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A Kinetic Study of the Reduction of Colloidal Manganese Dioxide by Oxalic Acid
JOURNAL OF COLLOID AND INTERFACE SCIENCE ARTICLE NO.
177, 288–297 (1996)
0034
A Kinetic Study of the Reduction of Colloidal Manganese Dioxide by Oxalic Acid JOAQUIN F. PEREZ-BENITO,* ,1 CONCHITA ARIAS,*
AND
ELISENDA AMAT†
*Departamento de Quimica Fisica, Facultad de Quimica, Universidad de Barcelona, Marti i Franques, 1, 08028 Barcelona, Spain; and †Unidad de Fisicoquı´mica, Facultad de Farmacia, Universidad de Barcelona, Avenida Diagonal, s/n, 08028 Barcelona, Spain Received May 15, 1995; accepted June 14, 1995
A kinetic study of the reaction between colloidal manganese dioxide and oxalic acid in aqueous acetate media (pH 4.3–5.1) is reported. The reaction is autocatalytic and, in order to determine the rate constants k1 and k2 corresponding, respectively, to the noncatalytic and autocatalytic reaction pathways, the Mn(III) formed as an intermediate was stabilized by addition of sodium pyrophosphate. Joint iodimetric and spectrophotometric studies indicated that the reduction of colloidal manganese dioxide follows the sequence R
Mn(IV) r Mn(II)
Mn ( IV )
R
Mn(III) r Mn(II),
where R stands for the reductant. Both reaction pathways exhibit acid catalysis, and the activation energies associated to k1 and k2 are 74.7 { 0.9 and 44.6 { 0.6 kJ mol 01 , respectively. The reaction is accelerated by addition of manganese(II) sulfate, and k2 increases with rising oxalate concentration, suggesting that MnC2O4 is the active autocatalyst. The possibility that Mn(III) might collaborate in the autocatalysis through a free-radical chain mechanism is also pointed out. q 1996 Academic Press, Inc. Key Words: autocatalysis; colloidal solution; kinetics; manganese dioxide; oxalic acid.
or reaction products in most permanganate oxidations (11, 12), being actively involved in the mechanism as autocatalysts in many cases (13, 14). On the other hand, although the kinetic aspects of the manganese–oxalate reacting systems have gained new interest in the last decade due to their involvement in some oscillating reactions (15–20), the reduction of colloidal manganese dioxide by oxalic acid has received much little attention (21–23) than its relative, the permanganate–oxalate reaction (24, 25). We have now undertaken a kinetic study of the reaction between colloidal manganese dioxide (in the form of a perfectly transparent, aqueous sol) and oxalic acid and found that it is in fact an autocatalytic reaction, whose stoichiometry may be written as MnO2 / H2C2O4 / 2H / Å Mn 2/ / 2CO2 / 2H2O.
[1]
We think that the results presented here for the title reaction might throw some light upon the common aspects of the autocatalytic behavior exhibited by both the MnO2 –H2C2O4 and MnO 40 –H2C2O4 (26) reactions. MATERIALS AND METHODS
INTRODUCTION
Preparation of the Colloidal Manganese Dioxide Sol Manganese dioxide is a substance of certain importance owing to its catalytic (1) and oxidizing (2–5) properties, but its applications are notably limited because of its insolubility under ordinary conditions (6). In recent years, however, the preparation of several forms of colloidal manganese dioxide stable for long periods as brown or yellow, perfectly transparent solutions, either in aqueous (7, 8) or organic (9) media, has made possible the kinetic study of its oxidizing behavior in redox reactions by easily implemented, conventional spectrophotometric methods (10). Moreover, the transparent sols of manganese dioxide are also of importance because of their widespread participation as intermediates 1
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/ 8MnO 40 / 3S2O 20 Å 8MnO2 / 6SO 20 3 / 2H 4 / H2O.
[2] Provided that the oxidant and reductant concentrations are used in stoichiometric ratio, a dark-brown solution is ob-
To whom correspondence should be addressed.
0021-9797/96 $12.00 Copyright q 1996 by Academic Press, Inc. All rights of reproduction in any form reserved.
To prepare the colloid, 10 ml of KMnO4 (0.100 mol dm03 ) and 20 ml of Na2S2O3 (1.88 1 10 02 mol dm03 ) aqueous solutions were mixed together in a 2-liter volumetric flask, and water was added until completion, the solution being homogeneized afterwards by gentle stirring. The reaction taking place under these conditions may be described by the stoichiometry
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tained, remaining perfectly transparent at least for several months (8, 10). Colloid Characterization The colloidal solution so prepared ([MnO2 ] Å 5.00 1 10 04 mol dm03 ) was characterized by means of a Malvern autosizer 2.C, the results indicating that 90% of the colloidal particles had a diameter comprised in the range 89–193 nm, with a peak at 120 nm. The zeta potential (measured with a Malvern zeta-sizer 4.C) was 047 { 3 mV. In addition, coagulation experiments performed in a previous work (8) indicated that the colloidal particles had a negative electrostatic charge, thus indicating that their stability in solution was probably due to the fixation of some anions on the colloid surface. The colloid’s UV–vis spectrum was recorded with a Hitachi U 2000 UV–vis double-beam spectrophotometer, and showed a large band covering the whole visible region of the spectrum, with absorbance uniformly decreasing with increasing wavelength, as well as a wide maximum at 300–400 nm. Other Reagents The solvent used was water previously purified by deionization followed by double distillation ( the first time from a potassium permanganate solution ) . The reductant used in the kinetic experiments was sodium oxalate ( always in a large excess with respect to MnO2 ) . To keep the pH constant during the runs, an acetic acid ( HAc ) – sodium acetate ( NaAc ) buffer mixture was used ( pH 4.3 – 5.1 ) , since the reaction either in aqueous HClO4 or in the presence of H3PO4 – NaH2PO4 buffer ( pH É 2 ) was too fast to be followed by conventional spectrophotometry. Manganese ( III ) was stabilized in the solution by addition of sodium pyrophosphate. When necessary, the ionic strength was kept constant with a large excess of sodium perchlorate. Other additives used in some of the experiments were manganese ( II ) sulfate, gum arabic and EDTA ( tetrasodium salt ) . All the chemicals used were purchased from Merck ( analytical quality ) . Iodimetric Determinations In some experiments, the average oxidation state (OSav ) of manganese during the MnO2 –H2C2O4 reaction was determined by taking successive 2-ml aliquots of the reaction mixture (buffered with HAc–NaAc and kept in a thermostatic bath at 25.07C) and introducing them into 5-ml volumetric flasks containing an excess of solid potassium iodide. After dilution with water until completion, the solutions were homogeneized and the iodine liberated was determined by UV–vis spectrophotometry [ lmax (I 30 ) Å 351 nm, 1351 Å 2.57 1 10 4 dm3 mol 01 cm01 ]. In accordance with many other previous works (27–29), it was assumed that both Mn(IV)
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FIG. 1. Average oxidation state of the manganese atom during the reduction of colloidal MnO2 (9.62 1 10 05 mol dm03 ) by Na2C2O4 (9.62 1 10 04 mol dm03 ), in the presence of HAc (0.504 mol dm03 ) –NaAc (0.481 mol dm03 ) buffer, at pH 4.57 and 25.07C. Open circles: [Na4P2O7 ] Å 0; filled circles: [Na4P2O7 ] Å 4.77 1 10 03 mol dm03 .
and Mn(III) were reduced by iodide ion to the Mn(II) state. Then, the average oxidation state of manganese could be obtained as OSav Å 2 /
5A351 , 1351l[Mn]T
[3]
where A351 and 1351 are the absorbance of the solution and the molar absorptivity of triiodide ion at 351 nm, respectively, l is the optical pathlength of the cuvettes (1 cm), and [Mn]T is the total manganese concentration present in the system during the reaction. Kinetic Measurements The progress of the reaction was followed by means of a Varian Cary 219 UV–vis spectrophotometer provided with a thermostatized cuvette holder. The wavelength chosen was 400 nm, taking into consideration that colloidal manganese dioxide had a strong absorption there, and that neither manganese(II) ion nor the oxalate and pyrophosphate complexes of manganese(III) showed any appreciable absorption. The pH’s of the solutions were determined by means of a Metrohm 605 pH-meter provided with a glass–calomel combination electrode. RESULTS
Influence of Pyrophosphate The average oxidation state of manganese during the MnO2 –H2C2O4 reaction both in the absence and in the presence of sodium pyrophosphate was determined by iodimetry. We can see in Fig. 1 that when Na4P2O7 was not present the
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FIG. 2. Fractions of Mn(IV) (open circles), Mn(III) (triangles), and Mn(II) (filled circles) during the reduction of colloidal manganese dioxide by sodium oxalate in the presence of sodium pyrophosphate. The experimental conditions are given in the legend of Fig. 1.
average oxidation state decreased uninterruptedly from 4 to 2, whereas in the presence of an enough amount of that additive the reduction took place until the oxidation state 3 in a first stage. At that moment, the UV–vis spectrum showed a small peak at 470 nm, thus confirming the presence in the solution of a manganese(III) –pyrophosphate complex (30). Under those conditions, the total reduction until the final Mn(II) state took several hours. Moreover, although the concentration of Mn(III) during the reaction in the experiment performed in the absence of pyrophosphate was too low to be determined, in the one performed in the presence of that additive it was possible to determine the fractions of the total manganese present in the system during the reaction in the oxidation states II, III, and IV ( fII , fIII , and fIV, respectively). To that end, the decay of Mn(IV) was followed by spectrophotometry at 400 nm for the same experimental conditions used in the iodimetric experiment (Fig. 1, upper curve). Since at that wavelength only Mn(IV) exhibits an appreciable absorption, the equations necessary to calculate the fractions of each manganese species present in the system at each particular instant t are fII / fIII / fIV Å 1
[4]
OSav Å 2 fII / 3 fIII / 4 fIV
[5]
fIV Å
[Mn IV ] A400 Å 0 [Mn]T A 400
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Determination of the Rate Constants The most simple way of describing mathematically the kinetic behavior of an autocatalytic reaction is to assume that two reaction pathways (one noncatalytic and the other autocatalytic) are involved, and that the kinetic orders of the reactant in defect (in both pathways) and of the autocatalyst (in the autocatalytic pathway) are the unity. If, in addition, we assume that the concentration of autocatalytic product at a particular instant t is directly proportional to the decrease of the reactant concentration from the beginning
[6]
where A400 and A 0400 are the absorbances of the solution at 400 nm at time t and at the beginning of the reaction, respectively. It can be seen in Fig. 2 that the fraction of Mn(IV) decreased with time according to a sigmoid curve typical of autocatalytic reactions. The fraction of Mn(III) increased first and then started to decrease very slowly, whereas that
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of Mn(II) increased, passed through a maximum, decreased, and finally showed a slow increase. It is interesting to notice that the final conversion of Mn(III) to Mn(II) seemed to occur when Mn(IV) was not present in the system any longer. When the reaction was followed spectrophotometrically, a small induction period was observed during which the reaction rate showed a fast decrease. After that, the rate started to increase in a typically autocatalytic manner, passing through a maximum and decreasing afterward (Fig. 3). Addition of sodium pyrophosphate to the solution at the beginning of the kinetic runs resulted in a clear acceleration of the reaction, since an increase of the reaction rate both at the minimum and the maximum of the rate vs. time plots was observed, whereas the half-life period of the reaction (t1 / 2 ) decreased (Fig. 4). On the other hand, an attempt to study the effect on the reaction of EDTA [a good complexing agent for both Mn(II) and Mn(III)] failed because in the media of our work colloidal manganese dioxide was reduced by that reactant even faster than by oxalic acid.
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FIG. 3. Dependence of the reaction rate on time during the reduction of colloidal MnO2 (1.00 1 10 04 mol dm03 ) by Na2C2O4 (1.00 1 10 03 mol dm03 ), in the presence of HAc (0.525 mol dm03 ) –NaAc (0.500 mol dm03 ) buffer, at pH 4.60 { 0.01 and 25.07C. Open circles: [Na4P2O7 ] Å 0; filled circles: [Na4P2O7 ] Å 4.96 1 10 03 mol dm03 .
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THE MnO2 –H2C2O4 REACTION
took place in the presence of a concentration of sodium pyrophosphate high enough, the rate-decreasing induction period was followed by another period in which the autocatalysis did not disappear and Eq. [8] clearly holds. This finding has allowed us to determine the rate constants k1 and k2 from the intercept and slope of the linear portion of the r/[MnO2 ] vs [MnO2 ]0 – [MnO2 ] plots, respectively. To that end, the experiments were performed in the presence of a pyrophosphate concentration high enough to stabilize the autocatalysis after the minimum of those plots (see Fig. 5, upper curve). In each kinetic run, the reaction rate was estimated at 5-s intervals as FIG. 4. Dependence of the rate/[MnO2 ] ratios at the minimum (open circles) and the maximum (filled circles) of the rate vs time plots, and of the half-life period (triangles), on the concentration of Na4P2O7 for the reduction of colloidal MnO2 (1.00 1 10 04 mol dm03 ) by Na2C2O4 (1.00 1 10 04 mol dm03 ), in the presence of HAc (0.525 mol dm03 ) –NaAc (0.500 mol dm03 ) buffer, at pH 4.60 { 0.01 and 25.07C.
of the kinetic run until that moment, we have the differential rate law rÅ0
d[MnO2 ] dt
Å k1[MnO2 ] / k2[MnO2 ]([MnO2 ]0 0 [MnO2 ]) [7]
or else r Å k1 / k2 ([MnO2 ]0 0 [MnO2 ]), [MnO2 ]
[8]
where r is the reaction rate at time t, [MnO2 ]0 and [MnO2 ] are the concentrations of colloidal manganese dioxide at the beginning of the reaction and at time t, respectively, k1 is the pseudo-first-order rate constant of the noncatalytic pathway, and k2 is the pseudo-second-order rate constant of the autocatalytic pathway. Differential rate laws similar to Eqs. [7] and [8] have been successfully used for the kinetic study of several autocatalytic permanganate reactions (31–33). Although, according to the latter equation, an increasing linear relationship would be expected when representing the ratio r/ [MnO2 ] against the difference [MnO2 ]0 – [MnO2 ], it can be seen in Fig. 5 that, in the absence of sodium pyrophosphate, once finished the rate-decreasing induction period, the ratio r/[MnO2 ] increased a little but reached soon a rather stationary value where the autocatalysis seemed to disappear, so that the kinetic behavior of the reaction can then be approximately described by a first-order rate law (horizontal stretch in the lower curve shown in Fig. 5). However, as can be seen in the same figure (upper curve), when the reaction
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rÉ0
1 D A400 , 1400l Dt
[9]
where D A400 is the change of absorbance of the reaction mixture at 400 nm between the instants t and t / Dt, with Dt Å 5 s, and 1400 is the molar absorptivity of colloidal manganese dioxide at that wavelength. Experimental Errors It should be noticed that both the accidental and systematic errors increase very rapidly as the number of rate constants to be determined in each kinetic experiment increases (34). In those reactions followed by spectrophotometry and where two rate constants have to be determined in each experiment, as is the case of autocatalytic reactions, one of the main sources of errors (both accidental and systematic) is the gradient between the temperature of the laboratory (external temperature, Te ) and that at which the kinetic runs are performed (internal temperature, Ti ). For instance, if Te õ Ti the entrance of cool air into the cuvette holder of the spectrophotometer at the beginning of the kinetic runs, and the
FIG. 5. Dependence of the rate/[MnO2 ] ratio on the decrease of MnO2 concentration during the kinetic runs in the absence (open circles) and presence (filled circles) of Na4P2O7 . The experimental conditions are given in the legend of Fig. 3.
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TABLE 1 Values of the Rate Constantsa at Various Initial Concentrations of Colloidal Manganese Dioxideb [MnO2]0c
k1d
1.00 1.20 1.40 1.60 1.80 2.00
7.1 6.6 5.9 5.2 5.1 4.6
{ { { { { {
k1e 0.2 0.2 0.1 0.1 0.2 0.1
4.5 4.0 3.5 3.3 3.2 3.1
{ { { { { {
k1f 0.1 0.1 0.1 0.1 0.1 0.1
12.7 11.4 12.2 11.5 12.2 11.1
{ { { { { {
k2d 0.3 0.1 0.8 0.8 0.1 0.7
26.8 24.4 23.4 22.5 20.9 20.1
{ { { { { {
k2e 0.5 0.6 0.1 0.2 0.1 0.1
19.2 17.9 17.2 15.9 14.7 13.8
{ { { { { {
k2f 0.6 0.2 0.1 0.1 0.2 0.1
88.1 81.1 72.5 67.5 60.1 57.5
{ { { { { {
1.3 0.5 1.0 1.8 0.5 0.7
Rate constant k1 is given in 1004 s01 and k2 in dm3 mol01 s01. [HAc] Å 0.525 mol dm03, [NaAc] Å 0.500 mol dm03, [Na4P2O7] Å 9.92 1 1003 mol dm03, pH 4.64, 25.07C. c In 1004 mol dm03. d [Na2C2O4] Å 2.00 1 1003 mol dm03; [gum arabic] Å 0. e [Na2C2O4] Å 2.00 1 1003 mol dm03; [gum arabic] Å 0.388 g dm03. f [Na2C2O4] Å 1.00 1 1002 mol dm03; [gum arabic] Å 0.388 g dm03. a b
subsequent increase of the temperature of that compartment to reach that of the thermostatic bath connected to it (Ti ) as the reaction progresses, result in considerable errors in the determination of both k1 (experimental value õ real value) and k2 (experimental value ú real value). On the contrary, when Te ú Ti the entrance of hot air into the cuvette holder at the beginning of the runs results in experimental values of k1 and k2 too high and too low, respectively. To minimize this problem, the temperature of the laboratory was controlled to keep the gradient ÉTe 0 TiÉ as low as possible (average differences: 0.57C for the experiments performed at 25.07C, and 3.07C for those corresponding to the determination of the Arrhenius parameters). Another source of systematic errors that should be taken into consideration is that related with the colloidal nature of one of the reactants (the oxidant); in that respect, since the reactivity of manganese dioxide is expected to change as the reaction progresses due to the decrease of the colloidal particles’ size, the values given in the present work for rate constant k2 have to be considered as mean values of that magnitude. Two independent determinations of each (k1 , k2 ) couple of values were done, the reproducibility of rate constant k1 being lower than that of k2 , although the accidental errors associated to the determination of both k1 and k2 were rather low, as indicated by the corresponding average standard deviations (2.9 and 1.2%, respectively). Some attempts were also made to measure the value of the initial rate corresponding to each kinetic run, but the sharp decrease of the reaction rate during the induction period (see Figs. 3 and 5) prevented us from extrapolating reproducible values of that magnitude. Kinetic Data Both rate constants k1 and k2 decreased as the initial concentration of colloidal manganese dioxide increased (Table 1). That decrease certainly reflects the existence of some
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systematic errors in the determination of the rate constants (see above), given that one of the basic tenets in chemical kinetics is that, when the isolation method is applied, the pseudorate constants obtained are independent of the initial concentration of the reactant in defect. To minimize the possible flocculation of the colloidal particles, the experiments were repeated in the presence of gum arabic, since this water-soluble polysaccharide is known to stabilize colloidal manganese dioxide in solution (35), but no improvement was observed. However, when the concentration of sodium oxalate was increased, k1 values notably independent of [MnO2 ]0 were obtained, although the decrease of k2 persisted. It should be noticed that rate constants depending on the initial concentration of the reactant in defect have been obtained in many permanganate reactions (36, 37), and in those with an autocatalytic character that dependence is especially notable in the case of k2 (14). Both rate constants increased with rising oxalate concentration in a nonlinear manner (Fig. 6), suggesting an approach toward a situation where the colloid surface is saturated by the reductant adsorbed on it. Addition of gum arabic resulted in a decrease of both k1 and k2 (see also Table 1), probably due to a competition with the reductant for the colloid surface. However, the inhibition effect caused by gum arabic was much less pronounced than the one found in a previous work for the reduction of colloidal manganese dioxide by formic acid (10). An important feature of the plots shown in Fig. 6 is that their extrapolation until zero concentration of reductant yields positive intercepts, suggesting that in the absence of oxalate some degree of reduction of colloidal manganese dioxide happened after addition of the buffering agents. By changing the concentrations of the two components of the buffer mixture, keeping constant both the total acetate concentration ([HAc] / [NaAc]) and the ionic strength (with excess NaClO4 ), it was observed that both rate constants exhibit acid catalysis (Fig. 7), but k1 was much more
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THE MnO2 –H2C2O4 REACTION
of the present work were made in the pH range 4.3–5.1 (Fig. 7), we can conclude that the reductant was present in the solution predominantly in the form of oxalate ion. Taking into consideration this fact along with the kinetic information found for rate constant k1 , we propose for the noncatalytic reaction pathway the following mechanism: / C2O 20 S HC2O 40 4 / H 0 4
/
[11]
HC2O / H S H2C2O4
[12]
(MnO2 )n / H2C2O4 S (MnO2 )n 0 H2C2O4
[13]
slow
(MnO2 )n 0 H2C2O4 FIG. 6. Dependence of rate constants k1 (open symbols) and k2 (filled symbols) on the concentration of Na2C2O4 for its oxidation by colloidal MnO2 (1.00 1 10 04 mol dm03 ), in the presence of HAc (0.420 mol dm03 ) – NaAc (0.600 mol dm03 ) buffer and Na4P2O7 (4.96 1 10 03 mol dm03 ), at pH 4.79 { 0.02 and 25.07C. Circles: [gum arabic] Å 0; triangles: [gum arabic] Å 0.232 g dm03 .
sensitive toward the solution pH than k2 . In a series of experiments where the concentration of NaClO4 was varied to study the effect of the ionic strength, k1 increased with rising perchlorate concentration to reach a maximum value and decrease later when that concentration was increased further, whereas k2 decreased uniformly (Table 2). Although addition of that electrolyte resulted in a decrease of the solution pH, once corrected the values of the rate constants to compensate for the pH variation, the dependences of k1 and k2 on the ionic strength at constant pH remained quite similar to those exhibited by the experimental, uncorrected values. Addition of manganese(II) sulfate (in concentrations not high enough to produce appreciable changes in the ionic strength of the solution) at the beginning of the kinetic runs resulted in an increase of k1 (Fig. 8, up) and a decrease of k2 (Fig. 8, down) but the former was much more pronounced, so that the net effect was an acceleration of the reaction (decrease of t1 / 2 ). The notable dependence of k1 on the initial manganese(II) concentration could be quantitatively described by a second-degree polynomial, k1 Å a / b[Mn 2/ ]0 / c[Mn 2/ ] 20 ,
(MnO2 )n01 0 MnO / 2CO2 / H2O
[14]
(MnO2 )n01 0 MnO / 2HAc r (MnO2 )n01 / Mn 2/ / 2Ac 0 / H2O.
[15]
Here (MnO2 )n stands for colloidal manganese dioxide. After adsorption of oxalic acid on the surface of the colloidal particles (Eq. [13]), the redox reaction takes place in the rate-determining step (Eq. [14]). Finally, in the last step (Eq. [15]) the acid component of the buffer mixture dissolves the manganese(II) oxide initially formed on the colloid surface. According to the mechanism proposed, the noncatalytic reaction pathway would be of first order in both colloidal manganese dioxide and adsorbed oxalic acid and, due to the saturation of the colloid surface at high reductant concentrations, the k1 vs [Na2C2O4 ] plots are expected to show a downward concavity (see Fig. 6). The assumption that molecular oxalic acid is the active reducing agent for Mn(IV) is in agreement with the results reported by Pimienta et al.
[10]
and the three experimental parameters involved (a, b, and c) increased with rising reductant concentration (Fig. 9). Finally, both rate constants fulfilled the Arrhenius law (Fig. 10), the corresponding activation energies (Table 3) indicating that the autocatalytic pathway (44.6 { 0.6 kJ mol 01 ) was more favorable from an energetic viewpoint than the noncatalytic one (74.7 { 0.9 kJ mol 01 ). DISCUSSION
Noncatalytic Reaction Pathway Given the dissociation constants of oxalic acid [pKa,1 Å 1.23 and pKa,2 Å 4.19 (38)], and that all the experiments
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FIG. 7. Dependence of rate constants k1 (open circles, slope 01.46 { 0.02) and k2 (filled circles, slope 00.67 { 0.03) on the pH for the reduction of colloidal MnO2 (1.00 1 10 04 mol dm03 ) by Na2C2O4 (2.00 1 10 03 mol dm03 ), in the presence of HAc–NaAc buffer ([HAc] / [NaAc] Å 5.04 1 10 02 mol dm03 ) and Na4P2O7 (4.96 1 10 03 mol dm03 ), at ionic strength 0.528 mol dm03 (NaClO4 ) and 25.07C.
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TABLE 2 Values of the Rate Constants at Various Concentrations of Sodium Perchloratea [NaClO4] (mol dm03)
pH
0.00 0.07 0.15 0.22 0.29 0.37 0.44 0.51 0.58 0.66
4.67 4.65 4.62 4.61 4.60 4.60 4.58 4.58 4.57 4.56
k1b (1004 s01) 3.66 4.86 5.19 4.99 4.91 4.60 4.82 4.38 4.23 3.97
{ { { { { { { { { {
k2b (dm3 mol01 s01)
0.21 0.02 0.05 0.10 0.05 0.12 0.15 0.17 0.01 0.01
33.8 31.5 30.3 29.5 28.5 27.9 26.8 26.1 25.7 25.0
{ { { { { { { { { {
0.2 0.2 0.1 0.1 0.1 0.1 0.5 0.1 0.3 0.1
k1c (1004 s01) 3.66 4.54 4.46 4.15 3.99 3.64 3.66 3.23 3.07 2.79
{ { { { { { { { { {
0.21 0.02 0.04 0.08 0.04 0.09 0.11 0.13 0.01 0.01
k2c (dm3 mol01 s01) 33.8 30.6 28.6 27.4 26.2 25.4 24.0 23.2 22.7 21.8
{ { { { { { { { { {
0.2 0.2 0.1 0.1 0.1 0.1 0.5 0.1 0.3 0.1
[MnO2]0 Å 1.00 1 1004 mol dm03, [Na2C2O4] Å 1.00 1 1003 mol dm03, [HAc] Å 0.105 mol dm03, [NaAc] Å 0.100 mol dm03, [Na4P2O7] Å 4.96 1 1003 mol dm03, 25.07C. b Experimental values. c Values corrected at constant pH (4.67). a
(23), and it can explain the strong acid catalysis found for rate constant k1 (Fig. 7). Moreover, the proposal of colloidal MnO2 being reduced first to MnO instead of directly to Mn 2/
may explain the fact that in neutral or alkaline solutions most reductants can provoke only a partial, superficial reduction of colloidal manganese dioxide (12, 39, 40), since, once covered by an insoluble, irreducible MnO monolayer, the colloid surface becomes passivated. It should be noticed that Pimienta et al. (23) have proposed that the reduction of Mn(IV) by oxalic acid yields Mn(III) in a first step, instead of Mn(II). However, our results indicate that at the beginning of the reaction Mn(II) was predominant over Mn(III), thus suggesting that Mn(IV) is directly reduced to Mn(II), and that Mn(III) is formed by reaction between Mn(IV) and Mn(II) once the latter becomes concentrated enough (see Fig. 2). Autocatalytic Reaction Pathway For the reaction pathway associated to rate constant k2 , we propose the following mechanism:
FIG. 8. Dependence of rate constants k1 (up) and k2 (down) on the initial concentration of MnSO4 for the reduction of colloidal MnO2 (1.00 1 10 04 mol dm03 ) by Na2C2O4 , in the presence of HAc (0.525 mol dm03 ) – NaAc (0.500 mol dm03 ) buffer, Na4P2O7 (4.96 1 10 03 mol dm03 ) and gum arabic (0.387 g dm03 ), at pH 4.61 { 0.01 and 25.07C. [Na2C2O4 ] Å 1.00 (open circles), 2.00 (filled circles), 3.00 (open triangles), 4.00 (filled triangles) and 5.00 (squares) 1 10 03 mol dm03 .
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FIG. 9. Dependence of parameters a (open circles, in 10 05 s 01 ), b (triangles, in dm3 mol 01 s 01 ) and c (filled circles, in 10 4 dm6 mol 02 s 01 ) on the concentration of sodium oxalate. The experimental conditions are given in the legend of Fig. 8.
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FIG. 10. Arrhenius plots for rate constants k1 (open circles) and k2 (filled circles) in the reduction of colloidal MnO2 (1.00 1 10 04 mol dm03 ) by Na2C2O4 (2.00 1 10 03 mol dm03 ), in the presence of HAc (0.525 mol dm03 ) –NaAc (0.500 mol dm03 ) buffer and Na4P2O7 (4.96 1 10 03 mol dm03 ), at pH 4.61 in the temperature range 14.7–34.87C.
Mn 2/ / C2O 20 4 S MnC2O4
[16]
(MnO2 )n 0 H2C2O4 / MnC2O4 S H2C2O4 0 (MnO2 )n 0 MnC2O4
[17]
slow
H2C2O4 0 (MnO2 )n 0 MnC2O4 (MnO2 )n01 / 2[Mn(C2O4 )] / / 2OH 0
[18]
[Mn(C2O4 )] / r Mn 2/ / CO2 / CO 2r0
[19]
2CO 2r0 r C2O 20 4 .
[20]
Here adsorption of manganese(II) oxalate on the surface of colloidal manganese dioxide (with oxalic acid previously adsorbed, Eq. [17]) is followed by the reaction between Mn(IV) and Mn(II) in the rate-determining step of this pathway (Eq. [18]) to yield a mono(oxalato)manganese(III) complex, whose unimolecular decomposition (Eq. [19]) results in the formation of the reaction products and a free radical. Finally, recombination of two of those radicals (Eq. [20]) leads to partial recuperation of oxalate ion. Our results seem to indicate that Mn(II) is the active autocatalyst (at least in the predominant autocatalytic reaction pathway), since Fig. 2 allows to infer the existence of a reaction between Mn(IV) and Mn(II) to yield Mn(III). Hence, the Mn(II) formed as reaction product contributes to accelerate the destruction of colloidal manganese dioxide. This was confirmed by the experiments performed in the presence of manganese(II) sulfate, since addition of Mn 2/ to the system resulted in an acceleration of the reaction (see Fig. 8). More difficult to explain is the quadratic term in the equa-
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tion describing the dependence of rate constant k1 on the initial concentration of manganese(II) ion (Eq. [10]). This might indicate that the existence of Mn 2/ in the electrical double layer surrounding the MnO2 colloidal particles favors the rate-determining step of the autocatalytic reaction pathway (Eq. [18]), probably by reacting with the two hydroxyl ions formed in that step to yield Mn(OH)2 (or hydrated MnO), thus preventing a decrease of the concentration of hydrogen-ion (a catalyst for the reaction, as shown by Fig. 7) in the surroundings of the colloidal particles. However, given the difficulty of obtaining precise values of the rate constants (especially of k1 ), the possibility that the deviation of the k1 vs [Mn 2/ ]0 plots (see Fig. 8, up) with respect to the linear relationship was due to the existence of systematic errors in the determination of k1 should not be excluded. The possible existence of a reaction of the type MnO2 – MnO in competition with the dissolution of MnO (Eq. [15]) should not be discarded either, although the known difficulty for the reaction between the oxides of manganese(IV) and manganese(II) to take place in the solid state (41) seems to indicate that the participation of the MnO2 –MnO reaction in the autocatalytic pathway might be of minor importance. In addition, according to the mechanism proposed, oxalate has an active participation in the autocatalytic reaction pathway (Eq. [18]). This is also in agreement with our experimental results, since we have observed that rate constant k2 increases with rising oxalate concentration (see Fig. 6) instead of being independent of that variable, as would happen if the autocatalytic pathway was due to a direct reaction between Mn(IV) and Mn(II) to yield Mn(III) without the participation of oxalate. Moreover, Fig. 9 confirms that the reductant plays an active role in the autocatalysis. On the other hand, Fig. 7 indicates that, although both reaction pathways present acid catalysis, the noncatalytic reaction pathway is much more influenced by the pH than the autocatalytic one. This might be a result of the participation of both oxalic acid and manganese(II) oxalate in the rate-determining step of the latter pathway (Eq. [18]), whereas in the one corresponding to the former (Eq. [14]) only oxalic acid is involved. It should be noticed that the existence of [Mn(C2O4 )] / (whose formation is proposed in Eq. [18]), as well as that of other complexes of Mn(III), is well known (42–44), TABLE 3 Arrhenius Parametersa Rate constant
ln Ab
Ea (kJ mol01)
k1 k2
22.2 { 0.4 21.6 { 0.2
74.7 { 0.9 44.6 { 0.6
a
The experimental conditions are given in the legend of Fig. 10. The Arrhenius preexponential factor (A) corresponding to rate constant k1 is given in s01 and that corresponding to k2 in dm3 mol01 s01. b
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and that there are many proofs on the involvement of free radicals (as the one whose participation is proposed in Eqs. [19] and [20]) in the manganese–oxalate reacting systems (45–49). Instead of recombination (Eq. [20]), other alternative for the destruction of those radicals would be (MnO2 )n 0 H2C2O4 / CO 2r0 r (MnO2 )n01 / [Mn(C2O4 )] / / CO2 / 2OH 0 .
[21]
Notice that, if this alternative was correct, Eqs. [19] and [21] would constitute a free-radical chain mechanism leading to the disappearance of the colloid. In that case, accumulation of Mn(III) in the system (see Fig. 2) would result in an increase (through Eq. [19]) of the concentration of free radicals available to reduce colloidal manganese dioxide, so that Mn(III) would also participate [along with Mn(II)] in the reaction as an autocatalyst. In fact, Jaky and Zrinyi (50) have suggested that Mn(III) might be the active autocatalytic agent for the reduction of an Mn(IV) –phosphate soluble complex by oxalic acid and other reductants. Induction Period Although it is difficult to ascribe the existence of a ratedecreasing induction period at the beginning of the kinetic runs (see Figs. 3 and 5) to a particular phenomenon, it might be related with the time necessary for the adsorptions of oxalic acid and manganese(II) oxalate on the colloid surface to reach an equilibrium situation (Eqs. [13] and [17]). Other possibility might be that the initial, sharp decrease of the reaction rate was due to the covering of the colloidal particles with an MnO monolayer (Eq. [14]) that would hinder the access of the reductant to their surface. Moreover, if Eq. [21] was correct, the induction period might be interpreted as the time necessary to elapse in order to reach an Mn(III) concentration high enough in the system (see Fig. 2). Stabilization of Mn(III) We have seen (Fig. 5, lower curve) that in the absence of pyrophosphate the acceleration of the reaction disappears very soon and a near-first-order situation is reached during a certain period. This situation might be the result of a compensation between the acceleration caused by the autocatalysis and the retardation caused by the dismutation of Mn(III), 2[Mn(C2O4 )] / / 2H2O r Mn 2/ / MnO2 / 2H2C2O4 , [22] since in that process the rate-monitoring species (MnO2 ) is partially recuperated. Given that pyrophosphate is known to be an excellent stabilizing agent for Mn(III) (51), when Na4P2O7 was added to the solution an acceleration of the reaction was observed (see Figs. 3 and 4), and the autocatal-
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ysis did not disappear (compare the two curves shown in Fig. 5). Comparison with Related Reactions In a previous work (10), we have reported a kinetic study of the reaction between colloidal manganese dioxide (prepared under identical conditions than the one used in the present work) and formic acid in perchloric acid aqueous media. The main differences found between the MnO2 – HCO2H and MnO2 –H2C2O4 reactions are that in the case of formic acid the reaction was much slower under similar conditions, did not show any appreciable autocatalysis and there was no rate-decreasing induction period. The apparent inexistence of autocatalysis in the MnO2 –HCO2H reaction can be taken as additional evidence that oxalate ions are necessary to facilitate the reaction between Mn(IV) and Mn(II) (Eq. [18]). However, it should be also noticed that the results shown in Fig. 7 seem to predict that if the acidity of the solution is high enough (as in the study corresponding to formic acid) the contribution of the autocatalytic reaction pathway would be negligible with respect to that of the noncatalytic one. The inexistence of a rate-decreasing induction period in the MnO2 –HCO2H reaction could be a consequence of the fact that the MnO initially formed on the colloid surface was rapidly dissolved due to the high acidity of the medium, so that the colloid was not protected by an irreducible MnO monolayer. However, both the MnO2 –HCO2H and MnO2 –H2C2O4 reactions exhibit a strong acid catalysis, and this situation contrasts with that found for the reduction of manganese dioxide by malonic acid (52), for which no dependence on [H / ] was observed. Jaky and Zrinyi (50) have reported the existence of a soluble Mn(IV) species in phosphoric acid aqueous media whose reactivity toward HCO2H and H2C2O4 is very similar to that of colloidal manganese dioxide, since only in the second case autocatalysis was found. It is interesting to notice that the UV–vis spectrum of that Mn(IV) –phosphate species (apparently noncolloidal, according to ultracentrifugal and electron microscopic studies) was very similar to that of colloidal MnO2 (12), and that phosphate ions are known to stabilize the MnO2 colloidal particles in aqueous solution by fixation on their surface (6, 12). Finally, mention should be made of the fact that there might be a close relationship between the autocatalysis found in the reactions of oxalic acid with permanganate ion (53) and colloidal manganese dioxide. To this respect, a parallelism can be established between the reactivities of the two oxidants, since the MnO 40 –H2C2O4 and MnO2 –H2C2O4 reactions are both markedly autocatalytic, whereas their related the MnO 40 –HCO2H and MnO2 –HCO2H reactions present, respectively, a slight autocatalysis (32) and no autocatalysis at all (10). This might suggest that the autocatalysis found
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THE MnO2 –H2C2O4 REACTION
in the first couple of reactions is caused predominantly by the reduction of either permanganate ion or colloidal manganese dioxide by MnC2O4 (see Eq. [18]), which would confirm the explanation proposed by Adler and Noyes (54) for the autocatalysis observed in the permanganate–oxalate reaction. ACKNOWLEDGMENT The authors thank Dr. J. Estelrich for the measurement of the zeta potential of the colloidal form of manganese dioxide used in this work.
REFERENCES 1. Abd El-Salaam, K. M., Z. Phys. Chem. (Frankfurt am Main) 95, 139 (1975). 2. Sharma, T. C., Lal, A., and Saksena, V., Bull. Chem. Soc. Jpn. 49, 2881 (1976). 3. Basak, B., and Malati, M. A., J. Inorg. Nucl. Chem. 39, 1081 (1977). 4. Kienzle, F., Tetrahedron Lett. 24, 2213 (1983). 5. Taniguchi, S., Bull. Chem. Soc. Jpn. 57, 2683 (1984). 6. Freeman, F., Rev. React. Species Chem. React. 1, 179 (1976). 7. Lume-Pereira, C., Baral, S., Henglein, A., and Janata, E., J. Phys. Chem. 89, 5772 (1985). 8. Perez-Benito, J. F., Brillas, E., and Pouplana, R., Inorg. Chem. 28, 390 (1989). 9. Perez-Benito, J. F., and Lee, D. G., Can. J. Chem. 63, 3545 (1985). 10. Perez-Benito, J. F., and Arias, C., J. Colloid Interface Sci. 149, 92 (1992). 11. Freeman, F., and Kappos, J. C., J. Am. Chem. Soc. 107, 6628 (1985). 12. Perez-Benito, J. F., and Arias, C., J. Colloid Interface Sci. 152, 70 (1992). 13. Lee, D. G., and Perez-Benito, J. F., Can. J. Chem. 63, 1275 (1985). 14. Perez-Benito, J. F., and Arias, C., Int. J. Chem. Kinet. 23, 717 (1991). 15. Dallison, A. M., Macer, D. R. J., and Rodley, G. A., Inorg. Chim. Acta 76, L219 (1983). 16. Cooke, D. O., J. Chem. Soc. Chem. Commun., 589 (1984). 17. Sevcik, P., and Adamcikova, L., Collect. Czech. Chem. Commun. 52, 2125 (1987). 18. Sevcik, P., and Adamcikova, L., J. Chem. Phys. 91, 1012 (1989). 19. Fujiwara, K., Kashima, T., Tsubota, H., Toyoshima, Y., Aihara, M., and Kiboku, M., Chem. Lett., 1385 (1990). 20. Keki, S., and Beck, M. T., React. Kinet. Catal. Lett. 44, 75 (1991). 21. Bradley, J., and Van Praagh, G., J. Chem. Soc., 1624 (1938). 22. Xyla, A. G., Sulzberger, B., Luther, G. W., Hering, J. G., Van Cappellen, P., and Stumm, W., Langmuir 8, 95 (1992).
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23. Pimienta, V., Lavabre, D., Levy, G., and Micheau, J. C., J. Phys. Chem. 98, 13294 (1994). 24. Ladbury, J. W., and Cullis, C. F., Chem. Rev. 58, 403 (1958). 25. Stewart, R., in ‘‘Oxidation in Organic Chemistry’’ (K. B. Wiberg, Ed.), Part A, p. 65. Academic Press, New York, 1965. 26. Pimienta, V., Lavabre, D., Levy, G., and Micheau, J. C., J. Phys. Chem. 96, 9298 (1992). 27. Simandi, L. I., and Jaky, M., J. Am. Chem. Soc. 98, 1995 (1976). 28. Freeman, F., Fuselier, C. O., Armstead, C. R., Dalton, C. E., Davidson, P. A., Karchefski, E. M., Krochman, D. E., Johnson, M. N., and Jones, N. K., J. Am. Chem. Soc. 103, 1154 (1981). 29. Perez-Benito, J. F., Brillas, E., and Arias, C., Can. J. Chem. 68, 79 (1990). 30. Jaky, M., and Simandi, L. I., J. Chem. Soc. Perkin Trans. 2, 1481 (1972). 31. De Andres, J., Brillas, E., Garrido, J. A., and Perez-Benito, J. F., J. Chem. Soc. Perkin Trans. 2, 107 (1988). 32. Perez-Benito, J. F., Arias, C., and Brillas, E., Int. J. Chem. Kinet. 22, 261 (1990). 33. Andres Ordax, F. J., Arrizabalaga, A., Peche, R., and Quintana, M. A., An. Quim. 88, 440 (1992). 34. Perez-Benito, J. F., Lamrhari, D., and Arias, C., J. Phys. Chem. 98, 12621 (1994). 35. Tompkins, F. C., Trans. Faraday Soc. 38, 131 (1942). 36. Senent, S., and Cuadrado, S., An. Quim. Ser. B 57, 11 (1961). 37. Wiberg, K. B., and Stewart, R., J. Am. Chem. Soc. 77, 1786 (1955). 38. Weast, R. C. (Ed.), ‘‘Handbook of Chemistry and Physics,’’ p. D150. CRC Press, Cleveland, 1977. 39. Ohashi, H., and Hirohara, K., Denki Kagaku 40, 610 (1972). 40. Ohashi, H., Denki Kagaku 40, 681 (1972). 41. Sigua, T. I., Svanidze, L. K., and Avaliani, A. T., Soobshch. Akad. Nauk Gruz. SSR 120, 129 (1985). 42. Taube, H., J. Am. Chem. Soc. 69, 1418 (1947). 43. Taube, H., J. Am. Chem. Soc. 70, 1216 (1948). 44. Adamcikova, L., Krizova, A., and Valent, I., Transition Met. Chem. 18, 218 (1993). 45. Konar, R. S., and Palit, S. R., J. Polym. Sci. Part A 2, 1731 (1964). 46. Mathur, P. C., and Bansal, O. P., J. Indian Chem. Soc. 43, 375 (1966). 47. Kolthoff, I. M., Meehan, E. J., and Kimura, M., J. Phys. Chem. 75, 3343 (1971). 48. Ganapathisubramanian, N., J. Phys. Chem. 92, 414 (1988). 49. Mukhopadhyay, S., Chaudhuri, S., Das, R., and Banerjee, R., Can. J. Chem. 71, 2155 (1993). 50. Jaky, M., and Zrinyi, M., Polyhedron 12, 1271 (1993). 51. Chandraju, S., Sherigara, B. S., and Made Gowda, N. M., Int. J. Chem. Kinet. 26, 1105 (1994). 52. Rao, I., Mishra, S. K., and Sharma, P. D., Transition Met. Chem. 18, 182 (1993). 53. Reckley, J. S., and Showalter, K., J. Am. Chem. Soc. 103, 7012 (1981). 54. Adler, S. J., and Noyes, R. M., J. Am. Chem. Soc. 77, 2036 (1955).
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177, 288–297 (1996)
0034
A Kinetic Study of the Reduction of Colloidal Manganese Dioxide by Oxalic Acid JOAQUIN F. PEREZ-BENITO,* ,1 CONCHITA ARIAS,*
AND
ELISENDA AMAT†
*Departamento de Quimica Fisica, Facultad de Quimica, Universidad de Barcelona, Marti i Franques, 1, 08028 Barcelona, Spain; and †Unidad de Fisicoquı´mica, Facultad de Farmacia, Universidad de Barcelona, Avenida Diagonal, s/n, 08028 Barcelona, Spain Received May 15, 1995; accepted June 14, 1995
A kinetic study of the reaction between colloidal manganese dioxide and oxalic acid in aqueous acetate media (pH 4.3–5.1) is reported. The reaction is autocatalytic and, in order to determine the rate constants k1 and k2 corresponding, respectively, to the noncatalytic and autocatalytic reaction pathways, the Mn(III) formed as an intermediate was stabilized by addition of sodium pyrophosphate. Joint iodimetric and spectrophotometric studies indicated that the reduction of colloidal manganese dioxide follows the sequence R
Mn(IV) r Mn(II)
Mn ( IV )
R
Mn(III) r Mn(II),
where R stands for the reductant. Both reaction pathways exhibit acid catalysis, and the activation energies associated to k1 and k2 are 74.7 { 0.9 and 44.6 { 0.6 kJ mol 01 , respectively. The reaction is accelerated by addition of manganese(II) sulfate, and k2 increases with rising oxalate concentration, suggesting that MnC2O4 is the active autocatalyst. The possibility that Mn(III) might collaborate in the autocatalysis through a free-radical chain mechanism is also pointed out. q 1996 Academic Press, Inc. Key Words: autocatalysis; colloidal solution; kinetics; manganese dioxide; oxalic acid.
or reaction products in most permanganate oxidations (11, 12), being actively involved in the mechanism as autocatalysts in many cases (13, 14). On the other hand, although the kinetic aspects of the manganese–oxalate reacting systems have gained new interest in the last decade due to their involvement in some oscillating reactions (15–20), the reduction of colloidal manganese dioxide by oxalic acid has received much little attention (21–23) than its relative, the permanganate–oxalate reaction (24, 25). We have now undertaken a kinetic study of the reaction between colloidal manganese dioxide (in the form of a perfectly transparent, aqueous sol) and oxalic acid and found that it is in fact an autocatalytic reaction, whose stoichiometry may be written as MnO2 / H2C2O4 / 2H / Å Mn 2/ / 2CO2 / 2H2O.
[1]
We think that the results presented here for the title reaction might throw some light upon the common aspects of the autocatalytic behavior exhibited by both the MnO2 –H2C2O4 and MnO 40 –H2C2O4 (26) reactions. MATERIALS AND METHODS
INTRODUCTION
Preparation of the Colloidal Manganese Dioxide Sol Manganese dioxide is a substance of certain importance owing to its catalytic (1) and oxidizing (2–5) properties, but its applications are notably limited because of its insolubility under ordinary conditions (6). In recent years, however, the preparation of several forms of colloidal manganese dioxide stable for long periods as brown or yellow, perfectly transparent solutions, either in aqueous (7, 8) or organic (9) media, has made possible the kinetic study of its oxidizing behavior in redox reactions by easily implemented, conventional spectrophotometric methods (10). Moreover, the transparent sols of manganese dioxide are also of importance because of their widespread participation as intermediates 1
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/ 8MnO 40 / 3S2O 20 Å 8MnO2 / 6SO 20 3 / 2H 4 / H2O.
[2] Provided that the oxidant and reductant concentrations are used in stoichiometric ratio, a dark-brown solution is ob-
To whom correspondence should be addressed.
0021-9797/96 $12.00 Copyright q 1996 by Academic Press, Inc. All rights of reproduction in any form reserved.
To prepare the colloid, 10 ml of KMnO4 (0.100 mol dm03 ) and 20 ml of Na2S2O3 (1.88 1 10 02 mol dm03 ) aqueous solutions were mixed together in a 2-liter volumetric flask, and water was added until completion, the solution being homogeneized afterwards by gentle stirring. The reaction taking place under these conditions may be described by the stoichiometry
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THE MnO2 –H2C2O4 REACTION
tained, remaining perfectly transparent at least for several months (8, 10). Colloid Characterization The colloidal solution so prepared ([MnO2 ] Å 5.00 1 10 04 mol dm03 ) was characterized by means of a Malvern autosizer 2.C, the results indicating that 90% of the colloidal particles had a diameter comprised in the range 89–193 nm, with a peak at 120 nm. The zeta potential (measured with a Malvern zeta-sizer 4.C) was 047 { 3 mV. In addition, coagulation experiments performed in a previous work (8) indicated that the colloidal particles had a negative electrostatic charge, thus indicating that their stability in solution was probably due to the fixation of some anions on the colloid surface. The colloid’s UV–vis spectrum was recorded with a Hitachi U 2000 UV–vis double-beam spectrophotometer, and showed a large band covering the whole visible region of the spectrum, with absorbance uniformly decreasing with increasing wavelength, as well as a wide maximum at 300–400 nm. Other Reagents The solvent used was water previously purified by deionization followed by double distillation ( the first time from a potassium permanganate solution ) . The reductant used in the kinetic experiments was sodium oxalate ( always in a large excess with respect to MnO2 ) . To keep the pH constant during the runs, an acetic acid ( HAc ) – sodium acetate ( NaAc ) buffer mixture was used ( pH 4.3 – 5.1 ) , since the reaction either in aqueous HClO4 or in the presence of H3PO4 – NaH2PO4 buffer ( pH É 2 ) was too fast to be followed by conventional spectrophotometry. Manganese ( III ) was stabilized in the solution by addition of sodium pyrophosphate. When necessary, the ionic strength was kept constant with a large excess of sodium perchlorate. Other additives used in some of the experiments were manganese ( II ) sulfate, gum arabic and EDTA ( tetrasodium salt ) . All the chemicals used were purchased from Merck ( analytical quality ) . Iodimetric Determinations In some experiments, the average oxidation state (OSav ) of manganese during the MnO2 –H2C2O4 reaction was determined by taking successive 2-ml aliquots of the reaction mixture (buffered with HAc–NaAc and kept in a thermostatic bath at 25.07C) and introducing them into 5-ml volumetric flasks containing an excess of solid potassium iodide. After dilution with water until completion, the solutions were homogeneized and the iodine liberated was determined by UV–vis spectrophotometry [ lmax (I 30 ) Å 351 nm, 1351 Å 2.57 1 10 4 dm3 mol 01 cm01 ]. In accordance with many other previous works (27–29), it was assumed that both Mn(IV)
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FIG. 1. Average oxidation state of the manganese atom during the reduction of colloidal MnO2 (9.62 1 10 05 mol dm03 ) by Na2C2O4 (9.62 1 10 04 mol dm03 ), in the presence of HAc (0.504 mol dm03 ) –NaAc (0.481 mol dm03 ) buffer, at pH 4.57 and 25.07C. Open circles: [Na4P2O7 ] Å 0; filled circles: [Na4P2O7 ] Å 4.77 1 10 03 mol dm03 .
and Mn(III) were reduced by iodide ion to the Mn(II) state. Then, the average oxidation state of manganese could be obtained as OSav Å 2 /
5A351 , 1351l[Mn]T
[3]
where A351 and 1351 are the absorbance of the solution and the molar absorptivity of triiodide ion at 351 nm, respectively, l is the optical pathlength of the cuvettes (1 cm), and [Mn]T is the total manganese concentration present in the system during the reaction. Kinetic Measurements The progress of the reaction was followed by means of a Varian Cary 219 UV–vis spectrophotometer provided with a thermostatized cuvette holder. The wavelength chosen was 400 nm, taking into consideration that colloidal manganese dioxide had a strong absorption there, and that neither manganese(II) ion nor the oxalate and pyrophosphate complexes of manganese(III) showed any appreciable absorption. The pH’s of the solutions were determined by means of a Metrohm 605 pH-meter provided with a glass–calomel combination electrode. RESULTS
Influence of Pyrophosphate The average oxidation state of manganese during the MnO2 –H2C2O4 reaction both in the absence and in the presence of sodium pyrophosphate was determined by iodimetry. We can see in Fig. 1 that when Na4P2O7 was not present the
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FIG. 2. Fractions of Mn(IV) (open circles), Mn(III) (triangles), and Mn(II) (filled circles) during the reduction of colloidal manganese dioxide by sodium oxalate in the presence of sodium pyrophosphate. The experimental conditions are given in the legend of Fig. 1.
average oxidation state decreased uninterruptedly from 4 to 2, whereas in the presence of an enough amount of that additive the reduction took place until the oxidation state 3 in a first stage. At that moment, the UV–vis spectrum showed a small peak at 470 nm, thus confirming the presence in the solution of a manganese(III) –pyrophosphate complex (30). Under those conditions, the total reduction until the final Mn(II) state took several hours. Moreover, although the concentration of Mn(III) during the reaction in the experiment performed in the absence of pyrophosphate was too low to be determined, in the one performed in the presence of that additive it was possible to determine the fractions of the total manganese present in the system during the reaction in the oxidation states II, III, and IV ( fII , fIII , and fIV, respectively). To that end, the decay of Mn(IV) was followed by spectrophotometry at 400 nm for the same experimental conditions used in the iodimetric experiment (Fig. 1, upper curve). Since at that wavelength only Mn(IV) exhibits an appreciable absorption, the equations necessary to calculate the fractions of each manganese species present in the system at each particular instant t are fII / fIII / fIV Å 1
[4]
OSav Å 2 fII / 3 fIII / 4 fIV
[5]
fIV Å
[Mn IV ] A400 Å 0 [Mn]T A 400
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Determination of the Rate Constants The most simple way of describing mathematically the kinetic behavior of an autocatalytic reaction is to assume that two reaction pathways (one noncatalytic and the other autocatalytic) are involved, and that the kinetic orders of the reactant in defect (in both pathways) and of the autocatalyst (in the autocatalytic pathway) are the unity. If, in addition, we assume that the concentration of autocatalytic product at a particular instant t is directly proportional to the decrease of the reactant concentration from the beginning
[6]
where A400 and A 0400 are the absorbances of the solution at 400 nm at time t and at the beginning of the reaction, respectively. It can be seen in Fig. 2 that the fraction of Mn(IV) decreased with time according to a sigmoid curve typical of autocatalytic reactions. The fraction of Mn(III) increased first and then started to decrease very slowly, whereas that
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of Mn(II) increased, passed through a maximum, decreased, and finally showed a slow increase. It is interesting to notice that the final conversion of Mn(III) to Mn(II) seemed to occur when Mn(IV) was not present in the system any longer. When the reaction was followed spectrophotometrically, a small induction period was observed during which the reaction rate showed a fast decrease. After that, the rate started to increase in a typically autocatalytic manner, passing through a maximum and decreasing afterward (Fig. 3). Addition of sodium pyrophosphate to the solution at the beginning of the kinetic runs resulted in a clear acceleration of the reaction, since an increase of the reaction rate both at the minimum and the maximum of the rate vs. time plots was observed, whereas the half-life period of the reaction (t1 / 2 ) decreased (Fig. 4). On the other hand, an attempt to study the effect on the reaction of EDTA [a good complexing agent for both Mn(II) and Mn(III)] failed because in the media of our work colloidal manganese dioxide was reduced by that reactant even faster than by oxalic acid.
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FIG. 3. Dependence of the reaction rate on time during the reduction of colloidal MnO2 (1.00 1 10 04 mol dm03 ) by Na2C2O4 (1.00 1 10 03 mol dm03 ), in the presence of HAc (0.525 mol dm03 ) –NaAc (0.500 mol dm03 ) buffer, at pH 4.60 { 0.01 and 25.07C. Open circles: [Na4P2O7 ] Å 0; filled circles: [Na4P2O7 ] Å 4.96 1 10 03 mol dm03 .
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THE MnO2 –H2C2O4 REACTION
took place in the presence of a concentration of sodium pyrophosphate high enough, the rate-decreasing induction period was followed by another period in which the autocatalysis did not disappear and Eq. [8] clearly holds. This finding has allowed us to determine the rate constants k1 and k2 from the intercept and slope of the linear portion of the r/[MnO2 ] vs [MnO2 ]0 – [MnO2 ] plots, respectively. To that end, the experiments were performed in the presence of a pyrophosphate concentration high enough to stabilize the autocatalysis after the minimum of those plots (see Fig. 5, upper curve). In each kinetic run, the reaction rate was estimated at 5-s intervals as FIG. 4. Dependence of the rate/[MnO2 ] ratios at the minimum (open circles) and the maximum (filled circles) of the rate vs time plots, and of the half-life period (triangles), on the concentration of Na4P2O7 for the reduction of colloidal MnO2 (1.00 1 10 04 mol dm03 ) by Na2C2O4 (1.00 1 10 04 mol dm03 ), in the presence of HAc (0.525 mol dm03 ) –NaAc (0.500 mol dm03 ) buffer, at pH 4.60 { 0.01 and 25.07C.
of the kinetic run until that moment, we have the differential rate law rÅ0
d[MnO2 ] dt
Å k1[MnO2 ] / k2[MnO2 ]([MnO2 ]0 0 [MnO2 ]) [7]
or else r Å k1 / k2 ([MnO2 ]0 0 [MnO2 ]), [MnO2 ]
[8]
where r is the reaction rate at time t, [MnO2 ]0 and [MnO2 ] are the concentrations of colloidal manganese dioxide at the beginning of the reaction and at time t, respectively, k1 is the pseudo-first-order rate constant of the noncatalytic pathway, and k2 is the pseudo-second-order rate constant of the autocatalytic pathway. Differential rate laws similar to Eqs. [7] and [8] have been successfully used for the kinetic study of several autocatalytic permanganate reactions (31–33). Although, according to the latter equation, an increasing linear relationship would be expected when representing the ratio r/ [MnO2 ] against the difference [MnO2 ]0 – [MnO2 ], it can be seen in Fig. 5 that, in the absence of sodium pyrophosphate, once finished the rate-decreasing induction period, the ratio r/[MnO2 ] increased a little but reached soon a rather stationary value where the autocatalysis seemed to disappear, so that the kinetic behavior of the reaction can then be approximately described by a first-order rate law (horizontal stretch in the lower curve shown in Fig. 5). However, as can be seen in the same figure (upper curve), when the reaction
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rÉ0
1 D A400 , 1400l Dt
[9]
where D A400 is the change of absorbance of the reaction mixture at 400 nm between the instants t and t / Dt, with Dt Å 5 s, and 1400 is the molar absorptivity of colloidal manganese dioxide at that wavelength. Experimental Errors It should be noticed that both the accidental and systematic errors increase very rapidly as the number of rate constants to be determined in each kinetic experiment increases (34). In those reactions followed by spectrophotometry and where two rate constants have to be determined in each experiment, as is the case of autocatalytic reactions, one of the main sources of errors (both accidental and systematic) is the gradient between the temperature of the laboratory (external temperature, Te ) and that at which the kinetic runs are performed (internal temperature, Ti ). For instance, if Te õ Ti the entrance of cool air into the cuvette holder of the spectrophotometer at the beginning of the kinetic runs, and the
FIG. 5. Dependence of the rate/[MnO2 ] ratio on the decrease of MnO2 concentration during the kinetic runs in the absence (open circles) and presence (filled circles) of Na4P2O7 . The experimental conditions are given in the legend of Fig. 3.
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TABLE 1 Values of the Rate Constantsa at Various Initial Concentrations of Colloidal Manganese Dioxideb [MnO2]0c
k1d
1.00 1.20 1.40 1.60 1.80 2.00
7.1 6.6 5.9 5.2 5.1 4.6
{ { { { { {
k1e 0.2 0.2 0.1 0.1 0.2 0.1
4.5 4.0 3.5 3.3 3.2 3.1
{ { { { { {
k1f 0.1 0.1 0.1 0.1 0.1 0.1
12.7 11.4 12.2 11.5 12.2 11.1
{ { { { { {
k2d 0.3 0.1 0.8 0.8 0.1 0.7
26.8 24.4 23.4 22.5 20.9 20.1
{ { { { { {
k2e 0.5 0.6 0.1 0.2 0.1 0.1
19.2 17.9 17.2 15.9 14.7 13.8
{ { { { { {
k2f 0.6 0.2 0.1 0.1 0.2 0.1
88.1 81.1 72.5 67.5 60.1 57.5
{ { { { { {
1.3 0.5 1.0 1.8 0.5 0.7
Rate constant k1 is given in 1004 s01 and k2 in dm3 mol01 s01. [HAc] Å 0.525 mol dm03, [NaAc] Å 0.500 mol dm03, [Na4P2O7] Å 9.92 1 1003 mol dm03, pH 4.64, 25.07C. c In 1004 mol dm03. d [Na2C2O4] Å 2.00 1 1003 mol dm03; [gum arabic] Å 0. e [Na2C2O4] Å 2.00 1 1003 mol dm03; [gum arabic] Å 0.388 g dm03. f [Na2C2O4] Å 1.00 1 1002 mol dm03; [gum arabic] Å 0.388 g dm03. a b
subsequent increase of the temperature of that compartment to reach that of the thermostatic bath connected to it (Ti ) as the reaction progresses, result in considerable errors in the determination of both k1 (experimental value õ real value) and k2 (experimental value ú real value). On the contrary, when Te ú Ti the entrance of hot air into the cuvette holder at the beginning of the runs results in experimental values of k1 and k2 too high and too low, respectively. To minimize this problem, the temperature of the laboratory was controlled to keep the gradient ÉTe 0 TiÉ as low as possible (average differences: 0.57C for the experiments performed at 25.07C, and 3.07C for those corresponding to the determination of the Arrhenius parameters). Another source of systematic errors that should be taken into consideration is that related with the colloidal nature of one of the reactants (the oxidant); in that respect, since the reactivity of manganese dioxide is expected to change as the reaction progresses due to the decrease of the colloidal particles’ size, the values given in the present work for rate constant k2 have to be considered as mean values of that magnitude. Two independent determinations of each (k1 , k2 ) couple of values were done, the reproducibility of rate constant k1 being lower than that of k2 , although the accidental errors associated to the determination of both k1 and k2 were rather low, as indicated by the corresponding average standard deviations (2.9 and 1.2%, respectively). Some attempts were also made to measure the value of the initial rate corresponding to each kinetic run, but the sharp decrease of the reaction rate during the induction period (see Figs. 3 and 5) prevented us from extrapolating reproducible values of that magnitude. Kinetic Data Both rate constants k1 and k2 decreased as the initial concentration of colloidal manganese dioxide increased (Table 1). That decrease certainly reflects the existence of some
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systematic errors in the determination of the rate constants (see above), given that one of the basic tenets in chemical kinetics is that, when the isolation method is applied, the pseudorate constants obtained are independent of the initial concentration of the reactant in defect. To minimize the possible flocculation of the colloidal particles, the experiments were repeated in the presence of gum arabic, since this water-soluble polysaccharide is known to stabilize colloidal manganese dioxide in solution (35), but no improvement was observed. However, when the concentration of sodium oxalate was increased, k1 values notably independent of [MnO2 ]0 were obtained, although the decrease of k2 persisted. It should be noticed that rate constants depending on the initial concentration of the reactant in defect have been obtained in many permanganate reactions (36, 37), and in those with an autocatalytic character that dependence is especially notable in the case of k2 (14). Both rate constants increased with rising oxalate concentration in a nonlinear manner (Fig. 6), suggesting an approach toward a situation where the colloid surface is saturated by the reductant adsorbed on it. Addition of gum arabic resulted in a decrease of both k1 and k2 (see also Table 1), probably due to a competition with the reductant for the colloid surface. However, the inhibition effect caused by gum arabic was much less pronounced than the one found in a previous work for the reduction of colloidal manganese dioxide by formic acid (10). An important feature of the plots shown in Fig. 6 is that their extrapolation until zero concentration of reductant yields positive intercepts, suggesting that in the absence of oxalate some degree of reduction of colloidal manganese dioxide happened after addition of the buffering agents. By changing the concentrations of the two components of the buffer mixture, keeping constant both the total acetate concentration ([HAc] / [NaAc]) and the ionic strength (with excess NaClO4 ), it was observed that both rate constants exhibit acid catalysis (Fig. 7), but k1 was much more
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THE MnO2 –H2C2O4 REACTION
of the present work were made in the pH range 4.3–5.1 (Fig. 7), we can conclude that the reductant was present in the solution predominantly in the form of oxalate ion. Taking into consideration this fact along with the kinetic information found for rate constant k1 , we propose for the noncatalytic reaction pathway the following mechanism: / C2O 20 S HC2O 40 4 / H 0 4
/
[11]
HC2O / H S H2C2O4
[12]
(MnO2 )n / H2C2O4 S (MnO2 )n 0 H2C2O4
[13]
slow
(MnO2 )n 0 H2C2O4 FIG. 6. Dependence of rate constants k1 (open symbols) and k2 (filled symbols) on the concentration of Na2C2O4 for its oxidation by colloidal MnO2 (1.00 1 10 04 mol dm03 ), in the presence of HAc (0.420 mol dm03 ) – NaAc (0.600 mol dm03 ) buffer and Na4P2O7 (4.96 1 10 03 mol dm03 ), at pH 4.79 { 0.02 and 25.07C. Circles: [gum arabic] Å 0; triangles: [gum arabic] Å 0.232 g dm03 .
sensitive toward the solution pH than k2 . In a series of experiments where the concentration of NaClO4 was varied to study the effect of the ionic strength, k1 increased with rising perchlorate concentration to reach a maximum value and decrease later when that concentration was increased further, whereas k2 decreased uniformly (Table 2). Although addition of that electrolyte resulted in a decrease of the solution pH, once corrected the values of the rate constants to compensate for the pH variation, the dependences of k1 and k2 on the ionic strength at constant pH remained quite similar to those exhibited by the experimental, uncorrected values. Addition of manganese(II) sulfate (in concentrations not high enough to produce appreciable changes in the ionic strength of the solution) at the beginning of the kinetic runs resulted in an increase of k1 (Fig. 8, up) and a decrease of k2 (Fig. 8, down) but the former was much more pronounced, so that the net effect was an acceleration of the reaction (decrease of t1 / 2 ). The notable dependence of k1 on the initial manganese(II) concentration could be quantitatively described by a second-degree polynomial, k1 Å a / b[Mn 2/ ]0 / c[Mn 2/ ] 20 ,
(MnO2 )n01 0 MnO / 2CO2 / H2O
[14]
(MnO2 )n01 0 MnO / 2HAc r (MnO2 )n01 / Mn 2/ / 2Ac 0 / H2O.
[15]
Here (MnO2 )n stands for colloidal manganese dioxide. After adsorption of oxalic acid on the surface of the colloidal particles (Eq. [13]), the redox reaction takes place in the rate-determining step (Eq. [14]). Finally, in the last step (Eq. [15]) the acid component of the buffer mixture dissolves the manganese(II) oxide initially formed on the colloid surface. According to the mechanism proposed, the noncatalytic reaction pathway would be of first order in both colloidal manganese dioxide and adsorbed oxalic acid and, due to the saturation of the colloid surface at high reductant concentrations, the k1 vs [Na2C2O4 ] plots are expected to show a downward concavity (see Fig. 6). The assumption that molecular oxalic acid is the active reducing agent for Mn(IV) is in agreement with the results reported by Pimienta et al.
[10]
and the three experimental parameters involved (a, b, and c) increased with rising reductant concentration (Fig. 9). Finally, both rate constants fulfilled the Arrhenius law (Fig. 10), the corresponding activation energies (Table 3) indicating that the autocatalytic pathway (44.6 { 0.6 kJ mol 01 ) was more favorable from an energetic viewpoint than the noncatalytic one (74.7 { 0.9 kJ mol 01 ). DISCUSSION
Noncatalytic Reaction Pathway Given the dissociation constants of oxalic acid [pKa,1 Å 1.23 and pKa,2 Å 4.19 (38)], and that all the experiments
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FIG. 7. Dependence of rate constants k1 (open circles, slope 01.46 { 0.02) and k2 (filled circles, slope 00.67 { 0.03) on the pH for the reduction of colloidal MnO2 (1.00 1 10 04 mol dm03 ) by Na2C2O4 (2.00 1 10 03 mol dm03 ), in the presence of HAc–NaAc buffer ([HAc] / [NaAc] Å 5.04 1 10 02 mol dm03 ) and Na4P2O7 (4.96 1 10 03 mol dm03 ), at ionic strength 0.528 mol dm03 (NaClO4 ) and 25.07C.
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TABLE 2 Values of the Rate Constants at Various Concentrations of Sodium Perchloratea [NaClO4] (mol dm03)
pH
0.00 0.07 0.15 0.22 0.29 0.37 0.44 0.51 0.58 0.66
4.67 4.65 4.62 4.61 4.60 4.60 4.58 4.58 4.57 4.56
k1b (1004 s01) 3.66 4.86 5.19 4.99 4.91 4.60 4.82 4.38 4.23 3.97
{ { { { { { { { { {
k2b (dm3 mol01 s01)
0.21 0.02 0.05 0.10 0.05 0.12 0.15 0.17 0.01 0.01
33.8 31.5 30.3 29.5 28.5 27.9 26.8 26.1 25.7 25.0
{ { { { { { { { { {
0.2 0.2 0.1 0.1 0.1 0.1 0.5 0.1 0.3 0.1
k1c (1004 s01) 3.66 4.54 4.46 4.15 3.99 3.64 3.66 3.23 3.07 2.79
{ { { { { { { { { {
0.21 0.02 0.04 0.08 0.04 0.09 0.11 0.13 0.01 0.01
k2c (dm3 mol01 s01) 33.8 30.6 28.6 27.4 26.2 25.4 24.0 23.2 22.7 21.8
{ { { { { { { { { {
0.2 0.2 0.1 0.1 0.1 0.1 0.5 0.1 0.3 0.1
[MnO2]0 Å 1.00 1 1004 mol dm03, [Na2C2O4] Å 1.00 1 1003 mol dm03, [HAc] Å 0.105 mol dm03, [NaAc] Å 0.100 mol dm03, [Na4P2O7] Å 4.96 1 1003 mol dm03, 25.07C. b Experimental values. c Values corrected at constant pH (4.67). a
(23), and it can explain the strong acid catalysis found for rate constant k1 (Fig. 7). Moreover, the proposal of colloidal MnO2 being reduced first to MnO instead of directly to Mn 2/
may explain the fact that in neutral or alkaline solutions most reductants can provoke only a partial, superficial reduction of colloidal manganese dioxide (12, 39, 40), since, once covered by an insoluble, irreducible MnO monolayer, the colloid surface becomes passivated. It should be noticed that Pimienta et al. (23) have proposed that the reduction of Mn(IV) by oxalic acid yields Mn(III) in a first step, instead of Mn(II). However, our results indicate that at the beginning of the reaction Mn(II) was predominant over Mn(III), thus suggesting that Mn(IV) is directly reduced to Mn(II), and that Mn(III) is formed by reaction between Mn(IV) and Mn(II) once the latter becomes concentrated enough (see Fig. 2). Autocatalytic Reaction Pathway For the reaction pathway associated to rate constant k2 , we propose the following mechanism:
FIG. 8. Dependence of rate constants k1 (up) and k2 (down) on the initial concentration of MnSO4 for the reduction of colloidal MnO2 (1.00 1 10 04 mol dm03 ) by Na2C2O4 , in the presence of HAc (0.525 mol dm03 ) – NaAc (0.500 mol dm03 ) buffer, Na4P2O7 (4.96 1 10 03 mol dm03 ) and gum arabic (0.387 g dm03 ), at pH 4.61 { 0.01 and 25.07C. [Na2C2O4 ] Å 1.00 (open circles), 2.00 (filled circles), 3.00 (open triangles), 4.00 (filled triangles) and 5.00 (squares) 1 10 03 mol dm03 .
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FIG. 9. Dependence of parameters a (open circles, in 10 05 s 01 ), b (triangles, in dm3 mol 01 s 01 ) and c (filled circles, in 10 4 dm6 mol 02 s 01 ) on the concentration of sodium oxalate. The experimental conditions are given in the legend of Fig. 8.
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THE MnO2 –H2C2O4 REACTION
FIG. 10. Arrhenius plots for rate constants k1 (open circles) and k2 (filled circles) in the reduction of colloidal MnO2 (1.00 1 10 04 mol dm03 ) by Na2C2O4 (2.00 1 10 03 mol dm03 ), in the presence of HAc (0.525 mol dm03 ) –NaAc (0.500 mol dm03 ) buffer and Na4P2O7 (4.96 1 10 03 mol dm03 ), at pH 4.61 in the temperature range 14.7–34.87C.
Mn 2/ / C2O 20 4 S MnC2O4
[16]
(MnO2 )n 0 H2C2O4 / MnC2O4 S H2C2O4 0 (MnO2 )n 0 MnC2O4
[17]
slow
H2C2O4 0 (MnO2 )n 0 MnC2O4 (MnO2 )n01 / 2[Mn(C2O4 )] / / 2OH 0
[18]
[Mn(C2O4 )] / r Mn 2/ / CO2 / CO 2r0
[19]
2CO 2r0 r C2O 20 4 .
[20]
Here adsorption of manganese(II) oxalate on the surface of colloidal manganese dioxide (with oxalic acid previously adsorbed, Eq. [17]) is followed by the reaction between Mn(IV) and Mn(II) in the rate-determining step of this pathway (Eq. [18]) to yield a mono(oxalato)manganese(III) complex, whose unimolecular decomposition (Eq. [19]) results in the formation of the reaction products and a free radical. Finally, recombination of two of those radicals (Eq. [20]) leads to partial recuperation of oxalate ion. Our results seem to indicate that Mn(II) is the active autocatalyst (at least in the predominant autocatalytic reaction pathway), since Fig. 2 allows to infer the existence of a reaction between Mn(IV) and Mn(II) to yield Mn(III). Hence, the Mn(II) formed as reaction product contributes to accelerate the destruction of colloidal manganese dioxide. This was confirmed by the experiments performed in the presence of manganese(II) sulfate, since addition of Mn 2/ to the system resulted in an acceleration of the reaction (see Fig. 8). More difficult to explain is the quadratic term in the equa-
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tion describing the dependence of rate constant k1 on the initial concentration of manganese(II) ion (Eq. [10]). This might indicate that the existence of Mn 2/ in the electrical double layer surrounding the MnO2 colloidal particles favors the rate-determining step of the autocatalytic reaction pathway (Eq. [18]), probably by reacting with the two hydroxyl ions formed in that step to yield Mn(OH)2 (or hydrated MnO), thus preventing a decrease of the concentration of hydrogen-ion (a catalyst for the reaction, as shown by Fig. 7) in the surroundings of the colloidal particles. However, given the difficulty of obtaining precise values of the rate constants (especially of k1 ), the possibility that the deviation of the k1 vs [Mn 2/ ]0 plots (see Fig. 8, up) with respect to the linear relationship was due to the existence of systematic errors in the determination of k1 should not be excluded. The possible existence of a reaction of the type MnO2 – MnO in competition with the dissolution of MnO (Eq. [15]) should not be discarded either, although the known difficulty for the reaction between the oxides of manganese(IV) and manganese(II) to take place in the solid state (41) seems to indicate that the participation of the MnO2 –MnO reaction in the autocatalytic pathway might be of minor importance. In addition, according to the mechanism proposed, oxalate has an active participation in the autocatalytic reaction pathway (Eq. [18]). This is also in agreement with our experimental results, since we have observed that rate constant k2 increases with rising oxalate concentration (see Fig. 6) instead of being independent of that variable, as would happen if the autocatalytic pathway was due to a direct reaction between Mn(IV) and Mn(II) to yield Mn(III) without the participation of oxalate. Moreover, Fig. 9 confirms that the reductant plays an active role in the autocatalysis. On the other hand, Fig. 7 indicates that, although both reaction pathways present acid catalysis, the noncatalytic reaction pathway is much more influenced by the pH than the autocatalytic one. This might be a result of the participation of both oxalic acid and manganese(II) oxalate in the rate-determining step of the latter pathway (Eq. [18]), whereas in the one corresponding to the former (Eq. [14]) only oxalic acid is involved. It should be noticed that the existence of [Mn(C2O4 )] / (whose formation is proposed in Eq. [18]), as well as that of other complexes of Mn(III), is well known (42–44), TABLE 3 Arrhenius Parametersa Rate constant
ln Ab
Ea (kJ mol01)
k1 k2
22.2 { 0.4 21.6 { 0.2
74.7 { 0.9 44.6 { 0.6
a
The experimental conditions are given in the legend of Fig. 10. The Arrhenius preexponential factor (A) corresponding to rate constant k1 is given in s01 and that corresponding to k2 in dm3 mol01 s01. b
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and that there are many proofs on the involvement of free radicals (as the one whose participation is proposed in Eqs. [19] and [20]) in the manganese–oxalate reacting systems (45–49). Instead of recombination (Eq. [20]), other alternative for the destruction of those radicals would be (MnO2 )n 0 H2C2O4 / CO 2r0 r (MnO2 )n01 / [Mn(C2O4 )] / / CO2 / 2OH 0 .
[21]
Notice that, if this alternative was correct, Eqs. [19] and [21] would constitute a free-radical chain mechanism leading to the disappearance of the colloid. In that case, accumulation of Mn(III) in the system (see Fig. 2) would result in an increase (through Eq. [19]) of the concentration of free radicals available to reduce colloidal manganese dioxide, so that Mn(III) would also participate [along with Mn(II)] in the reaction as an autocatalyst. In fact, Jaky and Zrinyi (50) have suggested that Mn(III) might be the active autocatalytic agent for the reduction of an Mn(IV) –phosphate soluble complex by oxalic acid and other reductants. Induction Period Although it is difficult to ascribe the existence of a ratedecreasing induction period at the beginning of the kinetic runs (see Figs. 3 and 5) to a particular phenomenon, it might be related with the time necessary for the adsorptions of oxalic acid and manganese(II) oxalate on the colloid surface to reach an equilibrium situation (Eqs. [13] and [17]). Other possibility might be that the initial, sharp decrease of the reaction rate was due to the covering of the colloidal particles with an MnO monolayer (Eq. [14]) that would hinder the access of the reductant to their surface. Moreover, if Eq. [21] was correct, the induction period might be interpreted as the time necessary to elapse in order to reach an Mn(III) concentration high enough in the system (see Fig. 2). Stabilization of Mn(III) We have seen (Fig. 5, lower curve) that in the absence of pyrophosphate the acceleration of the reaction disappears very soon and a near-first-order situation is reached during a certain period. This situation might be the result of a compensation between the acceleration caused by the autocatalysis and the retardation caused by the dismutation of Mn(III), 2[Mn(C2O4 )] / / 2H2O r Mn 2/ / MnO2 / 2H2C2O4 , [22] since in that process the rate-monitoring species (MnO2 ) is partially recuperated. Given that pyrophosphate is known to be an excellent stabilizing agent for Mn(III) (51), when Na4P2O7 was added to the solution an acceleration of the reaction was observed (see Figs. 3 and 4), and the autocatal-
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ysis did not disappear (compare the two curves shown in Fig. 5). Comparison with Related Reactions In a previous work (10), we have reported a kinetic study of the reaction between colloidal manganese dioxide (prepared under identical conditions than the one used in the present work) and formic acid in perchloric acid aqueous media. The main differences found between the MnO2 – HCO2H and MnO2 –H2C2O4 reactions are that in the case of formic acid the reaction was much slower under similar conditions, did not show any appreciable autocatalysis and there was no rate-decreasing induction period. The apparent inexistence of autocatalysis in the MnO2 –HCO2H reaction can be taken as additional evidence that oxalate ions are necessary to facilitate the reaction between Mn(IV) and Mn(II) (Eq. [18]). However, it should be also noticed that the results shown in Fig. 7 seem to predict that if the acidity of the solution is high enough (as in the study corresponding to formic acid) the contribution of the autocatalytic reaction pathway would be negligible with respect to that of the noncatalytic one. The inexistence of a rate-decreasing induction period in the MnO2 –HCO2H reaction could be a consequence of the fact that the MnO initially formed on the colloid surface was rapidly dissolved due to the high acidity of the medium, so that the colloid was not protected by an irreducible MnO monolayer. However, both the MnO2 –HCO2H and MnO2 –H2C2O4 reactions exhibit a strong acid catalysis, and this situation contrasts with that found for the reduction of manganese dioxide by malonic acid (52), for which no dependence on [H / ] was observed. Jaky and Zrinyi (50) have reported the existence of a soluble Mn(IV) species in phosphoric acid aqueous media whose reactivity toward HCO2H and H2C2O4 is very similar to that of colloidal manganese dioxide, since only in the second case autocatalysis was found. It is interesting to notice that the UV–vis spectrum of that Mn(IV) –phosphate species (apparently noncolloidal, according to ultracentrifugal and electron microscopic studies) was very similar to that of colloidal MnO2 (12), and that phosphate ions are known to stabilize the MnO2 colloidal particles in aqueous solution by fixation on their surface (6, 12). Finally, mention should be made of the fact that there might be a close relationship between the autocatalysis found in the reactions of oxalic acid with permanganate ion (53) and colloidal manganese dioxide. To this respect, a parallelism can be established between the reactivities of the two oxidants, since the MnO 40 –H2C2O4 and MnO2 –H2C2O4 reactions are both markedly autocatalytic, whereas their related the MnO 40 –HCO2H and MnO2 –HCO2H reactions present, respectively, a slight autocatalysis (32) and no autocatalysis at all (10). This might suggest that the autocatalysis found
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THE MnO2 –H2C2O4 REACTION
in the first couple of reactions is caused predominantly by the reduction of either permanganate ion or colloidal manganese dioxide by MnC2O4 (see Eq. [18]), which would confirm the explanation proposed by Adler and Noyes (54) for the autocatalysis observed in the permanganate–oxalate reaction. ACKNOWLEDGMENT The authors thank Dr. J. Estelrich for the measurement of the zeta potential of the colloidal form of manganese dioxide used in this work.
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AP: Colloid