J . Ph)s. Chem. 1984,88, 31 11-31 15
3111
Pulse Radiolysis Study of the Kinetics and Mechanisms of the Reactions between Manganese(I I ) Complexes and H02/02- Radicals. 1. Sulfate, Formate, and Pyrophosphate Complexes Diane E. Cabelli and Benon H. J. Bielski* Chemistry Department, Brookhaven National Laboratory, Upton, New York 1 1 973 (Received: November 1 , 1983)
The reactions between Mn2+and H02/02-were studied in the presence of pyrophosphate, sulfate,or formate ligands. Manganous pyrophosphate is oxidized to manganic pyrophosphate by both HOz and Oz- at pH-dependent rates where k, varies from 3 X lo5 to 2 X lo7 M-I s-’ in the pH range of 0.1-7.2. The sulfate and formate complexes both react to yield the transient species MnOz+/MnOOHZ+;MnO2+ eventually disappears by complex processes while MnOOHZt reacts with additional Mn2+to form a dinuclear species. M I I O O H M ~and/or ~ ~ . oxidizes to Mn3+. Reaction kinetics have been established from pH 1 to 7 for these two systems.
Introduction Despite numerous investigations of the interactions between manganous cations and superoxide/perhydroxyl radicals, a question still existed as to whether Mn2+reacts with 0, to produce a transient metal-oxy complex, MnOZ+,l4 whether oxidation to Mn3+ occurs d i r e ~ t l y , ~or- ~ whether MnO,+ is a transient in the oxidation of Mn2+to Mn3+.9 Preliminary studies in this laboratory1° indicated that all pathways are in fact correct and that the pathway is dependent upon the particular ligand associated with Mn2+ in a given system. This paper describes a detailed study, using the pulse radiolysis technique, of the MnZ+-HOz/O,- system in the presence of sufficient pyrophosphate, formate, or sulfate to form the respective metal complexes in aqueous solution. While previous studies were primarily carried out in a pH range between 4 and 7, studies over a much broader pH range are necessary for a complete understanding of the mechanism in these systems because Oz-is always in equilibrium with its conjugate acid
+
H 0 2 + 02- H+
pKl = 4.711
(1,-1)
Experimental Section Materials. All solutions were prepared by using water which, after distillation, had been passed through a Millipore ultrapurification system. Sodium formate and sodium pyrophosphate (both Baker analyzed reagents) were purified by repeated recrystallizations.I2 Anhydrous sodium sulfate (Baker analyzed) was used without further purification. Manganous sulfate is commercially available in very pure form (Alfa/Ventron puratronic, 99.998%); manganous perchlorate was synthesized from manganese metal (Johnson Matthey and Co., Ltd., 99.99%) and perchloric acid in order to attain the required degree of purity. The pHs of all solutions were adjusted by the addition of either ( I ) Gotz, F.; Lengfelder, E. “Oxy Radicals and Their Scavenger Systems”; Cohen, G., Greenwald, R. A,, Eds.; Elsevier Biomedical: New York, 1983; Vol. I, pp 207-17. (2) Bielski, B. H. J.; Chan, P. C. J. Am. Chem. SOC.1978, 100, 1920. (3) Pick-Kaplan, M.; Rabani, J. J. Phys. Chem. 1976, 80, 1840. (4) Braun, D. M.; Dainton, F. S.; Walker, D. C.; Keene, J. P. “Pulse Radiolysis”; Academic Press: New York, 1965; p 221. ( 5 ) Archibald, F. S.;Fridovich, I . Arch. Biochem. Biophys. 1982,214,452. (6) Curnotte, J. T.; Karnovsky, M. L.; Babior, B. M . J. Clin. Invest. 1976, 57, 1059. ( 7 ) Kono, Y.; Takahashi, M.-A.; Asada, K. Arch. Biochem. Biophys. 1976, 174, 459. ( 8 ) Lumsden, J.; Hall, D. 0. Biochem. Biophys. Res. Commun. 1975.64, 595.
(9) Lati, J.; Meyerstein, D. J . Chem. Soc., Dalton Trans. 1978, 1105. (IO) Bielski, B. H. J.; Arudi, R. L.; Cabelli, D. E. “Oxygen Radicals in Chemistry and Biology”; Bors, W., Saran, M., Tait, D., Eds.; Walter de Gruyter: New York, 1984; p 1. (1 1) Bielski, B. H. J. Photochem. Photobioi. 1978, 28, 645. (12) Bielski, B. H. J.; Richter, H. W. J . Am. Chem. SOC.1977, 94, 3019.
0022-3654/84/2088-3111$01.50/0
perchloric acid (double distilled from Vycor, GFS Chemical Co.) or sulfuric acid (“Aristar” from BDH Chemicals, Ltd.). Solutions could only be acidified as the addition of base caused the formation of a pre~ipitate.’~A filtered source of oxygen (UHP, Matheson Co., 99.99%) was used to oxygenate solutions. All rate values were the average of five runs and are good to &lo%. Equipment. All UV/vis spectra were recorded on a Cary 210 spectrophotometer thermostated to 24 “C. The pulse radiolysis experiments were carried out by using the BNL 2-MeV Van de Graaff accelerator as described previo~s1y.l~Doses were computed by using either the ferrous dosimeter or (SCN)z- as calibrants, assuming G values of 15.8 and 6.13, respectively. The spectral path length was either 2 or 6.1 cm, depending upon the nature of the experiment. Radical concentration was varied from 2 to 8 I~.Mof HOz/Oz-.
Results and Discussion Water, when radiolyzed, yieldsI5
--
HzO O H (2.75), eaq- (2.65), H (0.65), HzOz (0.70), H, (0.45) (I) where the values in parentheses are G values, that is, the number of radicals produced per 100 eV of energy dissipated in the system. In the presence of sodium formate and oxygen all primary radicals are converted to HOZ/O2-in the following manner:
+ O2 H +0 2
eeq-
+ HC02COz- + 0 2
OH
-
-
-
02-
HOz
Hz0
--+
COZ
+ COZ-
+ 02-
(2) (3) (4) (5)
The rate of reaction 4 (k4 = 3.5 X lo9 M-’ s-’ (r ef 16)) proved to be the limiting factor in the choice of systems to be studied as many ligands of interest, Le., fumarate, malate, etc., react as fast as if not faster than formate with O H radicals16 while having stability constants of the same magnitude as formate or sulfate with Mn2+.17 This investigation involved studies of the formate, sulfate, and pyrophospate liganded manganous systems. Stability constants (13) Davies, G. Coord. Chem. Rev. 1969, 4, 199. (14) Cabelli, D. E.; Bielski, B. H. J. J . Phys. Chem. 1983, 87, 1809. (15) Schwarz, H. A. J. Chem. Ed. 1981, 58, 101. (16) Farhataziz; Ross, A. B. “Selected Specific Rates of Reaction of Transients from Water to Aqueous Solution. 111. Hydroxyl Radical and Perhydroxyl Radical and Their Radical Ions”; National Bureau of Standards: Washington, DC; No. NSRDS-NBS 59. (17) Martell, A. E.; Smith, R. M., Eds. “Critical Stability Constants”; Plenum Press: New York, 1977; pp 200-2.
0 1984 American Chemical Society
3112 The Journal of Physical Chemistry, Vol. 88, No. 14, 1984
Cabelli and Bielski TABLE I: Rates of Reaction between Manganese(I1) Pyrophosphate and H02/02- as a Function of pH at 25 OC in Solutions Containing 10 m M HCOONa, 10 mM Na4PZ0,, 1.2 mM Oz, and Variable Amounts of MnZ+
T
7.18 6.50 5.77 5.22 4.87 4.37 4.04 3.70 3.36
i
-I 5
~
5
;4t P I
:c
formatecomDlex
[Mn2+l, mM k6, M-I
pH
pyrophosphate complex
t\
a
0.3 0.437 0.39 0.437 0.39 0.39 0.437 0.39 0.437
1.68 X 2.64 X 3.60 X 3.95 X 4.36 X 3.38 X 1.96 X 1.61 X 1.01 X
s'l
pH
lo7 lo7 lo7' lo7 lo7" 10'" lo7 lo7" lo7
2.97 2.69 2.39 2.07 1.68 1.29 1.10 0.70 0.14
[Mn2+l,
k6, M-I
mM 0.52 0.52 0.437 0.52 0.437 0.437 0.461 0.461 0.461
7.70 X 6.20 X 4.58 X 4.17 X 2.97 X 1.77 X 1.31 X 5.87 X 3.10 X
s-l
lo6' lo6"
lo6 lo6' lo6 lo6
lo6 lo5 lo5
1 mM NaCOOH used, other conditions unchanged.
3
IO8F-----T
i 210
230
250
270 290 X ,nm
310
330
I
*
0
350
Figure 1. Spectra of species produced upon reaction of H02/02-with MnZ+in pyrophosphate, sulfate, and formate solutions. ( 0 )Spectrum of manganese(II1) pyrophosphate: pH 7.18, 10 mM Na4P207,10 mM HCOONa, 0.437 mM Mn(ClO,),, 1.2 mM 02,4 WMHOz/O 450 pM. The change in absorbance of manganese(II1) pyrophosphate as a function of pH (see Figure 2, inset), determined at 260 nm, indicated that the manganic complex has a pK of approximately 4.0, in good agreement with a previously published result that indicated a pK of 4.25 for the following reaction: 22
+
I
I
I
I
I
II 0
Mn(H2P207)33-+ Mn(HP207)23- H4P2O7
(7) 0
Our results also indicate an additional pK below pH 1 in this system (Figure 2, inset). Manganese(II1) pyrophosphate is very stable in aqueous solution and, unlike many Mn3+ systems, this stability is not due to the presence of Mn2+. Therefore, a method could be developed to produce manganese(II1) pyrophosphate in almost 100%yield from Mn2+ by using 6oCoradiation according to reaction 6. The presence of H 2 0 2 , however, complicates the conversion since peroxide reacts with Mn3+ that is produced in reaction 6, in the radiolysis of water (eq 1) and in the spontaneous disproportionation of H02/02-. The rate of reaction of manganese(II1) pyrophosphate and H202has not yet been accurately measured. The addition of catalase to the solution obviates this problem as catalase reacts with peroxide to yield water and oxygen. Irradiation of an aqueous solution of Mn2+ (0.1 M formate, 0.02 M pyrophosphate, 250 p M Mn2+,3-5 pM catalase, 02,pH 6.5) for 60 min in a source that produces 14 pM (02-/H02)/min (approximately 2.2 h a d ) along with subsequent heating to 80 "C, cooling, and filtration leads to the conversion of 95-100% manganese(I1) pyrophosphate to manganese(II1) pyrophosphate. The solution did not change during 2 days under refrigeration, as monitored by the absorbance at 260 nm. This is a convenient method of preparing a manganic pyrophosphate solution that is virtually free of manganous pyrophosphate. Manganous SulfatelManganous Formate. The reactions between manganese(I1) sulfate or manganese(I1) formate and H02/0< at pH 4-6.5 result in the formation of transients having broad absorption bands with maxima at 270 nm. The corresponding molar absorbances of the sulfate and formate species are 2.2 X lo3 and 3.2 X lo3 M-' cm-' , respectively (see Figure 1). These studies were carried out with varying concentrations of MnL+,sodium sulfate, and sodium formate in order to establish that the kinetics were concentration independent and that the Mn2+ was fully liganded. This was particularly important in the manganous sulfate system where the addition of some formate was necessary in order to generate H02/02-(see reactions 2-5). The ionic strengths of the solutions were, therefore, quite high for kinetic studies and, whereas the rates of the reactions did vary somewhat with ionic strength, no systematic study of this effect was practical. It should be noted that since these studies are carried out in aqueous solutions, the Mn2+ is not only liganded but also hydrated. Although the mechanisms in the sulfate and formate systems are similar, the kinetics and the spectral characteristics vary with the ligand. The initial steps in the mechanism of these two systems are described by the following reactions:
+ H02 Mn2+ + 0;
Mn2+
MnOOH2+
-
MnOOH2+
+ MnOz+
M n 0 7 + + H+
I
(8)
formate complex sulfate complex
i
IO61
2
3
4
5
6
7
8
PH
Figure 3. Plots of kobdvs. pH for the reactions of manganese(I1) formate and sulfate with HOz/Oz-; see Tables I1 and 111for experimental conditions. Theoretical curve for the formate system calculated by using eq 11, k8 = 6.4 X lo6 M-l s-l a nd k9 = 4.6 X lo7 M-' s-l. TABLE II: Rates of Reaction between Manganese(I1) Formate and H02/0f as a Function of pH at 25 OC in Solutions Containing 0.4 M HCOONa, 1.2 mM 02,and Variable Amounts of Mn2+ pH 7.06 5.94 5.78 5.71 5.29 4.91 4.62 4.02
IMnZ+l, -mM3.57 3.57 7.5 15.0 3.57 3.57 3.57 3.57
M-' S-' 4.67 x 107 4.57 x 107 4.27 X lo7 4.18 X lo7 3.52 X lo7 2.70 X lo7 1.90 X lo7 1.90 X 10'
kobsd,
pH 3.75 3.37 3.07 3.02 2.90 2.81 2.21
IMn2+1, mM3.57 3.57 3.57 17.8 3.57 17.87 17.8
kobr,j,M-' S-' 9.80 X lo6 6.50 x i o 6 6.30 X lo6 5.00 X lo6 7.70 X lo6 4.60 X lo6 4.60 X lo6
TABLE III: Rates of Reaction between Manganese(I1) Sulfate and H02/0f as a Function of pH at 25 OC in Solutions Containing 0.1 M Na2S04, 1.2 mM Of, and Variable Amounts of Mn2+ uH IMnZ+l,mM INaCOOH1. mM knhOA. M-' s-l 5.60 1.o 1.o 5.4 x 107 5.0 x 1070 5.15 5.03 2.0 4.85 10.0 1.o 3.6 x 1070 1.o 3.36 x 107 4.72 1.o 4.43 1.o 1.o 2.83 x 107 4.11 1.0 0.5 2.36 x 107 2.17 x 107 1.o 3.94 1.o 3.65 1.0 0.5 2.00 x 107 20.0 3.39 20.0 6.85 X lo6 12.8 3.31 25.0 5.50 X lo6 3.11 10.0 5.0 7.00 X lo6 3.07 20.0 20.0 6.05 X lo6 2.82 10.0 1.o 8.00 X lo6 2.77 20.0 20.0 6.80 X lo6 "0.05 M NaSO., used.
a study of the dependence of the observed rate upon pH, by using a simple kinetic equation:
(9,-9) (10,-10)
where, as noted previously, H 0 2 and 02-are in equilibrium as are Mn2+ + 0; and Mn02+.3 All measurements of equilibrium 9,-9 were, therefore, carried out at [Mn2+]> 1 m M to ensure that the equilibrium was dominated by the foward reaction. The rate constants for reaction 8 and reaction 9 can be obtained from (22) Tarayan, V. M.; Elumyan, M. G. Izu. Akad. Nauk Arm. SSR,Khim. Nauki 1958, 11, 2 3 .
The observed rates, measured under pseudo-first-order conditions, are listed in Tables I1 and I11 and plotted in Figure 3. In the formate system these rates have been used along the pK1 and eq I1 to derive a theoretical curve, giving ka = 6.4 X IO6 M-I a nd k9 = 4.6 X lo7 M-' s-l . Values for k( and kg/ in the sulfate system can be obtained from the plateau regions in Figure 3 and are kg' 7.0 X lo6 M-I s-l a nd kg/ = 5.4 X lo7 M-' s-l. Observed rates of reactions in both systems can be measured only down to pH 2-3 and the values in the low pH region tend to have large ex-
-
3114
The Journal of Physical Chemistry, Vol. 88, No. 14, 1984
Cabelli and Bielski TABLE I V Rates of Formation of Manganese(II1) Sulfate in Solutions Containing 0.5 M Na2SOd,1.2 mM Os and Variable Amounts of MnZt and HCOON; at'pH 1-1.3 aid 25 OC
[Mn2+], mM
m
2 K
m
m
a
lo5 !
I
I
I
I
20
40
60
80
i
10 25 50 100
'"v 35
30
4
25 k
m
M-'
S-'
kcalcd,
8.23 X 10' 1.48 X lo2 3.06 X lo2 5.58 X lo2 9.38 X lo2
7.73 X 1.39 X 3.09 X 5.54 X 9.33 X
s-I 10' lo2 lo2 lo2 lo2
TABLE V Rates of Formation of Mn3+ in 0.5 M HCOONa Solutions Containing 1.2 mM Oz and Variable Amounts of MnZ+at pH 1-1.3 and 25 OC
5
40
kobsd,
1 5 10 25 50
1
45
0K
3.9 7.8 19.4 39.0 78.0
[MI?+], mM
TIME x 1 0 6
m
[NaCOOH], mM
knhdr S-'
7.52 X 1.33 X 1.97 X 3.53 x 5.85 X 9.25 X
10'
lo2 lo2 102 lo2 lo2
kraicdl
9.16 X 1.39 X 1.97 X 3.55 x 5.81 X 9.27 X
s-'
10' lo2 lo2 102 lo2 lo2
The observed rate of the process illustrated by the trace in Figure 4b can then be calculated in terms of k12,k13, K1,, and the total Mn2+ concentration according to the relationship
m
a
15 10
5 1 1
Od
I
I
IO TIME x ~ 0 3
5
15
i li
20
Figure 4. Traces produced from pulse radiolysis experiments of the reaction of manganese(I1) formate with H02: (a) initial buildup that represents the reaction of H 0 2 with manganese(I1) formate; (b) subsequent first-order buildup, see text for discussion (pH 2.8, 270 nm,0.4 M HCOONa, 18 mM Mn2+).
perimental uncertainties associated with them. The reason for this and the equilibrium constant Klobetween the protonated and unprotonated manganese dioxygen complex will be discussed later. The unprotonated complex, Mn02+, disappears, in both sulfate and formate systems, by mechanisms that involve 02-and are currently under investigation, At low pH (pH 1-3) the observed kinetics become complicated. The aforementioned buildup that led to the determination of k8 and k9 (Figure 4a) was followed by a strictly first-order process (Figure 4b). Although the latter buildup varies with Mn2+ concentration, it is not directly proportional to the cation concentration. In the sulfate system the spectrum of the transient obtained 10-20 ms after the pulse (i.e., at the end of the trace shown in Figure 4b) was identical with the well-established spectrum of manganese(II1) sulfate (from the permanganate oxidation of Mn2+ to Mn3+ in acidic solution where manganese(II1) sulfate is reasonably stable).23 These results can be explained in light of previous work describing the ferrous perchlorate/HO, system that suggests a mechanism involving an equilibrium between Fe3+-H02- and Fe3+-H02--Fe2+ with the subsequent oxidation of both monomer and dimer to Fe3+.24 In analogy, the Mn2+/H02system can be described by reactions 8-10 followed by reactions 12 and 13 and equilibrium (1 1,-1 1) MnOOH2+
+ MnZf
-
(MnOOHMn)4t (11,-11)
+ HOz( M I I O O H M ~ ) ~ + Mn3+ + product MnOOH2+
Mn3+
(12) (13)
(23) Fackler, J. P., Jr.; Chawla, I. D.Inorg. Chem. 1964, 3, 1130. (24) Jayson, G. G.;Parsons, B. J.; Swallow, A. J. J . Chem. SOC.,Faraday Trans. 1 1913, 69, 236.
Experiments were carried out a t pH 1-1.3 where the observed rate was measured while varying the concentration of Mn2+ by approximately 2 orders of magnitude (see Tables IV and V). The results led to values of k12= 7.93 X 10' s-l, k,, = 2.81 X lo3 s-I, and Kll = 4.5 M-' for the formate system and k l i = 1.27 X 10' s-', k13/ = 3.06 X lo3s-l, and Kll' = 5.5 M-' for the sulfate system. It should be noted that the rate of reaction 13, the dissociation of the dinuclear species, (MnOOHMn)4+, is considerably faster than the rate of reaction 12, the dissociation of the metal-perhydroxyl complex, MnOOH2+, in both formate and sulfate systems. This indicates that although an increase in Mn2' concentration tends to favor the formation of (MnOOHMn)4+ (equilibrium 11,-1 1), the lack of stability of this complex counteracts that effect. This makes direct measurements of the absorption spectra of either transient species untenable. The spectra of MnOOH2+ and (MnOOHMn)4+can be obtained through less direct methods, as described for the Fe(C104)2/H02systems.24 In analogy to eq 111, a relationship can be derived defining the molar absorbances of these species in terms of the observed molar absorbances seen in Figure 4 at the end of trace a or the beginning of trace b:
If the observed spectra obtained by varying MnZ+concentrations are used (spectra taken from the experiments described in Tables IV and V), the spectra of MnOOH2+-and (MnOOHMn)4+shown in Figure 5 for the formate and sulfate system can be obtained. The spectra of Mn3+in formate and sulfate solutions, also shown in Figure 5 , were measured directly and were invariant to Mn2+ concentration at [Mn2+] > 5 mM. The manganic cation has been studied extensively in aqueous solutionI2 and will therefore not be discussed here except to note that it disappears by a second-order process with a rate that increases as the concentration of Mn2+ decreases; Le., Mn3+ is stabilized by the presence of MnZ+as well as by the presence of strong acids. Also, the change in absorbance with pH for Mn02+ indicates that reaction 10,-10 has a pK in the range of 2.5-3.5 for both the formate and sulfate systems. A more accurate determination is not posssible due to the mechanism of disappearance of MnOOH2+ (reactions 11-13) and the disappearance of Mn02+ itself. Finally, it should be noted that formulas describing the distribution of electrons were avoided in denoting the mononuclear
J . Phys. Chem. 1984, 88, 3 115-3 118
(-)
in the presence of sufficient sulfate or formate, Mn2+reacts with H02/02- to form the metal dioxygen species Mn02+/Mn00H2+. MnOz+ disappears by complex mechanisms. The subsequent reactions of MnOOH2+are much more complex; MnOOH2+is in equilibrium with the dinuclear species (MIIOOHM~)~'.Both species then disappear with the concomitant appearance of Mn3+; these reactions occurring at relatively fast rates (tl12) of the order of 9-20 ms for the mononuclear species and 0.2 ms for the dinuclear species). The oxidized species Mn3+ disproportionates at a rate that varies with the Mn2+ concentration. The Mn02+/Mn00H2+pair has pK"s that are estimated to be in the range of 2.5-3.5 for both the formate and sulfate species. While complex formation between H02/0; and metal cations is well d ~ c u m e n t e d , ~ ~such ~ ~ ~systems ~ - * ' were usually not investigated over a wide pH range. As is apparent, the reaction mechanisms and reaction products can vary significantly with pH. Studies of this nature are very fundamental to the understanding of autoxidation processes in the presence of metals and metalcatalyzed oxidation/hydroxylation reactions.
SULFATE, pH 1.3
(-4 FORMATE, pH 1.0
230 240
250
260
270
280 290
300
310
3115
320
PH
Figure 5. Spectra calculated by using eq IV;see text for experimental conditions and discussion.
and dinuclear complexes, MnOOHZ+and (MnOOHMn)4+. A determination of the charges on the manganese centers was not within the scope of this work.
Summary and Conclusion In aqueous solution manganese(I1) pyrophosphate reacts with H02/O; to yield manganese(II1) pyrophosphate. In contrast,
Acknowledgment. We thank Prof. R. L. Arudi for many stimulating discussions. This research was carried out at Brookhaven National Laboratory under contract with the U S . Department of Energy and supported by its Office of Basic Energy Sciences. Registry No. Mn2+,16397-91-4; MnOOH2+,90342-66-8; Mn02+, 66460-00-2; H02, 3 170-83-0; 02-,11062-77-4; manganese(I1) sulfate, 7785-87-7; manganese(I1) formate, 3251-96-5; manganese(I1) pyrophosphate, 13446-44-1; manganese(" pyrophosphate, 64042-23-5.
(25) Samuni, A.; Czapski, G. J. Phys. Chem. 1970, 74, 4592. (26) Sellers, R. M.; Simic, M. G. J . Am. Chem. SOC.1976, 98, 6145. (27) Ilan, Y. A.; Czapski, G.; Ardon, M. Isr. J . Chem. 1975, 13, 15.
NO3 Quantum Yields from N,05 Photolysis Diane Swanson, Brian Kan, and Harold S. Johnston* Materials and Molecular Research Division, Lawrence Berkeley Laboratory and Department of Chemistry, University of California, Berkeley, California 94720 (Received: November 18, 1983)
The technique of laser flash photolysis/laser resonance absorption was used to directly measure the quantum yield for NO, production upon ultraviolet photolysis of NzO5. The average NO3 quantum yield was found to be 0.89 f 0.15 . There appeared to be a slight decrease in quantum yield as the concentration of N 2 0 Sincreased and as the pressure of carrier gas (N, or 0,) increased. The indicated NO3quantum yield at low N205concentrationand low gas pressure is 1.O, but there is uncertainty from the experimental error. A fast rise in the concentration of N03(u= 0) was observed, and the rise time was dependent on N2O5 concentration and on N2 pressure. These data are interpreted to give rate constants for the deactivation process N 0 3 ( u>0) + M NO,(u=O) + M where k(N,O,) = (4.0 f 1.0) X cm3 molecule-' s-I and k(N2) = (1.6 f 0.3) X io-', cm3 molecule-' s-'.
-
Introduction After N2O5 absorbs near-ultraviolet radiation, the possible primary photochemical products are 2N02+ 0 X C 401 nm NzO5 + hv (1)
- +
-
NO2
NO
+ O2
X C 1125 nm
(2)
NO2 + NO3 X C 1300 nm (3) A study' published in 1934 looked at pressure changes upon constant illumination and evaluated quantum yields for overall Three unpublished investigations formation of 2 N 0 2 '/*02.
+
(1)
H. H. Holmes and F. Daniels, J . Am. Chem. Soc., 56, 630 (1934). 0022-3654/84/2088-3115$01.50/0
are in Ph.D. these^.^-^ The first two used indirect methods and will not be discussed here. The third used laser flash photolysis between 290 and 300 nm and looked for nitric oxide and atomic oxygen by means of resonance fluorescence. N o nitric oxide was seen, which set an upper limit of 0.1 on the quantum yield of channel 2. The quantum yield of atomic oxygen was found to be 0.35 f 0.10, which presumably comes from channel 1. The (2) R. Murphy, Ph.D. Dissertation, University of California at Los Angeles, 1969. (3) P. S. Connell, Ph.D. Dissertation, University of California at Berkeley, 1979. (4) F. Magnotta, Ph.D. Dissertation, University of California at Berkeley, 1979.
0 1984 American Chemical Society
