J. Phys. Chem. 1992, 96, 265-272
265
Aqueous Solubility and Reaction Kinetics of Hydroxymethyl Hydroperoxide Xianliang Zbou and Yin-Nan Lee* Environmental Chemistry Division, Department of Applied Science, Brookhaven National Laboratory, Upton, New York 1 1 973 (Received: June 21, 1991)
The aqueous-phase equilibria and kinetics of the formation of hydroxymethyl hydroperoxide(HMP) and of bis(hydroxymethy1) peroxide (BHMP) from formaldehyde and hydrogen peroxide, i.e., H2C0+ H202a HOCH2O2H(1, -l), and HOCH2O2H H2C0e HOCH2OZCH20H(2) were studied using an amperometric technique which is highly sensitive and selective for H202. The equilibrium constants of reactions 1 and 2 between 5 and 35 OC were determined to be K 1 = 2.35 X exp(2610/7') M-' and Kz = 1.04 X exp(2780/T) M-I, respectively, both independent of pH between 4.0 and 8.4. The rate coefficients of (1) and (-1) determined at pH 7.07 0.02 between 5 and 35 "C are kl = 6.0 X 10" exp(-9450/T) M-I s-l and k-l = 1.O X 1015exp(-ll8OO/r) s-l. Both kl and k-l are hse-catalyzed and are linearly dependent on pH between 4.0 and 8.2, namely, k-, = (5.0 0.3) X l0-lo/[H+] s-I at 25.8 f 0.1 OC. The Henry's law constants of HMP and BHMP, determined by measuring their corresponding gas and aqueous concentrations at phase equilibrium, are 5.02,fX lo5 and 6:; X lo5 M atm-I at 22.0 f 0.1 OC, and 6.22,; X 10s and 20fi8 X lo5 M atm-' at 10.0 0.1 OC, respectively. The reaction kinetics of HMP with S(1V) was studied by a competition technique using the H202-S(IV) reaction as the reference; the reaction is acid-catalyzed, with an effective second-order rate constant of (2.2 X 107)[H+]f 15% M-I s-I for the pH range 3-4 at 22.0 f 0.1 "C. These results indicate that gas-phase HMP in the atmosphere is efficiently removed by wet scavenging processes and would be quantitatively detected by peroxide instruments involving gas-liquid scrubbers, provided that the scrubbed HMP is stabilized. The time constant of the dissociation of dissolved HMP to Hz02is fairly short, being 100 min at pH 5.5, shorter at higher pH. Consequently, HMP is expected to be stable and detected in atmospheric liquid water only at pH 1 5 . 5 .
+
*
*
-
Introduction Organic peroxides along with other photooxidants in the atmosphere are produced from oxidation of hydrocarbons initiated by OH radicals1 and 03,1-4 with intricate involvement of oxides of nitrogen as catalyst^.^ The detection and characterization of atmospheric organic peroxides therefore constitute an important task for gaining insight into mechanisms by which hydrocarbons are oxidized and free radicals are generated and destroyed. Among the several organic peroxides identified in the atmosphere from analysis of air and precipitation samples (e.g., CH302Hand C H 3 C ( 0 ) 0 2 H ) hydroxymethyl hydroperoxide (HMP, HOCH202H) has also been d e t e ~ t e d ,with ~ . ~ a maximum concentration approaching that of H202,Le., at parts per billion levels, under strong photochemical condition^.^ This peroxide has also been found to be present in leaves of isopreneemitting plants after exposure to 03,8 indicating a source in the reactions of O3and naturally emitted hydrocarbons. Because of its phytotoxicity and mutagenicity,+I' HMP is thought to be an important contributing factor to forest decline! In addition, since organic hydroperoxides are capable of oxidizing S(IV),I2the contribution of HMP in this regard must be quantified. In order to assess the budget and the atmospheric importance of HMP, the physical and chemical properties governing its atmospheric lifetime must first be determined. Since scavenging by cloud and precipitation is an important process by which gaseous species are incorporated into atmospheric water and removed from the atmosphereI3-l6 we report in this paper the de(1) Hanst, P. L.; Gay, B. w., Jr. Atmos. Enuiron. 1983, 17, 2259. (2) Martinez, R. I.; Herron, J. T.; Huie, R. E.J . Am. Chem. Soc. 1981, 103, 3807. (3) Gab, S.; Hellpointner, E.; Turner, W. V.; Korte, F. Nature 1985, 316, 535. (4) Becker, K.H.; Brockmann, K. J.; Bechara, J. Nature 1990,346,256. ( 5 ) Kasting, J. F.;Singh, H. B. J. Geophys. Res. 1986, 91, 13239. (6) Hellpointer, E.; Gab, S. Nature 1989, 337, 631. (7) Hewitt, C. N.; Kok, G. L. J . Atmos. Chem. 1991, 12, 181. (8) Hewitt, C. N.; Kok, G. L.; Fall, R. Nature 1990, 344, 56. (9) Sobel, F. H. Nature 1956, 177, 977. (10) Marklund, S. Acta Chem. Scand. 1971.25, 3517. (11) Marklund, S. Arch. Biochem. Biophys. 1973, 154,614. (12) Lind, J. A.; Lazrus, A. L.; Kok, G. L. J. Geophys. Res. 1987, 92, 4171. (13) Levine, S . 2.;Schwartz, S . E.Atmos. Enuiron. 1982, 16, 1725. (14) Chameides, W. L. J . Geophys. Res. 1984,89,4739.
0022-3654/92/2096-265$03.00/0
termination of the Henry's law solubility,the kinetics of its aqueous decomposition, and the reaction with S(1V) that influence its rate of removal from the atmosphere and its subsequent fate in the aqueous phase. These aqueous properties are also important for the evaluation of the collection efficiency and speciation accuracy of peroxide sampling techniques involving scrubbing of gases by aqueous solution^.'^ In aqueous solution, H M P decomposes into and reaches an equilibrium with hydrogen peroxide and formaldehyde (principally in its hydrated form, methanediol): k
HzCO + H2O25.1.HOCH202H k-1
(1, -1)
where kl and k-l are the rate constants of the forward and reverse reactions, respectively. In this chemical system, H M P reacts further with H2C0to reach another equilibrium producing bis(hydroxymethyl) peroxide (BHMP), Le.,
HOCH202H
+ H2CO K2 HOCH202CH20H
(2) where K2is the equilibrium constant. It may be pointed out that the formula HzCO is taken in this paper as the total analytical formaldehyde, Le., hydrated plus unhydrated. The aqueous kinetics and equilibria of reactions 1 and 2 have previously been s t ~ d i e d ' ~but J ~ Jat~fairly high reagent concentrations, Le., at tens of mM to 1 M, as necessitated by the approach taken where BHMP and mixtures of BHMP and H M P were employed as starting reagents. Consequently, these studies were complicated by the fact that at high concentrations BHMP undergoes further reactions to produce H C 0 2 H and H2. Since the concentrations of HzCOand H202found in cloudwater and rain are fairly low, typically tens of pM and less,w23studying the kinetics of reactions H
(15) Jacob, D. J . Geophys. Res. 1985, 90D, 5864. (16) Schwartz, S . E. In Chemistry of Multiphase Atmospheric Systems; Jaeschke, W., Ed.; Springer-Verlag: Berlin, 1986. (17) Lazrus, A. L.; Kok, G . L.; Lind, J. A.; Gitlin, S . N.; Heikes, B. G.; Schetter, R. E.Anal. Chem. 1986, 58, 594. (18) Dunicz, B. L.; Perrin, D. D.; Style, D. W. G. Trans Faraday SOC. 1951, 47, 1210. (19) Benner, W. H.; Bizjak, M. Armos. Enuiron. 1988, 22, 2603. (20) Kelly, T. J.; Daum, P. H.; Schwartz, S . E.J. Geophys. Res. 1985, 90, 7861. (21) Lee, Y.-N.; Shen, J.; Klotz, P. J. Water, Air Soil Pollut. 1986, 30, 143.
0 1992 American Chemical Society
266 The Journal of Physical Chemistry, Vol. 96, No. 1, 1992
Zhou and Lee
1 and 2 at micromolar concentrations not only offers the advantage that decomposition of BHMP other than reaction 2 is minimized, but also provides kinetics data that can be applied to atmospheric situations without great extrapolations. A major difficulty encountered in the study of reactions 1 and 2 arises from the speciation of the coexisting peroxides, Le., H,02, HMP, and BHMP. For example, Marklund'O used a colorimetric technique involving the complexation of Ti(1V) and H202in acidic medium for the determination of H202. HMP and BHMP were determined as H202by the same technique after conversion under basic conditions. This technique, however, suffers from its nonspecificity because both HMP and BHMP react with Ti(IV), albeit at lower rates. Corrections of the contributionsfrom these reactions and assumptions of extinction coeficient of the complexes of Ti(1V)-HMP were therefore necessary. This difficulty is circumvented in the present study by use of an amperometric technique, sensitive only to H202,to follow the concentration of H202continuously in real time. The high sensitivity and selectivity as well as the fast response of this technique facilitated the study of reactions 1 and 2 in the micromolar concentration range. Experimental and Methodology Section Materials. Reagent grade H202(30%)and H 2 C 0 (37%) were obtained from Mallinckrodt. Volumetric standard solutions of HCl and NaOH were obtained from Aldrich. Reagent-grade anhydrous Na2HP04,CH,CO,Na (both from MCB, Cincinnati, OH), and tris(hydroxymethy1)aminomethane (tris, Sigma Co.) were used for preparing pH buffers. Reagent grade KCI (Fisher) was employed for maintaining ionic strength. Na2S03was obtained from Aldrich. All of the reagents were used without further purification. Ultra high purity N2 (UHP, 99.999%) used in the solubility measurements was obtained either from Lind or Liquid Carbonic. A ca. 1 M formaldehyde stock solution, free of alcohol preservative, was prepared from paraformaldehyde (Eastman Kodak) following Benner and Bizjak.19 Millipore Milli-Q water (resistivity 118 Mohm at 25 "C) was used to prepare all the solutions. Concentration Calibrations. The concentration of the commercial stock formaldehyde solution was standardized by a sulfite-iodometric technique24and was found to be within 5% of the specified value. The concentration of the 30% H 2 0 2solution was determined by titrating against standard sulfite solution using the amperometric technique (see below). The concentrationsof sulfite solutions used in this regard were standardized by an iodometric technique.24 pH and Temperature Controls. The pH of reaction mixtures was controlled by typically 1-5 mM of the following buffers: acetate for pH 4-5.5,phosphate for pH 6-8,and tris for pH 8. HCl was used as a concentration buffer for pH below 3.3. The pH of the solutions was measured by a research-grade pH meter (Beckman, Model 4500) equipped with a combination electrode (Orion, Ross electrode), which was calibrated before use. The temperature of reaction mixtures was maintained constant within fO.l OC using a circulating water bath (Neslab, Model RTE-9). H202Measurement. An amperometric technique was applied to the measurement of H 2 0 2in the micromolar concentration rangesz5 The commercial instrument used,oxidase meter (Model 25, Yellow Springs Instrument, Inc.), consists of a concentric Pt-Ag electrode pair, a power supply, a current measurement circuitry, and a recorder output. The principle of operation involves the application of a 700-mV bias voltage across the electrodes, causing oxidation of H 2 0 2to 0,on the platinum anode and reduction of O2 to H20on the silver cathode. The current thus generated is linearly proportional to [H20,] up to -50 pM; the detection limit of the device is -0.1 pM ( S I N = 3). The
-
(22) Gunz, D. W.; Hoffmann, M . R. Atmos. Enuiron. 1990, 24A, 1601. (23) Sakugawa, H.; Kaplan, I. R.; Tsai, W.; Cohen, Y. Enuiron. Sci. Technol. 1990, 24, 1452. (24) Intersociety Committee. Methods of Air Sampling and Analysis; American Public Health Association: Washington, D.C., 1972. (25) Lee, Y.-N.; Shen, J.; Klotz, P. J.; Schwartz, S . E.; Newman, L. J . Geophys. Res. 1986, 910, 13264.
I
time
Figure 1. The kinetics of the formation of HMP from the reaction of H202with H2C0as monitored by following the decrease of H202as a function of time: [H2COl = 7.5 mM,[H202]~= 10.0 rM,pH = 7.07, = 0.02, and T = 35.0 "C. The first-order plot of the data is shown in the inset.
sensitivity to H 2 0 2is slightly dependent on pH, increasing by a factor of about 3 from pH 3 to pH 8;detailed calibration characteristics have been reported previously.26 The response of the instrument is fast, with time constant shorter than 0.1 s. The amperometric device responds also to S(IV) and H,CO, however, at much diminished sensitivities. The sensitivity to HzCO is about 3 orders of magnitude lower than that to H202. The sensitivity to S(IV) decreases with increasing pH with the relative sensitivity to H 2 0 2and to S(IV) varying from nearly 2:l at pH 3 to about 20:l at pH 5.26 The response of the amperometric device to HMP and BHMP was examined and found to be I 18 of that to H202as established by the fact that the signal responding to a 10 rM H 2 0 2solution (pH 7.2) diminished to -8% of the initial value when a major portion of the HzOz was converted to HMP and BHMP by the addition of HzCO (to yield 40 mM). This remaining signal was found to be due to the free H202in equilibrium with the organic peroxides, consistent with values calculated by the equilibrium constants of reaction 1 and 2 determined in this work (see the kinetics and equilibrium measurement sections for details). The concentrations of H M P and BHMP were determined as H,O, after decomposition of these organic peroxides by dilution ( S O 100-fold); H202becomes the predominant peroxide under the new equilibrium conditions. It may be pointed out that since the amperometric device responds only to H202and not to the organic peroxides, absolute calibration of the device against Hz02is often unnecasary. While this is certainly the case for studying kinetics that conform to first-order behavior, it also offers an added advantage in the equilibrium measurements where only the ratio of H2O2to total organic peroxides is needed, the latter quantity is equal to the lost H 2 0 2signal. Any minor changes in calibration arising from, for example, slight pH shifts, were internally cancelled out. Finally, the ability of the oxidase meter to monitor H 2 0 2continually in real time makes the kinetics studies in this work more feasible and reliable. KinetirsMesslaecllenta FornratiOnandDecompositioaofHMP. The reaction kinetics between H202and H 2 C 0 (in large excess) was studied under pseudo-first-order conditions by following the decrease of H 2 0 2with time after mixing of the two reagents. In (26) Shen, J.; Lee, Y.-N.Brookhauen Naiional Laboratory Report, BNL-52013, Brookhaven National Laboratory: Upton, NY, 1986.
The Journal of Physical Chemistry, Vol. 96, No. 1, 1992 267
Solubility and Kinetics of Hydroxymethyl Hydroperoxide
where [H2O2ITis the total peroxide concentration, a quantity known from the initial amount of H 2 0 2used and the extent of dilution involved. Reaction of HMP and S(IV). The kinetics of the reaction of HMP with S(IV) (= S02.H20 HSOY + S032-) was studied by a competition kinetics method using the H202-S(IV) reaction as a reference. When S(1V) is added into a mixture of H202 and HMP, the following reactions take place concurrently:
innc
l
-
+
$1
,
,
,
,
,
,
,
I
1
0
50
100
150
'i 200
H 2 0 2+ S(1V) and
250
HMP
t, s
time
Figure 2. The kinetics of the decomposition of HMP initiated by dilution (by X100) as monitored by following the inrease of H202as a function of time: pH = 7.07, p = 0.02, and T = 22.0 OC. The mother solution contained [H2COIT= 10.0 mM, [I1202]7.= 1.0 mM. The first-order plot of the data is shown in the inset.
a typical run, the Pt-Ag electrodes were first conditioned in a 20-mL portion of H 2 C 0 solution (ca. 2-20 mM, maintained at a desired pH) in a thermostated cone-shaped glass reaction vessel (volume -50 mL) under continuous stirring. When temperature equilibration was reached and a stable background signal registered, the solution was added with a small portion ( 50 pL) of H202stock to give an initial [H202]of ca.10 pM and the decrease of the signal corresponding to the loss of H 2 0 2was continuously recorded (Figure 1). The rate of H 2 0 2 disappearance under pseudo-first-order conditions, Le., [H2CO] remains unchanged, is given by: -d[H202]/dt = kl'[H202] - k-,[HMP] (3)
-
where k,' = kl[H2CO]. Upon integration, a first-order rate expression is obtained as:
In ([H2O2If- [H2O2lm)= -k,d
+C
(4)
where subscripts 0, t and m denote time 0, t , and infinity, C = In ( [H202]o-[H202]m), and kobs is the apparent first-order rate constant defined as: koh = k-I + kl[H2CO] (5) By plotting kobsagainst [H2CO], one obtains kl and kl as the slope and intercept, respectively. However, since the magnitude of klis small compared to k,' and is subject to large uncertainties under the experimental conditions employed, this technique was used primarily for the determination of k l . To determine k-l, we followed the decomposition of HMP (reaction -l), by monitoring the build up of H 2 0 2as a function of time. In this approach, HMP was prepared in a solution containing millimolar levels of H 2 0 2and H2CO: with [H2CO] = 5 mM, -50% of the total peroxide is present in HMP. To initiate a kinetic run, a small portion of this HMP stock solution (typically 200 pL) was added to the reaction cell containing a buffer solution (20 mL), resulting in a 100-fold dilution of the HMP stock. The increase of [H202]accompanying the decomposition of HMP in reaching a new equilibrium was continuously monitored. A typical trace of such a kinetic run is shown in Figure 2. Since BHMP constitutes only a few percent of the total peroxide in solutions containing [H2CO] I10 mM and the decomposition of HMP is slow compared to that of BHMP,l0 any BHMP initially present in the HMP stock would instantaneously dissociate to HMP and H 2 C 0 upon dilution. Furthermore, in the diluted solution where the concentrations of peroxides and H 2 C 0are in the range of tens of micromolar, decomposition of HMP goes essentially to completion and the reaction is effectively irreversible. Consequently, the rate law of HMP decomposition is simply d[HzO,]/dt = k-1 [HMP], = k-1([H2021T - [H2021f) (6)
+ S(1V)
k7
ks
products
(7)
products
(8)
At constant pH, the rates of reactions 7 and 8 can be expressed as 4[H2021 /dt = k7[H2021 [S(Iv)l
(9)
-d[HMP]/dt = ks[HMP][S(IV)]
(10)
and It may be noted that the assumptions of a pH dependence of the reaction kinetics and a fmt-order dependence on the concentration of HMP in eq 10 are based on the well established kinetic behavior of reaction 725327and the S(IV)-organic peroxide reactions.12 Dividing eq 10 by eq 9 followed by integration, we obtain
In ([HMPly/[HMPli) = ( k s / h ) In ([H202l~/[H202li) (1 1) where the subscripts i andfrepresent initial and final. The value of k8 is calculated from the slope of eq 11 and the known value of k7. In a typical experiment, a small portion of Na2S03stock solution was added to an acidic solution (pH 3-4, maintained by either HCl or acetate buffer) containing 20-200 pM total peroxides. The initial concentration of S(1V) was typically half of that of the total peroxides. [H202],and [HMP],were measured with the amperometric method as soon as the reactions were complete, typically within 10 min. At the low pH of the study, the rate of exchange between H 2 0 2and HMP was sufficiently slow (time constant -hours) that these two species are essentially independent of each other within the time of reaction (time constant 2 min). Equilibrium Measurement of the H202-H2C0System. Since both HMP and BHMP are produced in mixtures of H 2 0 2and H 2 C 0 and cannot be individually quantified, the determination of the equilibrium constants of reaction 1 and 2, KI and K2,was achieved by measuring an apparent equilibrium constant Kap which relates the equilibrium concentrations of H$O, H202, an8 the total organic peroxides, [OP], (=[HMP], + [BHMP], = [H2021T - [HZ021-)r by Kapp = [OPI =/([HZ021 m[HXOIm) (12)
-
and its dependence on [H2CO],. Substituting KI and K2, which are defined as KI = [HMPl=/([HzOzl m [H2COl=)
(13)
K2 = [BHMP],/ ([HMPJm[H2CO],)
(14)
and into eq 12, we obtain Kapp= KI + KlK2:[H2C0lm
(15)
The values of KI and K2can then be determined from the intercept and slope of the plot of Kappagainst [H2CO],. To determine Kapp,a small volume of H 2 0 2stock solution (typically 50 pL, to result in a pM concentration) was added to a formaldehyde solution (typically 20 mL, in mM concentration, maintained at a desired pH) identical to the procedure employed (27) McArdle, J. V.; Hoffmann,M. R.J . Phys. Chem. 1983.87, 5425.
268
The Journal of Physical Chemistry, Vol. 96, No. 1, 1992
Zhou and Lee
in the kinetics studies, and the reaction was allowed to proceed until the H z 0 2signal had stabilized (Figure 1). The equilibrium concentrations of H202,[H202],, and the total organic peroxides were determined from the final oxidase meter signal and the difference, [H2O2lT- [H202],, respectively. In fact, many of the experimental runs for the determination of Kappwere the very ones performed for the determination of k l . Henry's L a w Solubility Measurement. The Henry's law solubilities of H202,HMP, and BHMP were determined by measuring the corresponding gas- and aqueous-phase concentrations of these species under gas-liquid equilibrium, and relating them according to Hx = [Xl/PX
= (IHMP] + (BHMPl)/@HMP + PBHMP)
(17)
Substituting eqs 13, 14, and 16 into eq 17, we obtain HOP= HiHz(1
+ K2[H2COI)/(Hz + HiK*[HzCOl)
0 01
(16)
where H is Henry's law constant, p is partial pressure, and the subscript X denotes the chemical species. The experimental setup for Henry's law solubility measurement is similar to those reported before.28 Equilibrium gas-phase peroxides were generated by passing a N2flow through a bubbler gas-liquid reactor containing a solution of known concentrations of the peroxides (typically at millimolar level). The Pyrex bubbler gas-liquid reactor was equipped with a water jacket for temperature control and a fritted-disk bottom which allows gas to enter into the solution as finely divided bubbles. The large surface to volume ratio, a, gives rise to a highly efficient mixing between the two phases. The stochastic mixing time constant, 7 , (or l/kLa, kL is the liquid side mass transfer coefficientz9),ranged from 1 to 5 s (liquid volume VI= 10-50 mL and total gas flow rate F = 2 L/min) and were sufficiently short to assure that gas-liquic! .~~ equilibrium of a dissolved species was e s t a b l i ~ h e d . ~ ~The concentration of the gaseous peroxides in the N2effluent gas were then determined by scrubbing the gas using a scrubber containing a known amount of water attached downstream of the bubbler. The Pyrex scrubber (ASTM D1607-69)31consists of a delivery tube with a sintered frit at the end concentrically inserted to the bottom of a test tube (- 15 cm long) which has an enlarged mid-section to allow for bubble breaking. To improve the scrubbing efficiency, the temperature of the scrubbing solution was maintained at 0 OC using an ice bath. Since HMP cannot be isolated as a pure compound and is in equilibrium with H 2 0 2and BHMP, the determination of the Henry's law constant of HMP requires the delineation of the contribution of the other two peroxides, H z 0 2and BHMP. The concentrations of each individual peroxide, i.e., H202,HMP, and BHMP, in the source solution were determined by the technique described in the equilibrium measurement section. However, it was difficult to quantitate the concentrations of the individual gaseous peroxides collected in the scrubber because of the low concentrationscollected. As a result, the three. peroxides collected in the scrubber were analyzed as H 2 0 2after rapid conversion of the organic peroxides to H 2 0 2in pH 7 buffer solutions upon dilution (typically by 5-fold). The contribution of H 2 0 2to the total peroxides collected in the scrubbing solution was determined from H 2 0 2concentration of the source solution and its Henry's law constant independently determined in this work. The total partial pressure of the peroxides, pT,after correcting for pHrOl,is related to the total concentration of the organic peroxides of the source solution by an operational Henry's law constant by HOP
100
(18)
where H I and H 2 are the Henry's law constants of HMP and (28) Park, J.-Y.; Lee, Y.-N. J . Phys. Chem. 1988, 92, 6294. (29) Danckwerts, P. V. Cos-Liquid Reactions; McGraw-Hill: New York, 1970. (30) Lee, Y.-N.; Schwartz, S. E. J . Phys. Chem. 1981, 85, 840. (31) Public Health Service Publication, "Selection Methods for the Measurement of Air Pollutants," 1965, No. 999-AP-1l .
0.02
0.03
[HKOI-, M
Figure 3. The dependence of the apparent equilibrium constant Kappon [H2CO],. The straight line represents a least-squares best fit of the data. Conditions: [H2O2IT= 10.0 pM, pH = 7.07, p = 0.02, T = 22.0 "C.
TABLE I: Summpry of the Equilibrium Constants between H20b H,CO, HMP,and BHMP (pH 7.07 and p = 0.02)' T, OC K , , M-I K,,bM-' K,. M-' 5.0 10.0 22.0 25.0 35.0
275 24 1 163 157 111
266 256 149 163 124
26.5 16.6 12.7 10.8 9.6
a K I and K2 are equilibrium constants of reactions 1 and 2 as defined by eqs 13 and 14 and are obtained from eq 15 and Figure 3. n 1 3 for all values and the standard deviations 15% for K,, -10% for K2. bDetermined as k , / k _ , ,see the kinetics study section.
BHMP, respectively. The value of HI is determined by extrapolating HOPto [H2CO] = 0. Alternatively, eq 18 can be rearranged to (1 + K2[H2COI)/HOP = 1/H, + (K,/H,)[H2COI
(19)
By plotting (1 + K2[H2CO])/HoPagainst [H2CO],one obtains l/H1 as the intercept and K 2 / H 2as the slope, respectively. In a typical experimental run, a source solution was prepared by mixing H202and H 2 C 0 to yield 5-30 mM total peroxide and 5-30 mM total H2C0. The pH of the solution was maintained at 7.2 by 5 mM phosphate buffer to allow for a rapid equilibration among all the species (time constant -2 min). A portion (ca. 50 mL) of this source solution after standing for -30 min was introduced into the bubbler while a flow of prehumidified N2gas (1.OO f 0.05 L/min) had already been established. After temperature equilibration was reached (510 min) and the system conditioned for -30 min, a scrubber containing 10 mL of either water or 1 mM HCl solution was connected downstream of the bubbler to collect the gaseous peroxides. The typical collection time was 20-60 min. Because of the high solubilities of the peroxides, care was taken to ensure the wall of the tubings connecting the bubbler and the scrubber was free from any water droplets that might remove gaseous peroxides. The collection efficiency of the scrubber was examined by connecting a second scrubber in series. Since the amount of peroxides in the second scrubber was
