J. Phys. Chem. 1985, 89, 5779-5783
5779
Chemistry of Colloidal Manganese Dioxide. 2. Reaction with 02-and H20, (Pulse Radiolysis and Stop Flow Studies) S. Baral, C. Lume-Pereira, E. Janata, and A. Henglein* Hahn- Meitner-Institut fur Kernforschung Berlin, Bereich Strahlenchemie, 0-1000 Berlin 39, Federal Republic of Germany (Received: June 1 1 , 1985; In Final Form: August 29, 1985)
The reactions of 0,- and H 2 0 2with colloidal MnO, in alkaline solution were studied by using pulse radiolysis and stop flow techniques. In both reactions, Mn4+ions in Mn02 are reduced to a certain degree. The reduction processes are autocatalytic, and the rate can be increased by the addition of manganous ions. These effects are explained in terms of the reduction mechanism proposed in part 1. The autocatalysis involves (1) reduction of Mn3+centers at the surface of the colloidal particles by the substrate and (2) conproportionation of MnZ+and Mn4+centers. When Mn3+centers have been accumulated to a certain degree, the conproportionation reaction is slowed down. 02-then catalytically dismutes, and H202is catalytically decomposed.
1. Introduction When manganese dioxide is reduced in alkaline solution, Mn3+ centers are formed in the solid and play an important role in the course of the reaction. In part 1 it was shown that the reduction can be followed by using MnOz in the form of a transparent sol and measuring the changes in optical absorption and conductivity which accompany the reduction.’ In the present paper, the reactions of the superoxide anion, OF, and of hydrogen peroxide with MnO, are described. The reaction of H,02 has quite often been studied2-I2 by using powdered MnO, which does not allow one to detect directly the changes in the solid. In acidic solutions, HzOzessentially reduces MnO, to Mn2+,and in alkaline solutions, H 2 0 2is catalytically decomposed. 02-is an intermediate of the In the present Mn02-catalyzed decomposition of H202.10*11 studies, 02-. was produced radiolytically in alkaline MnO, sols and its reactions were followed by using the method of pulse radiolysis. The reaction of H202was investigated with the stop flow method. In both cases, MnOz is reduced to a certain degree and activated in this way. The mechanism of reduction which was described in part 1 is applied to explain the features of the 0,- and H 2 0 2reactions.
2. Experimental Section The preparation of the colloid and the pulse radiolysis set up with the optical and conductometric detection systems have been described in part 1. I A commercial stop flow apparatus, Model D-110 Durrum Dionex, was used in the studies of the H202 MnO, reaction. The following apparative improvements were made. Optical absorption vs. time curves were recorded and stored by a programmable digitizer (7D20T, Tektronix). The digitizer was connected to a computer (micro-1 1, DEC) via an IEEE-488 bus for data analysis and storage. In order to overcome the time consuming procedure of adjusting the photomultiplier light level and the baseline everytime the wavelength was changed, the buffer amplifier of the apparatus was replaced by back-off circuitry for automatic base-line compensation and measurement of the
+
( I ) Lume-Pereira, C.; Baral, S.; Henglein, A,; Janata, E. J. Phys. Chem., preceding paper in this issue. (2) “Gmelins Handbuch der Anorganischen Chemie”; Verlag Chemie: Weinheim, West Germany, 1973; 8. Auflage, Mangan 56, Teil C1, p 304. (3). “Gmelins Handbuch der Anorganischen Chemie”; (1966) Verlag Chemie: Weinheim; West Germany, 1966; 8. Auflage, Sauerstoff 3, p 2356-2359. (4) Broughton, D. B.; Wentworth, R. L. J . Am. Chem. SOC.1947,69,741. (5) Broughton, D. B.; Wentworth, R. L.; Laing, M. E. J . Am. Chem. Sor. 1947, 69, 744. (6) Scholder, R.; Kolb, A. Z . Anorg. Chem. 1949, 260, 231. (7) Fouinat, F.; Magat, M. J . Chim. Phys. 1950, 47, 514. (8) Mooi, J.; Selwood, P. W. J . A m . Chem. SOC.1952, 74, 1750. (9) Roy, C. B. J . Catal. 1968, 12, 129. (10) Ono, Y . ; Matsumura, T.; Kitajima, N.; Fukuzumi, %-I. J . Phys. Chem. 1977, 81, 1307. (1 1) Kitajima, N.; Fukuzumi, S.; Ono, Y . J. Phys. Chem. 1978,82, 1505. (12) Vasilenko, I. I . Russ. J . Phys. Chem. 1983, 57, 1642.
0022-3654/85/2089-5779!§01 S O / O
background current.I3 In the beginning of the recording, the solution of the preceding experiment was still in the cell. It was removed by the new mixture within a few milliseconds. During this removal the absorption of the solution changed, and the photomultiplier current varied correspondingly. The photomultiplier current was compensated to zero during all this time and then the compensation current was kept constant during the following recording. The compensating current was read by the computer and used for the calculation of absorbances or absorption coefficients.
3. Results 3.1. Reaction of 0,- with Colloidal MnO,. 0; radical anions were generated by irradiation of alkaline aerated solutions containing both 2-propanol and acetone.I4 The primary radicals from the radiolysis of water, i.e., hydrated electrons, hydroxyl radicals, and hydrogen atoms, are known to react according to H,O
-+ - + ea;,
eaq- + (CH3),C0 OH(H)
H+
+ (CH3),HCOH
(CH3)ZCOH
+02
-
H+, OH, H
(1)
(CH3),COH
H20(H2)
(2)
(CH3),C0H
(3)
OH-
(CH3)2C(OH)O2 (CH3),C0
+ 02-(4)
All these reactions are fast and complete at a few microseconds after the pulse. Note that six 0,- radicals are formed per 100 eV of absorbed radiation energy and that the molar conductivity of the solution is decreased by AA = 120 0-l cm2 as OH- ions are substituted by less mobile 02-ions. The observed changes in absorbance were divided by the concentration of the radicals generated in the pulse to obtain an apparent change Ac in the absorption coefficient of the solution. Figure 1 shows these changes for a pulsed 2 X M Mn02 sol containing 0.4 M 2-propanol, 0.2 M acetone, and air. The pH of the solution was 10. A concentration of 0,- radicals of 1.5 X lod M was generated per pulse. The solution was uniformly irradiated. The ciphers at the curves give the number of pulses with which the sample was preirradiated before the signals were recorded after the monitoring pulse. The prepulses were applied as a train, the time interval between two successive pulses being 20 s. Parts a and b of Figure 1 show the respective changes in Ae and AA upon preirradiation with 0-17 pulses, and parts a’ and b’ show what happens upon application of 33-62 prepulses. It is seen from part a that the absorbance at 380 nm, where MnO, absorbs,’ decreases after the pulse, the curves being slightly S-shaped, which indicates an induction period or the occurrence of an autocatalytic process. The final decrease in specific ab~~
~~
(13) Janata, E. patent application FRG P 3432091.1, 1984. (14) Ilan, Y . ;Rabani, J.; Henglein, A. J. Phys. Chem. 1976, 80, 1558.
0 1985 American Chemical Societv
5780
The Journal of Physical Chemistry, Vol. 89, No. 26, 1985
Baral et al.
t l c l
Figure 1. Time profiles of the specific absorption Ae at 380 nm and the molar conductivity A h of pulsed solutions containing 2 X lo4 M MnO,, 0.4 M 2-propanol, 0.2 M acetone, and air, and having pH 10. The numbers at the curves give the number of pulses with which the samples were preirradiated.
sorbance was about IO4 M-I cm-I which corresponds to the consumption of one M n 0 2 molecule per 0, formed. The conductivity is decreased immediately after the pulse as expected from the Occurrence of the fast reactions 1-4. After the pulse, the negative conductivity signal recovered which indicates that OH- was regenerated during the reaction of 0, with MnOz. The rate of these changes in Figure la,b increased with increasing preirradiation, i.e., the colloid became activated by the preceding 02attack. The half-life for corresponding curves of A€ and AA were practically equal. The curves observed for preirradiation with 17-33 pulses were M 02-radicals had acted rather similar. Note that (2-5) X on the 2 X M M n 0 2 in this range of preirradiation. New phenomena occur at stronger preirradiations (Figure 1, parts a' and b'). The optical absorption starts to increase within less than 1 s after the pulse, but then decreases, ultimately leading again to a bleaching of the solution. With further preirradiation the initial rise in absorbance increases in magnitude, while the ultimate bleaching toward the end of the 4-s time window in Figure 1 decreases. The conductivity seems to recover in two steps. The first step occurs in the time range where the absorption rises; the second step takes place at much longer times. The final recovery of the negative signal becomes smaller with increasing preirradiation. In the absence of colloid, the conductivity very slowly recovered, the first half-life of this recovery being longer than 10 s, as the 0,- radicals disappeared via the dismutation 2 H 2 0 + 20,-
-
H 2 0 , + O2 + 2 0 H -
(5) Curve 0 in Figure 1b was not much different from the curve obtained in the absence of colloid. This means that only a small fraction of the 0,- radicals produced in the first pulse reacted with MnO,. The hydrogen peroxide produced in the dismutation of 0,- (eq 5) could then react with M n 0 2 during the time which elapsed until the next pulse of radiation was given to the solution. The fact that different time profiles of A€ and AA were observed after application of only a few pulses shows that small amounts of 0; or H202produced in the preirradiation were able to activate the colloid. Experiments were also carried out in which the solution was strongly preirradiated with 6 2 pulses, then left to stand for 5 h until the monitoring pulse was applied. The signals obtained were quite different from those observed in Figure la',b'. In fact, the signals resembled the signals for preirradiation with 17 pulses (Figure la,b). This shows that the colloid was able to recover its activity which had been decreased by too high a preirradiation. When experiments of this type were carried out with smaller preiriadiations, no recovery effects were observed. It was found that the colloid could also be activated by Mn2+ ions or small amounts of hydrogen peroxide. Different amounts of these additives were given to the 2 X M M n 0 2 sol and the conductivity recorded after the first pulse (without preirradiation). In Figure 2 one finds a plot of the half-life of the recovery of the negative conductivity signal. It is seen that the reaction of 0,- with MnO, becomes faster with increasing concentration
I
"0
I
40
20
60
cant. 110-6M I
Figure 2. First half-life of the recovery of the conductivity signal as a function of the concentration of Mn2*and H202added (no preirradiation).
0.8 I
5 N
E 0.6 u " c 0
0.L 0
02
n 200
LOO
600
hinml
I
II
I
1
0 200
300
LOO
500
X lnml
Figure 3. (a) Absorption spectrum of a 4 X lo4 M MnO, sol before and after reaction with various amounts of added hydrogen peroxide; pH 10.6. (b) Absorption spectrum of a 4 X M MnO, sol before and after reaction with various amounts of added hydrogen peroxide. Argon was bubbled through the solution during reaction to flush out the oxygen produced.
of the additive, H20zbeing more efficient than Mn2+. At concentrations higher than 10 KM H 2 0 2 or 20 MMMn2+,an acceleration of the reaction is still noticeable, although it is no longer so pronounced as for the lower concentrations. 3.2. Absorption Spectrum of MnOz after Reaction of HZO2. Figure 3a shows the absorption spectrum of a 4 X M MnO, sol before and after reaction with various amounts of hydrogen peroxide. A small amount of a more concentrated H 2 0 2solution was rapidly mixed with the sol and the spectrum measured after 1 h, Le., when oxygen no longer was developed. The intensity of the spectrum decreased in a similar way as in the reduction of the colloid by an organic radical.' In the beginning, the decrease in absorption became more and more pronounced as the amount of added H 2 0 2 increased. However, at H 2 0 2 concentrations greater than lod3M, the decrease became smaller again. At H202 concentrations above 6 X M, the final spectrum was always the same. Figure 3b shows the results of similar experiments in which the solution was bubbled with argon before and during the reaction with H2O2. Under these conditions, any oxygen developed during
The Journal of Physical Chemistry, Vol. 89, No. 26, 1985 5781
Chemistry of Colloidal Manganese Dioxide
L
10-6
10-6
0
0 0
2.0
4,O
6.0 IH2021 [lVLM1
8.0
Figure 4. Final MnO, concentration after reaction of 4 X IO4 M MnO, colloid with various amounts of added H202. Dashed curve: ratio R as a function of the concentration of added H20, (see eq IO). tIsl 1
0
2
-
Figure 5. Decrease in MnO, concentration as a function of time. ConM MnO, and 1 X M centration of the reacting mixture: 4 X H202; pH 10.6.
2
t\ 2o 0
10-~ 104
io-)
io-1
IH202IIMI
10.0
Figure 7. Initial rate ro of MnO, reduction as a function of H20, concentration. 13
I
10
4 2
0
I ti202 I I 10-4M I
Figure 8. Time ti and decrease in Mn02 concentration -Aq at the turning point (see Figure 5) and final decrease -Ac, as functions of H202concentration. Initial M n 0 2 concentration: 4 X M; pH 10.6.
I
100
tlsl
Figure 6. Decrease in Mn0, concentration-a scale. Conditions same as in Figure 5 .
more compressed time
the reaction was flushed out and could not act as a reactant with reduced manganese centers. The comparison between the two figures reveals that the course of reaction was quite different in the absence of 0,. Much more MnO, was reduced than in the presence of 0,. In both types of experiments, the H202was completely consumed. In both types of experiments, H202also was added in two portions. 1 X loW3M H202was added and after completion of the reaction, a second portion of 1 X M H202was added to the solution. In both cases, the addition of the second portion caused further consumptions of Mn02, while the direct addition of 2 X M H202was accompanied by a weaker consumption. In the direct addition of 2 X M H 2 0 2 ,the O2 evolution was more violent, Le., 0, as a reactant for reduced manganese centers was present at a higher concentration. All these experiments clearly show that the oxygen which is developed in the reaction of H202with MnO, has a strong influence on the course of the reaction. The results obtained in the experiments in Figure 3a for H202 concentrations below M were used to draw the curve (full line) in Figure 4. The M n 0 2 concentration after the completion of reaction is plotted here vs. the concentration of H202added. The M n 0 2 concentration was calculated from the 340-nm absorbance of the solution after reaction, where the reduction product of MnO, has negligible absorption.’ 3.3. Stop Flow Measurements. Figures 5 and 6 show typical curves obtained in recording the 380-nm absorption of a solution after mixing the M n 0 2 sol with H 2 0 2solution. The M n 0 2 sol contained 8 X M manganese dioxide, the hydrogen peroxide solution was 2 X M. As equal volumes were mixed, both concentrations in the reacting mixture were halved. On the ordinate scale of the figure one finds the decrease in MnO, concentration, calculated from the decrease in 380-nm absorbance. The curve in Figure 5 is S-shaped which indicates that the reaction
20 40 60
80 100 120 140 160s
2-04
5 -0.8
Figure 9. Concentration of oxygen evolved and decrease in MnO, concentration as functions of time. 4 X M Mn02; M H202;pH 10. Total 0, produced: 5 X M.
between MnO, and H 2 0 2is autocatalytic. However, at longer times the reaction strives toward a stationary state where the M n 0 2 concentration no longer is decreased. ro, the initial rate of reaction, was calculated from the time at M. ro was which -AcMnO2had a given small value of 2 X measured for various H 2 0 2concentrations; a double logarithmic plot of the data is shown in Figure 7. The straight line with slope one indicates that the reaction follows the rate law:
ro = k[MnO2l[H2021
(6)
k was obtained as 7 X 10, M-’ s-’ . In Figure 8, one finds the time t, and concentration change -Acl at the inflection point of the S-shaped curve (see Figure 5) and the final decrease -Acfin M n 0 2 concentration, Le., [MnO2l0- [Mn021fas functions of the H202concentration. It is seen that strong changes in these entities M. At the occur at small H202concentrations below 1 X higher H202concentrations, -Acf was almost constant at 2.5 X = 1.5 X M or, in other words, M, Le., (4.0-2.5) X (1.5 X 10-5)/(4.0 X X 100 = 38% of the MnO, remained unchanged. The concentration change -Ac, at the inflection point was about 7 times smaller than -Acf. The question now arises as to whether most of the oxygen from the decomposition of hydrogen peroxide is formed during phase I of reaction when substantial reduction of M n 0 2 takes place or during phase I1 where almost no reduction occurs any more. In Figure 9 one finds again a curve from the stop flow experiments which gives the decrease in MnO, concentration as a function of
5782 The Journal of Physical Chemistry, Vol. 89, No. 26, 1985
-
0
0
2
4
6
8
1
0
tis1
Figure 10. O2 produced as a function of time after addition of M H20, to 4 X M Mn02sol in the absence and presence of Mn2+.The fast increase in the beginning occurs in the time range where -AcMn02 is fastest (Figure 5); pH 10.6.
time. In addition, the 0, concentration produced during the reaction is shown. It was measured with a Clark electrode. It is recognized that after 60 s, when 90% of the final reduction of M n 0 2 has been reached (phase I), only one-third of the oxygen was produced. In other words, two-thirds of the reaction of H,O, with MnO, occurred in phase 11, Le., under conditions where the concentration of Mn0, remained almost constant. During phase I, reduction of MnO, and oxidation of H202to form 0, + 2Hf was an important part of the overall decomposition of H202,while during phase I1 mainly disproportionation of H202to form 0, + H,O occurred. Experiments also were carried out in which small amounts of manganous perchlorate were added to the MnO, sol. The very first part of the reaction became much faster. At [Mn2+] = 1 X M, for example, the curve was no longer S-shaped. The 0, evolution was also faster in the beginning of the reaction, as can be seen from a comparison of the two curves in Figure 10. 4. Discussion 4.1. Reaction of 0,-with MnO,. Mn4+ ions in MnO, are consumed in the reaction with 02-as the absorption of the colloid is decreased (Figure la). Similar changes in absorbance were found in part I, where the reaction of a reducing organic radical with MnO, was investigated. It is concluded that 02- reduces MnO, via electron transfer.' During this reaction, an OH- ion is released from the colloid (Figure Ib) as it was found in the reaction described in part I. As pointed out in part I, a colloidal particle may be described by the formula (Mn02),(OH-)o.5,. The reaction with 02-radical anions therefore may be written as
This reaction is very slow in the beginning (curves 0 in Figure 1). However, after preirradiation with up to 17 pulses in Figure 1, the reaction becomes faster and faster. This shows that the reaction is autocatalytic. A similar effect was also found for the reaction of the 1-hydroxy-1-methylethyl radical with Mn02.' The autocatalytic nature of the reaction also explains the slight S-shape of the curves in Figure la. Autocatalysis can only be possible if active centers are accumulated in the colloid. In part 1 it was pointed out that Mn$+ centers are responsible for the activation of the colloid. These centers react with the radicals and are regenerated and even increased in their number via the following sequence of reactions: Mn$+
+ 0,
-
Mnc2++ 0,
(8)
Mn;+ +- Mn?+ 2Mn,3+ (9) The index c is to indicate that the species is part of the colloid, probably on its surface, and not in solution. OH- is released in reaction 8 and not in reaction 9. In the beginning of the multiple pulse experiment of Figure 1, where the preirradiation affects the colloid only slightly ( n
