1 Photodegradation of Dicloran in Freshwater and Seawater Emily N

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Photodegradation of Dicloran in Freshwater and Seawater Emily Vebrosky, Parichehr Saranjampour, Donald G. Crosby, and Kevin Armbrust J. Agric. Food Chem., Just Accepted Manuscript • DOI: 10.1021/acs.jafc.8b00211 • Publication Date (Web): 23 Feb 2018 Downloaded from http://pubs.acs.org on February 25, 2018

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Journal of Agricultural and Food Chemistry

Photodegradation of Dicloran in Freshwater and Seawater Emily N. Vebrosky,† Parichehr Saranjampour,†,‡ Donald G. Crosby,§ and Kevin L. Armbrust*,† †Department of Environmental Sciences, College of the Coast & Environment, Louisiana State University, Baton Rouge, LA 70803 ‡Current address: United States Environmental Protection Agency, 109 TW Alexander Drive, Durham, NC 27709 §Department of Environmental Toxicology, University of California at Davis, Davis, CA 95616

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ACS Paragon Plus Environment


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ABSTRACT

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Dicloran appears to be a model pesticide to investigate photodegradation processes in

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surface waters. Photodegradation processes are particularly relevant to this compound as

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it is applied to crops grown in proximity to freshwater and marine ecosystems. The

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photodegradation of dicloran under simulated sunlight was measured in distilled water,

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artificial seawater, phosphate buffer, and filter-sterilized estuarine water to determine the

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half-life, degradation rate, and photodegradation products. The half-life was

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approximately 7.5 hours in all media. There was no significant difference in the rate of

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degradation between distilled water and artificial seawater for dicloran. For the

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intermediate products, a higher concentration of 2-chloro-1,4-benzoquinone was

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measured in artificial seawater versus distilled water, while a slightly higher

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concentration of 1,4-benzoquinone was measured in distilled water versus artificial

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seawater. The detection of chloride and nitrate ions after two hours of light exposure

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suggests photonucleophilic substitution contributes to the degradation process.

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Differences in product distributions between water types suggest that salinity impacts on

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chemical degradation may need to be addressed in chemical exposure assessments.

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KEYWORDS: Dicloran, photodegradation, salinity, surface water

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INTRODUCTION

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Dicloran (Figure 1) is a fungicide used to prevent spore germination of Rhizopus

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spp., Sclerotium rolfsii, and other fungal diseases on various crops including stone fruits,

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sweet potatoes, tomatoes, lettuce, and celery throughout the western and southern United

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States.1 It is applied to crops by sprinkler irrigation, aerial spray, dip tank, and

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chemigation. With a Koc of 760-1062, dicloran is classified as having low mobility in

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soils; it still may be susceptible to transport to surface waters in areas where these crops

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are grown close to aquatic ecosystems.1,2 Dicloran is marketed under the trade name

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Botran™ in a variety of formulations.3 Although it has been registered for use in the

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United States since 1961 and is reported to be susceptible to photodegradation, very little

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published data is available outside of that reported in handbooks on its physical and

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chemical properties or degradation pathways or its persistence on crops.2-4 Boscá, et al.

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indicates that dicloran can induce photosensitization in workers exposed to this chemical

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during application. In this paper, the authors reported that dicloran (DCL, 2,6-dichloro-4-

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nitroaniline) undergoes photoreduction of the nitro group from its triplet state and

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homolytic rupture of the C-NH2 bond from its singlet state. These authors conducted

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irradiations in organic solvents and reported photoreduction of the nitro group as the most

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important photodegradation pathway. The photohemolysis assay, as an in vitro

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phototoxicity test, has demonstrated the involvement of radical-mediated cellular

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membrane damage in dicloran photosensitization.5

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Multiple components and environmental variables can influence chemical

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photodegradation including pH, dissolved inorganic carbon, dissolved organic carbon,

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nutrients, and salinity.6,7 Within marine and estuarine systems, salinity can influence the

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measured degradation rates, half-lives, and generation of intermediate products of

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chemicals that may be susceptible to photodegradation. Pentachlorophenol (PCP) and

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3,4-dichloroaniline (DCA) undergo photolysis reactions when degraded by sunlight; in

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distilled water, PCP has a reported half-life of 0.9 hours, while in artificial seawater the

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reported half-life is 2.3 hours and DCA has a half-life of 17.2 hours in distilled water and

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17.3 hours in artificial seawater.8-10 Photonucleophilic substitution is reported to be the

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principal mechanism responsible for the degradation of PCP while DCA is degraded

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through solvolysis of an aryl cation intermediate of the molecule’s excited state.9,10 Both

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PCP and DCA demonstrated media-dependent differences in measured photoproducts;

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the highest concentration of a PCP photoproduct, tetrachlorophenol, was 0.8% in distilled

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water versus 6.1% in seawater.8-11 Therefore, the trend does not apply uniformly to all

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chemicals and the apparent effects of seawater on a chemical’s photodegradation process

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indicates the importance of substituent effects which are unique to each chemical. 4-

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Hydroxychlorothalonil, a degradation product of the fungicide chlorothalonil, undergoes

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processes similar to PCP, degrading initially by photonucleophilic substitution to small

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fragments including aliphatic acids.12 Nitrofen (2,4-dichlorophenyl p-nitrophenyl)

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undergoes photodegradation with the replacement of the ring nitro group with a hydroxyl

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group; it further degrades into products including hydroquinone.13 With chlorine and

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nitro substituents, dicloran shares many of the same structural characteristics of PCP,

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DCA, nitrofen, and hydroxychlorothalonil, and would be expected to photodegrade

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similarly.

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Photodegradation is relevant not only to a chemical’s persistence but also its toxicity, potentially inducing toxic effects in organisms that would otherwise be

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unharmed by it at those concentrations.14,15 Many polycyclic aromatic hydrocarbons

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(PAH) have been reported to have phototoxic effects on aquatic organisms at

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concentrations lower than those effects observed in the absence of sunlight. Impacts on

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algae and aquatic vertebrates from PAH phototoxicity includes damage to cellular

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membranes.14 Both blue crab (Callinectes sapidus) and mahi-mahi (Coryphaena

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hippurus) embryos have been shown to be negatively impacted by phototoxic PAHs,

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such as anthracene and pyrene, from crude oil due to the Deepwater Horizon oil spill in

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2010; these impacts ranged from reduced fecundity to mortality.14 Dicloran has also

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demonstrated radical-mediated cellular membrane damage in a photohemolysis assay.5 It

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may be possible that dicloran or its photodegradates drive such impacts to organisms in

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aquatic environments. The purpose of this investigation was to gather evidence

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supporting a photodegradation pathway for dicloran in aqueous systems and determine if

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salinity impacted dicloran’s photodegradation rates or product distribution, as salinity

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does for PCP and DCA.

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MATERIALS AND METHODS

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Reagents and solvents.

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Analytical grade 2,6-dichloro-4-nitroaniline and 2,6-dibromo-4-nitroaniline

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(DBNA) were from Sigma-Aldrich (St. Louis, MO). Sodium nitrate, sodium nitrite,

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sodium chloride, sodium bromide, maleic acid disodium hydrate, succinic acid disodium

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salt, and oxalic acid were obtained from Sigma-Aldrich, and fumaric acid from Fluka (St.

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Louis, MO). The standards of the degradation products 2-chloro-1,4-benzoquinone, 1,4-

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benzoquinone, and 1,2,4-benzenetriol were from Sigma-Aldrich, and potassium

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phosphate dibasic (anhydrous) was from Mallinkrodt Chemical Company (Paric, KY).

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All solvents were HPLC grade; distilled water was from J. T. Baker (Center Valley, PA)

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and acetonitrile (ACN) from Fisher Scientific (Pittsburgh, PA). Instant Ocean®

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(Blacksburg, VA) was used for artificial seawater and mixed in distilled water according

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to directions on package. Natural estuarine water was collected from Lake Pontchartrain

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at Fontainebleau State Park (Mandeville, LA). Field and lab water samples were filtered-

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sterilized prior to use with 0.22 µm cellulose acetate membrane filters (Advantec MFS,

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Inc., Dublin, CA). Borosilicate glass 2 mL vials were obtained from Agilent

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Technologies (Santa Clara, CA), colorless vials for photodegradation experiments and

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amber for dark controls; borosilicate glass has been shown not to attenuate light in the

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actinic region of sunlight and has been used in other photodegradation studies.12

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Rate Experiments.

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Samples of dicloran and its brominated analogue, DBNA, were prepared from a 1000 mgL-1 stock solution in ACN, and later diluted for analysis in their proposed media.

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The degradation rate constant and half-life of DBNA was measured in comparison to

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dicloran to determine if the ring-halogen could potentially have a significant impact on

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the overall rate of degradation. Stock solutions were diluted to 1 mg L-1 for analysis in

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the various media including distilled water, artificial seawater (3.2%), filter-sterilized

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estuarine water (0.8%, pH 7.6), and 0.01 M phosphate buffer (pH 7); this concentration

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was below the reported solubility of dicloran (6.3 mg L-1 at 20°C) and the co-solvent was

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less than 1% of the total volume.16 Vials containing 1 mL of solution were irradiated for

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a maximum period of 24 hours in an Atlas SUNTEST XXL+ environmental chamber

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(Mount Prospect, IL) outfitted with a daylight filter. The environmental chamber

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simulated the irradiance of natural sunlight measured at wavelengths between 300-400

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nm with an energy output of 65 W/m2 and 20% relative humidity. Individual samples

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were removed at 0, 2, 4, 6, 12, and 24 hours for analysis. Each hour of exposure in the

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chamber at this intensity was calculated to be equivalent to approximately 1.8 hours of

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direct sunlight at 30°N latitude during the June summer solstice.17 Dark controls were

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run simultaneously. An Agilent 1260 Infinity High Performance Liquid Chromatograph

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(HPLC) (Santa Clara, CA) was used to measure the concentrations of dicloran and

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DBNA remaining in solution, as well as detect and quantify concentrations of

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intermediate photodegradation products and obtain UV-spectra for each compound. The

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samples were analyzed with photodiode array detection at 380 nm (for dicloran and

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DBNA) and 254 nm for observed intermediate products, using a water and acetonitrile

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gradient as the mobile phase and a ZORBAX C-18 Eclipse Plus Analytical 4.6x150 mm

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5-micron column.

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Intermediate Degradation Products.

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Two major intermediate photodegradation products were identified by matching

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the HPLC retention times of standards and by photodiode array UV spectra. These

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products as well as other intermediate products were additionally characterized using a

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Waters ACQUITY Ultra Performance Liquid Chromatography (UPLC) unit with a Triple

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Quad Detector (TQD) (Milford, MA), by direct injection, with Atmospheric-Pressure

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Chemical Ionization (APCI) in negative mode, a cone voltage of 30 V, and a scan range

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50-550 m/z, as well as their UV-spectra. A Varian Cary® 50 UV-Vis Spectrophotometer

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(Agilent Technologies, Santa Clara, CA) was also used to measure the UV-spectra of

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dicloran and two photoproducts at 0.01 mM from 200-500 nm.

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A Thermo Scientific Dionex ICS-5000+ Ion Chromatograph (IC) (Sunnyvale,

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CA) unit was used to measure the formation and decline of the concentrations of

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chloride, nitrite and nitrate ions. Analytes were separated on an AS-14 column with an

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isocratic mobile phase of water and sodium hydroxide and a 0.50 mL min-1 flow rate.

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A ZORBAX SB-Aq Rapid Resolution HT 1.8-µm 4.6x150 mm HPLC column

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was used to detect small aliphatic acids including maleic acid, fumaric acid, succinic

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acid, and oxalic acid, with a 99:1 isocratic gradient of 20 mM pH 2.0 phosphate buffer

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and ACN with a 1.0 mL min-1 flow rate at 210 nm.

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Data Analysis.

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At each sample point (t = 0, 2, 4, 6, 12, and 24 hours), the percent residual

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dicloran or DBNA remaining in solution was averaged across replicates and the pseudo-

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first order rate constant was calculated from a plot of the ln C/C0 where ln C/C0 = -kt (k =

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the pseudo-first order rate constant in hr-1 and t = time). The half-life was calculated

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from the rate constant, where t½ = (ln2)/k. One-way ANOVA was used for statistical

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analysis of the degradation rates and half-lives (α = 0.05) to determine any significant

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differences between media and to remove any outliers in the trials. Each trial was

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replicated in quadruplicate except for dicloran in distilled water and in artificial seawater,

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which were replicated 8 times.

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A small sample, one-sided paired t-test (α = 0.05) determined any significant

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statistical difference between the generation and subsequent degradation of 2-chloro-1,4-

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benzoquinone and 1,4-benzoquinone in distilled water and artificial seawater. The

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average of the sampling time points was used to plot the formation and degradation of the

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products. The limit of detection (LOD) was established based on the lowest, measured,

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quantifiable standard (LOD = 0.01 mg L-1). Measurements below the LOD were

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assumed to be zero and are depicted with asterisks.

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RESULTS AND DISCUSSION

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Degradation Rates and Half-lives.

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Dicloran has a strong UV absorption within the actinic region of sunlight with a

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peak at 365 nm, indicating a strong likelihood for photodegradation in aquatic

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environments (Figure 2). No degradation was observed in any dark control. Dicloran

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degraded rapidly in the simulated sunlight, t½ = 7.62 ± 0.094 hours in distilled water and

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7.37 ± 0.279 hours in artificial seawater; the rate constants and half-lives for dicloran and

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DBNA in all media are shown in Table 1.

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One-way ANOVA analysis (α = 0.05) indicated there was no significant

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difference between the degradation rate constants of dicloran in distilled water and

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artificial seawater 0.092 ± 0.0011 hr-1 and 0.094 ± 0.0036 hr-1, respectively (p = 0.19). A

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slight difference was observed between dicloran degradation in phosphate buffer and

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estuarine water compared to artificial seawater and distilled water, however the

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differences could have been skewed by the latest time points, as there was no significant

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difference between the rate constants in any media within the first half-life (Figure 3). A

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variety of different components that are found in estuarine water, including humic and

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fulvic acids and nutrients can induce reactions through indirect photochemical processes

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and may have been responsible for differences observed at later time points. Components

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of natural water have shown to enhance the photodegradation of other organic

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compounds such as 17α-ethynylestradiol and chlorpyrifos.6,7,11 DBNA behaved similarly

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to dicloran with a t½ in distilled water roughly one hour longer at 8.86 ± 0.257 hours.

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Although dicloran and DBNA appear to follow similar trends, the difference in t½ appears

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to be influenced by the ring-halogen. Chlorine is more electronegative and thus a

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stronger electron-withdrawing group than bromine, which likely explains why dicloran

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degrades slightly faster than DBNA. Dicloran and DBNA appear to behave more like

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DCA and nitrofen than either PCP or 4-hydroxychlorothalonil; salinity does not impact

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the photodegradation rate of dicloran.8-13

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Intermediate Degradation Products.

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Initial degradation products were detected after two hours of exposure to artificial

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sunlight. Two early intermediate products were 2-chloro-1,4-benzoquinone and 1,4-

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benzoquinone, identified by matching retention times, UV, and MS with those of

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standards. The rates of formation and degradation of both 1,4-benzoquinone and 2-

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chloro-1,4-benzoquinone are significantly different between DI water and artificial

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seawater (Figure 4). For 1,4-benzoquinone, a one-sided paired t-test produced p = 0.032

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(α = 0.05) and for 2-chloro-1,4-benzoquinone, p = 0.029; therefore, both photoproducts

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behave somewhat differently, statistically, in their generation and degradation between

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freshwater and seawater although this may not be visually apparent for 1,4-benzoquinone

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in Figure 4. For both products, the standard error appears larger for samples in seawater

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than those in distilled water. The statistical analysis of the photoproducts does not

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include time points where values fell below the LOD. Further MS analysis of solutions

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after two and four hours of irradiation showed ion clusters of products consistent with a

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proposed photodegradation pathway (Figure 5). Major ion clusters appeared at m/z 124-

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126 (box C), m/z 139-144 (boxes A and D), m/z 155-160 (boxes E and F), m/z 171-176

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(boxes G and H), and m/z 189-193 (box I). The proposed structures are consistent with

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these masses as well as photoproducts measured for similar compounds.8-10,12,18

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Bracketed compounds are assumed, but were not detected by HPLC or MS.

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In this pathway, dicloran undergoes an initial displacement of the nitro group by

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hydroxyl. The resulting 4-hydroxyaniline is then oxidized to the corresponding enimine,

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which is rapidly hydrolyzed to the chlorobenzoquinone. From this point, degradation is

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expected to proceed through a series of phenol oxidations and 1-4 addition reactions,

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ultimately resulting in ring cleavage. In this pathway, structures in boxes are illustrated

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as paired quinone-hydroquinone (quinhydrone) complexes.18 In aqueous media these

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often exist as electron donor-acceptor pairs and interchange through a semiquinone

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radical.18-21 They are also responsible for the production of singlet oxygen and superoxide

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radical in the presence of sunlight.22 Thus, it is reasonable to depict the progressive

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degradation of these intermediates as occurring through redox pairs. As all of these

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products appear to degrade more quickly than the parent compound, it is not surprising

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that they are detected at only trace levels over the course of the experiment.

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Based upon the products, the photodegradation of dicloran appears in many

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respects similar to that of nitrofen where a hydroxyl group replaces the nitro group, and

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this process is followed by formation of a quinone and chloroquinone.13 However, it is

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not apparent that the loss of chlorines for dicloran photodegradation is strictly the result

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of a photonucleophilic mechanism as a reductive mechanism also appears to be occurring

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as evidenced by the presence of the monochloroquinone and benzoquinone.

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Mechanistically this could be occurring through a reductive process similar to the loss of

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chlorines during the photodegradation of pentachlorophenol to tetrachlorophenol or 3,4-

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dichloroaniline to 3-chloroaniline.8 In the case of 3,4-dichloroaniline, the formation of 4-

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chloro-3-hydroxyaniline occurs through an aryl cation resulting the solvolysis of the

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excited state of the molecule supported by the high electron density at the 3-carbon

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position on the aromatic ring.10 This is likely different mechanistically then what appears

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to occur with both nitrofen and dicloran. The mechanism is apparently complicated with

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both reductive, photonucleophilic and hydrolytic processes occurring concurrently.

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Further evidence of this is the absence of 2,6-dichloro-1,4-benzoquinone in any sample in

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this investigation. Its presence would have been expected had photonucleophilic

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substitution of the nitro group for hydroxyl been the dominant process. The absence of

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2,6-dichloro-1,4-benzoquinone is possibly explained by an intermolecular oxidation-

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reduction occurring where the hydrogen from the amine replaces one of the chlorines,

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occurring simultaneously during the formation of the quinones.

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The detection of nitrate, nitrite and chloride ions provide additional lines of

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evidence for these processes. Nitrate ions (NO3-) were detected at two hours by ion

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chromatography (IC) suggesting photonucleophilic may contribute to the

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photodegradation of dicloran. Chloride ions were also detected after two hours of

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exposure to artificial sunlight and increased in concentration over time, finally accounting

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for 38% of the theoretical amount of chloride. After 12 hours of irradiation, maleic and

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fumaric acid were detected by HPLC, but at levels too low to be quantitatively measured.

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Nitrite could only be detected at 12 hours after sunlight exposure, but this suggests NO3-

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and NO2- are undergoing separate photolysis to form hydroxyl radicals or singlet

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oxygen.23-25 Zhou, et al. and Zafiriou, et al. suggest that NO3- and NO2- have the

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potential to regenerate as a result of photolysis to form both •OH and 1O2 in freshwater

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and marine environments.24-27 This provides additional evidence supporting the proposed

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photodegradation pathway of dicloran in Figure 4.

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In conclusion, salinity did not appear to impact the rate at which dicloran

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photodegrades. The detection of ions such as chloride and nitrate after two hours of

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irradiation as well as the structures within the proposed degradation pathway, suggests

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photonucleophilic substitution may be a contributing factor in the degradation process.

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Salinity did impact the generation of intermediate degradation products. The product 2-

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chloro-1,4-benzoquinone formed at nearly double the concentration after two hours of

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irradiation in seawater (0.10 ± 0.011 mg L-1) as opposed to distilled water (0.04 ± 0.003

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mg L-1). The product 1,4-benzoquinone was not as significantly impacted by salinity as

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2-chloro-1,4-benzoquinone but a difference in product distribution was observed;

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benzoquinone was measured at a slightly higher concentration in distilled water at both 4

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and 12 hours (0.044 ± 0.004 mg L-1 and 0.019 ± 0.002 mg L-1, respectively) opposed to

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seawater (0.041 ± 0.029 mg L-1 and 0.014 ± 0.010 mg L-1). A difference in the

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distribution of photoproducts between different waters, in this case freshwater and

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seawater suggests possible differences in toxicity in one media over another. As seawater

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has shown to elevate toxicities in previous studies, the potential for dicloran or its

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intermediate degradation products to be more toxic in seawater warrants further

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investigation.

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ABBREVIATIONS USED

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Dicloran (2,6-dichloro-4-nitroaniline), DBNA (2,6-dibromo-4-nitroaniline), PCP

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(pentachlorophenol), DCA (3,4-dichloroaniline), 2,4-D (2,4-Dichlorophenoxyacetic

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acid), EPA RED (Environmental Protection Agency’s Reregistration Eligibility Decision)

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ACKNOWLEDGEMENTS

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The authors thank the staff and other graduate students in the Department of

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Environmental Sciences at Louisiana State University for their support and assistance

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with this research, and also Andrea Warren and Amy Hernandez of Louisiana

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Department of Agriculture and Forestry for LC/MS analysis.

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REFERENCES 1. United States Environmental Protection Agency. Reregistration Eligibility Decision DCNA (Dicloran). (https://www3.epa.gov/pesticides/chem_search/reg_actions/reregistration/red_PC031301_14-Jun-06.pdf) (Jun. 14, 2016). 2. Roberts, T. R.; Hutson, D. H. Miscellaneous Fungicides. Metabolic Pathways of Agrochemicals, Part 2: Insecticides and Fungicides; The Royal Society of Chemistry, Cambridge, UK. 1999; 2., pp 1391-1394. 3. Gowan Company, LLC. (http://www.gowanco.com) (Feb. 26, 2016). 4. International Programme on Chemical Safety, INCHEM. Dicloran. (http://www.inchem.org/documents/jmpr/jmpmono/v074pr16.htm). (May 30, 2017). 5. Boscá, F.; Miranda, M. A.; Serrano, G.; Vargas, F. Photochemistry and Photobiological Properties of Dicloran, a Postharvest Fungicide with Photosensitizing Side Effects. Photochemistry and Photobiology. 1998, 67, 532537. 6. Hen, D.; Huang, B.; Xiong, D.; Pan, X. Effects of pH and dissolved oxygen on the photodegradation of 17α-ethynylestradiol in dissolved humic acid solution. Environ. Sci. Process Impacts. 2016, 18, 78-86. 7. Pinto, M. I.; Salgado, R.; Cottrell, B. A.; Cooper, W. J.; Burrows, H. D.; Vale, C.; Sontag, G.; Noronha, J. P. Influence of dissolved organic matter on the photodegradation and volatilization kinetics of chlorpyrifos in coastal waters. J. Photochem. Photobiol. A: Chem. 2015, 310, 189-196. 8. Miille, M. J.; Crosby, D. G. Pentachlorophenol and 3,4-dichloroaniline as models for photochemical reactions in seawater. Mar. Chem. 1983, 14, 111-120. 9. Wong, A. S.; Crosby, D. G. Photodecomposition of Pentachlorophenol in Water. J. Agro. Food Chem. 1981, 29, 125-130. 10. Miller, G. C.; Miille, M. J.; Crosby, D. G.; Sontum, S.; Zepp, R. G. Photosolvolysis of 3,4-dichloroaniline in water. Tetrahedron. 1979, 35, 17971800. 11. Crosby, D. G. The Photodecomposition of Pesticides in Water. Fate of Organic Pesticides in the Aquatic Environment; Los Angeles, American Chemical Society. 1972, 111, 173-188.

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12. Armbrust, K. L. Photodegradation of Hydroxychlorothalonil in Aqueous Solutions. Environ. Toxicol. Chem. 2001, 20, 2699-2703. 13. Nakagawa, M.; Crosby, D. G. Photodecomposition of Nitrofen. J. Agro. Food Chem. 1974, 22, 849-853. 14. Alloy M.; Baxter, D.; Stieglitz, J.; Mager, E.; Hoenig, R.; Benetti, D.; Grosell, M.; Oris, J.; Roberts, A. Ultraviolet Radiation Enhances the Toxicity of Deepwater Horizon Oil to Mahi-mahi (Coryphaena hippurus) Embryos. Environ. Sci. Technol. 2016, 50, 2011-2017. 15. McDonald, B. G.; Chapman, P. M. PAH phototoxicity – an ecologically irrelevant phenomenon? Mar. Poll. Bull. 2002, 44, 1321-1326. 16. California Farm Bureau Federation. (http://www.cfbf.com) (Feb. 29, 2016). 17. Pidwirny, M. 2006. Earth-Sun Relationships and Insolation. Fundamentals of Physical Geography, 2nd Edition. (http://www.physicalgeography.net/fundamentals/6i.html) (Mar. 14, 2016). 18. Streitwieser, A.; Heathcock, C. H. Aromatic Halides, Phenols, Phenyl Esters, and Quinones. Introduction to Organic Chemistry, 2nd Edition; Macmillan Publishing Co., Inc., New York, New York, 1981. pp 982-1029. 19. Nobuyuki, A.; Kudoh, S.; Nakata, M. UV Photolysis of 1,4-Diaminobenzene in a Low-Temperature Argon Matrix to 2,5-Cyclohexadiene-1,4-diimine via 4Aminoanilino Radical. J. Phys. Chem. 2003, 107, 6725-6730. 20. Beck, S. M.; Brus, L. E. Transient Raman Scattering Study of the Initial Semiquinone Radical Kinetics following Photolysis of Aqueous Benzoquinone and Hydroquinone. J. Am. Chem. Soc. 1982, 104, 4789-4792. 21. Land, E. J. Flash Photolysis and Pulse Radiolysis Studies of some Semiquinones in Relation to Cancer Induction and Therapy. J. Chem. Soc. 1986, 82, 2183-2188. 22. Alegria, A. E.; Ferrer, A.; Santiago, G.; Sepulveda, E.; Flores, W. Photochemistry of water-soluble quinones. Production of hydroxyl radical, singlet oxygen and the superoxide ion. J. Photochem. Photobiol. A: Chem. 1999, 127, 57-65. 23. Michalska, R.; Kurzyca, I. Determination of Nitrogen Species (Nitrate, Nitrite and Ammonia Ions) in Environmental Samples by Ion Chromatography. Pol. J. Environ. Stud. 2005, 15, 5-18. 24. Zafiriou, O. C.; True, M. B. Nitrite photolysis in seawater by sunlight. Mar. Chem. 1979, 8, 9-32.

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25. Zafiriou, O. C.; True, M. B. Nitrate photolysis in seawater by sunlight. Mar. Chem. 1979, 8, 33-42. 26. Mack, J.; Bolton, J. R. Photochemistry of nitrite and nitrate in aqueous solution: a review. J. Photochem. Photobiol. A: Chem. 1999, 128, 1-13. 27. Zhou, X.; Mopper, K. Determination of photochemically produced hydroxyl radicals in seawater and freshwater. Mar. Chem. 1990, 30, 71-88.

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FIGURE CAPTIONS Figure 1. The structure of 2,6-dicloro-4-nitroaniline (X = Cl) and 2,6-dibromo-4nitroaniline (X = Br). Figure 2. The ultraviolet absorption spectrum of dicloran (blue), 1,4-benzoquinone (red), and 2-chloro-1,4-benzoquinone (green) and the output of sunlight (purple) at 33.2°N latitude in mid-June, measured between 200-500 nm.10 The concentration of all chemicals was 10-5 M. Figure 3. The rate of degradation of dicloran in various media over the 24-hour photoperiod. = distilled water, = 3.2 % artificial seawater, = 0.01 M phosphate buffer (pH 7), = 0.8 % filter-sterilized estuarine water; error bars indicate standard deviation. Figure 4. The formation and degradation of 2-chloro-1,4-benzoquinone (A) and 1,4benzoquinone (B) in distilled water () and artificial seawater (), error bars indicate standard error. Concentrations with an asterisk (*) were below the LOD and not included in statistical analysis. Figure 5. The proposed photodegradation pathway of dicloran in distilled water. Blue boxed compounds were detected by HPLC and MS (A-B); red bracketed compounds are assumed, but were not detected; black boxed compounds were detected by MS only (CI); pink boxed compounds are assumed based upon degradation of similar compounds with the exception of maleic and fumaric acid which were detected below quantifiable values by HPLC only (J).

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TABLES Table 1. Photodegradation rates and half-lives of dicloran and DBNA in various media. rate (hr-1)

half-life (hours)

samples (n)a

dicloran distilled water

0.092 ± 0.0011

7.62 ± 0.094

5

3.2 % artificial seawater

0.094 ± 0.0036

7.37 ± 0.279

7

pH 7 10-5 M phosphate buffer

0.120 ± 0.0046

5.78 ± 0.216

3

filter-sterilized estuarine water

0.110 ± 0.0009

6.29 ± 0.048

3

distilled water

0.078 ± 0.0023

8.86 ± 0.257

4

3.2 % artificial seawater

0.083 ± 0.0014

8.31 ± 0.144

4

pH 7 10-5 M phosphate buffer

0.085 ± 0.0006

8.20 ± 0.059

3

filter-sterilized estuarine water

0.091 ± 0.0021

7.56 ± 0.181

3

DBNA

a

Number of samples after outliers removed, ANOVA used to determine outliers.

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Page 20 of 30

FIGURE GRAPHICS

Figure 1

0.25

0.6

Absorbance

0.4 0.15 0.3 0.1 0.2 0.05

0.1

0

0 200

250

300 350 400 Wavelength (nm)

Figure 2

20


450

500

Irradiance (W/ sq m)

0.5

0.2

Page 21 of 30


0 0

5

10

15

-0.5

ln C/Co

-1

-1.5

-2

-2.5

-3

Time (Hours)

Figure 3

21


20

25


Concentration (ppm)

A

Page 22 of 30

0.12 0.10 0.08 0.06 0.04 0.02

*

*

0.00 0

Concentration (ppm)

B

5

10 15 Time (Hours)

20

25

0.06 0.05 0.04 0.03 0.02 0.01

*

0.00 0

5

10 15 Time (Hours)

20

Figure 4

22


25

Page 23 of 30


Figure 5

23



Page 24 of 30

GRAPHIC FOR TABLE OF CONTENTS

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NO3-

NH2

!

Cl

Cl

Cl-

NH2 Cl

UPLANDS

%& '

%& '

Cl NO2

$

O

%& '

O

NO2

" " " "

Cl O

NO3

NH2 Cl

-

Cl

OCEANS

O

Cl-

# $ # $

# $ # $ O

O

NO2

# $ # $

Cl

# $ # $

O

O

24


NH

2 ofPage Agricultural 25 of 30and Food Ch

X

X

CS Paragon Plus Environmen

NO2

0.25

Page0.6 26 of 30


Absorbance

0.4

0.15

0.3 0.1

0.2

0.05

0.1

0

0 200

250

300Paragon Plus 350Environment 400 ACS Wavelength (nm)

450

500

Irradiance (W/ sq m)

0.5

0.2

Page 27 of 30


0 0

5

10

15

20

25

-0.5

ln C/Co

-1 -1.5 -2 -2.5 -3

Time (Hours)

Figure 3. The rate of degradation of dicloran in various media over the 24-hour photoperiod. u = distilled water,

1 Photodegradation of Dicloran in Freshwater and Seawater Emily N

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