Environ. Sci. Technol. 1999, 33, 2561-2565
pH Dependent Photoformation of Hydroxyl Radical and Absorbance of Aqueous-Phase N(III) (HNO2 and NO2-) T A K E M I T S U A R A K A K I , * ,† TAKAYUKI MIYAKE,‡ TSUYOSHI HIRAKAWA,† AND HIROSHI SAKUGAWA§ The Center for Forest Decline Studies, CREST, Japan Science and Technology Corporation (JST), Hiroshima Techno Plaza 308, 3-13-26 Kagamiyama, Higashi-Hiroshima, Hiroshima 739-0046, Japan, and Graduate School of Biosphere Sciences and Faculty of Integrated Arts and Sciences, Hiroshima University, 1-7-1 Kagamiyama, Higashi-Hiroshima, Hiroshima 739-8521, Japan
Ultraviolet-visible (UV-Vis) absorption spectra of aqueous-phase N(III) (HNO2 and NO2-) were studied for the typical pH ranges observed in the atmospheric waters. The molar absorptivity of HNO2 is larger than that of NO2- in UV-A regions. The N(III) molar absorptivity at a specific pH can be determined based on the molar absorptivities and acid-base equilibrium of HNO2 and NO2- (pKa ) 3.27). Photochemical formation of hydroxyl radicals (OH radicals) was also studied in aqueous solutions of N(III) for pH values between 1.9 and 6.2 (seven pH values). The OH radical photoformation rate constants showed a distinctive pH dependence, approximately 10fold higher at pH ) 1.9 than at pH ) 6.2. The pH-dependent OH radical photoformation also followed the speciation pattern of HNO2 and NO2-. Hydroxyl radical photoformation rate constants ((SE) were estimated to be (3.1 ( 0.08) × 10-4 s-1 for HNO2 and (3.2 ( 0.61) × 10-5 s-1 for NO2for vernal equinox solar noon conditions at 34° N (HigashiHiroshima), using the least-squares best-fit analysis. We further performed a model calculation to elucidate the significance of N(III) photolysis in the atmospheric hydrometeors, assuming that the dissolution of gaseous HONO was the only source of N(III). We found that photolysis of aqueous-phase N(III) could play a significant role in initiating oxidation reactions in atmospheric hydrometeors (i.e., dew) exhibiting higher pH and N(III) concentrations.
Introduction Gaseous HONO absorbs ultraviolet regions (98%), and phenol (99%) were obtained from Katayama Chemical. All solutions were prepared using Milli-Q pure water (Millipore, >18 MΩ cm, air saturated). Measurement of Absorbance of N(III) at Different pH Values. The pH of N(III) solutions was adjusted by adding appropriate amounts of sulfuric acid. No acid was used for pH ) 6.2. The solution pH was measured before and after the photochemical experiments using a glass electrode (Horiba 6366-10C) that was calibrated with pH ) 1.68, 4.01, and 6.86 standard solutions (Katayama Chemical). There was little pH change before and after the photochemical experiments. UV-Vis absorption spectra of N(III) solutions (0.00125-0.0100 M) at specific pH values were recorded by Shimadzu UV-2400PC (double-beam, 2 nm bandwidth) using a 1-cm quartz cuvette cell with a Teflon stopper at room temperature. For the absorbance measurements, the instrument was zeroed using the sulfuric acid solution at the pH of interest (in Milli-Q pure water), and the sulfuric acid solution was used as a reference. Thus, the absorption spectra do not contain the effects of sulfate ions in the N(III) solutions at any pH. No attempt was made to maintain a constant ionic strength in the solutions, and all the equilibrium constants were used without adjusting for ionic strength. All the absorption spectra were recorded within 5 min of sample preparations in the volumetric flasks. Considering the physical Henry’s law constant for HONO (KH ) 49 M atm-1) (1) and since the cuvette cell had a Teflon stopper and small headspace, little HONO was expected to escape from the solution even at pH ) 1.1 (which can be seen from eq 9 with high liquid water content). Photochemical Experiments. Photochemical experiments were carried out using a solar simulator (Oriel model 81160-1000, Oriel Corp.) with a 300-W Xe lamp (ozone free, model 6258, Oriel Corp.). To simulate actual solar irradiance, wavelengths less than 300 nm were filtered out by two optical filters (Oriel AM0 and AM1.0, Oriel Instruments). The photochemical reactions were performed in a custom-made quartz cuvette cell (Workshop for Advanced Techniques, Hiroshima University). The cylindrical quartz cell was 3 cm in diameter, 1 cm in thickness, and 6.8 mL in capacity. The cuvette cell was designed to be gas-tight by using a silicone cap with a Teflon syringe needle adapter (Z11731-5, Aldrich). During the photochemical experiments, the sample was stirred by a magnetic stirrer covered with Teflon, and the temperature was kept at 20 °C by a water circulator (RTE111, NESLAB). The daily actinic flux was determined by chemical actinometry (2-nitrobenzaldehyde, for simplicity “2-NB”) using the same quartz cuvette that was used for the photochemical experiments. The 2-NB photolysis rate constants for the solar simulator ranged from 0.00508 to 0.00656 s-1. The vernal equinox solar noon conditions of the 2-NB disappearance rate constant was theoretically calculated to be 0.00867 s-1 using quantum yields and molar absorptivities of 2-NB (21-23) and actinic flux at 34° N (24). Experimentally determined actinic flux at 36° N (ground level) was reported to be 0.0101 s-1 for autumnal equinox solar noon conditions (25), which agreed well with our estimated value. For comparison, all the photochemical data were normalized to the vernal equinox solar noon conditions of HigashiHiroshima (34° N), following the same procedure of Faust and Allen (21). It should be noted that actinic flux may significantly vary from dawn to sunset, thus our normalization to the solar noon conditions may overpredict photoformation of OH radicals. 2562
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FIGURE 1. Ultraviolet-visible (UV-Vis) absorption spectra of N(III) (HNO2 and NO2-) solutions at pH values between 2.14 and 6.45 (four pH values). The pH values were adjusted by adding appropriate amounts of sulfuric acid, and no buffer was used. Measurement of OH Radicals. For a detailed description of detection techniques for OH radicals and control experiments, please refer to the refs 16 and 21. In short, a photochemically formed OH radical is trapped by benzene (1.2 mM) added to the reaction solutions. The reaction between benzene and OH radical forms phenol as a product (16). In this study, phenol concentrations were determined by a high-performance liquid chromatography (HPLC) by detecting ultraviolet absorbance at 220 nm. For separation, a Superco Sil C-18 reverse phase 5 µm, 250 × 4.6 mm column was used. The mobile phase was 50% CH3CN/50% H2O (v/ v), and the flow rate was 1.0 mL min-1. For the reaction between benzene and OH radical, the yield of phenol was reported to be 75 ( 15% for a pH range of 2.0-5.0 (16), thus photochemically formed OH radical concentrations ([OH radical]) were determined by
[OH radical] ) [phenol]/0.75
(2)
It should be mentioned that the OH radical concentrations reported here are total concentrations of OH radical formed during the given illumination time and are not steady-state concentrations.
Results and Discussions Absorption of N(III) at Different pH Values. Figure 1 shows the ultraviolet-visible (UV-Vis) absorption spectra of N(III) (HNO2 and NO2-) at different pH values. Absorbance of N(III) shows significant pH dependence at wavelengths above 326.5 nm. At wavelengths below 326.5 nm, N(III) absorption is lower for lower pH values. Absorbance of NO2- near the UV region is ascribed to two electronic transitions, namely, n f π* and π f π* (26). Strickler and Kasha (26) studied solvent effects on N(III) absorbance and found that absorption spectra were quite different for polar and nonpolar solvents. Thus, absorbance of HNO2 and NO2- are expected to be quite different because of the proton in HNO2. Molar absorptivity of NO2- is well documented (26, 27), but only a few studies have reported molar absorptivity of aqueous-phase HNO2 (28, 29). Molar absorptivity of aqueous HNO2 was determined at pH ) 1.1 (adjusted with H2SO4, ionic strength ) 0.09 M). Since the equilibrium constant for nitrous acid dissociation (HNO2 S H+ + NO2-) is 10-3.27 M (30), 99.3% is HNO2 at pH ) 1.1. Thus, absorption spectra at pH ) 1.1 can be regarded as being due exclusively to HNO2, and we were able to determine the molar absorptivity of HNO2 at specific wavelengths. Molar absorptivity of HNO2 at specific wavelengths can be found in the Supporting
FIGURE 2. Observed and calculated molar absorptivity of aqueousphase N(III) at pH ) 2.1 and 3.9. Molar absorptivities at pH ) 2.1 and 3.9 were calculated based on molar absorptivities of HNO2 and NO2- and the acid dissociation constant of nitrous acid (pKa ) 3.27) (30). Equilibrium constants were not corrected for the ionic strength. See details in the main text. Information. Our study found that the shape of HNO2 absorbance was very similar to the data presented in Finlayson-Pitts and Pitts (1) but with some difference in magnitude. The maximum molar absorptivity was found to be 51.9 M-1 cm-1 at 371 nm in this study, which was approximately 70% of the value in Graedel and Waschler (28) and 120% of that in Fischer and Warneck (29). Molar absorptivity of N(III) at various pH values (N(III),λ) were fit to the following speciation equation:
N(III),λ ) fHNO2HNO2,λ + fNO2-NO2-,λ
FIGURE 3. Hydroxyl radical photoformation from aqueous-phase N(III) (HNO2 and NO2-) at different pH values. The slope of each curve indicates the OH radical formation rate constant (h-1). Solution pH was adjusted by H2SO4, and no buffer was added to avoid scavenging reactions of OH radical. Hydroxyl radical mainly reacted with added benzene. SO42- scavenged less than 0.1% of OH radical even at pH ) 1.9.
(3)
where fHNO2 and fNO2- are fractions of HNO2 and NO2- at a specific pH, and HNO2,λ and NO2-,λ are molar absorptivity (M-1 cm-1) of HNO2 and NO2- at λ nm, respectively. The fHNO2 and fNO2- are determined by the following equations (31):
fHNO2 ) [HNO2]/[N(III)] ) 1/(1 + Ka/[H+])
(4)
fNO2- ) [NO2-]/[N(III)] ) 1/(1 + [H+]/Ka)
(5)
where Ka ) 10-3.27 M (30) is an equilibrium constant for HNO2 dissociation and [N(III)] ) [HNO2] + [NO2-]. Figure 2 shows the observed and calculated absorbance of N(III) at pH 4.0 and pH 2.1, respectively. Molar absorptivity for NO2- is from Zuo and Deng (27). The calculated N(III) absorption spectra for the two pH values are very similar to the observed absorbance, indicating that the molar absorptivity of N(III) at specific pH can be determined by molar absorptivity and fractions of HNO2 and NO2-. OH Radical Photoformation from N(III) at Various pH Values. Figure 3 depicts examples of OH photoformation from aqueous-phase N(III) (HNO2 and NO2-) at three different pH values. As can be seen in Figure 3, OH photoformation rates are higher at lower pH values and are linearly dependent on concentrations of N(III). The slope of each curve in Figure 3 indicates the OH radical photoformation rate constants (h-1) at the given pH values. The OH radical photoformation rate constants at specific pH values were normalized to vernal equinox solar noon conditions of 34° N. It should be noted that solutions of N(III) and benzene (at pH ) 2.0 and 5.6) kept in the dark did not form any detectable phenol over 1 h, thus the OH radical formed in the quartz cell was not due to the reactions between N(III) and benzene.
FIGURE 4. Hydroxyl radical photoformation rate constants of aqueous-phase N(III) at pH between 1.9 and 6.2 (seven pH values). The O represents experimentally determined rate constants. The solid line was constructed using the acid dissociation constant of nitrous acid (pKa ) 3.27) (30) and OH radical photoformation rate constants ((SE) for HNO2 ) (3.1 ( 0.08) × 10-4 s-1 and for NO2) (3.2 ( 0.61) × 10-5 s-1, R 2 ) 0.99 (see details in the main text). No correction was made for ionic strength. Figure 4 shows the pH dependence (1.9-6.2) of OH photoformation from N(III). Hydroxyl radical photoformation rate constants (JN(III)) at specific pH values were fit to the following speciation equation (similar to molar absorptivity) to estimate the OH photoformation rate constants of HNO2 and NO2-:
JN(III) ) fHNO2 JHNO2 + fNO2-JNO2-
(6)
where fHNO2 and fNO2- are fractions of HNO2 and NO2- at specific pH, and JHNO2 and JNO2- are the OH radical photoformation rate constants (s-1) of HNO2 and NO2-, respectively. The JHNO2 and JNO2- were estimated by least-squares best-fit VOL. 33, NO. 15, 1999 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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analysis (r 2 ) 0.99, n ) 7) using eq 6. The best estimates of OH radical photoformation rate constants (( standard error) were (3.1 ( 0.08) × 10-4 s-1 for HNO2 and (3.2 ( 0.61) × 10-5 s-1 for NO2- for vernal equinox solar noon conditions at 34° N. The corresponding half-lifes of N(III) are 36.7 min for HNO2 and 6.0 h for NO2-. The estimated photoformation rate constants are similar to the tentatively determined values at pH ) 1.9 and 5.6 (13). Equation 6 is plotted in Figure 4 together with the experimentally determined rate constants. The close correlation between eq 6 and experimental data indicates that OH photoformation rate constants follow the speciation pattern of N(III). Hydroxyl radical photoformation rate constant at a specific wavelength (J(λ)) is expressed as
J(λ) ) ln (10)I(λ)(λ)Φ(λ)
(7)
where,I(λ) is the actinic flux, (λ) is the molar absorptivity, and Φ(λ) is the quantum yield of OH radical photoformation, all at λ nm. Integration of eq 7 for the solar spectral region yields the OH radical photoformation rate constants. On the basis of the reported solar irradiance of 34° N (24), molar absorptivity of NO2- (27), and quantum yields of OH radical photoformation (0.07 for 290-350 nm, 0.05 for 350-500 nm) (15), we calculated the OH photoformation rate constant of NO2- to be 4.4 × 10-5 s-1, which agreed well with our experimentally determined value (JNO2- ) 3.2 × 10-5 s-1). Similarly, using solar irradiance of 34° N (24), molar absorptivity of HNO2 (this study), and quantum yields of OH radical photoformation (0.35, independent of wavelength) (29), the OH radical photoformation rate constant for HNO2 was calculated to be 5.2 × 10-4 s-1, which also agreed well with our experimentally determined value (JHNO2 ) 3.1 × 10-4 s-1). It should be mentioned that Zafiriou and Bonneau (14) found that the quantum yield of OH radical photoformation from NO2- was independent of ionic strength (0.035-0.9 M), thus the increase in OH radical photoformation at lower pH was not caused by increases in SO42- concentrations. Zellner et al. (15) also found that the quantum yields of OH radical photoformation from N(III) were higher at lower pH. Implications. We examined the significance of N(III) photolysis in natural atmospheric hydrometeors, assuming that gaseous HONO was the only source of N(III) in the aqueous phase. As indicated in the Introduction, the dissolution of NO and NO2 is negligible because of the small physical Henry’s law constants. For the calculations, atmospheric concentrations of 0.1, 1, and 5 ppbv were used for gaseous HONO (PHONO). Field measurements found that gaseous HONO concentrations of up to 8 ppbv were present in polluted urban locations (1, 4, 6, 7, 12). Thus, the HONO mixing ratio used for this calculation could be observed in any urban location. We analyzed for open and closed systems (Figure 5). For the open system, gaseous PHONO was assumed to be constant, thus gas-to-liquid transfer did not deplete the gaseous HONO. For the closed system, the initial PHONO was fixed, and the reduction of gaseous HONO by dissolution in the aqueous phase was taken into account. In both cases, the atmospheric pressure was assumed to be 1 atm. Furthermore, acid-base equilibrium (HNO2 S H+ + NO2-) was assumed to be instantaneous in the aqueous phase. For the open system, aqueous-phase N(III) ([HNO2] + [NO2-]) concentration was calculated by the following equation:
[N(III)]open ) PHONOKH/fHNO2
(8)
where KH ) 49 M atm-1 is the physical Henry’s law constant for HONO (1) and fHNO2 is fraction of HNO2 defined in eq 4. 2564
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FIGURE 5. Model calculations of N(III) concentrations (A) and OH radical photoformation rates (B) in the aqueous-phase at different PHONO levels. The solid lines (a-c) indicate results for the closed system, and the dashed lines (a′-c′) indicate the results for an open system. The initial concentrations of HONO (PHONO) used were 5.0 (a and a′), 1.0 (b and b′), and 0.1 ppbv (c and c′), respectively. Hydroxyl radical photoformation rates are calculated for vernal equinox solar noon conditions at 34° N (Higashi-Hiroshima). Temperature is assumed to be 20 °C (293 K). See details in the main text. For the closed system, a mole fraction of HONO present in the aqueous-phase (χ) was calculated by (31)
χ ) 1/[1 + 1/(KHWRT(1 + Ka/[H+]))]
(9)
where W is a dimensionless liquid water content (assumed to be 3 × 10-7 m3 m-3) (32), R is the gas constant (0.08206 L atm mol-1 K-1), and T is temperature (K). Thus, N(III) present in the aqueous phase is
[N(III)]closed ) PHONOχ/(WRT)
(10)
Figure 5A shows the concentrations of N(III) at specific pH values. For instance, at an atmospheric HONO mixing ratio of 1.0 ppbv and pH ) 6.0 (χ ) 15.4%), the N(III) concentration is 25.2 µM for the open system and 21.3 µM for the closed system. But, when pH ) 7.0 (χ ) 64.5%), N(III) is 251 µM for the open system and 89.4 µM for the closed system. Hydroxyl radical photoformation was also estimated by combining eqs 6 and 8 for the open system and the eqs 6 and 10 for the closed system (Figure 5B). Similarly, when pH ) 6.0, OH radical photoformation rate is 2.4 µM h-1 for the open system and 2.0 µM h-1 for the closed system. When pH ) 7.0, the rate is 23.6 µM h-1 for open system and 8.4 µM h-1 for the closed system. The major difference between the open and closed systems is due to the significant depletion of gaseous HONO at higher pH for the closed system. It should be noted that the depletion of gaseous HONO decreases as the liquid water content decreases. Figure 5B clearly shows that OH radical photoformation in the aqueous phase is significant at higher pH. Our preliminary study found that natural dew waters had higher
pH (mean ( SD ) 6.1 ( 0.4 in nine summer samples) and contained higher concentrations of N(III) (9.1 ( 4.4 µM), which is almost exclusively NO2- (13). On the basis of the concentrations of NO2- and pH of dew waters, 99.3 ( 8.6% of OH radical was formed from photolysis of NO2- (13). Dew is well formed under stagnant air, thus the closed system appears to be more realistic for the dew waters than the open system. If N(III) in dew waters was derived solely from gaseous HONO, then PHONO ) 0.4 ppbv would account for the N(III) concentration in the dew waters at pH ) 6.1 for the closed system. The results from Figure 5 can be further extended to other atmospheric hydrometeors such as cloud, rain, and fog, which may exhibit higher PHONO and pH. However, at low pH, other OH radical sources (i.e., Fenton’s reaction) may be more important than N(III) photolysis. The results also indicate that nighttime formation of HONO and subsequent dissolution to the aqueous phase could have a significant impact on photochemistry in the aqueous phase in the morning. In particular, N(III) photolysis could act as an important oxidation reaction initiator in the aqueous phase by generating OH radicals, where PHONO is high. Furthermore, as compared to the predicted half-life of gaseous HONO (∼4.2 min) (33), the half-life of aqueous N(III) is relatively long (36.7 min for HNO2 and 6.0 h for NO2-) for vernal equinox solar noon conditions at 34° N, assuming that the disappearance of HONO is the same as photoformation of OH radical. N(III) in the aqueous phase could be one of the important sources of OH radical not only at dawn but also during the day if N(III) persists and pH does not decrease significantly in the aqueous phase.
Acknowledgments We would like to thank anonymous reviewers for their helpful suggestions on the manuscript, Dr. H. Wang for her helpful comments and some of the calculations, and Dr. L. Bartlett for editing our manuscript. This work has been supported by CREST (Core Research for Evolutional Science and Technology) of Japan Science and Technology Corporation (JST).
Supporting Information Available One table giving the decadic molar absorptivity of aqueous HNO2. This material is available free of charge via the Internet at http://pubs.acs.org.
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Received for review July 27, 1998. Revised manuscript received April 2, 1999. Accepted April 28, 1999. ES980762I
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