57 10.04' .... 8.22 3.54 19.6 1.45 47.5 1.05 65.7
I0.13(av.) 10.5 13.3 11 . 6 13.5 =k 11.0 11.43 (av.) =k
=
k(Fe(II))(Fe(III))
(1)
(3)
F c ; Z ~ ~Y- F~ ~~A= ,~ ( Fe.L,,Z-mo)/(Fe ++)(d-*)" (4)
where p and y are the over-all complex ion formation constants, and the exurhange processes by Fe*A,S--na
+
Fe*A,z--mo
+ FeX,3-*a
(5)
The rate of a process such as ( 5 ) is then given by R F ~ ~ " - F=~ .kFeA.-Fe4m I~ (Fehn3-no)(Fe~~mz-m.j (6)
The over-all rate of the electron-exchange reaction is given by ,?
(8)
Now under these circumstances (Fe++) c (FelII)) (Fe+++) = (Fe(II1)) (Fe;\n3-na)
(9)
A.n
=
(Fe(II1))
-
PFaA,,
(Fe+++)(h-a)n (10)
A.n
(Fe+++) = (Fe(III))/ 1
+
PF~\, A.n
hence
By comparison of equations 1 and 12 we see that the term in brackets in the latter is k. For the case of the oxalate catalysis of the iron(11)-iron(II1) electron-exchange reaction in perchlorate media the important terms in equation 8 are X = k p e (Fe++ k)(Fe++)+ ~ F ~ C (FeClOi++)(Fe++) I O ~
+ +
~ F , O E (FeOH +
kFdCPO4)Z
')( Oe
+
I)
+ kFecZo4(F e C O +)(Fe ++)
(13)
+ ...
(Fe((X4)2-)(Fe++)
The first two terms of equation 13 involve a controversial choice among the alternatives O ~0, ~ F ~ C I#O 0,~ and k~~ f 0 ( a ) P F ~ C If (b) /3werio4 # 0, k F e L i n l f 0, and k t , = 0 ( c ) PFerin4 = 0, t n d kre # 0
Hitherto alternative (e) has been preferred,i*12 although there now appears to be evidence for (a) or (b).2027-29 Fortunately, for purposes of analyzing the oxalate catalysis, the first three terms in equation 13 can be combined into n single term R = Xo' kFeClO4 (FeCdL+)(Fe++)
+
+
~ F ~ ( c(Te(C204)2-)(Fe++) ~ o ~ ) ~ i-
(11)
where, since the ratio of the concentration of uncomplexed to complexed iron(II1) is large ( e . g . , about 40 a t 10" with 10 X 10-5j total added oxalic acid), Ro' may be approximated by Ro ( 2 5 ) hI. Boftelsky, D. Chasson and S. F. Klein, .\nul. 460 (19.53).
c h i n Actu. 8 ,
(28) H. yon Stackelburg and H. yon Freyhold, Z. Elekfrochem , 4 6 , 120 (1940). (27) J. Button, Nature. 169, 7 1 (1952).
RFokn-FeAm
Am.n
k~,a.- ~
(Fe.An3-na)(Fe++)
fO.0
(FeA,a-"")/(Fe
+ mA-5
c
k~e.4~
+
Fe++
=
An
3.0
where R is the rate of the reaction, k the over-all specific reaction rate constant, and (Fe(I1)) and (Fe(II1)) are the total concentrations of the iron(I1) and iron(II1) species. The over-all specific reaction rate constant was evaluated, corrections being made for induced exchange when indicated, by using a McKay plotz4and the expression 12 = O.E93{ [(Fe(II)) (Fe(III))lt1/2]-1 (2) where t I / % is the half-life of the exchange. The formation of iron(I1) and iron(II1) complex ions can be represented by FeAns-no,P F ~ A = " Fc+++ + nA-0 and
R
12 . 3
media, the exchange is first order with respect to both of the iron valence states R
quite general, being valid for all complexing anions, A, that have been studied to date, and for n = 0, 1 and 2. The limits on the values of m are unknown. Fortunately for many complexing anions p >> y , and hence it may be possible to make the assumption that m = 0. Iron(II), for example, forms oxalate complexesz5but they are less stable than the corresponding iron(II1) complexes by a factor of 10'3 and make a negligible conlribution.Z6 When m = 0 equation 7 becomes simply
~ (FeAnS-n5)( 4 ~ FeAm2-m0)
A,m.n (24) H. A. C. McKay. Nature, 142, 097 (1938).
(7)
(28) K. A ' .' Sykes, et al., "The Kinetics and Mechanism of Inorganin Reactions in Solutions," Chem. SOC., London, 1954, pp. 04 et seq. (29) Cf. H. Call. R. V. Naumsn a n d P. W. West. J . A m . Chem. Soc.. 81; 1284
