Benon H. J. Bielski and Augustine 0. Allen
Mechanism of the Disproportionation of Superoxide Radicals' Benon H. J. Blelskl" and Augustine 0. Allen Chemistry Department, Brookhaven National Laboratory, Upton, New York 1 1973 (Received February 4, 1977) Publication costs assisted by Brookhaven Natlonal Laboratory
New measurements are reported on the rate of the spontaneous uncatalyzed second-order decay of the superoxide radical in aqueous solution at 23 "C. By extreme purification of reagents, and addition of small amounts of EDTA to sequester traces of catalytic metal ions, purely second-orderkinetics were found up to pH 13.2. The pH effects clearly show that the ionic form 0; never reacts with itself but only with the H02 in equilibrium with it, even when the ratio HOz/Oz- is as small as 2 X lo-'; the rate constant for 0; + 0; is less than 0.3 M-' s-'. The pK of HO2 was determined as 4.75 f 0.08, and the bimolecular rate constants were kl = (7.61 f 0.55) x 105 M-' s-1 for H02 + HOz and kz = (8.86 f 0.43) X lo' M-' s-' for HO2 + O;, all in good agreement with published values. The change in the spontaneous decay of superoxide radicals with pH follows the equation kobsd = [k,+ k~K~o,/(H')1/[1 + KHO~/(H+)]~. The optical extinction coefficients of hydrogen peroxide were determined over the pH range 6.1-13.8 at wavelengths from 280 to 230 nm.
In the course of developing techniques for studying the ihemical reactions of radiation-generated superoxide radicals, we have obtained considerable data on their spontaneous decay, which is due to their disproportionation to form oxygen and hydrogen peroxide. Our measurement techniques include pulse radiolysis and the recently described method of stopped-flow radiolysis;' but equally important are the techniques of purification of reagents. Although the disproportionation of these radicals should be a purely second-order reaction, previous workers3s4 have found that in an alkaline solution a first-order component, attributed to the presence of catalytic impurities, was superimposed on the second-order reaction. By reducing heavy metal impurities in our solutions to the order of a few parts per billion, and working always in the presence of ethylenediaminetetraacetic acid (EDTA) to sequester the remaining traces of metals, we have been able to carry the determinations of reaction rate to a higher pH than previous workers. The results indicate that the disproportionation of the ionic form 0 2 - in aqueous solution appears always to occur by reaction with the small quantity of HOz present in equilibrium, and there is no evidence that the ionic form 0; is capable of reacting directly with itself. Experimental Section Materials. Ordinary distilled water was purified by a Milli-Q reagent grade water system obtained from the Millipore Corporation of Bedford, Mass. This treatment, consisting of deionization followed by ultrafiltration, was found to produce water more nearly free of catalytic heavy metal impurities than the formerly used multiply distilled water. Reagent grade EDTA, purchased as the disodium salt, was dissolved in water and precipitated in the acid form by addition of perchloric acid. The acid form was then recrystallized directly from hot water. Analysis showed that the original product contained 10 ppm of Fe which was reduced to 0.037 ppm by two recrystallizations. The product was always recrystallized at least three times before use. Perchloric acid was the double vacuum distilled product from G. Frederick Smith Chemical Co., guaranteed to contain less than 0.01 ppm of iron or heavy metals. Trisodium phosphate and sodium formate were recrystallized from water containing EDTA and then twice more from pure water. Baker Analyzed Reagent grade sodium hydroxide was made into a 50% slurry in water, cooled in The Journal af Physical Chemistry, Val. 8 l , No. l l s 1977
an ice bath, and filtered through ultra-fine sintered glass, It was then diluted with purified water to a concentration of about 1.5 M and was then filtered through a "fluoropore" cartridge (Millipore Corporation) having a pore size of 0.2 pm. Water and all solutions were stored in carefully cleaned vessels of fused silica. Solutions Used. All irradiated solutions were saturated with air or oxygen and contained formic acid or sodium formate in concentration ranging from 0.01 to 0.1 M, together with EDTA to a concentration 0.01 of that of the formate. Presence of these solutes converts all the primary radicals, e-, H, and OH to the form HOz or 0;.394 A 100-fold excess of formate over EDTA is sufficient to protect the latter from attack by OH radicals. Procedures. Four different irradiation procedures were used depending upon the pH and the rates of radical decay to be measured. In every case the measurement consisted of determining the rate of decline of the optical density at 250-270 nm and 23 "C due to HOz or 0;. (a) Pulse Radiolysis. This was the standard pulse radiolysis method for H 0 2 detection described for example by Bielski and S c h w a r ~ . ~ (b) Stopped-Flow Radiolysis. This method (described by Bielski and Richter') uses a fast kinetics spectrophotometer (Durrham Instrument Co. Model D110) modified so that one of the flowing solutions passes through a 2-MeV electron beam produced by a Van de Graaff generator. This solution was always adjusted to pH 9.5 with trisodium phosphate. In the fast mixing chamber the solution was mixed with one containing acid or buffer to produce the desired pH, and then entered a spectrophotometer where, after the flow stopped, the decay of the radical was monitored. The optical cell in this setup has an internal diameter of only 2 mm, and it was found that for reactions lasting longer than 1min the kinetics became distorted by wall reactions which led to introduction of a first-order component superimposed on the second-order decay reaction. (c) Continuous-Flow Radiolysis. For systems in which the first half-life of the radicals was longer than 1 min, instead of stopping the flow the entire irradiated mixed solution was collected in a test tube and was then poured into an optical cell which was placed in a Cary 14 spectrophotometer, and the optical density was read out as a function of time on a recorder. The optical cells used in the spectrophotometer are 2 cm in diameter and no de-
1049
Disproportionation of Superoxide Radicals
i 400
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/
I
, , -240nm
1 i
4
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w,200nm L / / / I /
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-!
-
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PH Figure 1. Optical extinction coefficlentsfor alkaline aqueous hydrogen peroxide solutions.
tectable wall reaction occurs in these cells in the course of 2 h. (d) y Ray Radiolysis. When the decay of the radicals was very slow (at pH >lo) the solutions were irradiated with y rays which resulted in building up the concentration of radicals to a higher level than was obtained in the other methods. The solutions were allowed to stay in the y-ray source until an optical density of the order of unity due to the radicals had been built up. Depending on the concentration, cells of optical path lengths of 10,5, or 1.8 cm were used. These solutions contained 0.01 M HCOONa, lo-* M EDTA, and the pH was adjusted with NaOH. Extinction Coefficients. The extinction coefficients of HO, (t25 nm 800 M-' cm-'; t260nm 495 M-' cm-I; t270nm260 M-' cm-') and 0; ( t 50nm 1930 M-'cm-'; t260nm 1680 M-' cm-'; t270nm 1330 M- ? cm- ') were taken from the work of Behar et al.4 At any pH the ratio of 0; to HOz is determined by the equilibrium constant K = (H+)(O;)/ (HO,). We set K/(H+) = X so that log X = pH - pK. The ratio (O13 without finding any deviation from eq C. The effect of reaction 3 would be to increase points at high pH, with a hobad leveling off at a value equal to ka. No such effect is seen, and we can say definitely that k is less than 0.3 M-' s-'. Behar et a1.4 found k3 < 100 M-a s-l; a more recent article' claims k3 = 6 M-l s-' at pH >12. These results must have been affected by traces of impurity in the solutions used. Certainly the direct reaction of 0; with itself plays no role in the behavior of the radical ion in aqueous solution under any ordinary conditions, and 0; disappears only by reaction with the H 0 2which exists in equilibrium, or by reaction with catalysts. We suggest that the selfdisproportionation cannot occur at all, as shown by the reported stability of 0; in aprotic solvents. Thus, a 0.15 M solution of KO2 in dimethyl sulfoxide, stabilized by a crown ether; showedg less than 10% decomposition in 1 day at room temperature. Discussions" of the supposed self-reaction of 0; and the state of the O2 formed from it have little relevance to real situations. The weak acid catalysis shown by the rise in at pH values below 1 is real and not due to ionic strength effects.
References and Notes Research carried out at Brookhaven National Laboratory under contract with the US. Energy Research and Development Administration and supported by Its Division of Physical Research. B. H. J. Bielski and H. W. Richter, J . Am. Chem. Soc., submitted for publication. J. Rabani and S. 0.Nielsen, J . Phys. Chem., 73, 3736 (1969). D. Behar, G. Czapski, J. Rabanl, L. M. Dorfman, and H. A. Schwarz, J . Phys. Chem., 74, 3209 (1970). B. H. J. Belski and H. A. Schwarz, J. phys. Chem., 72, 3836 (1968). B. H. J. Blelski and A. 0.Allen, Inf. J . Radht. Phys. Chem., 1, 153 (1969). W. C. Schumb, C. N. Satterfield, and R. N. Wentworth, "Hydrogen Peroxide", Reinhold, New York, N.Y., 1955, p 287. J. Divikk and B. Kastening, J. Electroanal. Chem., 65, 603 (1975). J. S. Valentine and A. B. Curtis, J. Am. Chem.Sm., 97, 224 (1975). A. U. Khan, J . Phys. Chem., 80, 2219 (1976).
