FLASH PHOTOLYSIS STUDY OF
THE
2361
MECHANISM OF OZONIDE IONDECAY
Flash Photolysis Study of the Mechanism of Ozonide Ion Decay in Basic Aqueous Hydrogen Peroxidel
by Vincent Russell Landi and Lawrence J. Heidt Department of Chemistry, Massachusetts Institute of Technology, Cambridge, Massachusetts (ReceQed December 12, 1968)
02139
A study has been made at 27 i 2' of the rate of decay of the 4300-A absorption peak of the ozonide ion produced by the flash photolysis of hydrogen peroxide in aqueous alkaline solutions. The observed decay in every case followed the course of a first-order reaction. The rate constants increased with increase in Hi02 concentration and decreased with increase in the concentration of base and dissolved molecular oxygen. A longer lived species was formed in part at the same rate as the ozonide decayed. Spectral and kinetic evidence kl suggests that this longer lived species is ozone. The main reactions explaining the results are 08- 02 0-, kz HOH 02 0- 0 3 - , 0HOZ- 02- OH-, and 0 3 02- +0 3 HOZOH- for which kl = 2.81 f 0.14 x lo3 sec-1 and k& = 3.6 (h0.3). The marked increase in stability of Os- at high (NaOH) can be explained on the basis of the formation of the ion pair: Na+ 03- = Na03, the equilibrium constant for which is 2.24 i 0.11 M - l .
+
--.f
+
2
+
+
+
+
+
+
Introduction The absorption spectrum of the ozonide ion, Os-, has been observed2-7 in pulse radiolysis and flash photolysis studies of certain basic aqueous systems. The ion is believed to be formed by the reaction 0-
+
0 2
+0 3 -
where 0-, the basic form of hydroxyl radica1,O OH, is only important at pH greater than 9. I n addition there appears to be formed under the same conditions as the ozonide ion another species which absorbs light in a different region of the spectrum and decays at a much slower rate than the 0 3 - . Presented here is a detailed flash photolysis study of these transients.2
Experimental Section The flash photolysis equipment has been described in its main essentials elsewhereUa Four cylindrical flash lamps arranged coaxially around a reaction vessel were filled with 30 mm of oxygen and operated around 14 kV. A capacitance of 36 pF was employed and yielded l / e flash times of 15 psec. The cylindrical quartz reaction vessel had an inner diameter of 1.25 cm and 20 cm length; a concentric jacket of 0.5-cm annulus held the light-filtering solutions. Water-free carbon tetrachloride was used as the light filter to eliminate light shorter than 2600 to provide more homogeneous photolytic light absorption at the largest HzOz concentrations. A 450-W Osram xenon arc lamp powered by batteries was used as the spectrophotometric light source. Changes in optical density with time were analyzed at selected wavelengths with a Zeiss MM 12 double monochromator, EM1 6256 B photomultiplier tube, and Tektronix 545A oscilloscope with Type D preamplifier. The oscilloscope trace was 8 9
triggered with a second phototube responding to light from the flash. Highly pure 50% HzOzwithout keeper was a gift from the Becco Co. Rates of the thermal decomposition of the basic hydrogen peroxide stock solutions used in the photolysis studies were from 0.5 to 2%/hr as monitored by changes in optical density. For most experiments, the high-purity 50% sodium hydroxide solutions were tapped directly from the electrolytic cell in which they were prepared to minimize the carbonate content; the solutions were kindly donated by the Hooker Chemical Corp. Water was distilled and redistilled over alkaline permanganate. Other chemicals were reagent grade. In experiments at 1 atm 02, oil-free oxygen was bubbled through the solutions for an hour. This was done in a glove bag which had been filled and flushed twice with oxygen. A suitably calibrated Cary 11 spectrophotometer wms used for monitoring suitably slow changes in optical density. (1) This is publication No. 81 of the MAT. Cabot Solar Energy Fund. (2) Submitted to the Graduate School of the Massachusetts Institute of Technology in partial fulfillment of the requirements for the degree of Doctor of Philosophy, 1965, for V. R. Landi. (3) L. J. Heidt and V. R. Landi, J . Chem. Phys., 41, 176 (1964); L. J. Heidt, J . Chem. Educ., 43, 623 (1966), and J . Solar Energy, 11, 123 (1967). (4) (a) G. Caapski and L. M. Dorfman, J . P h m Chem., 68, 1169 (1964): (b) G. Caapski, ibid., 71, 1683 (1967). (5) W. D. Felix, B. L. Gall, and L. M. Dorfman, ibid., 71, 384 (1967). (6) G. E. Adams, J. W. Boag, and B. D. Michael, Proc. Roy. Soc., Ser. A , 289, 321 (1966). (7) (a) E. Hayon and J. J. McGarvey, J . Phys. Chem., 71, 1472 (1967); (b) L. Dogliotti and F. Hayon, ibid., 71, 2511 (1967). (8) J. Rabani and M. S. Matheson, J . Amer. Chem. Soc., 86, 3176 (1964). (9) L. Lindqvist, Rev. Sei. Instr., 35, 993 (1964). Volume 73, Number 7
July 1969
VINCENT RUSSELL LANDIAND LAWRENCE J. HEIDT
2362
Results The flash photolysis of the basic hydrogen peroxide solutions produced two observable transient absorption spectra: the one peaking at 4300 +& was due to ozonide and decayed much faster than the longer lived species that peaked around 2600 A. The ozonide decay was investigated mainly at 4300 A in M to 10 M NaOH and to M HzOz. I n every experiment the decay followed the first-order rate law (Figure 1) at rate constants koa- (Table I), which varied with changes in the initial solution composition. Qualitatively, koa- increased with increase in hydrogen peroxide and decreased with increase in oyxgen (Table IV) and sodium hydroxide concentrations.
-1.0
-1.2
'1.4 0
0
-v-1.6 -1.8
TableoI: Typical First-Order Decay Rates at 4300 and 2800 A in Air-Saturated Solutions at 27 + 1.5" -2.0
1.28
12.8 128 100 1.36 13.6 136 2.07 20.7 207 1.66 16.6 166 3.99 19.9 199 3.19 12.8
9.5 *
.
I
...
7 4.75 .
I
.
... 0.95 ,
I
.
... 0.0475
... .
.
I
0.0095
...
... pH
=
10,8
128
2.17 6.46 10.29 63.8 48.7 140 358 80.8 771 2430 166 974 3640 656 1543 4360 1186 2392 4725
2
I
0.00092
Figure 1. First-order decay of ozonide absorption at 4300 in flash-photolyzed 5 x M HzOz,1.9 M NaOH air-saturated solution; optical density, O.D., that of 20-cm path length.
0.028
2.03 4.10
o'z 0.1
8
k
0.05
W
r\l u
6.29
n 0
0.02
At wavelengths progressively shorter than 3000 light absorption measurements became increasingly difficult due to decreasing response of the spectrophotometric system and to increasing light absorption of the hydrogen peroxide ion, H02- (the predominant form of HzOz in most of our work'"). Measures taken to increase sensitivity also increased the interference of stray light from the flash. Nevertheless good measurements of the long-lived absorption could be made at 2800 A. The determination of the absorption spectrum of the peak region, however, was difficult to obtain with the flash unit. The decay at 2800 R is presented in Figure 2. The decrease in rate with increase in the sodium hydroxide concentration is presented in Table I. Figure 3 shows this spectrum produced by photolysis of hydrogen peroxide in 7 M sodium hydroxide. A correction was made for absorption due to the HOz-. The JOUTnal of Physical Chemistry
4
3 mSec
\ \\ '
0.0 I
0
,
I
I
/
I
I
l
1
2
3
4
5
6
7
T I M E ( u n i t s bc'ow)
Figure 2. First-order decay of optical density of 20-cm path length at 2800 in flash-photolyzed basic hydrogen peroxide air-saturated solutions: 1, 0,05 M NaOH, 1.7 X lo-' M HaOa, 1 time unit = 0.2 sec; 2, 0.01 M NaOH, 6.7 X M H102, 1 time unit = 0.1 sec.
The buildup of the 2800-A absorption could be observed to take place in part well after the effects from stray light had ceased. The time dependence of the increased absorption at 2800 A corresponded to the first-order decay of the Os- species. Table I1 shows (10) J. Jortner and G. Stein, Bull. Res. Coicncil Israel, Sect. A , 6, 239 (1957).
FLASH PHOTOLYSIS STUDY OF
THE
MECHANISM OF OZONIDE IONDECAY
MB
-
200
-
100
-
2363
y s -
b
I 2
1
I
5
I
I
I
10 20 ( ~ 2 ~ x2 )105 ( m o l e r h i t ~ r )
% .
1
100
.
1
200
Figure 4. Points are psuedo-first-order rate constants of ozonide decay, koa -, as a function of hydrogen peroxide concentration in air-saturated 1.9 M NaOH solution at 27 f 1.5". Curve calculated from expression I. I
1
2133
1
l
2600
l
l
I
2603
I
30JO
I
l
3%05
Table I11 : Effect of Sodium Hydroxide Concentration on
h (1)
Parameters in Rate Constant Equation for Ozonide Decay
Figure 3. Uv spectra of: 1, product of photolysis of hydrogen peroxide in 7 M NaOH air-saturated solution; 2, the same 7 A4 NaOH (without hydrogen peroxide) through which ozone had been passed.
that the first-order decay constants calculated for the precursor to the 2800-A absorbing species correspond closely to the decay rates found for ozonide under the same conditions.
9.5 4.75 1.9 0.95 0.475 0.1 0.0475 0.0095 (pH10.8)
a X 10-8, d X 104, b/a X 108,
(HzOz)
AY
range,a
dev,
800-1
M
M
M x io6
%
0.0120 0.449 1.75 3.093 3.73 5.07 5.09 5.46 5.26
1.36 4.01 4.63 6.29 7.60 8.97 7.47 6.03 2.22
14.1 32.1 8.2 5.4 0 12.4 9.7 39.4 22.9
1.3-130(3) 1.4-140(5) 1.4-140(6) 20-500(8) 20-200(5) 0.81-800(7) 1 . 7 X 170(5) 4.0-500(6) 3.2-128(5)
1.8 1.6 2.3 1.1 3.0 4.7 4.7 4.6
I n parentheses are the number of different concentrations over indicated range of (HtOz).
Table 11: First-Order Rates of Decay of Ozonide and Precursor to 2800
(NaOH), M
Absorption in Air-Saturated Solutions
(NaOH), M x 101
VM" M X 106
k(4300 A), seo -1
0.95 0.95 47.5
16.6 41.4 5.06
1100 3880 234
k(pre2800 A),5 sec-1
1070 3730 258
Calculated from rate of increase of light absorption a t 2800 h;.
small but was necessary t o fit the data at low hydrogen peroxide concentrations. The effect of increasing oxygen concentration (Table IV) was to decrease the rate of ozonide decay and increase the amount of ozonide formed by the flash. The relative changes in these quantities depended upon the concentration of hydrogen peroxide. Our results can be explained by
osThe dependence of ozonide decay rate on hydrogen peroxide concentration, holding other known variables constant, was found to fit the equation
00-
-0 2
+0 2
+ 0-
(1)
Oa-
(2)
+ Hot- + + HO02-
(3)
Copious evidence exists for reactions 2 and 3,4-7111-18and Figure 4 shows the fit of eq I to the experimental values. The parameters a, b, and d were obtained by a procedure which minimized squares of the fractional deviations of the experimental values from the values predicted by eq I. Values of the parameters a, b, and d at different NaOH concentrations are summarized in Table 111. Column 5 gives the number of concentrations of Hz02 employed (regularly spaced in a logarithmic sense) and the range in H202 concentration covered. Column 6 shows the average % deviation of the experimental values from expression I. Parameter b was relatively
(11) E. J. Hart, 8. Gordon, and D. A. Hutchinson, J . Amer. Chem. SOC., 7 5 , 6165 (1953); 74, 5548 (1952). (12) 0.L.Forchheimer and H. Taube, ibid., 76, 2099 (1954). (13) C. R. Giuliano, N. Schwartz, and W. K. Wilmarth, J. Phys. Chem., 63, 346 (1959). (14) A. D. McLachlan, M. C. R. Symons, and M. G. Townsend, J . Chem. Soc., 952 (1959). (15) K. Schmidt, Z.Naturforsch, B , 16, 206 (1961). (16) E. Saito and B. H. J. Bielski, J . Amer. Chem. SOC.,83, 4467 (1961). (17) G. Czapski and B. H. J. Bielski, J. Phys. Chem., 67, 2180 (1963). (18) H.A. Schwartz, ibid., 66, 255 (1962). Volume 73, Number 7 July 1969
VINCENTRUSSELL LANDIAND LAWRENCE J . HEIDT
2364 the dependence of yields of 0 8 - on oxygen and hydrogen peroxide concentrations (Table IV) is consistent wit’h competition by (2) and (3) for 0- produced by the photolysis of HOz-.
tions Czapski and Dorfman observed a species having a of similar lifetime to the longpeak of 24004or 2500 lived species observed by us. I n their observations of the changes in absorption at 2537 A they noted a slowdown in the decay of transient absorption which occurred above pH 9, a t about the time 03- begins to appear.” In neutral solutions they observed a fast initial decay (presumably due to OH 023 which occurred before reaction 7 predominated. Above p H 10 there was no fast initial decay, but instead they observed a fast initial increase in absorption at 2537 8 followed by the slow decay. Adams, Boag, and Michaels also studied the species produced in oxygenated basic solutions by pulse radiolysis and found an absorption that peaked at 2600 A. When NzO (a scavenger for eaq- which diminishes the formation of Oz- caused by the pulse) was added, the 2600-f absorption exhibited a rapid decay followed by a very slow decay. The rate of the rapid decay was about that of the decay of Os- found in the same solutions. Among the conclusions of their work was that “08- reacts with 02- and the product probably absorbs light at 2600 8.jJ If 0 2 - is consumed by reaction 4
A5
+
Table IV ODa
koa-,
(HzOz),
ko8-,
ODa
(4300)d
(2800)’
om-’ X
om-’ X
pol,&
obsd,*
M X 106
atm
see - 1
see-1
loa
10’
2.70 2.70 10.8 10.8 28.6 28.6 108 108
0.2 1 0.2 1
169 35.3 431 125 1086 280 1880 781
153 31,6 476 108 985 263 11960 795
1.71 1.82 5.45 5.68 10.7 16.3 16.1 37.2
0.58 0.65 2.28 1.95 6.60 6.60 22.8 23.2
0.2
1 0.2 1
a po2 at 0.2 atm refers to air-saturated solutions. At 27 It 0.5’; NaOH = 0.95 M . See text. Optical density (OD) at 4300 per flash; at zero time after flash by extrapolation. Maximum OD at 2800 d (after buildup, before decay). L&
A prior indication of the existence of 08- in aqueous solution has been the isotopic exchange of oxygen from water with dissolved molecular oxygen in systems in which hydroxyl radical are p r e ~ e n t . l l - ~The ~ chain length of the exchange was found to increase dramatically above p H 8 from the small constant value when the pH was less than 9. It was deduced that the exchange reaction 0Oz* -+ 0-* O2 must go by way of Os- formation, in the light of the evidence for reaction 2. The high values of the chain length suggested the dissolution of 0 3 - into oxygen and 0- by reaction 1. The chain length was found to be greatly reduced by the addition of hydrogen peroxide, l 1 as would be expected by reactions 1, 2, and 3. Inclusion of the reaction 0803- = products6j6 was not consistent with our results, since no sign of second-order decay was found even at high 0 3 - concentrations. The apparent second-order component of 03- decay reported in the absence of hydrogen peroxide could well be accounted for by reactions of 0- with Os- or with itself. For systems in which an OH scavenger such as HOn- is not present, any concentration of Os- should be associated with a higher concentration of 0- than in the present work. The 02-produced by reactions 1, 2, and 3 might be expected to react by (7). *(The numbering of this
+
+
08-
+ 02- HOH 03 + HOz- + OH-
(4)
production of ozone would account for a long-lived species having a similar spectrum to that found. Conditions for reaction 4 to compete successfully with (7) are discussed below. Since an asymptotic limit for ozonide decay with increase in hydrogen peroxide has been found in our work, reaction 1 must be relatively slow and rate controlling; hence, steady-state approximations for (0-) and (02-) may be made.
+
+ 02-
02-
H02-
+ O2 + OH-
(7)
reaction is out of order but it seems less awkward than to give it the number 4 and spoil the sequence of reactions 1 through 6 as given later in this article). I n pulse radiolysis studies of oxygenated basic soluThe Journal of Physical Chemistry
By this approximation one calculates a first-order decay of ozonide whose rate constant is
Carbonate ion was an uncontrolled impurity in our experiments, since the distilled water used in the experiments was stored under air. Carbonate ion is known6to react with 0-.
0-
+COP
Cos-
+ 20H-
(5)
The spectrum of COa- was not found in our hydrogen peroxide system, but it was found by us using similar methods and materials2 in flash-photolyzed basic solutions of S2082-, another system in which 0- (and
FLASH PHOTOLYSIS STUDYOF
THE
MECHANISM OF OZONIDE IONDECAY
2l
Os-) is p r o d u ~ e d . ~ Reaction ~ ~ ~ ’ ~ 6, therefore, seems possible. COS-
+ HOz-
C032-
+ 02-+ HOH
2365
(6)
/
20
Inclusion of reactions 5 and 6 and a steady-state approximation for (C03-) in the above scheme leads to
16
3 12
7.
H
P ” Except for a factor of 2 in the numerator exactly the same expression is obtained by replacing reaction 4 by reaction 7 without making the steady-state approximation for 02-. Comparing expressions I and I11
e
4
a = 2kl 0
I
I
I
I
I
1
2
3
4
5
(Na+)(?&)a
(in moles/liter)
Figure 5. Test of expression I V of text; data at 27 =t1.5’ in air-saturated solutions.
k3
of ozonide decay. For example when a solution 0.1 M in NaOH and 7,5 X 10-4 M in H202was also 0.004, 0.02, and 0.1 M in Na2COs, the value of koS-increased from 2230 to 2690, 3220, and 3890 sec-l, respectively. The limited amount of data in which the carbonate concentration varied appears to be fit by expression 111,and k5/k8 = 6 X The ratio kz(O$/k3 is affected in several ways by the large changes in the p H and ionic strengths used in this work, namely: (1) At higher ionic strengths dissolved molecular oxygen is a t lower concentrations due to “salting out,” thereby decreasing parameter d. Calculations were carried out using the empirical relationship log (02)o/(Oz)= P(Na0H) where ( 0 ~ is )the~zero ionic strength concentration and P is a constant which depends on the ionic species.2o I n the absence of a value of P for NaOH and 02,that of P for NaCl and O2 was taken. The values of kZ/k3, so corrected, plateaued from 10-l to 2 M NaOH (Table V) where ]c2/k3 was Na+ 03- = (Na+03-) calculated to be 3.6 f 0.3. (2) At p H less than 12 the species OH and Ha02 (pK of HzOz = pK of HO = 12) become larger a t the expense of 0- and Hot- and so (Values of yi for high (NaOH) are pub1i~hed.l~) affect rates of reaction along paths 2,3, and 5 and this is ) the probable reason for the decrease in d ( 0 2 ) o / ( 0 2at Thus we add reaction 1‘ to reaction 1. 0.01 M (Table V). (3) At p H much (NaOH) less than (Na+Oa-) -+ negligible decomposition of 03- (1’) greater than 12 the salt effect on reaction rates of and parameter a of expression I becomes charged speciesz1 becomes increasingly important because of increasingly greater ionic strengths (5 and 10 M NaOH). which approximates kz/k3(0z) a t low c03’-. A test of the predicted oxygen dependence was made by calculating ICos- for solutions saturated with 1 atm of oxygen using a, b, and d obtained in air-saturated solutions and assuming the parameter d to be proportional to (02). Table IV shows good agreement between calculated and observed rates. The observed values of koa- a t 0.2 atm of oxygen shown in Table IV were obtained by the same independent experiments used t o predict the results when the pressure of oxygen gas was 1 atm. The decrease in the rate of ozonide decay as sodium hydroxide concentration is increased primarily results from a decrease in the parameter a of expression I (Table 111). The effect can be explained by assuming the formation of stable ion pairs between the ozonide ion and sodium ion.
+
Figure 5 shows the linear fit of l/a to (ri)2(Na+) as predicted. The new value calculated for kl is (2.81 i 0.14) X lo3 sec-’ and K1 is 2.24 f 0.11 M-l. The effect of sodium carbonate is to increase the rate
(19) R. A. Robinson and R. H. Stokes, “Electrolyte Solutions,” 2nd ed, Butterworths, London, 1959, pp 492, 504. (20) T. Morrison, J . Chem. Soc., 3814 (1962). (21) G . Scatchard, Chem. Rev., 10, 229 (1932). Volume 73, Number 7 July 1969
VINCENTRUSSELLLANDIAND LAWRENCE J. HEXDT
2366 ~
~~~~~~~
~~
Table V : Parameter d Corrected for Salting Out of Oxygen (NaOH), M
9.5 4.75 1.9 0.95 0 475 I
0.10 0.0475 0.0095 (PH 1 0 . 8 )
d(Oz)o/(Oz), M
29.1 18.5 8.52 8.55 8.86 9.25 7.58 6.05 2.22
(Yi)
a.4
2.84
1.03 0.705 0.678 0 690 0.766 0.798 0.899 0.964
The assumption that reaction 4 excludes reaction 7 requires evaluation of the ratio, R = k4(0a-)/k7(02-), which is equivalent to the amount of reaction taking place by (4) relative to (7). By the steady-state approximations previously made
0.3
I
i.
I
0
20
I
I
40
I
66
I
100
(~~02 x )105
I
120
(molss/utor)
Figure 6. Dependence of maximum in 2800-b absorption on hydrogen peroxide concentration at: 1, X, pH 10.8; 2, 0, 1.9 M NaOH; 3, a, 4.75 M NaOH. Optical density, O.D., for 20-cm path. length through air-saturated solution.
I n a typical experiment (0.95 M NaOH, 2.07 X 10-3 M H202) koa- = 2430 sec-I, and the starting optical density of 0 3 - at 4300 8 was 0.023 per centimeter of the solution. At one-tenth this initial value of O.D., first-order kinetics were still found. Assuming” the molar extinction coefficient of 03- to be 2000 M-1 cm-I at 4300 A, the concentration of Os- was roughly 10+ M . Then, since‘l k7 = 1,5 X lo7M-l sec-l, R = k4Z X 0.5 X 10-lB at M 03-.If IC4 = lo* M-1 sec-I, R = 50 or only 2% of the reaction goes by path 7. Since diffusion-controlled second-order rate constants are greater than 1Olo M-’sec-I, a magnitude of IC4 > 109 M-1 sec-l is feasible.22 The intermediate absorption at 2800 can be treated as a product of reactions taking place during the ozonide decay, since it builds up during the 08- decay and it decays a t a negligible rat: on this time scale. For this reason the yield at 2800 A per flash can be taken as the maximum in the light absorption-time curve. This yield was found to be directly proportional to hydrogen peroxide concentration (Figure 6 ) and independent of O2 concentration (Table IV), that is, it appeared to have a constant quantum yield. The yield of OS(Table IV) is a more complicated function of H202and increases with oxygen concentration. The over-all quantum yield of hydrogen peroxide photolysis has been reported by a number of authors to be independent of pH, H202 concentration, and light intensity28*24 H202 h’, H2O
+ 0.502
Through a scavenger technique, Baxendale and Wilson showed that the primary photolytic scission of hydrogen peroxide was one-half the over-all quantum yield; hence, also independent of (H2Oz).z4 These data have The Journal of Physical Chemistry
suggested a photolytic scission of H202into hydroxyl free radicals followed by reactions 3 and 7
H202 h’,2H0 The exchange of oxygen in photolyzed HzOzsolutions, however, cannot be explained by this series of reactions.2SJe An alternate suggestion stemming from the isotopic exchange work that atomic oxygen is produced has also been disputedm27 The presence of 08- shows that free-radical pairs are produced in the photolysis. I n the present work, the constant quantum yield of the long-lived product, despite varying yields of Os-, suggests that it is the ultimate trap for the free-radical pairs. If 03-is produced in excess of 02-by the very fast steps following the light absorption, reactions 1 through 6 predict the Os produced is equal to the HZOZdestroyed. Then, if, as has been r e p ~ r t e d , ~ ~ ozone -~O reacts with an excess of hydrogen peroxide by the stoichiometry 08
+ HSO2
--t
202
.+ Ha0
(8)
(22) M. S. Matheson, “Solvated Eleotron,” Advances in Chemistry Series, No. 50,American Chemical Society, Washington, D.C., 1965, p 61. (23) M. C. R. Symons, “Peroxide Reaction Mechaniams,” J. 0. Edwards, Ed., Interscience, New York, N.Y.,1962,p 137. (24) (a) J. H.Baxendale and J. A. Wilson, Trans. Faraday Soc., 5 3 , 344 (1957); (b) L. J. Heidt, J . Amer. Chem. Soc., 54, 2840 (1932). (26) M.Anbar, Trans. Faraday Soc., 57,971 (1961). (26) J. P.Hunt and H. Taube, J . Amer. Chem. Soc., 74,6999 (1962). (27) G. Buxton and W. K. Wilmarth, J . Phys. Chem., 67, 2836 (1963). (28) H. Taube and W. C. Bray, J . Amer. Chem. Soc., 62, 3357 (1940). (29) W. C. Bray, ibid., 60, 82 (1938). (30) V. Rothmund and A. Burgsteller, Monatah., 38, 295 (1917).
FLASH PHOTOLYSIS STUDYOF THE MECHANISM OF OZONIDE ION DECAY an over-all quantum yield twice that of the photolytic scission is predicted in accord with Baxendale and Wilson's rcsults. If 02-is produced in excess of Oa', this scheme breaks down, since it predicts that then no 03-would have existed long enough to be observed by our technique. A conceivable mechanism for the photolysis and subsequent reactions which take place early in our solutions involves an excited HOZ- ion. HOz- h', HOz-*
+ O2 % 03-+ 0- + HOH HO2-* + H02+ eaqeaq- + HO2- +HO- + 0eaq- + 02
HOz-*
--f
03-
-
02-
(9)
(10) (11) (12) (13)
(also reactions 2, 3, and 4). This scheme predicts that 03-produced in initial steps is always in excess to 0 2 which is consistent with the Os- yield data found in this worka2 Studies of the spectrum of ozonized alkaline solutions have shown3 an increase in the stability of ozone as basicity is increased. This is consistent with the identification of ozone as the species responsible for the long-lived 2800-.& absorbing product of the present study which is also stabilized as the basicity is increased. Reactions 1-4 and the slower decay of 0 3 at higher NaOH concentrations may well be connected with this stabilization. Ozone is believed to decay in water by a chain mechanism.12~28 Adapting this
2367
mechanism to alkaline solutions in which OH and HOB are chain carriers
+ 008 + 020- + 0Oa
--f
+ 202 + 0 -
0 2
02-
4
02-
I
+ 02- chain ending
Through reaction 2,O- is converted into the less reactive species 0 3 - , which should accumulate to a much higher concentration than 0-. By reaction 4, which should have a high rate constant, two chain carriers are consumed. This would slow the decomposition of ozone. One could speculate that 0 d 2 - (from reaction 7) or Ob2- (from reaction 4) is the long-lived species, but it is difficult to explain the pH dependence of the life times of such products. Also, scatter in the half-life data, which we2 and others4 have found, is not to be expected for a simple first-order decay of these species into permanent products.
Acknowledgments. The authors wish to give special thanks to Dr. Lars Lindquist and Dr. Gunnar Wettermark for their recommendations concerning the design and use of the flash photolysis equipment. Thanks are due to Professor Nichols Milas for letting us use his ozonizing equipment. Grateful acknowledgment is given to the National Science Foundation, the Corning Glass Foundation, the National Institutes of Health, and the Godfrey Lowell Cabot Solar Energy Fund of the Massachusetts Institute of Technology for financial assistance.
Volume 79, Number 7
July 1969
