Quantum Yield of Nitrite from the Photolysis of Aqueous Nitrate above

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Quantum Yield of Nitrite from the Photolysis of Aqueous Nitrate above 300 nm Katherine Beem Benedict, Alexander S McFall, and Cort Anastasio Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.6b06370 • Publication Date (Web): 24 Mar 2017 Downloaded from http://pubs.acs.org on April 5, 2017

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Environmental Science & Technology

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Quantum Yield of Nitrite from the Photolysis of Aqueous Nitrate above 300 nm

9 10 Katherine B. Benedict1, Alexander S. McFall and Cort Anastasio*

11 12 13

Department of Land, Air, and Water Resources

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University of California Davis

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Davis, CA 95616

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*Corresponding Author, Department of Land, Air, and Water Resources, University of California Davis; Tel: 530-754-6095; Email: [email protected]; Fax: 530-752-1552

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1

now at Department of Atmospheric Science, Colorado State University, Fort Collins, CO 80523

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Submitted in Revised Form to Environmental Science and Technology on March 22, 2017.

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1

ACS Paragon Plus Environment


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ABSTRACT: Photolysis of nitrate (NO3–) produces reactive nitrogen and oxygen species via

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three different channels, forming: (1) nitrogen dioxide (NO2) and hydroxyl radical (•OH), (2)

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nitrite (NO2-) and oxygen atom (O(3P)), and (3) peroxynitrite (ONOO–). These photoproducts are

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important oxidants and reactants in surface waters, atmospheric drops, and snowpacks. While the

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efficiency of the first channel, to form NO2, is well documented, a large range of values have

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been reported for the second channel, nitrite, above 300 nm. In part, this disagreement reflects

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secondary chemistry that can produce or destroy nitrite. In this study, we examine factors that

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influence nitrite production and find that pH, nitrate concentration, and the presence of an •OH

31

scavenger can be important. We measure an average nitrite quantum yield (Φ(NO2–)) of (1.1 ±

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0.2)% (313 nm, 50 µM nitrate, pH ≥ 5), which is at the upper end of past measurements and an

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order of magnitude above the smallest – and most commonly cited – value reported for this

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channel. Nitrite production is often considered a very minor channel in nitrate photolysis, but our

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results indicate it is as important as the NO2 channel. In contrast, at 313 nm we observe no

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formation of peroxynitrite, corresponding to Φ(ONOO–) < 0.26%.

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2


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TOC/Abstract Art

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Introduction

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Nitrate (NO3–) is an important, photochemically active contaminant in surface waters,

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atmospheric water drops and particles, waste water, and snow. As shown in Figure 1, the longer-

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wavelength absorption maximum of aqueous nitrate is at 302 nm, which enables direct

47

photolysis in sunlight. This photolysis proceeds via three main channels1–3 (Figure 2): (1)

48

formation of NO2 and •OH (which is formed via protonation of •O–), (2) formation of NO2– and

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O(3P), and (3) at least at wavelengths below 300 nm, formation of peroxynitrite, ONOO–.

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Nitrate photolysis in channel 11,4,5 can be a major source of •OH in atmospheric and surface

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waters as well as a source of nitrogen oxides.5–7 There is good agreement regarding the quantum

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yield of this channel, with an average value of Φ(•OH) of (1.35 ± 0.3)% at 298 K for the 302-nm

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absorption band.4,5,8,9 This quantum yield is higher at short wavelengths, e.g., Φ(•OH) is 9% at

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254 nm,3 but decreases steadily with increasing wavelength to a nearly constant value of about

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1% above 300 nm.4,9,10

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The second nitrate photolysis channel, which forms nitrite, has been examined in several

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studies recently10–12 and has been a subject of investigation for decades.4,5,9,13 Nitrite and its

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protonated form, nitrous acid (HONO), are important photochemical sources of •OH and nitric

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oxide (NO) in a variety of environments. Measurements of Φ(NO2–) in aqueous solutions near

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room temperature at tropospherically relevant wavelengths (> 290 nm) vary by nearly an order of

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magnitude, from 0.5% to 4%.8,10,11,13,18,19 The wide range in literature values make it difficult to

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model this channel and understand the environmental impacts of nitrite and O(3P) formation

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from nitrate photolysis. While these studies measured the photoformation of nitrite, the other

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product from channel 2, O(3P), should have the same quantum yield. However, the one study

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that monitored O(3P) reported a much lower quantum yield of 0.11%,8 which is commonly cited 4


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to suggest the nitrite channel is unimportant.12,20–22 We summarize the wide variation in

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conditions and results from previous studies in Table S1 of the Supplement.

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The large range in previously reported quantum yields for channel 2 might be due to the

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different conditions used in past studies. For example, the pH values of illumination solutions

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ranged from 3 to 11, with some studies not reporting a pH.19 The pH can affect the observed

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quantum yield due to HONO formation and volatilization to the gas phase.12 In addition, most

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studies used an organic compound, often formate, to scavenge •OH (reaction 4 in Figure 2),

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suppressing its concentration and thus protecting nitrite from •OH oxidation (reaction 5).

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However, reaction of •OH with organic compounds can form superoxide (•O2–; reaction 4),

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which can convert NO2 to NO2– (reaction 6) and interfere with determinations of Φ(NO2–), as

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observed recently by Sharko et al.12

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concentrations, from 10 mM to 5 M;8,10–13,19 although aerosol concentrations can be in this high

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range, nitrate concentrations in most atmospheric and surface waters are much lower.7,23–25

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Nitrite quantum yields from Goldstein and Rabani10 were independent of nitrate concentration,

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but yields from Roca et al.11 varied by over a factor of two with nitrate concentration (Table S1).

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In short, the influence of experimental conditions, including nitrate concentration, pH, and

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organic scavenger, has not been fully investigated to clarify which conditions are best used to

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determine Φ(NO2–).

Past studies have also used a wide range of nitrate

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Finally, the third channel of nitrate photolysis (Reaction 3), to form peroxynitrite, is poorly

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understood at environmental wavelengths. ONOO– is difficult to measure due to its rapid

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isomerization to nitrate (at pH < 6.6), decomposition,26,27 secondary formation from the reaction

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of NO with •O2–,28 and because it absorbs light in the same region as nitrate and nitrite (Figure

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1). ONOO– is formed from nitrate photolysis at 254 nm (Φ(OONO–) = 6.5 - 10%),3,10 but its 5



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importance as a product during illumination in the 300-nm absorption band is uncertain.1

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Goldstein and Rabani10 observed no ONOO– formation during 300-nm illumination (Φ(ONOO–)

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< 0.2%), but we are aware of no other studies of peroxynitrite formation at environmentally

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relevant wavelengths.

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Our overarching goals are to understand the impacts of several environmental variables on

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the nitrite quantum yield and use our results to assess past reports of Φ(NO2–). To do this, we

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explore the impacts of pH, the presence of organic scavenger, nitrate concentration, and the

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nitrate salt cation on the nitrite quantum yield. We use a nitrate concentration, 50 µM, that is the

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lowest concentration ever used in a reported Φ(NO2–) determination and that is much more

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similar to typical environmental levels. Finally, we also examine peroxynitrite formation from

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nitrate illumination at 254 and 313 nm.

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Experimental Methods

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Materials

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We used ACS-certified sodium nitrite, sodium nitrate, hydrochloric acid (trace metal grade),

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and sodium formate from Fisher. 2-nitrobenzaldehyde (98%), potassium nitrate (ACS Reagent

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>99%), magnesium nitrate (ACS Reagent >99%), ammonium nitrate (>99%), sulfanilamide

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(>99%), N-1-napthylethylene diamine (ACS Reagent 98%), sodium hydroxide (>98%), and L-

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cysteine (>98%) were from Sigma-Aldrich. Calcium nitrate was from ACROS (99%) and

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potassium phosphate monobasic was from EM Science (99%). Solutions were prepared in air-

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saturated, deionized water from a Milli-Q Plus system (>18.2 MΩ). Stock solutions of nitrite

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(0.010 M) and nitrate (0.10 M) were prepared annually and stored in an amber glass bottle in a

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refrigerator. Absorbance spectra were taken periodically to check the stock concentrations using 6


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a UV-vis spectrophotometer (Shimadzu). Dilutions for standards and illuminations were made

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daily.

114 115

Measurement of Nitrite

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Within 1 minute of stopping or pausing illumination, we initiated nitrite determination using 29–32

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the Griess Method (e.g., refs.

), a colorimetric method that forms a strongly absorbing azo-

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dye complex with nitrite. To 1.0 mL of sample (or nitrite standard) we added 25 µL of 1.0%

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sulfanilamide in 10% (w/v) HCl solution, allowed the sample to react for 10 min, then added 25

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µL of 0.10% N-1-napthylethylene diamine solution, and waited 10 min. We then measured light

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absorption at 540 nm using a TIDAS II spectrophotometer (World Precision Instruments,

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Sarasota, FL) with a liquid waveguide capillary cell (LWCC; length of 100 cm, effective path

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length of 94 cm, 250 µL volume), and tungsten lamp. The TIDAS contains two lamps but the

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deuterium lamp (200-350-nm) causes an artifact, so we turned it off during our measurements

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(see section S1 in the supporting information for details). Absorption was measured from 350 to

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700 nm to correct for any baseline shifts. The peak height between 530 and 550 nm was

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determined as the difference between the maximum absorbance in this wavelength range relative

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to a baseline determined from local absorption minima between 400 and 500 nm and between

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550 and 700 nm. With the long pathlength of the TIDAS, we were able to measure nitrite very

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sensitively, with a detection limit based on replicate blank analyses of 3 nM. We made fresh

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standards of sodium nitrite each week and calibrated the spectrophotometer using concentrations

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from 0 to 100 nM (see supplement section S2c). Samples and other solutions were pulled

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through the LWCC at approximately 1 mL min-1 using a peristaltic pump. Manual syringe

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injections can also be used, but we found that pulling with the pump gave more consistent results 7



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and fewer air bubbles in the LWCC. We cleaned the LWCC between each measurement with a

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full cell volume of three separate cleaning solutions (1 M NaOH, 1 M HCl, and 50%

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methanol/50% water), each separated by an air bubble and followed by rinsing with Milli-Q

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water.

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While NO2 can also react in the Griess method to form the same azo dye as NO2–,33 our

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control experiments with N2 purging to remove NO2 during illumination show no statistically

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significant effect on the nitrite quantum yield, as discussed in section S3 of the supplemental

142

information. In addition, calculations of partitioning for the sparingly soluble NO2 in our vials

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indicate that the dissolved fraction of NO2 is small (20%) and should have at most a modest

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effect on our nitrite quantum yields, in agreement with the control experiments. It is also

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possible that NO2 is undergoing hydrolysis during illumination to make NO2–: Scharko et al.

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recently suggested that NO2 produced from nitrate photolysis is more easily hydrolyzed to nitrite

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compared to NO2 bubbled into solution.12 However, because our relatively rapid N2 flow (15 mL

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min–1) during purging did not cause a statistically significant decrease in the NO2– quantum yield

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(Section S3), it seems unlikely that NO2 hydrolysis is significantly interfering with our nitrite

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determination.

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Sample Illumination

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Illumination solutions generally contained 50 µM NaNO3. Solutions with a pH of 6 or above

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were adjusted using a potassium dihydrogen phosphate-sodium hydroxide buffer, while sulfuric

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acid was used for pH values below 5.5. When used, the concentration of organic scavenger was

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generally 100 µM for 50 µM nitrate solutions, and was at least 10% of the nitrate concentration.

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Illuminations were carried out with 313-nm light from a 1000-W Hg/Xe lamp in a

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monochromatic illumination system (Spectral Energy). Solutions were contained in stirred quartz

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cuvettes (1-cm or 5-cm path length) or 2.0-mL HPLC vials (low impurity Type I Class A

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borosilicate glass, 12 × 32 mm; Shimadzu P/N 228-45450-91) where the entire 1 mL of solution

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was illuminated but not stirred. The sample temperature was controlled at 25°C using a custom

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Peltier-cooled copper chamber (Paige Instruments). Quartz cuvettes, which were capped during

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illuminations, were only opened to remove an aliquot at each time point and deliver it to a vial

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already containing the first Griess reagent. Vials (each representing one time point) were capped

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during illumination and were only opened to add the Griess reagents. Dark controls contained the

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same solution as the illuminated sample but were not exposed to light and were analyzed

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periodically throughout the illumination series for an experiment; there was no production of

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nitrite in the dark. Periodic controls of deionized water were illuminated for 1 hour; no nitrite

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was formed. Based on reported molar absorptivities9,34 and measured photon fluxes (see below),

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the lifetimes of nitrite and aqueous HONO were 1.2 days and 4.5 hours, respectively, in our

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illuminated solutions.

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Detection of Peroxynitrite

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As in several past studies,10,28,35,36 we measured light absorption at 302 nm to determine the

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concentration of peroxynitrite. While peroxynitrite has a published molar absorptivity of 1670

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M–1cm–1 at 302 nm,37 there is some uncertainty in the value due to the instability of ONOO– in

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solution and the difficulty in making peroxynitrite. In addition, the absorption spectra of nitrate,

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nitrite, and peroxynitrite overlap at 302 nm (Figure 1), which confounds light absorption

179

measurements. We tried two different methods to determine the concentration of ONOO– formed 9



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during 254-nm and 313-nm illumination in a pH > 10 solution. A high pH was used to prevent

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peroxynitrite from protonating and forming peroxynitrous acid (pKa = 6.6), which undergoes

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rapid isomerization to HNO3.27 In the main ONOO– method, we measured light absorption in the

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illuminated sample at 302 nm, removed the expected absorbance contributions from nitrite and

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nitrate, and used the ONOO– molar absorptivity to calculate the amount of peroxynitrite at each

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time step (Supplemental section S4). We also tried the peroxynitrite method of Plumb and

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Edwards,38 but had no success getting this method to work (Supplemental section S5).

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Calculation of Quantum Yields

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We measured the photon flux (Iλ) daily (or more often if experimental conditions changed

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significantly) using the chemical actinometer 2-nitrobenzaldehyde (2NB)35 under the same

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conditions (container, light, and temperature) as the nitrate photolysis experiment; for

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experiments with high nitrate and scavenger concentrations we also added these species to the

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actinometry solution. Aliquots of a 10 µM 2NB solution were illuminated for varying lengths of

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time, then the concentration of remaining 2NB was measured using HPLC (C18 Column,

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60%/40% acetonitrile/water eluent at 0.7 mL min–1, absorbance at 258 nm).39 Under the low

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light-absorbing conditions of our actinometer, the measured rate constant for 2NB loss (j2NB,λ) is

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equal to:40

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j2NB,λ =2.303 × 10 Iλ × , ,

(1)

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where Iλl is the surface-area-normalized photon flux (mol-photon cm-2 s-1 nm-1), , ,

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(= 640 M–1 cm–1)40 is the product of the base-10 molar absorptivity and quantum efficiency for

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2NB at 313 nm, and 2.303 converts from base-10 to base-e. Average j2NB values were 0.018 s–1

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for vials and 0.005 s–1 for quartz cuvettes under typical conditions. Under the low light-absorbing 10


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conditions of our nitrate samples (absorbance at 313 nm < 0.16), the rate constant of nitrite

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formation is:

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(2) NO → NO = 2.303 × 10 ΦNO ,

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where ΦNO is

is the quantum yield of nitrite formation from nitrate photolysis and ,

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the base-10 molar absorptivity of nitrate at the illumination wavelength ( , =5.29 M-1cm-1,

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, !" = 3.52 M–1cm–1)9.

209 210

The rate of nitrite formation during illumination is equal to: #[ %] #'

= NO → NO [NO ]

(3)

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where t is time. We illuminate samples over short times where the change in nitrite concentration

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is linear and the formation rate is determined by a simple linear regression (see supplemental

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section S1d). This is in contrast to some past methods where experiments were sometimes

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performed over long time scales, resulting in non-linear plots of nitrite versus time.13,18

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Combining equations 1 through 3 we solve for the quantum yield of NO2-: ΦNO

=

#[ %] #'

(%)*,+ ,%-.,+

×/

[ %-.,+ ()0 ] ,+

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Similarly, the quantum yield for peroxynitrite is

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ONOO =

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#[ ] #'

×/

(%)*,+ ,%)*,+

[ %)*,+ ()0 ] ,+

(4)

(5)

For a given experiment, the relative standard error ranges from 1-10% for

#[ %] #'

and 1-7%

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for , . In comparison, the relative error for , , is 8% and for , is 1%.

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Errors on individual quantum yields (e.g., in figures) represent 1 standard error determined by

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propagating measured or reported errors in all terms in equation 4. Errors on means were

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determined from multiple individual experiments and are 1 standard deviation. 11



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Results and Discussion

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Influence of Organic Scavengers and pH

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While organic compounds can alter the apparent quantum yield of nitrite, presumably by

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scavenging •OH and thus minimizing NO2– oxidation, it is unclear if they can also lead to

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secondary formation of nitrite from the NO2 + •O2– reaction (Figure 2). To investigate this issue,

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we first examine the expected fate of •OH in a solution containing nitrite and one or more

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scavengers, either formate, cysteine, or bicarbonate (HCO3–)/carbonate (CO32–) from the

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dissolution of atmospheric CO2 (supplemental section S2). We examine the calculated fates of

233

•

234

to the product of the second-order rate constant of scavenger with •OH and the scavenger

235

concentration (e.g., for formate (Fo) this is k’OH = kFo+OH[Fo]). With no added organics, HCO3–

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/CO32– and NO2– are the dominant •OH scavengers (Figure 3a). The strength of the nitrite sink

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for

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bicarbonate/carbonate sink depends on pH, which governs CO2 partitioning from the gas phase

239

and its aqueous speciation. Between approximately pH 6 and 7 there is a transition zone where

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the dominant •OH sink changes depending on the amount of NO2– in solution. At higher nitrite

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concentrations (500 nM), bicarbonate and carbonate are the dominant sinks for •OH above

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approximately pH 7, while at a low nitrite concentration (20 nM), bicarbonate and carbonate are

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the dominant sinks for •OH above pH 6.2. Understanding where the experimental conditions fall

244

on Figure 3a is important for determining if •OH is an important sink for nitrite, which would

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(artificially) reduce the nitrite quantum yield.

OH by quantifying its pseudo-first-order rate constant with each scavenger, k’OH, which is equal

•

OH depends on the amount of nitrite formed from nitrate photolysis, while the

12


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To explore how the theoretical calculations compare to our experimental observations, we

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performed experiments determining the quantum yield of nitrite, Φ(NO2–), for pH 2 to 10.5

248

solutions containing either 50 µM or 10 mM nitrate (Figure 3b). By the end of illumination, we

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formed approximately 50 and 500 nM of NO2– in the 50 µM and 10 mM NO3– solutions,

250

respectively. At high pH, where we expect bicarbonate/carbonate to be the dominant •OH

251

scavenger, nitrite quantum yields were approximately constant, though noisy, with a higher value

252

for 50 µM NO3– (Figure 3b). In more acidic solutions, below approximately pH 5, Φ(NO2–)

253

decreases with decreasing pH, but there is quite a bit of scatter in the data. This decrease likely

254

has two causes in solutions with no added organic: (1) increased •OH reaction with NO2– in more

255

acidic solutions as bicarbonate/carbonate scavenging of •OH is less effective (Figure 3a) and, (2)

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at pH values near and below the pKa of HONO (2.8),41 protonation of nitrite to form HONO,

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which is likely lost to the gas phase.12 The latter effect has been observed by Scharko et al.,12

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where HONO production from nitrate photolysis increased with solution acidity. Thus, our

259

values at higher pH are probably most representative of the primary quantum yield for nitrite:

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average values for pH ≥ 5 in Figure 3b are 1.0% and 0.93% for 50 µM and 10 mM nitrite,

261

respectively. Since these solutions did not contain an organic •OH scavenger, it is possible that

262

our values are lower bounds; we address this issue below. Results from the literature for similar

263

conditions without scavenger (illumination wavelength above 300 nm, pH near 4, 0.01-5 M

264

sodium nitrate, room temperature) are less than our 50 µM nitrate quantum yield, ranging from

265

0.25% to 0.6% 11. Nitrite quantum yields measured at pH values near 4 without a scavenger are

266

likely underestimated because of NO2– scavenging by •OH and volatilization of HONO.

267

To determine if the decrease in Φ(NO2–) in acidic solutions was more influenced by •OH

268

scavenging of nitrite or by evaporation of HONO, we next performed experiments with an 13



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added organic scavenger. Experiments were performed at 25 °C and pH 3-8 in 50 µM NaNO3

270

with three scavenger conditions: (1) no added scavenger, (2) 50 µM sodium formate, or (3) 50

271

µM L-cysteine. As shown in Figure 4, for solutions with a pH of 5 or above, the average (± 1 σ)

272

value of Φ(NO2–) is similar for each of the three conditions: no scavenger (1.11 ± 0.17)%,

273

formate (1.18 ± 0.14)%, and cysteine (1.16 ± 0.04)%. The similar results for the three conditions

274

suggest that carbonate/bicarbonate scavenging of •OH is sufficiently protective of nitrite at these

275

pH values; it is also possible that trace organic contaminants in the solutions or containers are

276

helping to scavenge •OH.42 The similarity in the no scavenger, formate, and cysteine results also

277

suggest there is no significant secondary production of nitrite via reduction of NO2 by •O2– at

278

these low concentrations of nitrate and scavenger. To further test this idea, we performed a set of

279

experiments where we attempted to suppress •O2– and/or NO2 in order to prevent the secondary

280

formation of nitrite via reaction 6 (Figure 2). As described in supplemental section S3, these

281

attempts made no significant difference in Φ(NO2–), suggesting that secondary formation of

282

nitrite is negligible in our experiments with 50 µM NO3–.

283

Although there is significant scatter in the results of Figure 4, at lower pH values (< 4.5)

284

quantum yields decrease both with and without scavenger, consistent with loss of HONO to the

285

gas phase rather than a change in scavenging. Warneck and Wurzinger8 also found that the nitrite

286

quantum yield decreases with increasing acidity, starting somewhere between pH 9 and 5.6.

287

Based on our results, experiments should be performed at pH 5 or above so that HONO

288

evaporation is negligible. In contrast, at least two studies have used pH values below this (Table

289

S1).

290 291

Impact of Nitrate Concentration on Φ(NO2–) 14


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To further explore the difference in Φ(NO2–) between 50 µM and 10 mM nitrate solutions

293

(Figure 3b), we examined the impact of nitrate concentration for solutions with and without

294

organic scavengers at pH 7. The concentration of organic scavenger varied from 100 to 1000 µM

295

depending on the initial nitrate concentration; organic:nitrate ratios were 0.1 to 2, similar to

296

previous studies.10,11 The results of these experiments are shown in Figure 5 along with results

297

from several past studies. At our lowest nitrate concentration of 50 µM, Φ(NO2–) is the same

298

with and without an organic scavenger, consistent with the results in Figure 3. We are the first to

299

show this, likely because we are working at much lower nitrate concentrations than past work. In

300

contrast to the scavenger having no effect at low nitrate concentrations, our quantum yields with

301

and without scavenger diverge at higher nitrate concentrations. In 5 and 10 mM NO3– solutions

302

containing formate or cysteine, Φ(NO2–) is up to 40% larger than at 50 µM NO3– (with or

303

without scavenger). This enhancement is likely because of secondary formation of nitrite via

304

reaction of nitrogen dioxide and superoxide (reaction 6 in Figure 2), consistent with recent

305

observations by Scharko et al.12

306

scavenger at 5 or 10 mM NO3– is 20 – 60% smaller than the 50µM NO3– value. Experiments at

307

50 µM NO3– form approximately 40 nM of nitrite over the course of illumination (10 min) while

308

10 mM nitrate solutions form approximately 500 nM of nitrite after 1 minute. Thus the lower

309

nitrite quantum yield in the high concentration nitrate solutions without scavenger is consistent

310

with our hypothesis that the dominant •OH sink changes from HCO3–/CO32– to NO2– as the nitrite

311

concentration increases (Figure 3a).

In contrast, the nitrite quantum yield in the absence of

312

At 10 mM NO3–, our no-scavenger experiments agree well with results from Roca et al.11

313

while our with-formate results are significantly higher, suggesting losses of nitrite in their pH 4

314

solutions (Figure 5). Roca et al.11 did not observe a trend in Φ(NO2–) over their nitrate 15



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concentration range (0.01 - 5 M), although the quantum yields vary significantly. Goldstein and

316

Rabani10 report an average Φ(NO2–) of 0.94% for an unknown number of 0.02 - 1 M nitrate

317

solutions containing 10 mM formate at pH 4, somewhat lower than our results. Older work from

318

Daniels et al.13, using solutions containing 1.0 M nitrate without a scavenger at neutral pH,

319

observed a quantum yield of 4%, which is far outside the range of other results (see Table S1).

320

Our results are most similar to Warneck and Wurzinger,8 who examined Φ(NO2–) in the 300-nm

321

region (room temperature, pH 5.6, 10 mM NaNO3, 2-propanol scavenger) and reported an

322

average quantum yield of (1.02 ± 0.07)%.

323 324

Influence of Nitrate Salt Cation

325

We next examined the effect of nitrate salt on Φ(NO2–). In all experiments up to this point,

326

we used sodium nitrate, but here we present Φ(NO2–) measurements for four additional salts:

327

calcium nitrate (Ca(NO3)2), magnesium nitrate (Mg(NO3)2), ammonium nitrate (NH4NO3), and

328

potassium nitrate (KNO3) at 25°C, pH 7, and 50 µM NO3¯. Values of Φ(NO2–) range from (0.9 ±

329

0.1)% for NH4NO3 to (1.16 ± 0.13)% for KNO3 (Figure 6), but none of the results are

330

statistically different at p < 0.05 and the average quantum yield is (1.13 ± 0.15)% for our data.

331

Two past studies have reported nitrite quantum yields from nitrate salts other than NaNO3

332

(Figure 6 and Table S1): Roca et al.11 reported values for Φ(NO2–) from Ca(NO3)2 that are

333

significantly higher (with scavenger) or lower (without scavenger) than our results, while Alif

334

and Boule19 reported a quantum yield for nitrite from KNO3 that is approximately half our value.

335

Finally, our finding that the partner cation of nitrate has no effect on Φ(NO2–) in solution is

336

consistent with recent results showing the •OH quantum yield (reaction 1 in Figure 1) in solution

337

is also independent of cation.43 In contrast, this same work found that nitrate photolysis in thin 16


Page 17 of 35


338

aqueous films is affected by cation, with rates of NO2 release for KNO3, Mg(NO3)2, and NaNO3

339

approximately 3.5, 2.3, and 2.3 times higher, respectively, than Ca(NO3)2.

340

Wavelength Dependence

341

Since previous explorations of the nitrite quantum yield above 300 nm have used differing

342

wavelengths, we also tested for any wavelength dependence of Φ(NO2–). We conducted a series

343

of experiments over one month, varying only the illumination wavelength. As shown in Figure

344

S5, while there is some noise in our results, the quantum yield is essentially the same at 302, 305,

345

313, and 325 nm.

346

wavelengths, but the molar absorptivity of nitrate at 334 nm is very small (0.4 M–1 cm–1; Figure

347

S5) and difficult to measure; we hypothesize this parameter might be leading to an incorrectly

348

large quantum yield. We also performed experiments to confirm that the nitrite quantum yield is

349

independent of light intensity, as shown in Figure S6.

Our result at 334 nm is approximately 65% higher than at the other

350 351

Peroxynitrite Formation

352

To test the peroxynitrite (ONOO–) detection technique described in the methods section, we

353

first performed experiments at 254 nm to compare with reported quantum yields. Example

354

absorbance spectra for 254-nm illuminations at t = 0 min (black line) and t = 70 min (dashed

355

blue line) are shown in Figure 7a. The peak at time zero is from the initial 10 mM nitrate in

356

solution. After illumination we observe an increase in absorbance at all wavelengths (250 to

357

approximately 420 nm), which could be a result of photoformed nitrite or peroxynitrite (see

358

Figure 1). To determine the concentration of ONOO–, we subtracted the contributions of both

359

nitrate and nitrite from the absorbance spectrum, and assumed the residual absorbance was only

360

from ONOO– (Supplemental Section S4). The time profiles of the residual absorbance at 302, 17



Page 18 of 35

361

370, and 390 nm during 254-nm illumination are shown in the filled symbols in Figure 7b. The

362

ratio of the residual absorbance at the two wavelengths where ONOO– absorbs (for example,

363

A370nm:A390nm) is equal to the ratio of molar absorptivities between the same wavelengths (εONOO-

364

,370nm:εONOO-,390nm),

365

four experiments with 1 to 10 mM nitrate at pH 12 and illumination at 254 nm: the average (± 1

366

σ) quantum yield for ONOO– is (6.5 ± 1.5)%, which agrees well with measurements by

367

Goldstein and Rabani10 (6.5%) but is lower than that from Mark et al.3 (10%).

consistent with the residual absorbance being due to ONOO–. We performed

368

We then performed the same experiment with 313-nm illumination. In comparison to the

369

254-nm results, the main peak of absorbance (250 to 330 nm) did not change after 130 minutes

370

of illumination at 313 nm (Figure 7a). There is an increase in the absorbance between 350 and

371

390-nm, corresponding to photoformed nitrite. The residual absorbances at 302, 370, and 390 nm

372

show no increase during 313-nm illumination (Figure 7b). Instead, there is a slight decrease in

373

the residual absorbance, which is likely due to the loss of nitrate, since we assume the nitrate

374

concentration is constant over the course of illumination. (For comparison, we calculate that 4%

375

of the initial nitrate should have been lost during the 130 min of illumination, corresponding to

376

decreases in the residual absorption coefficient of 0.0003, 0.002, and 2.6E-6 cm–1 at 302, 370,

377

and 390 nm, respectively.) We performed additional experiments with 50 µM NaNO3 and

378

obtained similar outcomes. Based on these results, and the good agreement between our

379

peroxynitrite quantum yield at 254 nm and the literature, it appears there is no significant

380

ONOO– formation during 313-nm illumination. Based on our limit of detection, we calculated

381

Φ(ONOO–) to be less than 0.26%, similar to results from Goldstein and Rabani,10 who report an

382

upper bound of 0.2%.

383 18


Page 19 of 35


384

Implications and Environmental Impacts

385

Our recommended value for Φ(NO2–) is (1.1 ± 0.2)% at 25°C, which is the average of our 61

386

experiments with 50 µM NO3–, pH ≥ 5, with and without scavengers. This value is 17% higher

387

than the recently reported average from Goldstein and Rabani.10 In contrast, our recommended

388

value is 10 times higher than the quantum yield for channel 2 determined from O(3P)

389

measurements by Warneck and Wurzinger.8 This is significant because a number of studies refer

390

to this low quantum yield (e.g., refs 12,20–22,44,45) as the nitrite yield. Perhaps because of this

391

perception, a number of modeling studies include only the production of NO2 – and not N(III)

392

(i.e., NO2– + HONO) production – from nitrate photolysis.46–49 However, our results, along with

393

those from past work,8,10–12,19 indicate that the rate of production of NO2–and HONO from nitrate

394

photolysis is comparable to the rate of NO2 formation.

395

This often neglected source of NO2– might help reconcile model underpredictions of daytime

396

concentrations of gaseous HONO in the atmosphere, which is an important source of •OH.50–53

397

Several explanations have been offered for this underprediction, including heterogeneous

398

chemistry54 and surface photolysis.55–57 Including multiple HONO formation channels has

399

improved the ability of models to reproduce observations,58–60 but discrepancies remain. The

400

photolysis of particulate nitrate to form nitrite is often not included when HONO sources are

401

analyzed58 or only the NO2 production channel is considered as a source.46 There is evidence that

402

aerosol nitrate is associated with an increase in HONO59,61 but many attribute the HONO source

403

to heterogeneous chemistry,60,62,63 even though others suggest the unknown source of nitrate is

404

correlated with other photolysis terms.33,61,64–67 In some cases, the unknown photolytic HONO

405

source is the dominant daytime term.68,69

19



Page 20 of 35

406

To estimate the contribution of aerosol nitrate photolysis to HONO(g) production, we assume

407

an acidic, aqueous aerosol such that every NO2– formed leads to release of HONO. Under Davis,

408

CA summer solstice conditions with an aerosol nitrate concentration of 0.5 µg/m3 (8 nmol/m3),

409

photolysis will form approximately 0.3 ppt/hr of HONO. In winter in the Central Valley of

410

California, aerosol nitrate concentrations can be much higher,70 up to 30 µg/m3, corresponding to

411

a HONO source of 6 ppt/hr under winter solstice radiation. However, this is not enough to

412

provide closure to atmospheric measurements, which tend to differ several hundred ppt/hr64,69,71

413

or more,65–68,72, but it is a step in the right direction. Other studies suggest that photolysis of

414

nitrate on surfaces can be faster than in bulk solution. For example, it has been reported that the

415

molar absorptivity of surface-adsorbed nitrate is approximately 75 times greater than that of

416

aqueous nitrate73 and that photolysis of nitrate (measured by •OH production) occurs 4 orders of

417

magnitude faster in urban grime than in aqueous solution.74 However, the type of surface may be

418

important: Ye et al.56 found that while the photolysis rate for surface HNO3/nitrate was up to 3

419

times larger than gas-phase HNO3, the major product released was different for natural surfaces

420

(where HONO dominated) and artificial surfaces (where NOx dominated). Properly accounting

421

for these surface reactions in models will require measurements of both molar absorptivities and

422

quantum yields and an understanding of how these values vary with sample environment.

423 424

Acknowledgements

425

We thank Emily Lucic and Alex Funderburk for assistance with the experiments. This work

426

was funded by the Arctic Natural Sciences Program of the National Science Foundation (ANS-

427

1204169). 20


Page 21 of 35

428 429


Supporting Information. Additional experiment details, calculation details, effect of some experimental variables, Tables S1-S5, Figure S1-S6

430

21



431

Page 22 of 35

Figures

432 1800

28

ONOO–

Illumination

εNO3–, εNO2– (M–1cm–1)

24

NO2–

20

1400 1200 1000

16

800

12

600

8 4

1600 εONOO– (M–1cm–1)

32

400

NO3–

200 0

0 250 275 300 325 350 375 400 λ (nm)

433 434

Figure 1. Base-10 molar absorptivity of nitrate (NO3–)9, nitrite (NO2–)34, and peroxynitrite

435

(ONOO–)35. The spectral distribution of the illumination system output at 313 nm, in arbitrary

436

units, is included in purple.

437

22


Page 23 of 35


438 NO3– + hν 3 2 1

3

ONOO–

NO2– + O( P) hν

NO2 + •OH

NO + •OH org.

5

+

H

4 6 •

NO2 + OH–

HONO (aq)

O2– HONO (g)

439 440 441

Figure 2. Primary and secondary reactions occurring during nitrate photolysis. The secondary formation of nitrite from reaction of NO2 with •O2– is indicated by the dashed lines (reaction 6).

442 443 444

23



Page 24 of 35

a) 1.0E+07 Cysteine

1.0E+06

k'OH (s–1)

1.0E+05 Formate

–

2–

HCO3 + CO3

1.0E+04 –

500 nM NO2

1.0E+03

–

100 nM NO2

–

20 nM NO2

1.0E+02 1.0E+01 1 445

3

5

7

9

11

pH

b) –

–

H2CO3 ↔

HONO ↔ NO2

1.6%

HCO3

2–

↔ CO3

1.4%

Φ(NO2‾)

1.2% 1.0% 0.8% 0.6% 10 mM

0.4%

50 uM

0.2% 0.0% 1 446

3

5

7

9

11

pH

447 448

Figure 3: a) Calculated pseudo-first-order rate constants for •OH scavenging by nitrite, 50 µM

449

formate (orange line), 50 µM cysteine (purple line), and bicarbonate/carbonate (green line) in

450

equilibrium with atmospheric CO2 at a given pH. The three horizontal black lines represent the

451

rate constants for •OH scavenging for three different nitrite concentrations (20, 100, and 500 24


Page 25 of 35


452

nM). Nitrite will be protected from •OH under conditions where k’OH for nitrite is significantly

453

below the value for another scavenger at a given pH.b) Results from initial experiments to

454

examine the relationship between Φ(NO2–) and pH for both 50 µM and 10 mM NO3– solutions

455

(313 nm, 25°C, no organic scavenger, pH adjusted with H2SO4 or a phosphate sodium hydroxide

456

buffer). The three thicker vertical lines represent the pKa values for HONO (2.8)41, H2CO3

457

(6.3)75, and HCO3– (10.33)75.

458 459 460 1.6%

1.4

1.4%

1.2 1.0

1.0%

0.8

0.8%

0.6

0.6%

Formate Cysteine 0.4 No Org (Vial) No Org (Cuvette) 0.2 Average pH > 5 Mole Fraction

0.4% 0.2% 0.0%

0.0

2 461

Mole Fraction of NO2-

Φ(NO2–)

1.2%

3

4

5

6

7

8

9

pH

462

Figure 4. Quantum yield of nitrite (50 µM NaNO3, 313 nm, 25°C) as a function of pH

463

without and with 50 µM organic scavenger. The open circles and squares represent experiments

464

performed in glass HPLC vials and quartz cuvettes, respectively. The dashed line is the mole

25



Page 26 of 35

465

fraction of N(III) that is present as nitrite. pH was adjusted with either sulfuric acid (pH ≤ 5.6) or

466

potassium dihydrogen phosphate (pH ≥ 6).

467

26


Page 27 of 35


468 469 This Work

470

Roca et al.

Warneck & Wurzinger

471

1.6%

Φ(NO2‾)

1.2%

0.8%

0.4%

0.0% 0.00001

0.0001

0.001 0.01 0.1 Initial Nitrate Concentration (M)

1

10

472 473 474

Figure 5. Quantum yield of nitrite, Φ(NO2–), as a function of nitrate concentration at pH 7

475

(313 nm, 25°C) with and without cysteine or formate (50-1000 µM, depending on [NO3–]) .

476

Triangles represent data from Roca et al.11 for pH 4 solutions with (orange filled) and without

477

(open) 10 mM formate (Fo), illuminated at 310 nm and 25°C. Squares represent data from

478

Warneck and Wurzinger8 for pH =5.6 solutions (305 nm and 0.13 M acetone (open) or 0.001M

479

formate (solid)) where the quantum yield for reaction 2 was determined either from NO2– (dark 27



Page 28 of 35

480

green) or oxygen atom, O(3P) (light green line). The blue line is the average of data from

481

Goldstein and Rabani10 for pH 4.2-4.5 solutions containing 10 mM formate at 24°C with 300-nm

482

illumination. The dark purple diamond is the average of data from Alif and Boule19 for solutions

483

of 0.1 M KNO3 with formate at 310 nm (no temperature or pH are reported). The red “x” is the

484

average of two data points from Dubowski et al.18 for solutions of 10 mM NaNO3 illuminated at

485

313 nm and 20°C. Lines are qualitative fits to our data.

486 487

28


Page 29 of 35


1.6%

This Study

Roca et al. no Fo

Roca et al. w/Fo

Goldstein and Rabani

Warneck and Wurzinger

Dubowski et al

Alif and Boule

Φ(NO2‾)

1.4% 1.2% 1.0% 0.8% 0.6% 0.4% 0.2% 0.0%

488

NaNO3 NaNO 3

Ca(NO3)2 Ca(NO3)2

Mg(NO3)2 Mg(NO3)2

NH4NO3 NH4NO3

KNO3 KNO 3

489 490 491

Figure 6. The influence of nitrate salt on the quantum yield of nitrite for experiments from

492

this work (solid black bars: 313 nm, 25°C, 50 µM NO3–, no scavenger, pH > 7), Roca et al.11

493

(310 nm, 25°C, 10 mM NO3–, no formate (white bar) or 10 mM formate (light orange bar)),

494

Goldstein and Rabani10 (striped blue bar; 24°C, 300 nm, 0.02-1.0 M NO3–, 10 mM formate, pH

495

4.2-4.5), Warneck and Wurzinger8 (green bar; 305 nm, 22°C, 0.01 M NO3¯, 0.13 M 2-propanol,

496

pH 5.6) Dubowski et al.18 (red-hatched bar, 313 nm, 0.01 M NO3¯, 293K, no pH reported) and

497

Alif and Boule19 (purple bar; 310 nm, 0.1 M KNO3, 0.5 M formate, no temperature or pH

498

reported). Error bars for our data are ± 1 σ, while bars for other data are the reported errors.

499 500 501 29



Page 30 of 35

502 a)

0.50

Absorbance

TBefore = 0 Illumination 0.40

λ=313nm 313 nm, t t=130min =130 min

0.30

λ=254nm 254 nm, tt=70min =70 min

0.20

0.10

0.00 250

300

503 b)

350 Wavelength (nm)

400

450

Residual Abs. Coefficient (cm-1)

0.030 λ(ill): 254 λ(msmt): 302

0.025 0.020 0.015

λ(ill): 254 λ(msmt): 370

0.010 0.005

λ(ill): 254 λ(msmt): 390

0.000 λ(ill): 313

-0.005

0 504

50

100

150

Time (min)

505

Figure 7. Determination of peroxynitrite by UV absorbance measurements in solution (10

506

mM NaNO3, pH 13, 25 °C) illuminated with either 254 or 313 nm radiation. Panel a) shows UV

507

absorbance scans (5-cm pathlength) taken before illumination (t = 0; black line), after 70 min of

508

254-nm illumination (dashed blue line), and after 130 min of 313-nm illumination (dashed gray

509

line). Panel b) shows residual UV absorption coefficients (cm–1) as a function of time during 30


Page 31 of 35


510

illumination (λ(ill)) at 254 nm (filled symbols) or 313 nm (open symbols). The residual

511

absorption coefficient is the pathlength-normalized absorbance after subtracting contributions

512

from nitrate and nitrite (see supplemental section S4 for calculation details). Absorption

513

coefficient values are shown for three different measurement wavelengths, λ(msmt): 302 nm

514

(circles), 370 nm (diamonds), and 390 nm (squares).

515 516 517 518 519 520 521 522 523 524 525 526 527 528 529 530 531 532 533 534 535 536 537 538 539 540 541 542 543 544 545 546 547 548 549

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Quantum Yield of Nitrite from the Photolysis of Aqueous Nitrate above

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