J . Phys. Chem. 1993,97, 6664-6669
6664
Formation and Decay of Peroxynitrous Acid: A Pulse Radiolysis Study T. Lagager and K. Sehested' Section of Chemical Reactivity, Environmental Science and Technology Department, Rise National Laboratory, DK-4000 Roskilde, Denmark Received: February 22, 1993; In Final Form: April 13, 1993
Peroxynitrous acid and peroxynitrite anion have been studied using pulse radiolysis of nitrite and nitrate solutions. The formation rate constant is determined to be k(OH+N02) = (4.5 f 1.0) X lo9 M-l s-l, and the rate constant for the OH radical reaction with nitrite is determined to be k(OH+NOz-) = (6.0 f 1.0) X lo9 M-l s-l. In nitrate solutions, the competingreaction between O H and N O g i s found to have a rate constant of k(OH+N032-) = (3.0 f 1.0) X 109 M-1 s-l. The intermediate species in the nitrate system, Nos2-, HN03-, and H2NO3, decay into NO2 according to the first-order rate constants: (5.6 f 0.5) X lo4, (2.0 f 0.5) X lo5, and (7.0 f 2.0) X 105 s-1, respectively. The rate constants k(H+NO3-) = (1.0 f 0.3) X lo7 M-l s-l and k(H+N02) = (1.0 f 0.2) X 1010 M-I s-l were also determined. The PKa of ONOOH is found to be 6.5 f 0.1 by absorption measurements, and the maximum extinction coefficient at 240 nm is c z ~ ( O N 0 0 H = ) 770 f 50 M-l cm-l. The decay of peroxynitrous acid is determined to proceed through the first-order isomerization of ONOOH to HNOs according to the rate equation k& = &/( 1 Ka/ [H+]) with rate constants ki, = 1.O f 0.2 s-l and Ka = (1 .O f 0.3) X 10-7. A comparison of all available literature values for the PKa and the decay rate is reported.
+
Introduction Peroxynitrous acid (ONOOH) and its anion (ONOO-) are becoming increasingly acknowledged as major components in aqueous atmospheric chemistry. There are several methods of forming peroxynitrous acid in water, the most widely used are as follows: (1) reaction of hydrogen peroxide with nitrous acid at low pH and quenching of peroxynitrite with (2) reaction of NO with 02-,10.11 (3) flash photolysis of nitrate solutions,l2J6 and (4) pulse radiolysis of nitrate solutions.17.~*In this study we use pulse radiolysis of nitrate and nitrite solutions, the latter of which has been studied before, however, without special emphasis on peroxynitrous acid. Our initial interest was to study the great discrepancy between the decay and pK, value found using pulse radiolysisI7compared to the values found using other techniques.l-16 A closer literature examinationgave a great deal of information on both the decay rates and the pKa values, although sometimes very diverging and confusing. Lastly, only two literature rate constants for the formation of ONOOH from OH and NO2 were available: 1.3 X IO9 M-l s-l 17 and 1.2 X 1010 M-I s-1.l6 As the first is based on already diverging information and the determination of the latter was not described, this seemed to leave a gap in the peroxynitrous acid literature. Thus our aim was to study all these aspectscomprehensively over the whole pH range.
Experimental Section Materials and Solutions. All solutions were prepared using triply distilled water. NaN02 (Riedel-de HaBn p.a.), NaN03, Na2B407, Na2HP04, NaH2P04, NaOH (Merck, p.a.), HC104 (Frederick Smith Co, doubly distilled from Vycor), and N20 (N40) and Ar (N48) were used as supplied. None of the compounds contained more than lO-5% trace metals (Fe, Cu). NzO-saturated (0.05-1.0) X M NaN02 solutions and Arsaturated (1-5) X 10-3 M NaN03 solutions adjusted to the designated pH using 0.01 M NaOH or 0.01-1.0 M HC104 were studied in the wavelength region 240400 nm in the time range 0-8 ms. For the pKa determinations, the solutions at pH 5-8
* Address correspondence to this author. 0022-3654/93/2097-6664%04.00/0
were buffered with 1 X 10-3 M phosphate buffer, and those at pH 9 were buffered with 1 X l k 3 M borax buffer. The yields of ONOOH and ONOO- were measured at 250 and 300 nm, respectively, 8 ms after the pulse, as all other transient species have decayed at this time. The ONOOH and ONOO- decays were followed optically to completion at 250, 300, or 310 nm, respectively, either in the pulse radiolysis setup where 230-, 280-, or 305-nm filters were used to eliminate solute and peroxynitrite anion photolysis in the time scales < 0.5 s or, for slower reactions, in a 5-cm quartz cell to which the pulsed solution was transferred and followed optically using a SP8-400 UV/VIS spectrophotometer. Apparatus. The pulse radiolysis setup consists of a IO-MeV HRC Linac accelerator delivering 0.5-3.0-ps electron pulses into a quartz cell.19 The optical detection system comprises a 150-W Varian xenon lamp, a 5.1-cm light-path cell, a Perkin-Elmer double quartz prism monochromator, a 1P28 photomultiplier, and a LeCroy 9400 digital oscilloscope. An IBM PC/AT3 computer was used for data processing. As the formation of peroxynitrous acid is a second-order reaction (OH NO2) a fairly high dose per pulse is administered to enhance the yield. A dose of 16-20 krad in a 1.2-ps pulse is usually applied, but 40-50 krad in a 3-ps pulse has also been used. The dose was measured using the hexacyanoferrate(I1)dosimeter, G = 5.9 and E420 = 1000 M-I cm-l, containing (2-5) X M Fe(CN)6'. All experiments were carried out at ambient temperature. To expand our experimental possibilities, we have used our high-pressure, high-temperature cell to increase HZsolubility by applying Hz pressure.20 Modeling of the experimentalresults was carried out using CHEMSIMUL,21a program developed at RIS0 for the numerical simulation of chemical systems.
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Results and Discussion Two systems were used to study peroxynitrous acid formation and decay using pulse radiolysis: (A) N2O-saturated nitrite solutions and (B) Ar-saturated nitrate solutions. The formation of either ONOOH or ONOO- is described according to the reaction mechanism of each system as 0 1993 American Chemical Society
The Journal of Physical Chemistry, Vol. 97, No. 25, 1993 6665
Formation and Decay of Peroxynitrous Acid
TABLE I: Reaction Mechanism and Rate Constants Derived from the Experiments and Used for AU Computer Modeling of Both the Nitrite and Nitrate Systems no. 1
2 3
4 5 6 7 8
9 10 11
rate constant (M-1 s-1) e-, + N ~ O N* + OH + OH- 9.1 x 109 (6.0 1.0) x 109 OH NO,NO2 + OH(4.5 1.0) x 109 OH NO2 ONOOH %- NO,-+ NO,% 9.7 x 109 (5.6 t 0.5) X 104 s-l N0:2- Hz0 -+ NO2 + 20HpK,, see Table 111 ONOOH a H+ + ONOO5.5 x 109 OH + OH HzOz k7 4.5 X lo8 NO2 + NO2 + Nz04 k-7 = 7 x 103 s-1 N204 H20 HNOi + HNOJ 18.0 (7.0 2.0) x 105 S-I (4.6 0.3) X 10'
-
reaction
+ + +
+
+
+
+
+
* *
+
12 HNO:--NOz + OH1 3 OH + NO:%+ NO:- + OH14 %-+ H+ H + Hz0 HN0:15 H + NO:16 OH + H z + H + Hz0 17 H + NO2 HNOz 18 ONOOH HNO, 19 HNOz s H+ + NOz-
(2.0 t 0.5) x 105 S-l (3.0 t 1.0) x 109 2.3 X 10'0 (1.0 t 0.3) x 107 3.4 x 107 (1.0 0.2) x 1010 kls, see Figure 6 pK, = 3.2
+
+
*
+
+
e,,
+ N,O
-
-
+ NO; O H + NO,
OH
+ H,O OH + NO,
NO,"
+ OH + OH-
N,
NO,
-
-
+ OH-
ONOOH
NO,
+ 20H-
ONOOH
reference 22 this work this work 22 17, this work
22 24 24 24 this work this work
and N,04
+ H,O
-
HNO,
+ HNO,
-
6 -
I
0 0
z
0
4 -
Y
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1
2
3
4
5
6
7
8
9
10
Figure 1. Expeximental(0) and modeled (solid line) yields of ONOOmeasured at 300 nm, in NPsaturated nitrite solutions, pH 9.5, 8 m after the pulse, as a function of nitrite concentration. A 1.2-pa pulse of 18.9 krad was used.
this work this work 22 this work 20 this work
23
(1) (2) (3)
(5)
(3)
* +
-
8 -
3 v
[NO-*] (M x lo')
ONOOH H+ ONOO(6) As the formation is a second-orderreaction, each system is affected by other second-order radical-radical reactions
+ OH H,O, NO, + NO, ~iN,04
5
0
and in both systems according to the acid-base equilibrium of the peroxynitrous acid.
OH
-
10 m
(7) (8)
(9)
and by reactions of the H radicals (see Table I) which make up 10% of the total G value at neutral pH. Furthermore, the ionization state of the N032-intermediate affects the mechanism of ONOO- formation in the nitrate system. Formation and Yields of O N W and ONOOH in the Nitrite System, Determinationof k2 and kj. In the alkalinenitrite system, the formation of peroxynitrite anion is a competition between reactions 2,3,7, and 8. The yield of ONOO- was studied at 300 nm in N20-saturated solutions (pH 9.5), and the absorption 8 ms after a 19-krad pulse was measured as a function of nitrite concentration (Figure 1). The absorption was converted into ONOO- concentration using the extinction coefficient ~300-
(ONOO-) = 1670M-1cm-1.2Computer modelingoftheON00 yields, using the program Chemsimul and the reaction mechanism in Table I, gives rate constants for reactions 2 and 3, kz = (6.0 f 1.O) X lo9and k3 = (4.5 f 1.0) X lo9M-1 s-l, respectively, with a good fit throughout the nitrite concentration range used. Literature values for k2 range from 0.7 X 10'0 to 1.4 X 1010 M-1 s-1.22 The formation and the yield of ONOOH in acidic nitrite solution (pH 4-5) are assumed to be the same as those in alkaline solution as long as the pH is kept one pH unit above the pK. = 3.2 of nitrous acid.23 Two values for the formation rate of peroxynitrous acid, k3, exist in the literature, 1.3 X 109 M-1 s-1 l7 and 1.2 X 10'0 M-l s-1.16 It is not possible to findinformationabout thedetermination of the latter, and to comment on the value found by Griitzel et al.,17it isnecessary tonote that they havebased their determination of k3 on the absorption of ONOO- at 300 nm, 250 ps after the pulse in the pH range 4-6. At 250 ps after the pulse, NO2 and NzO4 are still present under the conditions used.24 and thus the determination of Griitzel et al. of [ONOOH]250 is contaminated with NO2 and NzO4 resulting in an incorrect determination of [NO2]zso = [OHIO- [ONOOHI25~ used for their determination of k3(OH+N02). Furthermore, their k3 is determined in the nitrate system where intermediate species like NO3,- complicate the mechanism (see below). Formation and Yields of O N O - and ONOOH in the Nitrate System, Determination of ks 40,k12, k1k krs and k17. The formation kinetics of peroxynitrite anion is more complex in the nitrate system due to the formation of an intermediate species (NO+-) formed from the hydrated electron (reaction 4). This species and its protonated forms decay into NOz, whereby the formationof peroxynitrous acid isdelayed(reaction 3). Therefore the formation rate of NO2 from these intermediate species was studied over the whole pH range. The intermediate has pK, values of 4.8 and 7.5,17 and the three intermediate forms (H2NO3, HN03-, and N032-) have decreasing decay rates with increasing pH. We have measured fint-order decay rateconstants of H2NO3 and NO3,- directly at 250 and 300 nm in unbuffered Ar-saturated 5 X 10-3 M nitrate solutions, pH 3 and 11, respectively, using a 2-3-had dose to minimize second-order reactions 3, 7, and 8 H,N03-NO, NO:-
+ H,O
+ H20-NO, + 20H-
4,, = 1.2 ps
rl,, = 10.0 ps
(10) (5)
givingk5=7.0X lO4s-landklo- (7.0h2.0) X loss-1. However, the decay of HN03- could only be studied in buffered solutions and the decay depends on the phosphate buffer concentration.
Lprgager and Sehested
6666 The Journal of Physical Chemistry, Vol. 97,No. 25, 1993 1.6
TABLE Ik Experimental and Modeled Yields of ONOOH in the Nitrate System as a Fuactim of Acidib exptl modeled pH dose [ONOOH] X 106 M [ONOOH] X 106 M
1.4 1.2
1 .o
0
1 .o 1.0 2.16 2.25 2.95 3.85
0.8 0.6
0.4 0.2
0.0 I 0.0
I
I
I
I
1
0.5
1.0
1.5
2.0
2.5
(M
[H,PO,-]
x 10')
*
Thus the protonation reaction with phosphate,
+ H,PO,
-
H,N03
+ HPOt-
(1 1)
was studied directly as a peudo-first-order decay of HN03- at 280 nm using a 4-krad pulse, At-saturated lCr3 M nitrate, pH 6, as a function of (0.3-3.0) X 10-3 M phosphate buffer (Figure 2). It was found to have a rate constant kll = (4.6 f 0.3) X lo8 M-1 s-I, which is in accordance with kll = 5 X IO8 M-l s-l.17 The first-order decay of HN03- into NO2 was found as the y-axis intercept of this relationship to be k12 = (2.0 f 0.5) X los s-l (Figure 2).
HNO,--
NO,
+ OH-
r1,, = 3.5 ps
(12)
Grtitzel et al.17 have found 12.5- and 3.0-ps half-lives for N032and HN03-,respectively, and the half-life of HzNO3 was assumed to be the same as that for HN03-. In their measurements they scavenged the OH radical with tert-butyl alcohol. In our study, however, the OH radical is free to react according to OH
+ NO-:
-
NO;
+ OH-
1.01 1.77 1.97 4.57 6.12 12.2 16.3
-
satisfactorily with a rate constant for reaction 15 of kls = (1 .O 3.0
Figure 2. Pseudo-fint-order decay rate constant of HNOr, measured at 280 nm using a 4-had pulse, as a function of phosphate buffer concentration at pH 6.0; the rate constant kll = (4.6 0.3) X 108 M-l s-l, and kl2 = (2.0 f 0.5) X lo5 s-l.
HNOL
0.8 1.7 1.9 4.8 6.0 12.3 15.9
17.8 17.8 23.3 17.8 23.3 23.3 23.3
(13)
Assuming a 12.5-ps half-life for N032-decay" (reaction 5), our observed decay rate (tip = 10 MS) suggests a rate constant for reaction 13 of k13 = (2.5 f 0.5)X lo9 M-1 s-l to compensate for the rate difference in half-lives. Additionally, the yield of ONOO- in alkaline solution, as measured at 300 nm using an extinction coefficient of 1670 M-l cm-1, revealed that reaction 13 is needed to decreasethe modeled yield at pH 1 10. From modeling the observed ONOO- yield at pH 10.7 using k3 = (4.5 f 1.0) X lo9 M-l S-I we get a rate constant kl3 = (3.0 f 1.0) X 109 M-l s-1, which agrees well with the rate constant determined directly from the decay of NO3,above. Our determination of the competing reactions 5 and 13 is thus in accordancewith the tllz = 12.5ps determined by Grtitzel et al.,I7 so that ks = 5.6 X 104 9-1. As the decay of the protonated forms HN03- and H2NO3 to NO2 is much faster than the N032- decay, their reaction with OH (analogous to reaction 13) has no appreciable effect on the yield of ONOO- (ONOOH). The yield of ONOOH in acidic solution, pH 0-4,measured at 250 nm with an extinction coefficient ~zJo(ONOOH) = 700 M-l cm-l (see the following section), is mainly determined by the differencein reactivity of eaq-and H atoms toward nitrate, as e, in acid is converted into H atoms by reaction 14, k14 = 2.3 X 1010 M-l s-l.22 The experimentalyields, Table 11, could be simulated
H
+ NO3-
HN03-
(15)
f 0.3) X 107 M-1 s-1. This is in agreement with the value found by Navon and SteinZSbut not with the kls = 1.4 X 106 M-l s-l cited by Buxton et aLzz To verify the rate constant for reaction 15, 140 atm of H2 pressure was added (dissolved after 20 min stirring to 0.1 M Hz) to an Ar-saturated, pH 0, 5 X 1O-, M nitrate solution in our pressure cell20 and the NO2 yield was monitored at 400 nm. The H2 converts all the OH to H according to
OH + H,
-+ H
H,O
k,, = 3.4 x 10'
M-l
s - ~22 (16)
and since H+ scavenges all electrons at pH 0 (reaction 14) all radicals are converted into H atoms. The rate constant for reaction 15 could thus be found in the Hz pressureexperimentby modeling the buildup and yield of NO2 using ~ w ( N 0 2 = ) 180 M-l cm-1.26 The experimentalyield of NO2 was 2.2 X M, and themodeled yield using kls = (1.0 f 0.3) X lo7 M-l s-l was 2.4 X l e 5 M, while when using the literature value kls = 1.4 X 106 a yield of only 6.0 X 1 V M NO2 was obtained. This system however also demanded that the rate constant of the reaction H
+ NO,
-
HNO,
(17) was k17 = (1.0 f 0.2) X 1010 M-l s-l, otherwise the NO2 yield was too high or the buildup too slow. ONOO- and ONOOH Spectrum and e(ONO0H). The spectrum of ONOO- is reproduced in both the nitrite and nitrate systems using 2 X 10-4 nitrite and M nitrate, respectively at pH = 9.3 (Figure 3). The absorption measurementsare taken 8 ms after a 19-krad pulse, at which time the peroxynitriteanion is the only absorbing speciespresent, as all other absorbing species have decayed (reactions 2-9). ONOO- has a maximum absorption at 300 nm with an extinctioncoefficient of c300(ONOO-) = 1670 M-1 cm-1.2 The spectrum of ONOOH is measured in Ar-saturated 10-3 M nitrate, pH 3 and 4, with 18- and 40-krad doses, respectively (Figure 3). A 230-nm light filter was used to minimize nitrate photolysis. The spectrum was taken 4 ms after the pulse and five pulses were averaged to improve the signal to noise ratio. The spectrum of ONOOH was also found in N20-saturated 5 X 1 V M nitrite, pH 2, with a 40-krad dose. Although about half of the electrons react with H+ rather than N20, H atoms react slowly with HN02 and calculations show only minor interference from these reactions, 8 ms after the pulse. The normalized spectra, corrected for loss in nitric acid, nitrate, and H202 produced, are shown in Figure 3. The ONOOH spectrumappears to have a maximum absorption at 24Cj245 nm, but it is difficult to determine accurately, because of the strong UV absorption of both nitrous acidZ7and nitrate.14 In the literature, the spectrumof ONOOH has been determined from flash photolysis of nitrate13J4 and from reaction of H202 with HN02.6.8 Our spectrum agrees well with the spectra of Shuali et al.,l3 Drexler et a1.,* and Barat et ale," however, it is different from what Benton and Moore6 have been observing. None of the literature studies have attempted to assign extinction coefficients to the spectrum; however, Shuali et al.13 show the
Formation and Decay of Peroxynitrous Acid
The Journal of Physical Chemistry, Vol. 97, No. 25, 1993 6667
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250
300
350
1500
1000
500
0
Wavelength (nm) Figwe 3. Spectra of ONOOH and ONOO-, normalized at 250 and 300 nm, respectively, found in both nitrite and nitrate systems, 4-8 ms after a 1840-krad pulse. (v)5 X 10-4 M nitrite, pH 2,40 krad; (A) M nitrate, pH 3, 18 had; (0) 10-9 M nitrata, pH 4,40 krad; ( 0 )2 X 10-4 M M nitrate, pH 9.3, 19 had; (0) M nitrate, pH 11, 21 krad. nitrite, pH 9.3, 19 krad; (+)
relative absorption between ONOOH and ONOO- to be 0.72 at 245 and 300 nm, respectively, assuming equal yields at the two pHs. We have determined the extinction coefficient of ONOOH at 250 nm by comparing the absorptions at 250 and 300 nm in NzO-saturated 10-4 M nitrite, pH 5 and 10, respectively, using a 16-krad pulse. As the mechanism for ONOO- and ONOOH formation and G values are aseumed to be identical at the two pHs, we assume identical yields. If e3w(ONOO-) = 1670 M-l cm-l is applied, the extinction coefficient for ONOOH is 6250(ONOOH) = 700 f 50 M-1 cm-1, which gives e z ~ ( 0 N O O H = ) 770 f 50 M-1 cm-1 for the maximum extinction coefficient. The extinction coefficient cannot be found in the nitrate system as the mechanism in nitrate and thus the yields are different in acid and alkaline solution due to the formation of the intermediate NO32-/HNO3-. However, if we in our calculations of the yields in the nitrate system adopt the extinction coefficient from the nitrite system, the calculated absorptions of ONOOH in acid (pH 3 4 ) at 250 nm and ONOO- in alkaline (pH 9-10) at 300 nm agree well with the absorptions obtained experimentally. Determination of pK,(ONOOH). The pH-dependent peroxynitrite anion absorption was studied in NzO-saturated 10-4 M nitrite solutions using a 16-krad pulse (Figure 4). The solutions were buffered using M phosphate buffer in the pH range 5-8. The yield was measured at 300 nm, 8 ms after the pulse. A fresh solution was used for each measurement. A pKa(ONOOH) = 6.5 f 0.1 is evaluated from the S-shaped absorption curve. By computer simulations of the system it was found that absorption corrections for nitrite used and nitrate and Hz02 produced were insignificant at 300 nm. The pKa curve in the nitrate system turned out to be quite different from the smooth pKa curve found in the nitrite system (Figure 5). Ar-saturated 5 X 10-3 M nitrate solutions were studied at 300 nm, 8 ms after a 22-krad pulse. Five first pulses were averaged. 1 X 10-3 and 5 X 10-3 M phosphate buffer was used in the pH range 5-8, and 1 X 10-3 and 5 X 10-3 M borax buffer was used at pH 9. The yield of ONOO- at a designated pH is mainly determined by the decay rate (reactions 5, 10, and 12) of the type of intermediate formed in reaction 4. Therefore the yield changes
0.08 I
0.07
0.06
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1
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1
I
I
I
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1
t
1
0.00 1 2
I
4
6
8
I
10
I
I 12
PH Figure 4. The yield of ONOO- at 300 nm, observed in NzO-saturated 10-4M nitrite, 8 ms after a 16-had pulse, as a function of pH. pH 5-8 were buffered with l t 3 M phosphate, and pH 9 was buffered with 10-3 M borax. A pK.(ONOOH) = 6.5 0.1 was evaluated from the curve.
throughout the pH range. This feature is observed directly at the characteristic maximum absorption around pH 7 which increases with increasing buffer concentration (Figure 5 ) . Due to a faster protonation of HN03- to H2NO3 by the buffer with increasing buffer concentration (reaction 11) and thereby a faster decay into NOz, a higher yield of ONOO- is obtained (reaction 3). The experimentally obtained absorptions at 300 nm of ONOO- in the nitrate system as a function of pH are perfectly reproduced by model calculations using pKa = 6.5 and the rate constants determined previously in this work. Table I11 shows the literature pKa values for peroxynitrous acid. The values have been found either by absorbance mea-
L~gagerand Sehested
6668 The Journal of Physical Chemistry, Vol. 97, No. 25, 1993 0.08
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Generally, determination of pK, values based on kinetic measurements of ONOOH and O N W decay will always be more uncertain than pK, curves baaed on absorption measurements in a clean system. Two reasons for this include (1) potential photolysis of peroxpitrite anion during decay (which we have observed, see later) would result in pK, values that are too high and (2) the decay mechanism may be dependent on the system used, such that reactions that accelerate the rates observed might be present (e.g., reactionwith buffersor catalysts),againresulting in high pK, v a l ~ e s . ~Thus ~ ~ -we ~ consider the pK, = 6.5 f 0.1 obtained from our absorbancemeasurementsin the nitrite system to be more accurate than the pK, = 7.0 f 0.3 found from our kinetic measurements(seebelow). Keith and Powell5 have found a pK, of 6.6 which is in agreement with our value, using a kinetic determination which is in accordancewith our rate equation (see below). They only give decays up to pH 9 due to buffer complications. The pK, value of 5.3, determined by GrHtzel et al,,l7 may be rejected for the followingreasons. Although it is based on O N W absorption as a function of pH, their absorption measurement is taken 250 ps after the end of the pulse, at which time both NO2 and N2O4 are still pressnt. Their concentrations are far greater than the concentration of ONOO-, which in spite of the much smaller extinction coefficients26 make the NO2 and N204 absorption a substantial part of the total absorption at 300 nm, 250 ps after the pulse. Additionally,they use the nitrate system for their determination, and as we have shown, this system is much more complicatedbecauseoftheN03%termediate, which gives pH-dependent yields. Wagner et al.16 base their determination of pK, on kinetic measurements of the isomerizationof ONOOH in the pH range 2.4-4.0. Their pK, = 5.3, which lies well out of the studied pH range, is found by fitting the data to a rate equation. Since their data lie on the plateau part of the pH decay curve (Figure a), it is pertinent to say that their determination of pK, is very uncertain. The two determinations of pK, = 6.0 are both observed in systems using flash photolysis of nitratel3J4 and both are found by absorption measurements of ONOOH and/or O N W (at 245,290, and 320 nm, respectively). The absorptionsare found "after flashing"l3 and, in the case of Barat et al.,l4 the time after
0.07
0.06 u
0
2 2
0.05
2
0.04
B
x
c .CI (
E a
0.03
2
0.02
0.00
I
PH Figure 5. The yield of ONOO- at 300 nm, observed in Ar-saturated 5 x 10-3 M nitrate, 8 rrm after a 22-krad pulse, as a function of pH. pH 5-8 were buffered with (0)lk3M and ( 0 )5 X 10-3 M phosphate, and pH 9 was buffered with 5 X lW3 M borax. (V) 5 X 10-3 M buffer,with a 3-r(s, 40-krad pulse, normalized to pH 3. Solid and dashed lines are modeled (see text).
TABLE I& Literature Valuea for p&(ONOOH) .ad the Teehdquea .ad System Used for Each Detemhtioa PK
system
reference pK 17 6.6 16 6.8 13 6.8 6.0 FP,nitrate 14 7.49 6.5 PR, nitrite (nitrate) this work
