7652
J . Phys. Chem. 1989, 93, 7652-7654
Nonmnetonaus Temporal Dependence of the Concentration of Ferroin as a Reactant Milan MelicherElk and hdovlt Treindl* Department of Physical Chemistry, Comenius University, 842 15 Bratislava, Czechoslovakia (Received: July 28, 1988; In Final Form: April 3, 1989)
In the course of a chemical reaction between Ce(S04)2and ferroin in a solution of sulfuric acid, the temporal dependence of the ferroin concentration has a nonmonotonous character. At the beginning of the reaction, the ferroin concentration decreases rapidly, then increases, attains its maximum, and finally decreases, tending to some limiting value. The proposed mechanism is based on the existence of tetravalent iron as an intermediate that is one of the products of the femin dismutation. An analogue computation of the corresponding kinetic differential equations, based on our reaction scheme, leads to the curve of temporal dependence of the ferroin concentration, which is qualitatively in accordance with experimental curves.
Introduction Phenanthroline-iron complexes have a quite rich phenomenology and an unusually developed ability of adapting themselves to different enviromental conditions.' Although the redox couple ferroin-ferriin as a catalyst of the Belousov-Zhabotinskii oscilonly scant work has been latory reaction was used in some ca~es,2~ done on the detailed kinetics of the ferroin-catalyzed oscillations.M This is unfortunate since most experimental studies of the Zhabotinskii chemical waves are made on the ferroin-catalyzed system. Showalter' has described trigger waves in the acidic bromate oxidation of ferroin. In an unstirred thin film of solution, a single trigger wave of chemical reactivity may develop and subsequently propagate, converting the reaction mixture from the reduced state to the oxidized state. Because the reaction is nonoscillatory, the trigger wave behavior in the ferroin-bromate system is a spatial analogue of bistability in the CSTR experiments and thus the coexistence of kinetic states and spatial bistability are considered. We report here a systematic investigation of an unusual character of the temporal dependence of a reactant concentration. The ferroin concentration exhibits a nonmonotonous temporal dependence induced by chemical or electrochemical oxidation of Fe(phen)32+to F e ( ~ h e n ) ~ions. ~ + We believe the study of this unusual temporal dependence of ferroin as a reactant could help us to elucidate the complex dynamic behavior of systems containing the ferroin-ferriin redox couple.
TABLE I: Rate of the Fe(I1) Anation by Phenantbroline with 1.25 X M FeSO, and 3.75 X IW3M Phenanthroline, and at 30 OC concn of H2S04, concn of H2SO4, mol dm-3 k X lo4, s-' mol dm-' k X lo4, s-' 1.000 0.01 0.250 14.29 0.625 2.27 0.187 36.08 0.500 2.85 0.125 118.51 0.375 4.63
Experimental Section Spectrophotometric measurements were carried out on a Specord UV-VIS-type apparatus (Carl Zeiss, Jena) in a 1- or 0.5-cm cuvette placed in a tempered block connected with a T M 150 type ultrathermostat (Medingen). The reaction courses were indicated by the change of spectra of Fe(phen)?+ and F e ( ~ h e n ) ~ ~ + ions in solutions of sulfuric acid in the presence of oxygen. There was no stirring of solutions in the cuvettes. The results refer to the temperature of 30 OC if nothing else is mentioned in the text. The electrooxidation of ferroin to ferriin was performed at the platinum electrode with the potential of +2.0 V in the presence of nitrogen of analytical grade. The reactions were followed always with the freshly prepared solution of ferriin. All reagents used were of analytical reagent grade. The corresponding differential equations were solved on a MEDA Model 42 TA type analogue computer (Czechoslovakia).
the value of the rate constant at 25 OC is k = 1.41 X lo5 dm3 s-I mol-I. In the first stage of the reaction, Fe(phen)32+ions are oxidized by Ce(IV) ions to Fe(phen),,+ ions, but the reaction in this solution is not yet finished. As we can see in Figure 1, the absorbance of F e ( ~ h e n ) ~ions ~ +with the value of, X = 590 nm decreases and the absorbance of Fe(phen)32+ions with the value of A-, = 510 nm increases, then attains its maximum, and finally decreases. During this complex reaction, the concentration of Fe(phen)$+ ions as a reactant behaves nonmonotonously. After its minimum is reached at the end of the first stage of reaction, this concentration increases to its maximum and then decreases, tending to some limiting value (Figure 2). Although the shape of this curve depends on the concentration of sulfuric acid and on the temperature (Figures 2 and 3), its nonmonotonous character persists. We can also observe this nonmonotonous behavior if we use instead of the chemical oxidation by Ce(IV) an electrochemical one. Since Fe(phen),*+ ions are substitution-inert species, we have focused our attention also on the anation of Fe(I1) ions by phenanthroline and have measured its rate in solutions of sulfuric acid spectrophotometrically (Table I). The anation of Fe(I1) ions by phenanthroline Fe(I1) + 3phen * Fe(phen),*+ (A) is influenced strongly by the variation of sulfuric acid concentration. The rate of ferroin aquation in this solution can also be measured in the same way (Table 11). Since this reaction is
Results and Discussion The oxidation of Fe(phen)$+ ions by Ce(IV) ions in the solution of 0.25 M H2S04proceeds very fast. According to Sutin et aL,* (1) Gutmann, V. Chem. Int. 1988, 10, 5. (2) Smoes, M. L. J . Chem. Phys. 1979, 71, 4669. (3) Smoes, M. L. In Instabilities, Bifurcations, and Fluctuations in Chemical Sysfems;Reichl, L. E., Schieve, W. C., Eds.; University of Texas Press: Austin, TX, 1982; pp 83-125. (4) Mrlkavovl, M.; Treindl, i.Collect. Czech. Chem. Commun. 1986,51, 2693. (5) Rovinsky, A. B. J . Phys. Chem. 1987, 91, 4606. (6) Rovinsky, A. B. J . Phys. Chem. 1987, 91, 5113. (7) Showalter, K. J . Phys. Chem. 1981, 85, 440.
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TABLE 11: Rate of the Ferroin Aquation with 1.2 X IO-' M Fe(phen),*+ and at 30 OC concn of phen, concn of phen, mol dm-' k X lo4. s-I mol dm-' k X IO4, 0 1.6 6.7 X lo-' 0.58 3.6 X IO4 1.6 1.3 X 0
s-I
TABLE 111: Rate of the Ferriin Decreasing, 1.25 X M Fe(~hen),~+. as Product of the Ferroin Oxidation bv C e W ) at 30 OC concn of concn of k x lo4, concn of concn of k x lo4, H2S04, M phen, M s-l H2S04, M phen, M s-l 0.25 0 1.2 0.125 0 1.64 0.25 1.76 0.5 0 0.84 0.25 4.8 X 3.71 0.75 0 0.69 3.8 0 0.12 1 .oo 0 0.57 3.8 4.8 X 0.34
(8) Dulz, G.;Sutin, N. Inorg. Chem. 1963, 2, 917.
0 1989 American Chemical Society
Temporal Concentration Dependence of Ferroin as Reactant
The Journal of Physical Chemistry, Vol. 93, No. 22, 1989 7653
a A
1.2
0.8
1.0
0.6
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0.4
0.0 0
350
400
500
I
I
1 \\
\I
600 h/nm
I
I
W-27,
1.0
0.8 0.6 0.4
0.2
0.0 L 350
Figure 1. Time dependence of an absorption spectrum 0.25 M H2S04, 1.25 X IF3M Fe(phen):+, 1.25 X lo-' M Ce(S04)*,temperatureof 30
OC, I-cm cuvette. The time interval between each two successive curves A = IO min: (a) 1-17 successive curves; (b) 18-46 successive curves.
4 0.60 - 0.45 - 0.30 - 0.15 I
20
t/ms
.
I
I
I
I
I
I
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18
24
30
36
0.00 42
L
5 6 t .10-'/5
7
8
9
10
cuvette.
1.2
0
3
2
Figure 3. Temporal dependence of the Fe(phen)32+ions absorbance at X = 510 nm on temperature ("C): 0,20; a, 30; A, 40. Conditions: 0.25 M Fe(phen)?+; 1.25 X IO" M Ce(SO&; 0.5-cm M H2S04;1.25 X
0.0
0.0
1
1,"s
0.2
A
20
48
t .163/5
Figure 2. Temporal dependence of the Fe(phen)32+ions absorbance at X = 510 nm. H2SO4 concentration (M): 0, 0.125; 0 , 0.25; A, 0.75. M Fe(phen)?+; 1.25 X M Ce(S04)2; Conditions: 1.25 X temperature, 30 OC; 0.5-cm cuvette.
reverse, we can see the rate of ferroin aquation to be decreased or even stopped by the presence of an appropriate amount of added phenanthroline. The rate of the decrease of ferriin concentration also depends on the presence of added phenanthroline (Table 111), but in this
case, the value of the rate constant slightly increases with increasing concentration of added phenanthroline. As a rule, the conversion of reactants in the course of a chemical reaction takes place monotonously. Sometimes, however, a nonmonotonous evolution may arise in a homogeneous reacting system evolving far from equilibrium. This nonmonotonous evolution has to do for example with intermediates and redox catalysts during an oscillatory reaction. In general, we expect the reactant concentration in closed systems to decrease in a smooth, monotonous manner. In this paper, we report a nonmonotonous change of ferroin as a reactant concentration. Since the ferroin-ferriin redox couple seems to play a very important role in Zhabotinskii chemical waves, we try to explain this unusual behavior, e.g., the described nonmonotonous temporal dependence of ferroin as a reactant. We believe this explanation might be useful not only as far as this reaction is concerned but also in connection with a detailed interpretation of the Zhabotinskii chemical waves in the presence of ferroin, as well as of the single trigger wave in the oxidation of ferroin by bromates.' Our explanation is based on the existence of tetravalent iron as an intermediate. Its existence has been postulated already in 1932 by Bray and G ~ r i n . The ~ formation of iron(1V) in the oxidation of iron(I1) has also been postulated by Sutin et al.IO The reality of iron(1V) complexes formation in water solutions has been confirmed in several and especially the existence of iron(1V) in the solid state has been proved by Warren et al.13 Therefore, we are encouraged to propose the following reaction mechanism. After either the chemical oxidation of ferroin Fe(phen):+ Ce(1V) F e ( ~ h e n ) , ~ + Ce(II1) (B) or its electrochemical oxidation is completed, the dismutation of ferriin in solution of 0.125-0.75 M H2S04 can proceed as follows: 2Fe(phen),,+ + Fe(phen)32++ F e ( ~ h e n ) , ~ + (C) The species F e ( ~ h e n ) , ~or + their products of aquation Fe( ~ h e n ) , ( H ~ O )as~ ~intermediates, + denoted as Fe(IV), oxidize water molecules (OH- ions) very probably in a subsequent step Fe(1V) + 2 H 2 0 Fe(I1) + 2H+ + Y2O2 H,O (D) Fe(I1) reacts with phenanthroline to form ferroin Fe(I1) 3phen + Fe(phen)?+ (E)
-
+
-
+
+
+
(9) Bray, W. C.; Gorin, M. H. J. Am. Chem. SOC.1932, 54, 2124. (10) Conocchioli, T. J.; Hamilton, E. J.; Sutin, N. J . Am. Chem. Soc.. 1965, 87, 926. (1 1) Melicherdk, M.; Svec, A.; Treindl, i.To be published. (12) Felton. R. H.: Owen. G. S.:DolDhin. D.: Faier. _ .J. J . Am. Chem. Soc. 1971, 93, 6332. (13) Warren, I. F.; Bennett, M. A. J . Am. Chem. SOC.1974, 96, 3340. (14) Pasek, E. A.; Straut, D.K. Inorg. Chem. 1972, 1 1 , 259. (15) Hazelden, G . S.; Nyholm, R. S.;Parish, R. V. J . Chem. Soc. A 1968, 162. (16) King, J.; Davidson, N. J . Am. Chem. SOC.1958, 80, 1542.
J. Phys. Chem. 1989, 93, 7654-7659
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(Table 11). As a parallel step to step D, the oxidation of water molecules by ferriin may proceed as follows:
+
2 F e ( ~ h e n ) , ~ + HzO
-
+
2 F e ( ~ h e n ) , ~ + 2H+
+ f/202
(G)
Oxygen molecules may have the role of a radical scavanger, and still other reaction steps may be taken into account. We formulated various alternative reaction schemes, based on the proposed mechanism, and were trying to find a solution of their corresponding rate equations. Finally, we selected the following scheme k,
A+B-C+D
(H)
2C&G+A k-2
k3
G+E-A+F
tlme
Figure 4. Computed c = A?)curves of A-C, G, and E, a qualitative simulation of a nonmonotonous temporal dependence of ferroin (A) as a reactant.”
where (E) is the sum of successive steps of anation. The reaction step C or its simpler analogue 2Fe(III)
Fe(I1)
+ Fe(1V)
(F)
might play a key role in this mechanism. This assumption is supported on one hand by the fact that Fe2(S04), in a solution of 0.25 M H2S04in the presence of added phenanthroline leads to the formation of ferroin and on the other hand by the fact that the formation rate of ferroin is inversely proportional to the sulfuric acid concentration (Table I). The reaction step E can be verified by experiments of FeSO, anation by phenanthroline molecules
where A = ferroin, B = Ce(IV), C = ferriin, D = Ce(III), G = Fe(IV), E = H 2 0 (OH-), W = water as ligand, I = inactive products of aquation, P = phenanthroline, and F = oxidation products. As we can see in Figure 4, the solution of the corresponding kinetic equations by an analogue computation” yields the temporal dependence of the ferroin concentration (of A) that is qualitatively in accordance with our experimental curves (Figures 3 and 4). Any other alternative reaction scheme does not lead to such an accordance. Unfortunately, for the existence of Fe(1V) species as intermediates, we have only such kinetic evidence. Registry No. Ce4+, 16065-90-0; ferroin, 14708-99-7.
Solvent Effects in the Reactions of Peroxyf Radicals with Organic Reductants. Evidence for Proton-Transfer-Mediated Electron Transfer P. Nets,*.* R. E. Huie,’ P. Maruthamuth~,~*~ and S. Steenken*32 Chemical Kinetics Division, National Institute of Standards and Technology, Gaithersburg, Maryland 20899, and Max- Planck- Institut fur Strahlenchemie. 0-4330 Mulheim, West Germany (Received: February 6, 1989; In Final Form: May 31, 1989)
Absolute rate constants for the reactions of substituted methylperoxyl radicals with ascorbate, urate, trolox (6-hydroxy2,5,7,8-tetramethylchroman-2-carboxylicacid), and TMPD (N,N,”,”-tetramethyl-p-phenylenediamine) have been determined by pulse radiolysis in different solvents. In wateralcohol or water-dioxane solutions, the rate constants for trihalomethylperoxyl radicals generally increase with increasing water content. The rate constant for reaction of CC1302’radicals with trolox was measured in water, MeOH, i-PrOH, 1-BuOH, ethylene glycol, diethyl ether, dioxane, acetone, acetonitrile, formamide, dimethylformamide, pyridine, and CCI,. The rate constants were found to correlate well with a two-parameter equation that includes the dielectric constant of the solvent and the coordinate covalency parameter, a measure of the proton-transfer basicity of the solvent. Kinetic isotope effects in H20/D20of about 2 and activation entropies of about -10 eu for reduction of R02’ by the organic reductants indicate that electron transfer to the peroxyl radical is concerted with the transfer of proton from the solvent to the incipient hydroperoxide anion.
Introduction Rate constants for reactions of peroxyl radicals with organic substrates have been determined for a wide variety of system^.^ Because of the involvement of the CCI3O2*radical in the toxicity of cc14,5 rate constants have been measured for reaction of
CC1302’ with several molecules of biological importance such as fatty acids, iron porphyrins, and natural antioxidant^.^ Although many of the rate constants refer to the reactions in aqueous solutions, the medium usually contained varying amounts of alcohols to help in the dissolution of CCI4 and of the organic substrate. Recently, it has been shown that the rate of oxidation of
(1) National Institute of Standards and Technology. (2) Max-Planck-Institut fur Strahlenchemie. (3) Visiting scientist from the University of Madras. (4) For a compilation see: Neta, P.;Huie, R. E.; Ross, A. B. J. Phys. Chem. Ref. Data, submitted for publication.
( 5 ) Slater, T. F. In: Biochemical Mechanisms of Liver Injury; Slater, T . F.,Ed.; Academic Press: London, 1978; p 1. Recknagel, R. 0.; Glende, E. A., Jr. CRC Crit. Rev. Toxicol. 1973, 2, 263. Brault, D. EHP, Environ.
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Health Perspect. 1985, 64, 53.
0 1989 American Chemical Society
