Multifactor Statistical Analysis of H - ACS Publications

Mar 23, 2009 - School of CiVil Engineering, Purdue UniVersity, 550 Stadium Mall DriVe, West ... Ultraviolet light combined with chemical oxidants such...
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Ind. Eng. Chem. Res. 2009, 48, 3955–3963

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Multifactor Statistical Analysis of H2O2-Enhanced Photodegradation of Nicotine and Phosphamidon Amanda M. Nienow,† Inez Hua,* Irene C. Poyer,‡ Juan Ce´sar Bezares-Cruz, and Chad T. Jafvert School of CiVil Engineering, Purdue UniVersity, 550 Stadium Mall DriVe, West Lafayette, Indiana 47906

Direct and indirect photolysis (λ ) 254 nm) of aqueous nicotine and phosphamidon were studied. A central composite design was used to explore the effects of initial [H2O2] (0-10 mM), pH (2.8-11.2), and ionic strength (Ic ) 0.02-0.71 M) on the rate constants of nicotine and phosphamidon separately dissolved in a surface water surrogate matrix. Five levels of each factor were included in the design. For nicotine, the fastest predicted photochemical oxidation rate constant (k ) 0.43 min-1) occurred under the following conditions: pH ) 7.5, [H2O2] ) 7.5 mM, and Ic ) 0.02 M. This rate constant predicts that 90% of the nicotine will react within 5.4 min under these conditions. In general, the photochemical oxidation of nicotine is more rapid at lower ionic strength and near-neutral pH values. For phosphamidon, the fastest predicted oxidation rate constant (k ) 0.65 min-1) occurred at a pH of 7.1 and [H2O2] of 5 mM. Under these conditions, 90% of the phosphamidon would react within 3.5 min of treatment. Like nicotine, the photochemical oxidation of phosphamidon is more rapid at near-neutral pH values. Ionic strength has no significant effect on the photochemical oxidation of phosphamidon. Introduction Ultraviolet light combined with chemical oxidants such as hydrogen peroxide is a free-radical-based method that has been applied to the treatment of a variety of organic and inorganic water pollutants.1,2 It is well-established that a number of solution and reactor characteristics (e.g., pH, oxidant dosage, photon flux) will influence the reaction rate and therefore treatment effectiveness, and the optimal combination of conditions will vary depending on the chemical compound to be treated. In this paper, we apply a design of experiment (DOE) methodology to optimize a bench-scale UV/H2O2 system for photochemical oxidation of nicotine and phosphamidon separately dissolved in surrogate surface water. The experiments were completed within the framework of a central composite design to explore the effects of initial oxidant (H2O2) concentration, pH, and ionic strength (Ic) on the photochemical reaction rate constants of the two compounds. Nicotine and phosphamidon were chosen for this study because of their contrasting physicochemical properties and because both are chemical compounds of concern. In addition, use of multiple compounds allows for an analysis of the DOE methodology. Nicotine is a diprotic acid with pKa1 ) 3.37 ( 0.02 and pKa2 ) 8.07 ( 0.02.3 At pH values below 3.37, diprotonated nicotine (Figure 1a) is the principal species. Between pH values 4 and 8.07, monoprotonated nicotine (Figure 1b) is dominant. At pH values greater than 8.07, the neutral form of nicotine (Figure 1c) is present in the greatest quantity. These three forms of nicotine can undergo direct and indirect photolysis at different rates. In addition, the charged nature of the diprotonated and monoprotonated forms can lead to changes in chemical activity as a function of ionic strength. The acid-base speciation of nicotine and the changes in chemical activity provide a basis to explore how physicochemical properties can affect photochemical oxidative efficiency. * Author to whom correspondence should be addressed. Telephone: (765) 494-2409. Fax: (765) 496-1988. E-mail: [email protected]. † Present address: Gustavus Adolphus College, 800 W. College St., St. Peter, MN 56082. ‡ Present address: ACT I Research and Analytical Laboratory, 316 Kelly Dr., Waco, TX 76710.

There are numerous sources of nicotine to the environment; according to information in the Toxic Release Inventory,4 the estimated release from various U.S. industrial facilities was over 2.2 million pounds between 1995 and 1997 and ∼6.5 million pounds between 1998 and 2005. Pure nicotine is extremely toxic; the probable oral lethal dose in humans is less than 5 mg/kg.5 There are increasing concerns about nicotine as a threat to food and water security. In 2003, the Centers for Disease Control and Prevention (CDC) reported a case of deliberate contamination of ground beef with nicotine.6,7 There has been recent testing and installation of nicotine sensors in drinking water plants,8,9 and nicotine and its metabolite cotinine have been detected in drinking water sources10 and other surface waters in the United States.11 In a recent study, nicotine was detected in surface waters at concentrations ranging from 0.2 to 0.9 µg/L.12 In addition, military exposure guidelines (MEGs) exist for nicotine in drinking water.13 The MEG for nicotine for an adult drinking 5 L of water for less than 7 days is 0.4 mg/L; and a 2-5 mg total dose will cause nausea in adults. Clearly, there is concern about mitigating human exposure to nicotine. However, the technology to treat nicotinecontaminated water is lacking. There are reports that describe nicotine oxidation or photochemistry;14,15 however, there is little contemporary literature on the reactivity of nicotine in aqueous systems. Phosphamidon is an organophosphorus herbicide with cis and trans isomers (Figure 2). Unlike nicotine, phosphamidon does not ionize in the range of pH values examined here (∼3-11). Also in contrast to nicotine, the structure of phosphamidon includes some oxygenated functional groups, and the compound

Figure 1. Acid-base species of nicotine: (a) diprotonated, (b) monoprotonated, and (c) unprotonated.

10.1021/ie801311f CCC: $40.75  2009 American Chemical Society Published on Web 03/23/2009

3956 Ind. Eng. Chem. Res., Vol. 48, No. 8, 2009

Figure 2. Cis (Z) and trans (E) isomers of phosphamidon. Table 2. Nicotine Experimental Conditions and Corresponding Rate Constants

Table 1. Experimental Factors factor levelsa factor (units)

-2

-1

0

1

2

x1 ) H2O2 (mM) x2 ) pH x3 ) ionic strength Ic (M)

0 2.8 0.02

2 4.5 0.16

5 7 0.36

8 9.5 0.57

10 11.2 0.71

a

Coded factors used by Design-Expert program.

undergoes hydrolysis. Both differences may lead to different results in the DOE analysis. Commercial formulations of phosphamidon were banned from use in the United States in 1998.16 However, development of treatment technologies is relevant because of phosphamidon’s continued availability in many countries outside the United States,17 its current status on the World Health Organization’s (WHO) 1a inventory of extremely toxic chemicals,18 and its high risk ranking as acutely toxic to deployed military personnel if ingested via contaminated drinking water.19 The destruction of organophosphorus (OP) compounds as a class has been explored, and free-radical based methods are effective against some types of OP compounds.20-32 However, with the exception of published works by Rabindranathan et al.31 in 2003 and Rahman et al.32 in 2005, there is limited recent information about the photochemical oxidation of phosphamidon. The primary objectives of the research were to demonstrate rapid decomposition of nicotine and phosphamidon in surrogate natural water and to determine the combination of initial conditions (H2O2 dose, pH, and ionic strength) that would result in the fastest reaction rates. Therefore, we employed a central composite design to build a second-order model for the response variable, and we build upon previous applications of multifactor analysis to systematically explore solution variables.33-35 In this study, the response variable is the observed reaction rate constant of nicotine or phosphamidon during exposure to UV/H2O2 under various combinations of initial conditions. The factors in the model are H2O2 dose, pH, and ionic strength. Five levels of each factor were included in the design. The data provide guidance for process optimization and also provide insight into how chemical activity and acid-base speciation substantially determine the efficacy of photochemical oxidation. Experimental Methods Experimental Design. A central composite design was used to explore the effects of initial [H2O2], pH, and ionic strength (Ic), the three factors in the design, on photochemical rate constants.36 Five levels of each factor were included in the design. Table 1 summarizes the factors and levels used in the design. There were six replicates of the center point of the design (i.e., 5 mM H2O2, pH 7, and Ic ) 0.36 M); the axial and factorial points were examined only once. With this design, a list of 20 experiments with 14 initial conditions was generated by use of Design-Expert, Version 7.0.2, Stat-Ease, Inc. The list of nicotine experiments is presented in Table 2; the corresponding phos-

run

block

H2O2 (mM)

pH

Ic (M)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

1 1 2 2 2 2 2 2 2 2 2 2 2 2 3 3 3 3 3 3

5 10 5 8 2 5 5 2 8 2 5 0 5 8 5 5 5 8 2 5

7 7 7 4.5 9.5 7 7 9.5 9.5 4.5 7 7 2.8 4.5 7 11.2 7 9.5 4.5 7

0.36 0.36 0.36 0.57 0.16 0.36 0.36 0.57 0.57 0.16 0.71 0.36 0.36 0.16 0.36 0.36 0.36 0.16 0.57 0.02

a

ka (min-1) 0.34 0.45 0.19 0.17 0.24 0.30 0.28 0.18 0.27 0.22 0.28 0.10 0.06 0.23 0.29 0.15 0.37 0.37 0.21 0.41

( ( ( ( ( ( ( ( ( ( ( ( ( ( ( ( ( ( ( (

0.04 0.06 0.03 0.05 0.06 0.04 0.03 0.09 0.14 0.02 0.08 0.03 0.03 0.04 0.04 0.01 0.08 0.06 0.02 0.03

Error bars are at the 95% confidence interval.

Table 3. Phosphamidon Experimental Conditions and Corresponding Rate Constants run

block

H2O2 (mM)

pH

Ic (M)

ka (min-1)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

1 1 1 1 1 2 2 2 2 2 2 2 2 3 3 3 3 3 3 3

5 10 5 8 2 5 5 2 8 2 5 0 5 8 5 5 5 8 2 5

7.1 7.1 7.1 4.3 9.5 7.1 7.1 9.5 9.5 4.5 7.0 7.0 2.3 4.6 7.1 11.2 7.0 9.5 4.6 7.1

0.37 0.37 0.37 0.57 0.16 0.37 0.37 0.57 0.57 0.16 0.70 0.37 0.37 0.16 0.37 0.37 0.37 0.16 0.57 0.002

0.57 ( 0.15 0.520 ( 0.10 0.50 ( 0.16 0.42 ( 0.09 0.30 ( 0.08 0.65 ( 0.10 0.57 ( 0.08 0.295 ( 0.09 0.24 ( 0.31 0.46 ( 0.17 0.58 ( 0.14 0.20 ( 0.02 0.22 ( 0.09 0.33 ( 0.07 0.32 ( 0.13 nab 0.50 ( 0.23 0.27 ( 0.11 0.29 ( 0.05 0.35 ( 0.26

a Error bars are at the 95% confidence interval. b Phosphamidon decomposed in solution before UV irradiation commenced.

phamidon experiments are shown in Table 3. As noted in Tables 2 and 3, the 20 experiments for each compound were separated into three blocks. This allowed for experiments to be run on different days while any systematic changes were taken into account during the statistical analysis. After the experiments were conducted, the Design-Expert software was used for data analysis to build a statistical model. Materials. All solutions were prepared with Barnstead Nanopure 18 MΩ water. Neat (-)-nicotine, 99%, was purchased

Ind. Eng. Chem. Res., Vol. 48, No. 8, 2009 3957 Table 4. Composition of Natural Water Surrogate compd

mg/L

mM

MgSO4 NaHCO3 FeSO4 · 7H2O Ca(NO3)2 · 4H2O KHCO3 CaCO3 IHSS NOM (1R101N)

200.17 125.10 0.66 3.93 80.10 180.73 9.99

1.663 1.489 2.4 × 10-3 0.0166 0.800 1.806 na

from Sigma-Aldrich and was used as received. Neat phosphamidon (Z + E isomer mix) was purchased from Supelco Inc. at 98% purity (71.2% Z isomer and 27.8% E isomer, 2% production byproducts). Dichloromethane solvent (GC Resolvgrade) used for extraction and to prepare stock and working standards was purchased from Fisher Scientific. A 30% weight by volume hydrogen peroxide solution, used to catalyze the reaction process, was obtained from Mallinckrodt. Suwannee River natural organic matter (NOM) 1R101N was purchased from International Humic Substances Society (IHSS) and was used as received. ReagentPlus-grade magnesium sulfate (g99.5%), ACS reagent-grade sodium bicarbonate, ACS reagentgrade iron(II) sulfate heptahydrate (99+%), USP-grade potassium bicarbonate, ACS reagent-grade calcium carbonate (99+%), and ACS reagent-grade calcium nitrate tetrahydrate (99%), used in preparing the natural water surrogate, were all purchased from Sigma-Aldrich and were used as received. Sodium hydroxide, hydrochloric acid, and sodium chloride, used to adjust the pH and ionic strength of the nicotine and phosphamidon solutions, were purchased from Mallinckrodt (NaOH, ACS reagent-grade) and Sigma-Aldrich (HCl 37% and NaCl g99.5%). Natural Water Surrogate. Photochemical experiments were conducted in an aqueous matrix representative of natural surface water; this matrix will be referred to as natural water surrogate (NWS). The inorganic composition was chosen by considering the occurrence and concentrations of relevant inorganic compounds in surface waters.37 Table 4 gives the recipe for NWS (Ic ) 0.02 M; pH ) 8.1). The organic constituent in the natural water surrogate consisted of International Humic Society Substances (IHSS) natural organic matter (NOM). Typically, there is 0-10 mg of NOM/L present in surface waters.38,39 Screening tests with NOM and nicotine or phosphamidon in deionized water showed that increasing NOM concentrations in the solutions resulted in a decrease in rate constants. Thus, NOM is a potential factor to investigate. However, in the experiments presented here, this factor was held constant at 10 mg/L. As increasing NOM concentrations always caused a decrease in degradation rates in the screening tests, the factor was not varied in order to reduce the overall number of experiments to be conducted. NWS was prepared in a large carboy and stored at room temperature. The carboy was covered with black fabric, and the experiments were conducted within 3 weeks of preparation. These precautions limited any fluctuations in the water quality and any biological growth. Photochemical Experiments. Photochemical reactions were carried out in a Rayonet RPR-100 photochemical reactor (Southern New England Ultraviolet Co.) equipped with eight 254-nm lamps. Photochemical oxidation experiments with phosphamidon were conducted over a period of 3 days to minimize variations in method or laboratory conditions. Nicotine experiments were conducted in three sets (1-2 days each) over a period of 2 weeks. When analyzed with the Design-Expert software, the data were thus separated into three blocks based on the day of experiment. As mentioned above, this allowed

the software to take any systematic changes into account during the statistical analysis. In each experiment, samples were collected at a series of time points (t ) 0, 1, 2, 3, 4, 6, 8, and 10 min for nicotine experiments and t ) 0, 0.5, 1, 1.5, 2, 4, 7, 10, and 15 min for phosphamidon). Immediately prior to photochemical experiments with nicotine, aqueous nicotine solutions were prepared in NWS with the designated ionic strength and pH. When necessary, the ionic strength of the solution was adjusted by adding NaCl and/or the pH was adjusted by adding small amounts of concentrated HCl or 0.5 M NaOH. The pH was measured with an Orion pH probe, and the ionic strength was calculated from the known concentrations of ions in the solution. Carbonate and bicarbonate concentrations were determined from the pH and the acidity constants of the different species. Eight microliters of neat (-)nicotine was added to every 500 mL of prepared solution, and the solutions were allowed to equilibrate for approximately 2 h. The final concentration of nicotine in each experiment was ∼98 µM. Prior to each experiment with phosphamidon, an aliquot of phosphamidon standard (∼3 mL of ∼16.7 mM), previously prepared in dichloromethane (DCM), was transferred into a 500 mL amber volumetric flask and plated onto the glass surface by use of nitrogen gas to evaporate the solvent. To the flask, 500 mL of NWS and a Teflon-coated magnetic stir bar were added. The flask was sealed with an amber glass stopper and Parafilm, and the solution was mixed for approximately 24 h on a magnetic stir plate. As with the nicotine experiments, the ionic strength of the solutions was adjusted by adding NaCl and/or the pH was adjusted by adding small amounts of concentrated HCl or 0.5 M NaOH prior to irradiation. Prior to irradiation, H2O2 (0-10 mM) was added to the equilibrated nicotine or phosphamidon solutions. A portion of these mixtures (∼50 mL) was set aside for monitoring any concentration changes occurring in the dark. The remaining solution was placed into a quartz reaction vessel (660 mL) inside the Rayonet RPR-100 photochemical reactor equipped with eight 254-nm lamps. The 660-mL quartz reaction vessel was held in the center of the RPR-100, and a glass magnetic stir bar mixed the solution throughout irradiation. These solutions were irradiated for up to 15 min with samples taken at the time points outlined above. At each time point, one or two 5-mL aliquots were removed via volumetric pipettes. The first 5 mL of solution removed was used for pH measurement (nicotine experiments only) and the second 5 mL was used for quantification of the remaining nicotine or phosphamidon. An aliquot of 0.5 M NaOH (5 mL) was added to the nicotine samples to raise the pH to ∼11 (thereby increasing extraction efficiency because this pH favors the neutral nicotine species), and the sample (total 10 mL) was then extracted with 3 mL of dichloromethane (DCM). The extraction was completed by rotating the samples on a GlasCol rotator for ∼5 min. The phosphamidon samples were added to individual amber vials containing 3 mL of DCM as the quenching reagent. Vials were placed on a Glas-Col rotator and extracted for 24 h at 20 rpm rotation speed. All nicotine sample extracts were analyzed on a Shimadzu 17A GC equipped with a flame ionization detector (GC-FID) equipped with an AOS-20S Shimadzu autosampler and an AOC20i Shimadzu autoinjector. Compound separation was accomplished on a J&W Scientific DB1 1- m by 0.25-mm i.d. column with a 0.10 µm film thickness. The GC-FID was run splitless with a sampling time of 0.80 min and a purge split flow at 25 mL/min. The injector and FID temperatures were set at 280 and 290 °C, respectively. One microliter of sample

3958 Ind. Eng. Chem. Res., Vol. 48, No. 8, 2009

was injected. The following oven temperature program was used: start at 50 °C and hold for 2 min, ramp to 100 at 15 °C/ min and hold 0.1 min, and ramp to 230 at 15 °C/min and hold 0.2 min. The total run time was 10.1 min. The phosphamidon DCM extracts were analyzed on a Thermo Finnigan Polaris Q gas chromatograph-mass spectrometer (GC/ MS) equipped with a Varian Factor 4 VF4 30 m × 0.32 mm × 0.25 µm film thickness analytical column. System temperatures were as follows: injection port, 150 °C; transfer line, 300 °C; and ion source volume, 200 °C. The column temperature program initiated at 35 °C, held for 4 min, andthen ramped to 280 at 10 °C/min with a final hold of 1 min. Mass spectra acquisition was in full scan mode of 50-600 m/z with quantitation of the parent compound based on the base peak ion of 127 m/z (126-128 m/z inclusive). The extracts were also analyzed on a Varian CP3800 gas chromatograph equipped with a pulsed flame photometric detector (GC/PFPD) in phosphorusspecific mode and a Varian Factor 4 VF4 (15 m × 0.25 mm × 25 µm film thickness) analytical column. GC temperatures were maintained at PFPD (310 °C) and injection port (150 °C); and the column oven program set initially to 80 °C and ramped to 240 °C at 10 °C/min. In addition to the photochemical experiments conducted to fulfill the DOE study, we also completed some initial photochemical experiments on nicotine (pH ) 1, 5.5, 7, and 11.3) and phosphamidon (pH ) 7) separately dissolved in phosphatebuffered reagent-grade water. The experiments were in the same photochemical reactor as previously described, and the reaction quenching, extraction, and analysis were also the same. Actinometry and Reactor Characteristics. The photon flux of the lamps in the Rayonet photoreactor was determined via chemical actinometry (with potassium ferrioxalate).40 The specifics of these measurements for the Rayonet reactor at Purdue can be found in Nienow et al.41 All experiments reported in this paper were completed with a 660-mL quartz reaction vessel held in the center of the Rayonet reactor with eight 253.7nm lamps surrounding the vessel (photon flux ) 7.6 × 10-6 einstein · s-1). Results and Discussion Twenty photochemical experiments were conducted with each compound. All degradation kinetics were approximated as pseudo-first-order, with kobs as the pseudo-first-order rate constant calculated from weighted linear least-squares analysis of the experimental data.42 The weights used for the regression were 1/σy2 (σy2 is the variance of the data), which are proportional to the squares of the sample concentration. The use of this weight effectively reduces the effect of the later time points (i.e., where concentrations are low and measurements likely have greater variance and error) on the fit of the regression line, yielding more accurate reaction rate constants. Phosphamidon results are presented as a combined rate constant (kobs) for the Z + E isomer mixture. Figure 3 gives degradation plots for the center point of the design (pH 7, 5 mM H2O2, and Ic of 0.37 M) for both nicotine and phosphamidon. The figure includes error bars at the 95% confidence interval and is representative of the degradation plots for the entire set of experiments. In all cases, the R2 values of the pseudo-first-order weighted linear regression fits were 0.91 or greater. Tables 2 and 3 indicate the rate constants that were input into DesignExpert as the response factor of the design. The contributions of each factor (x1 ) H2O2, x2 ) pH, and x3 ) Ic) to k was evaluated by statistical analysis. No data transform was used as was suggested by the Design-Expert

Figure 3. Degradation plots of (2) phosphamidon and ([) nicotine at pH 7 with 5 mM H2O2 and an ionic strength of 0.37 M. The lines are weighted least-squares linear regression fits to the data and errors bars represent the 95% confidence interval. Table 5. ANOVA Results for Full Quadratic Model for Nicotine parameter

factor

value

SE

β0 β1 β2 β3 β12 β13 β23 β11 β22 β33 lackoffit

intercept [H2O2] pH Ic [H2O2]-pH [H2O2]-Ic pH-Ic [H2O2]-[H2O2] pH-pH Ic-Ic

0.3122 0.0390 0.0207 -0.0202 0.0167 0.0113 0.0016 -0.0123 -0.0653 0.0243

0.0241 0.0157 0.0172 0.0153 0.0256 0.0218 0.0256 0.0157 0.0173 0.0151

t-value

p>t

6.1475 1.4408 1.7305 0.4268 0.2699 0.0037 0.6128 14.2092 2.5933 0.6580

0.0381 0.2643 0.2248 0.5319 0.6175 0.9528 0.4563 0.0055 0.1460 0.6822

Table 6. ANOVA Results for Full Quadratic Model for Phosphamidon parameter

factor

value

SE

β0 β1 β2 β3 β12 β13 β23 β11 β22 β33 lackoffit

intercept [H2O2] pH Ic [H2O2]-pH [H2O2]-Ic pH-Ic [H2O2]-[H2O2] pH-pH Ic-Ic

0.5248 0.0490 -0.0700 -0.0181 0.0091 -0.0390 -0.0199 -0.0734 -0.1313 -0.0035

0.0396 0.0297 0.0331 0.0303 0.0385 0.0445 0.0428 0.0273 0.0314 0.0271

t-value

p>t

2.7258 4.4865 0.3578 0.0563 0.7684 0.2166 7.2429 17.4644 0.01700 4.1443

0.1427 0.0719 0.5686 0.8193 0.4098 0.6558 0.0310 0.0041 0.8999 0.2057

software (the ratio of response maximum to response minimum was less than 10 for both the nicotine and phosphamidon experiments). Initially, a full quadratic model (eq 1) was fitted to the data. This model includes an intercept, β0; linear terms, β1x1, β2x2, and β3x3; quadratic terms, β11x12, β22x22, and β33x32; and cross-product terms to test for possible factor interactions, β12x1x2, β13x1x3, and β23x2x3. k ) β0 + β1x1 + β2x2 + β3x3 + β12x1x2 + β13x1x3 + β23x2x3 + β11x12 + β22x22 + β33x32 (1) Coefficients (βx) for each term were generated and tested for significance with an analysis of variance (ANOVA). Coefficients, standard error, corresponding t-statistics, and probability factors (i.e., p-values) for the full quadratic analyses for nicotine and phosphamidon are given in Tables 5 and 6, respectively. Further statistical analysis and results for nicotine and phosphamidon will be discussed separately below. The physical significance of the statistical analysis will also be presented for each compound.

Ind. Eng. Chem. Res., Vol. 48, No. 8, 2009 3959 Table 7. ANOVA Results for Reduced Quadratic Model for Nicotine parameter

factor

value

SE

β′0 β′1 β′2 β′3 β′22 β′33 lackoffit

intercept [H2O2] pH Ic pH-pH Ic-Ic

0.303 0.041 0.020 -0.019 -0.065 0.024

0.018 0.014 0.015 0.013 0.015 0.013

t-value

p>t

8.711 1.903 2.138 17.957 3.314 0.476

0.0121 0.1929 0.1624 0.0012 0.0937 0.8297

Nicotine. The full quadratic model for nicotine from the statistical analysis (see eq 1) predicted k with an R2 value of 0.82 for all experimental conditions. Only two factor effects were found to be significantly different from zero at the 95% confidence interval: β1x1and β22x22. These factors correspond to the linear effect of H2O2 and the quadratic effect of pH. Backward elimination, the elimination of terms with p > 0.05 one term at a time, was used to produce a more accurate model, in this case a reduced quadratic model. Through this process, the coefficients and p-values for the remaining terms will be altered slightly. However, if the reduced model is capturing the major effects, the R2 term will not be altered significantly. Terms eliminated from the full quadratic model were β12x1x2, β13x1x3, β23x2x3, and β11x12. β33x32 had a p-value of 0.0937; because the term is marginally significant (i.e., nearly significant at the 95% confidence interval and significant at the 90% confidence interval), it was kept in the reduced model. β3x3 was therefore kept in the model for hierarchy. The remaining terms are part of a reduced quadratic model given in eq 2: k′obs ) β′0 + β′1x1 + β′2x2 + β′3x3 + β′22x22 + β′33x32

(2)

The recalculated coefficients, β′x, and the corresponding t-statistics and p-values are found in Table 7. The R2 for the reduced quadratic model fit, 0.77, is relatively unchanged from the R2 for the full quadratic model fit. This suggests that the reduced model is indeed capturing the major effects. In addition, the RSD on k was ∼20%. The effects remaining include an intercept (the average rate constant at the center point), a firstorder dependence on [H2O2], and a first- and second-order dependence on both pH and ionic strength. The Design-Expert software can do a statistical optimization based on the ANOVA results. Here, the rate constant was maximized as a function of the three factors: [H2O2], pH, and ionic strength. The software found 33 solutions to this optimization problem, with six solutions giving maximum rate constants between 0.426 and 0.427 min-1. These six solutions had an ionic strength of 0.02 M, pH values between 7.66 and 7.81, and H2O2 concentrations ranging from 7.71 to 8.86 mM. The weaker dependence of the statistical model on H2O2 concentration is clearly illustrated by the wider variety of concentrations giving essentially the same rate constants, as compared to the much tighter variation in ionic strength and pH. Alternatively, at high ionic strengths (i.e., 0.70 M), the maximum rate constants predicted by Design-Expert were between 0.414 and 0.416 min-1. In this case, there were nine solutions with H2O2 concentrations always at 10 mM and a variation in pH from 7.63 to 8.14. Thus, in near-neutral solutions, with both high and low ionic strengths, nicotine can be photodegraded fairly rapidly (t1/2 ) 1.63-1.67 min) with the addition of ∼8-10 mM H2O2. These predictions can be explained by considering •OH generation, nicotine speciation, and variations in activity coefficients over a range of ionic strengths. Nicotine Photochemistry. Prior to the experiments pertaining to the central composite design, a few experiments with nicotine

Figure 4. (a) Pseudo-first-order rate constants, kobs, for direct photolysis of nicotine as a function of pH. (b) Speciation of nicotine as a function of pH.

were conducted to test the influences of UV light, H2O2, and the UV/H2O2 combination on nicotine degradation. These experiments involved measuring the concentration change of nicotine as function of time when (a) nicotine was exposed to 254-nm UV light, (b) nicotine and H2O2 were combined in the dark, and (c) nicotine and H2O2 were irradiated at 254 nm together. The majority of these experiments were conducted at pH 7, where the monoprotonated form of nicotine is the primary species present (91%). The experiments in which nicotine was irradiated at 254 nm without hydrogen peroxide show that nicotine does degrade via direct photolysis (k ) 0.19 ( 0.03 min-1). In experiments with 5-20 mM H2O2 added to the nicotine solutions in the dark, no reaction occurred. However, all reactions with UV/H2O2 and nicotine were 2-3 times faster than direct photolysis of nicotine. For example, at pH 7 with an initial H2O2 concentration of 1 mM, the degradation rate of nicotine was 0.34 ( 0.03 min-1. With 5 mM H2O2, the degradation rate was 0.57 ( 0.03 min-1. The enhancement in the degradation of nicotine is most likely due to reaction with •OH which forms when H2O2 is irradiated at 254 nm. The reported rate constant of •OH attack on nicotine is ∼4.5 × 108 M-1 s-1.43,44 Thus, it can be concluded that the dominant pathway of nicotine degradation in the DOE experiments will be reaction with •OH. To determine which nicotine species exhibits the greatest photodegradation rate, experiments were conducted at pH 1, 5.5, and 11.3, where the diprotonated, monoprotonated, and unprotonated nicotine species dominate, respectively. Figure 4 overlays the observed pseudo-first-order rate constants, kobs, and the speciation of nicotine as a function of pH. As can be seen in the figure, the fastest nicotine observed reaction rates were under conditions where the monoprotonated nicotine species is the most abundant. This outcome was true regardless of whether UV or UV/H2O2 was applied to the nicotine solution. These preliminary experiments help one frame the statistical results from the central composite design in terms of a mechanistic explanation. This will be more fully explored below. The initial H2O2 concentration plays a very important role in the observed nicotine photochemical oxidation rates because • OH is generated from H2O2. However, at high H2O2 concentra-

3960 Ind. Eng. Chem. Res., Vol. 48, No. 8, 2009

Figure 6. Nicotine reaction rate constant (k) distribution as a function of the initial [H2O2] and Ic conditions (pH ) 7.5). Contours indicate combinations of ionic strength and H2O2 that the model predicts will yield a constant value of k. At higher initial [H2O2], rate constants decrease moderately when Ic increases from 0.02 to 0.57 M. However, at higher ionic strengths, rate constants increase significantly as initial [H2O2] increases. Figure 5. Nicotine reaction rate constant (k) distribution as a function of the initial pH and [H2O2] conditions (Ic ) 0.02 M). Contours indicate combinations of pH and H2O2 that the model predicts will yield a constant value of k.

tions, inner filter effects due to light absorbance by H2O2 can reduce the uniformity and amount of light passing through the sample. This effect and rapid •OH recombination decreases the overall rate of formation of •OH at high H2O2 concentrations.41 Thus, one would expect the nicotine degradation rate constants to increase with H2O2 concentration until some critical concentration is reached. The statistical models show a linear dependence of rate constants on the concentration of H2O2. Due to the positive sign of β1′ , the highest rate constants will occur at the highest concentrations of H2O2. This suggests that the optimal concentration of H2O2 lies close to or above 10 mM for the solutions prepared and tested here. However, as can be seen in Figures 5 and 6, at low ionic strength and pH ∼8.3, there is a slight drop in rate constants at the highest H2O2 concentrations. Only under these conditions is the critical concentration of H2O2 reached. The ionic strength of the solution will play a significant role in nicotine speciation and observed reaction rates (Figures 6 and 7). The increase or decrease of ionic strength produces changes in the nicotine species concentrations when compared to a reaction in an ideal solution (e.g., low ionic strength). Activity coefficients corrections were performed by employing a modified Davies equation (eq 3) for the relevant conditions (Ic ) 0.02-0.71 M).45,46 -log γNz ) Az2

√Ic 1 + √Ic

- 0.3Ic

(3)

where γNz ) activity coefficient of nicotine ion with charge z, A ) 0.51 at 25 °C, and Ic ) ionic strength. This equation indicates that as the ionic strength increases, the activity coefficient of the diprotonated species will decrease much more rapidly than that of the monoprotonated species. For example, at Ic ) 0.36 M, γN1 ) 0.5 whereas γN2 ) 0.13. A decrease in the magnitude of the activity coefficient indicates a lower “effective” concentration, in this case due to shielding effects (the negatively charged chloride ions will shield the

Figure 7. Nicotine reaction rate constant (k) distribution as a function of the initial pH and Ic conditions ([H2O2] ) 7.5 mM). Contours indicate combinations of ionic strength and pH that the model predicts will yield a constant value of k. Regardless of the value of Ic, a near-neutral solution pH results in faster nicotine decomposition.

positively charged nicotine ions). Because free radical attack is a second-order process, a decrease in effective concentration of one reactant (e.g., the nicotine) will result in an overall decrease in the observed reaction rate. In addition to the changes in effective concentration of each species at a particular ionic strength, the change in concentration of each nicotine species relative to the other two also needs to be considered, because each species exhibits a different reaction rate when UV/H2O2 is applied. In nonideal solutions, the infinite dilution equilibrium constants (Ka1 and Ka2) must be corrected. The full equations for Ka1 and Ka2 are shown in eqs S1 and S2 in Supporting Information. In the experiments described in this paper, the solutions were maintained at a constant ionic strength over the course of the reactions, and therefore the appropriate equilibrium constants are cKa1 and cKa2. Functional relationships between cK and K are shown in eqs S3 and S4 in Supporting Information.47 The magnitudes of cKa1 and cKa2 diminish as the ionic strength increases, indicating that both the diprotonated and monoprotonated forms of nicotine become weaker acids (see Figure S1 in Supporting Information).

Ind. Eng. Chem. Res., Vol. 48, No. 8, 2009 3961 Table 8. ANOVA Results for Reduced Quadratic Model for Phosphamidon parameter

factor

value

SE

β′0 β′1 β′2 β′11 β′22 lackoffit

intercept [H2O2] pH [H2O2]-[H2O2] pH-pH

0.520 0.044 -0.067 -0.069 -0.130

0.0028 0.0023 0.0027 0.0022 0.0025

t-value

p>t

3.52 6.23 9.84 25.56 2.45

0.0850 0.0281 0.0086 0.0003 0.3248

Using the corrected equilibrium constants, we generated a series of log C versus pH diagrams for the nicotine species. The log C versus pH diagrams for a series of ionic strengths from 0 to 0.71 M are presented as Figures S2-S7 in Supporting Information. Clearly, the concentrations of each species are strongly dependent on the pH, and because the monoprotonated species reacts most quickly, it would be expected that the pH range in which the monoprotonated species is predominant (pH ) 4-7) would result in the fastest overall degradation of nicotine. Figures S2-S7 also show that the relatiVe concentration of each nicotine species with respect to the other two does not change substantially over the range of ionic strength conditions in our experiments, suggesting that ionic strength has a lesser effect on observed rate constants than the pH of the solution. In addition to speciation of nicotine, the pH also affects speciation of other •OH scavengers: the carbonates. The carbonate ion reacts more quickly with •OH than the bicarbonate ion (k ) 3.8 × 108 M-1 s-1 versus 8.5 × 106 M-1 s-1) and competes with nicotine for •OH.48 Thus, the observed reaction rates of nicotine would also decrease at higher pH values because carbonate ion would interfere with •OH attack of nicotine. Another •OH scavenger is the peroxy anion (HO2-). The reaction rate constant for HO2- scavenging of •OH (k ) 7.5 × 109 M-1 s-1) is quite fast, and it is comparable to the rate constant for •OH attack on the target compounds.48,49 The pKa of H2O2 is 11.6.50 Thus, for the few experiments conducted at pH > ∼11, the decreased reaction rate constant may be due to the scavenging of •OH by both CO32- and HO2-. However, for the majority of the experiments pH < 11, and the predominant species of hydrogen peroxide is H2O2. Phosphamidon. Statistical analysis of the phosphamidon data was first modeled with a full quadratic model, and then backward elimination of the insignificant terms was conducted. The full quadratic model predicted k with an R2 value of 0.80 for all experimental conditions. After backward elimination, the factors remaining were the intercept and first- and second-order [H2O2] and pH terms (i.e., all cross-product terms and the independent ionic strength terms were eliminated). The R2 value after elimination was 0.76 and the RSD was again ∼20%, suggesting that all significant factors are accounted for in the reduced model. The reduced quadratic model is described by eq 4: kobs ) β′0 + β′1x1 + β′2x2 + β′11x12 + β′22x22

(4)

where x represents the relevant input variables [H2O2] and pH. Elimination of the ionic strength terms (x3 and x32) is not unexpected because the parent compound phosphamidon is not an ion, and therefore its activity is not influenced significantly by the ionic charge of the matrix. The coefficients, β′x, and the corresponding t-statistics and p-values for the reduced quadratic model are found in Table 8. Numerical optimization of the phosphamidon rate constant was completed by use of the reduced quadratic model and the

Figure 8. Three-dimensional model distribution of the DOE design data and input parameters for phosphamidon. Contours indicate combinations of pH and initial [H2O2] concentration that the model predicts will yield a constant value of k.

Design-Expert software. On the basis of the reduced quadratic model, the rate constant was maximized as a function of just two factors: [H2O2] and pH. Unlike the optimization completed with the nicotine data, the software found only one solution to the optimization problem: the optimal oxidant concentration required to degrade phosphamidon in NWS was calculated to be 5.94 mM H2O2 at a pH of 6.34. The predicted rate constant (kobs) was 0.534 min-1. This rate constant is slightly higher than the average experimental rate constant (kobs ) 0.520 ( 0.30 min-1) of the six center-point experiments (runs 1, 3, 6, 7, 15, and 17 in Table 3). In these experiments, the H2O2 concentration was 5 mM and the pH was 7. Phosphamidon Photochemistry. The same series of preliminary experiments conducted with nicotine was repeated with phosphamidon to elucidate the dominant reaction pathway in the UV/H2O2 experiments in the central composite design. It was found that, like nicotine, reaction with •OH will be the major reaction pathway for phosphamidon. Phosphamidon reacts when exposed directly to UV light (254 nm); at pH 7, the pseudofirst-order rate constant for this reaction is 0.27 ( 0.04 min-1. Adding 1 mM H2O2 to this reaction increases the rate constant to 0.37 ( 0.04 min-1; adding 5 mM H2O2 increases the rate even further to 1.1 ( 0.2 min-1. This enhancement in rate constants can be attributed to the formation of •OH. Like nicotine, phosphamidon does not react with H2O2 in the dark. Initial H2O2 concentration is a significant factor for the photochemical reactions of phosphamidon for the same reasons discussed for nicotine above. In essence, the H2O2 serves as a chemical source of oxidant in the form of •OH. Increasing the H2O2 concentration increases the concentration of oxidant, thereby leading to an increase in rate constants. As discussed above, at high H2O2 concentrations, the rate constants start to decline. Statistically, this is modeled by including both linear and quadratic terms. Graphically, the effect can be seen in Figure 8, a 3D model distribution of the DOE design data and input parameters. Here the curvature in the 3D surface as a function of initial H2O2 concentration can be easily observed. Although phosphamidon does not ionize within the pH range in our studies, pH will certainly influence the observed

3962 Ind. Eng. Chem. Res., Vol. 48, No. 8, 2009

photochemical rate constants. The pH affects speciation of the carbonates, as described earlier in this paper, and the observed reaction rates of phosphamidon would also decrease at higher pH values because carbonate ion would interfere with •OH attack of phosphamidon. In addition, phosphamidon undergoes hydrolysis over a wide pH range (4-12) with half-lives ranging from hours to days; the hydrolysis reaction appears to be basecatalyzed.51 Note, however, that the hydrolysis rate constants are orders of magnitude slower than the photochemical rate constants, especially at near-neutral pH. Conclusions Nicotine and phosphamidon, two organic compounds with quite different functional groups, undergo rapid photochemical reaction when treated with UV/H2O2. In NWS, the photochemical oxidation of nicotine is more rapid at lower ionic strength and near-neutral pH values. This is attributable to acid-base speciation and the ionic nature of two of the nicotine species. For phosphamidon, the rate constants were also fastest at nearneutral pH and there was no dependence on ionic strength. The model constructed from experimentally determined rate constants predicts that nicotine and phosphamidon in NWS will be 90% destroyed within 3-5.5 min of treatment in this reactor system. Although the reactor is a batch system, the optimized solution conditions are likely applicable to a larger-scale reactor. In addition, DOE methods are valuable for comparing optimal reactor conditions for treating compounds with distinct structures and functional groups. Supporting Information Available: Full equations for Ka1 and Ka2 and functional relationships between cK and K, and figures showing cKa1 and cKa2 versus ionic strength and log C versus pH diagrams for a series of ionic strengths from 0 to 0.71 M. This material is available free of charge via the Internet at http://pubs.acs.org. Literature Cited (1) Crittenden, J. C.; Trussel, R. R.; Hand, D. W.; Howe, K. J.; Tchobanoglous, G.; Harza, M. W. Water Treatment: Principles and Design, 2nd ed.; John Wiley and Sons: New York, 2005. (2) Letterman, R. D. Water Quality and Treatment: A Handbook of Community Water Supplies, 5th ed.; McGraw-Hill Inc.: New York, 1999. (3) Turyan, Y. I.; Moskatov, I. I. Potentiometry of the stability constant of nicotinic complex of nickel(II) and acid-base dissociation-constants of nicotine. Zh. Neorg. Khim. 1980, 25 (10), 2742–2745. (4) U.S. Environmental Protection Agency, Toxic Release Inventory Explorer. http://www.epa.gov/tri/. (5) Gosselin, R. E.; Gleason, M. N. Clinical toxicology of commercial products: acute poisoning, 4th ed.; Williams & Wilkins: Baltimore, MD, 1976. (6) Boulton, M.; Stanbury, M.; Wade, D.; Tilden, J.; Bryan, D.; Payne, J.; Eisenga, B. Nicotine Poisoning After Ingestion of Contaminated Ground Beef -- Michigan. Morbidity and Mortality Weekly Report, Centers for Disease Control and PreVention 2003, 52, 413–416. (7) Dasenbrock, C. O.; Ciolino, L. A.; Hatfield, C. L.; Jackson, D. S. The determination of nicotine and sulfate in supermarket ground beef adulterated with black leaf 40. J. Forensic Sci. 2005, 50 (5), 1134–1140. (8) Ripley, B. D.; Hall, J. A.; Chau, A. S. Y. Determination of fenitrothion, phosphamidon, and dimethoate in natural water. EnViron. Lett. 1974, 7 (2), 97–118. (9) van der Schalie, W. H.; James, R. R.; Gargan, T. P. Selection of a battery of rapid toxicity sensors for drinking water evaluation. Biosens. Bioelectron. 2006, 22 (1), 18–27. (10) Hua, W. Y.; Bennett, E. R.; Letcher, R. J. Ozone treatment and the depletion of detectable pharmaceuticals and atrazine herbicide in drinking water sourced from the upper Detroit River, Ontario, Canada. Water Res. 2006, 40 (12), 2259–2266.

(11) Metcalfe, C. D.; Miao, X. S.; Koenig, B. G.; Struger, J. Distribution of acidic and neutral drugs in surface waters near sewage treatment plants in the lower Great Lakes, Canada. EnViron. Toxicol. Chem. 2003, 22 (12), 2881–2889. (12) Huerta-Fontela, M.; Galceran, M. T.; Ventura, F. Stimulatory drugs of abuse in surface waters and their removal in a conventional drinking water treatment plant. EnViron. Sci. Technol. 2008, 42 (18), 6809–6816. (13) Chemical Exposure Guidelines for Deployed Military Personnel; USACHPPM Technical Guide 230, Version 1.3; U.S. Army Center for Health Promotion and Preventive Medicine, Aberdeen Proving Ground, MD, May 2003-January 2004, Addendum, 2004; p 307. (14) Rayburn, C. H.; Harlan, W. R.; Hanmer, H. R. The effect of ultraviolet radiation on nicotine. J. Am. Chem. Soc. 1941, 63 (1), 115–116. (15) Weil, L. Photochemical oxidation of nicotine in the presence of methylene blue. Science 1948, 107 (2782), 426–427. (16) U.S. Environmental Protection Agency. www.epa.gov/. (17) TOXNET Toxicology Data Network. toxnet.nlm.nih.gov. (18) World Health Organization. www.who.int/. (19) Hauschild, V. D.; Bratt, G. M. Prioritizing industrial chemical hazards. J. Toxicol. EnViron. Health A 2005, 68, 857–876. (20) Doong, R. A.; Chang, W. H. Photoassisted titanium dioxide mediated degradation of organophosphorus pesticides by hydrogen peroxide. J. Photochem. Photobiol., A 1997, 107 (1-3), 239–244. (21) Doong, R.-a.; Chang, W.-h. Photoassisted iron compound catalytic degradation of organophosphorous pesticides with hydrogen peroxide. Chemosphere 1998, 37 (13), 2563–2572. (22) Harada, K.; Hisanaga, T.; Tanaka, K. Photocatalytic degradation of organophosphorus insecticides in aqueous semiconductor suspensions. Water Res. 1990, 24 (11), 1415–1417. (23) Hua, Z.; Zhang, M. P.; Xia, Z. F.; Low, G. K. C. Titanium-dioxide mediated photocatalytic degradation of monocrotophos. Water Res. 1995, 29 (12), 2681–2688. (24) Huston, P. L.; Pignatello, J. J. Degradation of selected pesticide active ingredients and commerical formulations in water by the photoassisted Fenton reaction. Water Res. 1999, 33 (5), 1238–1246. (25) Katagi, T. Abiotic hydrolysis of pesticides in the aquatic environment. ReV. EnViron. Contam. Toxicol. 2002, 175, 79–261. (26) Konstantinou, I. K.; Sakellarides, T. M.; Sakkas, V. A.; Albanis, T. A. Photocatalytic degradation of selected s-triazine herbicides and organophosphorus insecticides over aqueous TiO2 suspensions. EnViron. Sci. Technol. 2001, 35 (2), 398–405. (27) Krosley, K. W.; Collard, D. M.; Adamson, J.; Fox, M. A. Degradation of organophosphoric acids catalyzed by irradiated titanium dioxide. J. Photochem. Photobiol., A 1993, 69, 357–360. (28) Lartiges, S. B.; Garrigues, P. P. Degradation kinetics of organophosphorus and organonitrogen pesticides in different waters under various envrionmental conditions. EnViron. Sci. Technol. 1995, 29, 1246–1254. (29) O’Shea, K. E.; Beightol, S.; Garcia, I.; Aguilar, M.; Kalen, D. V. Photocatalytic decomposition of organophosphonates in irradiated TiO2 suspensions. J. Photochem. Photobiol., A 1997, 107, 221–226. (30) Pignatello, J. J.; Sun, Y. C. Complete oxidation of methoxychlor and methyl parathion in water by the photoassisted Fenton reaction. Water Res. 1995, 29 (8), 1837–1844. (31) Rabindranathan, S.; Devipriya, S.; Yesodharan, S. Photocatalytic degradation of phosphamidon on semiconductor oxides. J. Hazard. Mater. 2003, 102 (2-3), 217–229. (32) Rahman, M. A.; Muneer, M. Photocatalysed degradation of two selected pesticide derivatives, dichlorvos and phosphamidon, in aqueous suspensions of titanium dioxide. Desalination 2005, 181 (1-3), 161–172. (33) Lam, M. W.; Tantuco, K.; Mabury, S. A. PhotoFate: A new approach in accounting for the contribution of indirect photolysis of pesticides and pharmaceuticals in surface waters. EnViron. Sci. Technol. 2003, 37 (5), 899–907. (34) Walse, S. S.; Morgan, S. L.; Kong, L.; Ferry, J. L. Role of dissolved organic matter, nitrate, and bicarbonate in the photolysis of aqueous fipronil. EnViron. Sci. Technol. 2004, 38 (14), 3908–3915. (35) Hefner, K. H.; Fisher, J. M.; Ferry, J. L. A multifactor exploration of the photobleaching of Suwannee River dissolved organic matter across the freshwater/saltwater interface. EnViron. Sci. Technol. 2006, 40 (12), 3717–3722. (36) Oehlert, G. W. A First Course in Design and Analysis of Experiments; W. H. Freeman: New York, 2000. (37) Faust, S. D.; Aly, O. M. Chemistry of Natural Waters; Ann Arbor Science Publications: Ann Arbor, MI, 1981. (38) de Wuilloud, J. C. A.; Wuilloud, R. G.; Sadi, B. B. M.; Caruso, J. A. Trace humic and fulvic acid determination in natural water by cloud point extraction/preconcentration using non-ionic and cationic surfactants with FI-UV detection. Analyst 2003, 128 (5), 453–458.

Ind. Eng. Chem. Res., Vol. 48, No. 8, 2009 3963 (39) Michalowski, J.; Halaburda, P.; Kojlo, A. Determination of humic acid in natural waters by flow injection analysis with chemiluminescence detection. Anal. Chim. Acta 2001, 438 (1-2), 143–148. (40) Kuhn, H. J.; Braslavsky, S. E.; Schmidt, R. Chemical actinometry. Pure Appl. Chem. 2004, 76 (12), 2105–2146. (41) Nienow, A. M.; Bezares-Cruz, J. C.; Poyer, I. C.; Hua, I.; Jafvert, C. T. Hydrogen peroxide-assisted UV photodegradation of Lindane. Chemosphere 2008, 72, 1700–1705. (42) Green, J. R.; Margerison, D., Statistical Treatment of Experimental Data; Elsevier: New York, 1977; pp 198-293. (43) Ferger, B.; Spratt, C.; Earl, C. D.; Teismann, P.; Oertel, W. H.; Kuschinsky, K. Effects of nicotine on hydroxyl free radical formation in vitro and on MPTP-induced neurotoxicity in vivo. Naunyn-Schmiedeberg’s Arch Pharmacol 1998, 358 (3), 351–359. (44) Wang, S. L.; Wang, M.; Sun, X. Y.; Li, W. Z.; Ni, Y. M. Study on the chemical activity of nicotine by pulse radiolysis. Spectrosc. Spect. Anal. 2003, 23 (3), 481–483. (45) Butler, J. N. In Ionic Equilibrium: Solubility and pH Calculations; John Wiley and Sons: New York, 1998; p 49. (46) Morel, F. M. M.; Hering, J. G. In Principles and Applications of Aquatic Chemistry; John Wiley and Sons: New York, 1993; pp 76-78.

(47) Pankow, J. F. Aquatic Chemistry Concepts; CRC Press: Boca Raton, FL, 1991; p 683. (48) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. Critical review of rate constants for reactions of hydrated electrons, hydrogen atoms and hydroxyl radicals (•OH/•O-) in aqueous solution. J. Phys. Chem. Ref. Data 1988, 17 (2), 513–886. (49) Kochany, J.; Bolton, J. R. Mechanism of photodegradation of aqueous organic pollutants. 2. Measurement of the primary rate constants for reaction of OH radicals with benzene and some halobenzenes using an EPR spin-trapping method following the photolysis of H2O2. EnViron. Sci. Technol. 1992, 26, 262–265. (50) Ulanksi, P.; von Sonntag, C. The OH radical-induced chain reactions of methanol with hydrogen peroxide and with peroxodisulfate. J. Chem. Soc., Perkin Trans. 2 1999, 2, 165–168. (51) Gunther, F. A.; Gunther, J. D. Phosphamidon Residue ReViews; Springer-Verlag: New York, 1971; Vol. 37.

ReceiVed for reView August 29, 2008 ReVised manuscript receiVed February 3, 2009 Accepted February 13, 2009 IE801311F