DECAYKINETICSOF PERHYDROXYL RADICAL
2263
stein8 used the diffusion flame method to measure the rate of oxidation of ?\lg vapor. He reports a value of 4 X lo8 l./mol see for a bimolecular rate constant. O2 + RIgO 0 is 22 The endothermicity of hIg kcal/mol. This implies a second-order rate constant three orders of magnitude less than the observed value. A third-order rate constant calculated from Markstein's data is 5 X 10l2L2/mo12see. The two diffusion flame results are higher by about a factor of lo3 than those obtained by other methods. Possibly the rate constant for the Mg 0 2 reaction should not be used in this comparison in view of the experimental difficulties encountered in that work. One is left with the possibility that the rate constant for the Na-02 reaction measured by the diffusion flame method may be in error. The fast-flow method can be used to determine the rate constant for this reaction. The result would provide an independent check on the
+
+
+
diffusion flame method and the lean H2-02-N2 flame method. Such a check would be of general interest in that comparison between rate constants calculated from molecular beam cross-section values and classical chemical rate constants are now being made for alkali metal-alkyl halide reaction^.^^'^ Rate constants for such reactions have been measured only by the diffusion flame method. Evaluation by another method would put these comparisons on firmer ground.
Acknowledgment. The quartz reactors were made by RIr. R. F. Brennan and Mr. J. Wolford, whose careful work and constructive suggestions are gratefully acknowledged. (8) G. H. .Markstein, S y m p . (Int.) Combust. [Proc.], 9, 137 (1963). (9) J. P.Toennies, Ber. Bunsenges. P h y s . Chem., 72, 927 (1968). (10) K. R. Wilson and D. Herschbach, J . Chem. Phys., 49, 2676 (1968).
Deuterium Isotope Effect on the Decay Kinetics of Perhydroxyl Radical'
by Benon H. J. Bidski* Chemistry Departement, Brookhaven National Laboratory, Upton, ,Vew York
11978
and Eiichi Saito Centre #Etudes ~VuclCairesde Saclay, Departement de Recherches et Analyse, B.P. n o 2, 91-Gif-sur-Yvette, France (Received December $1, 1970) Publication costs assisted by Brookhaven iVational Laboratory
Rate constants for the disproportionation of HOz in HzO and DOz in DzO have been determined as a function of acidity and temperature. The radicals were generated by pulse radiolysis and by oxidation of peroxide with ceric sulfate. It was found that the isotope effect on the decay rates of the two radicals varies with pH. The pK values for HOZaudDOz were 4.8 and 4.9, respectively. Activation energies for the decay in 0.8 N HzS04 and 0.8 N D2S04 were 5.8 kcal/mol for H02 and 7.3 kcal/mol for D20.
Introduction Relatively high concentrations of the perhydroxyl radical can be generated in aqueous solutions by either high-energy ionizing radiations or by the oxidation of hydrogen peroxide by ceric sulfate. I n the field of radiation chemistry of aqueous solutions the perhydroxyl radical is of considerable interest, since it is probably the most important secondary product of the interaction between ionizing radiation and oxygenated water. Its formation can be described by the following equations. H20 -w+ 0 2
H, cas-, OH, HzOZ,H2
+ H +HOz
(1) (2)
+ eaQ--+ + H+ E HOz
(3) (4) Formation of HOZas an intermediate in the oxidation of hydrogen peroxide by ceric sulfate was originally postulated by Baer and Stein2 and later demonstrated by the studies of Saito and BielskL3r4 The following reactions account for the process. Ce(1V) H ~ O Z Ce(II1) H+ HOZ ( 5 ) 0 2
02-
+
02-
+
+
(1) Research performed under the auspices of the U. S. Atomic Energy Commission. (2) 9. Baer and G. Stein, J . Chem. SOC.,3176 (1953). (3) E. Saito and B. H. J. Bielski, J . A m e r . Chem. SOC.,83, 4467 (1961). (4) B. H. J. Bielski and E. Saito, J . P h y s . Chem., 66, 2266 (1962).
The Journal of Physical Chemistry, Vol. 76, N o . 16,1971
2264
BENONH. J. BIELSKIAND EIICHISAITO Ce(1V)
+ HOz -+ Ce(II1) + H f + O2
(6)
Bielski and Allen5 have shown more recently that, in the range below p H 1.0, the perhydroxyl radical generated by either method was kinetically indistinguishable. Despite this agreement it was not clear whether in the Ce(IV)-HzO2 system the H 0 2 was present as a free radical or complexed as Ce(II1)-OOH, as suggested by Anbar.6 Samuni and Czapski' solved this question recently by showing that HOz can exist either as a free radical or as a complex, depending upon the concentration of cerous ion in the system. Ce(II1)
+ HO2 eCe(II1)-OOH
(71
Their results suggest that Bielski and Allen studied the decay of the free form of the radical. I n the present study a comparison is made between the kinetics of HO2 in HzO and DO2 in DzO in order to establish the magnitude of the isotope effect on the decay rates of the light and heavy radicals.
For the esr studies, experimental equipment consisted of a Varian V-4502-13 electron spin resonance spectrometer, a 9-in. magnet with Fieldial, and a thermostated flow system. I n order to be able to carry out corrections for variations in the microwave power in the cavity over a period of several hours, a double cavity was employed. With this equipment the signal of interest is constantly monitored against a standard KC1-pitch sample and corrections can easily be made. The flow system consisted of a double-jet mixing chamber and a flow tube with an inner diameter of 0.86 mm. The flow tube itself was sealed into an 11-mm diameter quartz tube, and the space between them was evacuated. Tubes leading to the mixing chamber were jacketed by wider pipes through which the coolant was pumped from a thermostat bath by a circulating pump. Flow velocities were determined for each run. Temperature was monitored by thermometers embedded 15 ern above the mixing chamber. Average fluctuation during a run was about 0.2".
Experimental Section Chemicals. DzO, 99.7% in D , was further purified to ensure the removal of organic impurities. Two alternate methods of pretreatment were used. The D20 was mixed with oxygen in the vapor phase and passed through a silica tube at 800°, or it was preirradiated with 6OCo y rays. I n either case the heavy water was subsequently triple-distilled from acidic dichromate and alkaline permanganate solutions in an all-quartz distilling apparatus. D2S04was prepared by the addition of SO3 (Kuhlman, Paris) to DzO. DC104 was prepared by addition of Ba(C10& to D2S04. After removal of excess BaS04, pure De104 was obtained by distillation. D202and Hz02 solutions were prepared by the addition of Na202(Prolabo Product R. P.) to D2O or H2O and subsequent acidification. Concentrations of the peroxide solutions were determined by the ceric sulfate method. Ceric sulfate solutions mere prepared by dissolving the required amount of the anhydrous salt and checking the concentrations spectrophotometrically at 320 nm. The molar extinction coefficient in 0.8 N sulfuric acid was taken as 5580 M-l em-'. Apparatus. The apparatus for the pulse radiolysis experiments has been described in detail in an earlier p u b l i ~ a t i o n . ~The ~ 1.9-NeV electron pulses used in this study were of 2-20-msec duration with currents of 0.1 to 2 mA. The electron beam current striking the cell and cell holder was monitored and related to ferrous dosimetry. Total doses delivered to samples ranged from 0.3 to 2.5 krads. The corresponding perhydroxyl radical concentrations observed were of the order of 2 to8pM. The pH measurements were talien on an ORION Research Ionalyzer Model 801. The Journal of Physical Chemistry, Vol. 7 5 , N o . 15, 1071
Results Pulse Radiolysis Studies. Pulse radiolysis was used in the study of the decay rates in the pII range from 0 to 5.5. The change in concentration of the radical was followed spectrophotometrically at 241 nm. The molar extinction coefficients used for H 0 2 and 02at this wavelength were l l O O g a and 1950 M-1 cm-l,9b respectively. It was assumed that the values in D20 are the same. Since under pulse radiolytic conditions a considerable quantity of hydrogen sesquioxide or deuterium sesquioxide, which decays by first-order kinetics, also forms, the contributions of the two decay processes were resolved by a computer program which was used in an earlier study by Bielslii and S c h ~ a r z . ~ ~ The numerical results of the observed second-order rate constants are given in Figure 1. The acidity scale was normalized by the Glasoe-Long relationship, lo since pH measurements in D20 solutions are 0.4 pH units lower than in H 2 0 solutions of comparable acidity. The temperature dependence of the rate constants kHOzand k D O z mas determined in 0.8 N sulfuric acid solutions only. The corresponding activation energies, EHOz = 5.8 0.5 lical/mol and E m , = 7.3 f 0.4
*
( 5 ) B. H. J . Bielski and A. 0 . Allen, Proceedings of the Second Tihany Symposium on Radiation Chemistry, Akad. Kiodo, Budapest, 1967. (6) M. Anbar, J . A m e r . Chem. SOC.,83, 2031 (1961). (7) A. Samuni and G. Czapski, Israel J . Chem., 7, 375 (1969). (8) C . J. Hochanadel and J. A. Ghormley, J . Chem. Phys., 21, 880 (1953). (9) (a) B. H . J. Bielski and H. A. Schwarz, J . Phys. Chem., 72, 3836 (1968); (b) D. Behar, G. Czapski, L. M. Dorfman, J . Rabani, and H. A. Schwarz, ibid., 74, 3209 (1970). (10) P. K . Glasoe and F. A . Long, ibid., 64, 188 (1960).
2265
DECAYKINETICSOF PERHYDROXYL RADICAL
32
PH
Figure 1. Plot of log kabsd as a function of acidity for the disproportionation of HOZin HzO and DO1 in DzO (A,HOZ, pulse radiolysis data; A, HOZ,esr data; 0, DOz, pulse radiolysis data; 0, DO,, esr data).
-
* r
-
3.2
3.3
3.4 (IIT)
3.5 x
3.6
3.7
3.8
lo3
Figure 2. Plot of log kabad for the disproportionation of HOZ in HzO and DOz in DzO as a function of reciprocal temperature (A,HO,, pulse radiolysis data; A, HOz, esr data; 0 , DO,, pulse radiolysis data; 0, DO,, esr data).
kcal/mol, were calculated from the curves shown in Figure 2 . Electyon Spin Resonance Studies. The generation of perhydroxyl radicals in the Ce(IV)-H202 system is limited to the strong acid region, since above pH 1.5 precipitation of the cations as hydroxides interferes with the experiment. For this reason, the esr experiments were carried out only in 0.8 N sulfuric acid. The HO, and DO2 radicals generated in a mixing chamber were transported to the spectrometer cavity by a flow apparatus. Flow velocities were about 380 cm/sec for HzO and 210 cm/sec for D,O solutions at 23". The concentration of the Ce(1V) was 1.5 X M , while the HzOz and DzOz were always present in excess at a concentration of 0.3 M . Under these conditions the steady-state concentration of the free radical
O C
Figure 3. Change in line width for HOz in 0.8 A- HzS04and for DO2 in 0.8 N D2SOawith temperature at 9.5 kMc.
was measured as a function of flow time from the mixing chamber to the center of the esr cavity. Calibrations were carried out after each run by replacing the ceric peroxide solutions by AIn(I1) sulfate solutions (in HzO and DzO) in order to determine the absolute number of spins and hence the concentration of the free radicals. The computation of the number of spins is based on the correspondence of the integrated surface areas of the spectra of Mn2+,HOz, and DOZat a given temperature. These calibrations had to be carried out at each temperature, as seen in Figure 3. The g values for HOz and DOz radicals are 2.016 and 2.017, respectively. These values were computed wlth reference to the g = 2.0036 for @,a-diphenyl-P-picrylhydroxylfree radical. Plots of reciprocals of concentrations of H02 and DOn against time gave in all cases straight lines, indicating that second-order kinetics were obeyed. The corresponding rate constants for the radical decay are average values from five experiments. The numerical values are given in Figure 1 as a function of pH and in Figure 2 as a function of temperature.
Discussion It has been shown by several research teams that the perhydroxyl free radical disproportionates by secondorder kinetics to oxygen and p e r o ~ i d e . ~ ~ > The "-~~ (11) J. H. Baxendale, Radiat. Res., 17, 312 (1962). (12) G. Czapski and B. H . J. Bielski, J. P h y s . Chem., 6 7 , 2180 (1963). (13) G. Czapski and L. M.Dorfman, (bid., 68, 1169 (1964). (14) 2. P. Zagorski and K. Schested, Riso Report No. 114 (1965). (16) G. E. Adams, J. W. Boa& and B. D. Michael, P T O CRou. . SOC., Ser. A , 289, 321 (1966). (16) J. Rabani and 9 . 0. Nielsen, J. P h y s . Chem., 73, 3736 (1969). (17) D. Behar and G. Czapski, Israel J . Chem., in press.
The Journal of Physical Chemistry, Vol. 75, No. 16, 1971
BENONH. J. BIELSKIAND EIICHISAITO
2266 overall decay scheme, described by equations
+ HOz + 02-5Hz02+ O2 + OHOa- + Oz-2 H20a + O2 + 20HHOz
+ H02 -%-
H202
0 2
(8) (9)
(10)
postulates a pH dependence based on the existence of two forms of the radical KHO~
HOz
e02- + H+
(11)
where small k's are the various rate constants and ~ H O is , the dissociation constant of HOz. Contrary to some earlier report^,^^^^^^^^ the rate constant for reaction 10 is now believed to be many orders of magnitude smaller (kla < lo2 M-' sec-l). Hence, neglecting reaction 10 for the pH range investigated in this study, the decay mechanism can be described by the following equation. ks
kbsd =
[1
+
+]I2
+ 1
The solid curves in Figure 1 were calculated from this equation using the following rate constants, For HOz
ks = 7.0 X lo5 1W-l sec-I k 9 = 8.50 X lo7 M-I see-l
KHoa= 1.6 X
9b
or ~ K H =o 4.8 ~
For DO2
ks = 1.0 X lo5 M-l sec-l k s = 2.06 X 107M-1sec-1 or ~ K D o=, 4.9 K D O=~ 1.26 X These values are in close agreement with figures given recently in two independent report^.^^"^
The Journal of Physical Chemistry, Vol. 76, No. 16, 1971
3t1
20
1
2
3
4
5
6
PH
Figure 4. Plot of kHol/knop as a function of acidity.
Deuterium isotope effect upon the overall decay of the perhydroxyl radical is illustrated in Figure 4, where the solid line represents the ratio of [kHOz/kDOz]obsd computed from the rate constants and e q I . The experimental points are the ratios of the data shown in Figure 1. As expected, the isotope effect is most pronounced in the strong acid region ( ~ H O ~ / ~ Dabout O ~ 7.0),since the decay then takes place almost completely via reaction 8. With a shift in p H toward the neutral region this ratio decreases t o a value of 3.5 at pH -5.0. A smaller isotope effect in the less acid region could be explained by the fact that reaction 9, which is rate controlling, requires the breaking of one less bond than reaction 8. The difference in the activation energies of 1.5 kcal/ mol for the decay of HOz and DO2 in 0.8 N HzS04 and 0.8 N D2SO4(Figure 2) is of reasonable magnitude. Acknowledgment. We wish to thank Drs. J. Gebicki, H. Schwarz, and G. Czapslii for stimulating discussions
and constructive criticism. Thanks are also due t o Mr, G. Pereto and Mr, D. Comstocli for their excellent technical assistance. (18) K. Schmidt, 2. Naturforsch. B , 16, 206 (1961)
