Temperature dependence of hydroxyl radical reactions with chloramine species in aqueous solution

Abstract

The absolute temperature-dependent kinetics for the reaction between hydroxyl radicals and the chloramine water disinfectant species monochloramine (NH₂Cl), as well as dichloramine (NHCl₂) and trichloramine (NCl₃), have been determined using electron pulse radiolysis and transient absorption spectroscopy. These radical reaction rate constants were fast, with values of 6.06 × 10⁸, 2.57 × 10⁸, and 1.67 × 10⁸ M⁻¹ s⁻¹ at 25 °C for NH₂Cl, NHCl₂, and NCl₃, respectively. The corresponding temperature dependence of these reaction rate constants, measured over the range 10–40 °C, is well-described by the transformed Arrhenius equations:

$$\begin{gathered} \ln \mathrm{k}(•\mathrm{OH}\&{\mathrm{NH}}_2\mathrm{Cl})=(23.68\pm 0.23)-\{(8570\pm 580)/({8.314}^{\ast }\mathrm{T}/\mathrm{K})\} \\ \ln \mathrm{k}(•\mathrm{OH}\&{\mathrm{NHCl}}_2)=(21.96\pm 0.56)-\{(6110\pm 400)/({8.314}^{\ast }\mathrm{T}/\mathrm{K})\} \\ \ln \mathrm{k}(•\mathrm{OH}\&{\mathrm{NCl}}_3)=(21.26\pm 0.29)-\{(5770\pm 720)/({8.314}^{\ast }\mathrm{T}/\mathrm{K})\} \end{gathered}$$

giving activation energies of 8.57 ± 0.58, 6.11 ± 0.40, and 5.77 ± 0.72 kJ mol⁻¹ for these three chloramines, respectively. These data will aid water utilities in predicting hydroxyl radical partitioning and chemical contaminant removal efficiencies under real-world advanced oxidation process treatment conditions.

1. Introduction

For over a century the disinfection of drinking water and wastewater effluent has been performed through the addition of chlorine prior to distribution for human consumption and reuse applications (CRWQCB, 2002; CDPH, 2012; Stone et al., 2009). However, the application of chlorine leads to the formation of halogenated disinfection by-products (DBPs) that are toxic to the environment and show evidence of human toxicity through chronic exposure (Rook, 1974; Krasner et al., 2006; Richardson et al., 2007; Krasner, 2009; Krasner et al., 2009a, 2009b). Consequently, many water utilities are now considering the use of monochloramine as an alternative water disinfectant because its reaction with dissolved organic matter (DOM), in general, promotes lower formation of regulated DBPs such as haloacetic acids and trihalomethanes (EPA, 1999; Diehl et al., 2000; Vikesland et al., 2001; Hua and Reckhow, 2008; Zha et al., 2014; Wang et al., 2016).

However, the simultaneous presence of chloramines and chlorine reacting with DOM can increase overall DBP production (Wang et al., 2016), and the reaction of monochloramine with DOM and other trace chemical contaminants can produce N-nitrosodimethylamine and other species such as aromatic halogenated DBPs (Pan and Zhang, 2013; Hua et al., 2015; LeRoux et al., 2016; Pan et al., 2016; LeRoux et al., 2017; Jiang et al., 2017; Tian et al., 2017) that can be more toxic than trihalomethanes and haloacetic acids, but are currently not as widely regulated (Krasner et al., 2013; Pan et al., 2013; Gong et al., 2016; Guo et al., 2016; Nihemaiti et al., 2016; Zeng et al., 2016; Spahr et al., 2017). The specific timing of chloramination in the water treatment train is thus very important; chloramines need to be generated at specific times that allows for maximum microbial inhibition while minimizing harmful DBP formation (Carlson and Hardy, 1998; Hua and Reckhow, 2007; Wang et al., 2016). In post-treated wastewaters, levels of monochloramine (NH₂Cl) can exceed 2 mg L⁻¹ (measured as equivalent Cl₂) (Johnson et al., 2002).

At the Orange County Water District (OCWD) Advanced Water Purification Facility (AWPF) in Fountain Valley, CA, USA chloramines are deliberately generated by adding sodium hypochlorite (NaOCl) to a secondary-treated wastewater effluent prior to microfiltration (MF) to prevent reverse osmosis (RO) membrane biofouling downstream and to avoid the damaging effects of free chlorine on the thin-film composite polyamide RO membranes. The hypochlorite reacts with residual ammonia (~2.5 mg L⁻¹) present in the secondary effluent to produce a mixture of chloramines. At the injection point, the initial high OCl⁻:NH₃ ratio and pH (~12.5) likely favours the formation of trichloramine (NCl₃); however, as the NCl₃ is eventually diluted in the source water to the approximately neutral system pH, formation of both NH₂Cl and dichloramine (NHCl₂) will occur:

$${\mathrm{NCl}}_3\rightleftharpoons {\mathrm{NHCl}}_2\rightleftharpoons {\mathrm{NH}}_2\mathrm{Cl}$$ (1)

A large fraction of the residual chloramines pass through the MF and RO membranes and are detected in the RO permeate. Analysis of the chloramines by membrane introduction mass spectrometry indicated an approximate ratio of 48% NH₂Cl: 48% NHCl₂: 2% NCl₃ (unpublished data), consistent with expectations at this pH of 5.5–5.7 (Palin, 1950).

As the final treatment step before groundwater recharge for indirect potable reuse, OCWD incorporates an advanced oxidation process (AOP) for disinfection and removal of any trace organic contaminants that are not effectively removed by the RO process. AOPs are characterized by their in-situ generation of highly oxidizing hydroxyl radicals (•OH). In the AWPF, UV light (254 nm) from low pressure mercury lamps is utilized with added hydrogen peroxide (H₂O₂) to produce •OH radicals (von Gunten, 2003; von Gunten, 2007; Wert et al., 2010). The efficiency of the AOP is highly dependent upon the water quality, and consequently the introduction of chloramines into this system could significantly impact the efficiency of the desired reaction of •OH with chemical contaminants. All chloramines react with H₂O₂ (McKay et al., 2013), but these reactions are too slow to impact the AOP chemistry. In contrast, literature measurements for the room-temperature reaction of •OH with NH₂Cl (Poskrebyshey et al., 2003; Johnson et al., 2002) show that this reaction is fast, and thus this radical scavenging pathway could significantly impact the AOP efficiency:

$$\mathrm{OH}+{\mathrm{NH}}_2\mathrm{Cl}\to {\mathrm{H}}_2\mathrm{O}+•\mathrm{NHCl}{\mathrm{k}}_2=5.2\times {10}^8{\mathrm{M}}^{-1}{\mathrm{s}}^{-1}$$ (2)

Given the relative concentrations of chloramine species in the OCWD treatment system, the contribution of NHCl₂ and to a lesser extent, NCl₃ may also impact AOP efficiency. However, no temperature-dependent kinetic data for this reaction, nor any product species for the analogous reactions with NHCl₂ and NCl₃ have been reported:

$$•\mathrm{OH}+{\mathrm{NHCl}}_2\to {\mathrm{H}}_2\mathrm{O}+•{\mathrm{NCl}}_2∕\mathrm{HOCl}+•\mathrm{NHCl}{\mathrm{k}}_3$$ (3)

$$•\mathrm{OH}+{\mathrm{NCl}}_3\to \mathrm{HOCl}+•{\mathrm{NCl}}_2{\mathrm{k}}_4$$ (4)

Therefore, in this study, absolute temperature-dependent rate constants for •OH radical reaction with NH₂Cl, NHCl₂, and NCl₃ were measured to allow a quantitative understanding of their impacts under AOP conditions.

2. Experimental

2.1. Materials and methods

Ammonium chloride (NH₄Cl), sodium hypochlorite (NaOCl), sodium tetraborate decahydrate (Na₂B₄O₇·10H₂O), and hydrochloric acid (HCl) were obtained from VWR International. Potassium thiocyante (KSCN) was purchased from the Sigma-Aldrich Chemical Company. All chemicals were >99% purity and used as received. Solutions were prepared in Milli-Q water (≥18.2 MΩ cm⁻¹). Concentrations of individual chloramine solutions were determined spectrophotometrically using literature (Yiin and Margerum, 1989) absorption coefficients (ε(NH₂Cl) = 461 M⁻¹ cm⁻¹ at 243 nm, ε (NHCl₂) = 272 M⁻¹ cm⁻¹ at 294 nm, and ε (NCl₃) = 195 M⁻¹ cm⁻¹ at 336 nm). Fresh chloramine solutions were prepared daily as described below, kept in the dark, and were used as soon as possible at their buffered pHs. The thermal decay of the chloramines was found to be less than 10% over a 5 h period.

2.2. Chloramine preparation

The concentration of sodium hypochlorite in the stock solution was spectrophotometrically determined at λ_max = 290 nm based on the absorption coefficient ε = 350.4 M⁻¹ cm⁻¹ (Morris, 1966). NH₂Cl was prepared in 2.00 mM sodium tetraborate decahydrate solution buffered to pH 8.8 ± 0.1 by adding 20.0 mL of 6.25 mM sodium hypochlorite dropwise to 5.00 mL of 25 mM NH₄Cl with vigorous stirring over a period of ~15–30 min (Jafvert and Valentine, 1992). This allowed the reaction:

$${\mathrm{NH}}_3+\mathrm{HOCl}\rightleftharpoons {\mathrm{NH}}_2\mathrm{Cl}+{\mathrm{H}}_2\mathrm{O}{\mathrm{K}}_5=3.8\times {10}^{10}$$ (5)

(Margerum et al., 1994) to occur (Armesto et al., 1998). The solutions were kept stirring for at least 30 min after complete NaOCl addition. Typical efficiencies of NH₂Cl production were >95%, as based on initial reactant concentrations.

Dichloramine solutions were prepared by adjusting a monochloramine solution to pH = 3.5 using hydrochloric or perchloric acid. This allowed the disproportionation reaction (Vikesland et al., 2001; Vikesland and Valentine, 2002):

$$2{\mathrm{NH}}_2\mathrm{Cl}+{\mathrm{H}}^{+}\to {\mathrm{NH}}_4^{+}+{\mathrm{NHCl}}_2$$ (6)

to occur. The pH was continuously adjusted until it became stable (constant to 0.01 pH units for 10 min), and this solution was also kept in the dark until use. Dichloramine production efficiencies were 90–95% by this method. Trichloramine was produced by mixing a pH 6.5 ± 0.1 solution of NaOCl with NH₄Cl in a 3:1 mol ratio. Upon complete addition, the pH was then adjusted to 3.5 ± 0.1, and left stirring for at least 30 min before experimental use (Yiin and Margerum, 1989). The efficiency of NCl₃ production was 85–90% by this method.

2.3. Kinetics experiments

The nanosecond pulse radiolysis transient absorption measurements were performed using the linear electron accelerator system at the University of Notre Dame’s Radiation Research Laboratory. A detailed description of the irradiation procedure and transient absorption detection system has been published previously (Whitman et al., 1996; Hug et al., 1999).

The electron pulse radiolysis of water generates the following suite of radicals and molecular products in the stoichiometry (Buxton et al., 1988):

$${\mathrm{H}}_2\mathrm{O}⇝[0.28]•\mathrm{OH}+[0.06]\mathrm{H}•+[0.27]{\mathrm{e}}^{-}(\text{aq})+[0.05]{\mathrm{H}}_2+[0.07]{\mathrm{H}}_2{\mathrm{O}}_2+[0.27]{\mathrm{H}}_3{\mathrm{O}}^{+}$$ (7)

The values in brackets are the absolute yields (G-values, μmol J⁻¹) of generation of each species. In order to isolate the •OH radical, all chloramine solutions were individually pre-saturated with N₂O before mixing and irradiation, which allows the quantitative conversion of radiolytically-produced hydrated electrons and some hydrogen atoms to •OH (Buxton et al., 1988):

$$\mathrm{H}•+{\mathrm{N}}_2\mathrm{O}\to •\mathrm{OH}+{\mathrm{N}}_2{\mathrm{k}}_8=2.1\times {10}^6{\mathrm{M}}^{-1}{\mathrm{s}}^{-1}$$ (8)

$${\mathrm{e}}_{(\mathrm{aq})}^{-}+{\mathrm{N}}_2\mathrm{O}+{\mathrm{H}}_2\mathrm{O}\to {\mathrm{N}}_2+{\mathrm{OH}}^{-}+•\mathrm{OH}{\mathrm{k}}_9=9.1\times {10}^9{\mathrm{M}}^{-1}{\mathrm{s}}^{-1}$$ (9)

Absolute energy deposition measurements (dosimetry, Buxton and Stuart (1995)) were performed using N₂O saturated solutions of 1.00 × 10⁻² M KSCN at λ_max = 475 nm, (Gε = 5.2 × 10⁻⁴ m² J⁻¹) with doses of 3–5 Gy per 2–3 ns pulse, giving initial •OH concentrations in the range 2–4 μM.

During kinetic measurements, NH₂Cl solutions were continuously sparged with the minimum amount of N₂O necessary to prevent air ingress. The more volatile NHCl₂ and NCl₃ solutions were not sparged, instead they were kept in glass vessels whose head space was filled with N₂O gas. All transient absorption measurements were made using a quartz flow cell, with the flow adjusted such that each irradiation occurred on a fresh sample. Sample solutions were flowed through a temperature-controlled condenser into the cell, and the solution temperature was measured immediately after the irradiation. The temperature stability of the system was ±0.3 °C. Typically, 10–15 individual measurements were averaged to generate a single kinetic trace. The quoted errors for the reaction rate constants of this study are a combination of the measurement precision and concentration errors.

3. Results and discussion

3.1. Monochloramine reaction with hydroxyl radical

The hydroxyl radical has been demonstrated to react with monochloramine by hydrogen atom abstraction (Reaction (2)), yielding water and the •NHCl radical (Poskrebyshey et al., 2003). To determine the temperature-dependence rate constants for this reaction, we used benzoate competition kinetics, as no significant transient absorption change was found in the UV–visible region (200–800 nm) for the concentration of radicals produced under our experimental conditions. The reaction of •OH with benzoate (${\mathrm{C}}_6{\mathrm{H}}_5{\mathrm{CO}}_2^{-}$) gives the transient adduct $(\mathrm{HO}-{\mathrm{C}}_6{\mathrm{H}}_5{\mathrm{CO}}_2^{-})•$:

$$•\mathrm{OH}+{\mathrm{C}}_6{\mathrm{H}}_5{\mathrm{CO}}_2^{-}\to (\mathrm{HO}-{\mathrm{C}}_6{\mathrm{H}}_5{\mathrm{CO}}_2^{-})•$$ (10)

which has a strong absorption at 350 nm (Hug, 1981). The temperature dependence of Reaction (10) has been well established, and over the range 10–70 °C is given by (Ashton et al., 1995; Poskrebyshey et al., 2002):

$$\ln {\mathrm{k}}_{10}=(26.682\pm 0.133)-\{(10,150\pm 350)∕({8.314}^{\ast }\mathrm{T}∕\mathrm{K})\}$$ (11)

Upon addition of NH₂Cl to a benzoate solution, the competition for the •OH radical results in a decrease in the absorption intensity of the benzoate adduct transient, as shown in Fig. 1a. The competition between Reactions (2) and (10) is readily solved to give the following relationship:

$$\frac{{\mathit{Abs}}^O{\left(\mathit{HO}-C_6{H}_5{\mathit{CO}}_2^{-}\right)}^{•}}{\mathit{Abs}{\left(\mathit{HO}-C_6{H}_5{\mathit{CO}}_2^{-}\right)}^{•}}=1+\frac{k_2[{\mathit{NH}}_2\mathit{Cl}]}{k_{10}\left[C_6{H}_5{\mathit{CO}}_2^{-}\right]}$$ (12)

where ${\mathit{Abs}}^O(\mathit{HO}-C_6{H}_5{\mathit{CO}}_2^{-})•$ is the yield of the transient adduct absorbance in the absence of NH₂Cl, and $\mathit{Abs}(\mathit{HO}-C_6{H}_5{\mathit{CO}}_2^{-})•$ is the reduced yield in the presence of NH₂Cl. By plotting this ratio of intensities against the concentration ratio $[{\mathrm{NH}}_2\mathrm{Cl}]∕[{\mathrm{C}}_6{\mathrm{H}}_5{\mathrm{CO}}_2^{-}]$ (Fig. 1 (b)) an excellent straight line with a slope of k₂/k₁₀ is obtained. Based on the calculated rate constant of k₁₀ = 6.2 × 10⁹ M⁻¹ s⁻¹ at this temperature, this gives a calculated absolute second-order rate coefficient of (6.10 ± 0.08) × 10⁸ M⁻¹ s⁻¹. However, this calculated value includes the contribution of Reaction (2) as well as that of residual ammonia with •OH (Buxton et al., 1988):

$$•\mathrm{OH}+{\mathrm{NH}}_3\to {\mathrm{H}}_2\mathrm{O}+•{\mathrm{NH}}_2{\mathrm{k}}_{13}=9.9\times {10}^7{\mathrm{M}}^{-1}{\mathrm{s}}^{-1}$$ (13)

Fig. 1.

(a) Growth kinetics of $(\mathrm{HO}-{\mathrm{C}}_6{\mathrm{H}}_5{\mathrm{CO}}_2^{-})•$ at 350 nm for N₂O saturated solutions of 102.5 μM benzoate at pH 9.2 in aqueous solution at 22.2 °C with 0(□), 0.52 (O), 1.01 (Δ), and 1.48 mM (∇) NH₂Cl (b) Transformed competition kinetics plot using maximum intensity data (shown as horizontal lines) in (a). Solid line is weighted linear fit, with slope value of (0.984 ± 0.019) and an intercept of 0.998 ± 0.003, R² = 0.999. This slope corresponding to a total k = (6.10 ± 0.12) × 10⁸ M⁻¹ s⁻¹ under these conditions.

Residual NH₃ is presumed to exist as the formation of NH₂Cl (Reaction (5)) is not quantitative (measured at ~95% using a theoretical yield based on NaOCl and chloramine absorption coefficients). By taking this small side reaction into account, a final value of k₂ = (5.99 ± 0.12) × 10⁸ M⁻¹ s⁻¹ is obtained, in very good agreement with the 22 ± 1 °C reported value (Poskrebyshey et al., 2003) of (5.1 ± 0.6) × 10⁸ M⁻¹ s⁻¹. Unfortunately, there is a paucity of reported rate constants for •OH with single chlorine-containing chemicals in aqueous solution; however, our measured room-temperature rate constant for NH₂Cl is also comparable to the reaction of the hydroxyl radical with chloroethane, determined as 5.5 × 10⁸ M⁻¹ s⁻¹ (Milosavljevic et al., 2005), but significantly higher than measured for chloroacetamide ((8.1 ± 0.3) × 10⁷ M⁻¹ s⁻¹, Xu et al., 2014) or chloronitromethane ((1.9 ± 0.3) × 10⁸ M⁻¹ s⁻¹, Mincher et al., 2010; Mezyk et al., 2006).

Similar benzoate competition kinetics measurements were made for NH₂Cl reaction over the temperature range of 10–41 °C, the second-order rate constants obtained are summarized in Table 1. These NH₂Cl reaction data are well described by the transformed Arrhenius equation:

$$\ln {\mathrm{k}}_2=(23.68\pm 0.23)-\{(8570\pm 580)∕({8.314}^{\ast }\mathrm{T}∕\mathrm{K})\}$$ (14)

corresponding to an activation energy of E_a2 = 8.57 ± 0.58 kJ mol⁻¹ (see Fig. 2). No comparable literature data were found for this measured activation energy.

Table 1.

Summary of temperature-dependent rate constants measured in this study.

Species	Temperature °C	10−8 k value M−1 s−1	Ea kJ mol−1	ln (A)
NH2Cl	10.8	(5.06 ± 0.04)
	20.8	(5.72 ± 0.12)
	22.2	(5.99 ± 0.12)	8.57 ± 0.58	23.68 ± 0.23
	30.2	(6.53 ± 0.63)
	40.8	(7.13 ± 0.35)
NHCl2	11.6	(2.28 ± 0.06)
	22.1	(2.54 ± 0.10)	6.11 ± 0.40	21.96 ± 0.56
	29.9	(2.70 ± 0.08)
	38.1	(2.84 ± 0.18)
NCl3	7.1	(1.43 ± 0.18)
	19.9	(1.63 ± 0.18)	5.77 ± 0.72	21.26 ± 0.29
	31.2	(1.71 ± 0.14)
	39.0	(1.88 ± 0.23)

Fig. 2.

Arrhenius plot for the reaction of ^•OH with NH₂Cl (▪), NHCl₂ (●), and NCl₃ (▴). Solid lines correspond to weighted linear fits, corresponding to activation energies of Ea₂ = 8.57 ± 0.58, Ea₃ = 6.11 ± 0.40, and Ea₄ = 5.77 ± 0.72 kJ mol⁻¹, respectively.

3.2. Determination of di- and trichloramine spectra and kinetics

Conversely to NH₂Cl the pulse irradiation of an aqueous solution of NHCl₂ gave a measurable transient absorption spectrum peaking at 310 nm (see Fig. 3). The same, but more intense, absorption spectrum was observed for hydroxyl radical reaction with NCl₃. Absorption coefficients for these spectra were calculated based upon our measured dosimetry (which gave absolute initial •OH concentrations), utilizing the scavenging radical yield corrections of LaVerne and Pimblott (1993).

Fig. 3.

Absolute transient absorption spectra obtained for NCl₂• radical from reaction of hydroxyl radical reaction with NHCl₂ (□) and NC1₃ (O).

By assuming that Cl-atom abstraction is the only possible mechanism for •OH reaction with NCl₃, we can attribute these measured transient absorptions to the •NCl₂ radical. For NH₂Cl, there are two possible products of the hydroxyl radical-induced oxidation:

$$•\mathrm{OH}+{\mathrm{NHCl}}_2\to {\mathrm{H}}_2\mathrm{O}+•{\mathrm{NCl}}_2{\mathrm{k}}_{3\mathrm{a}}$$ (3a)

$$•\mathrm{OH}+{\mathrm{NHCl}}_2\to \mathrm{HOCl}+•\mathrm{NHCl}{\mathrm{k}}_{3\mathrm{b}}$$ (3b)

Through normalization of the NHCl₂ measured product spectra to that obtained for NCl₃ (at 310 nm) we can readily calculate the branching ratio for the overall NHCl₂ reaction to be 20.3% for Reaction (3a) and 79.7% for Reaction (3b). These fractions can be used to calculate individual reaction rate constants for these two reactions (see later). It is also important to note that while the •NHCl radical also absorbs in this wavelength region (Poskrebyshey et al., 2003) this previous study showed that the •NHCl radical intensity was similar to the aqueous •NH₂ radical peak (ε = 80 M⁻¹ cm⁻¹, Hug, 1981), and thus at our product radical concentrations, the absorbance of produced •NHCl in Reaction (3b) would not significantly contribute to the overall •NCl₂ intensity seen in our study.

3.3. Dichloramine reaction with the hydroxyl radical

The kinetics of hydroxyl radical reaction with NHCl₂ were directly measured at 310 nm. Typical kinetic data at 22.1 °C are given in Fig. 4. These measurements were conducted at pH 3.3, to maximize the stability of dichloramine in aqueous solution. By fitting the measured transient absorptions (Fig. 4a) to exponential growth kinetics, pseudo-first-order rate coefficients were attained. Plotting these fitted values against the dichloramine concentration (Fig. 4b) yielded a total second-order rate coefficient of k₃ = (2.54 ± 0.12) × 10⁸ M⁻¹ s⁻¹ under these conditions. From the spectral analysis performed for this compound, this overall rate constant consists of the sum of Reaction (3a) and (3b) pathways, from which we can calculate k_3a = (5.2 ± 0.2) × 10⁷ M⁻¹ s⁻¹ and k_3b = (2.0 ± 0.1) × 10⁸ M⁻¹ s⁻¹ at 22.1 °C. The overall rate constants for this reaction measured over the range 11–38 °C, were well described by the equation:

$$\ln {\mathrm{k}}_3=(21.96\pm 0.56)-\{(6110\pm 400)∕({8.314}^{\ast }\mathrm{T}∕\mathrm{K})\}$$ (15)

corresponding to an activation energy of E_a3 = 6.11 ± 0.40 kJ mol⁻¹ for both pathways. All the measured temperature-dependent kinetic data for NHCl₂ are summarized in Table 1, and shown in Fig. 2. Again, no comparable Arrhenius parameters for analogous species were found in the literature. However, the decrease in hydroxyl radical reactivity between NH₂Cl (5.99 × 10⁸ M⁻¹ s⁻¹) to NHCl₂ (2.54 × 10⁸ M⁻¹ s⁻¹) at ~22 °C is consistent with the decrease seen for this radical reaction with chloromethane (5.5 × 10⁸ M⁻¹ s⁻¹) and 1,1-dichloroethane (1.3 × 10⁸ M⁻¹ s⁻¹) in aqueous solution (Milosavljevic et al., 2005). In contrast, an increase in reaction rate constant was seen in going from chloroacetamide ((8.1 ± 0.3) × 10⁷ M⁻¹ s⁻¹) to dichloracetamide ((1.25 ± 0.02) × 10⁸ M⁻¹ s⁻¹; Xu et al., 2014), and from chloronitromethane ((1.9 ± 0.3) × 10⁸ M⁻¹ s⁻¹) to dichloronitromethane ((5.1 ± 0.8) × 10⁸ M⁻¹ s⁻¹; Mezyk et al., 2006).

Fig. 4.

(a) Transient absorption kinetics determined at 310 nm for reaction of hydroxyl radical with dichloramine in N₂O-saturated solution at pH 3.3 and 22.1 °C for 1.34 (□), 1.06 (O), 0.40 (Δ), and 0.20 mM (∇) NHCl₂. Fitted lines are pseudo first-order growth kinetics, with rate constants of (4.57 ± 0.09) × 10⁵, (3.68 ± 0.09) × 10⁵, (1.92 ± 0.05) × 10⁵, and (1.39 ± 0.10) × 10⁵ s⁻¹, respectively (b) Transformed second-order kinetic plot from kinetic data. Solid line corresponds to weighted linear fit, with slope corresponding to k = (2.54 ± 0.12) × 10⁸ M⁻¹ s⁻¹, R² = 0.994.

3.4. Trichloramine reaction with hydroxyl radical

The kinetics of hydroxyl radical reaction with NCl₃ were measured as for NHCl₂, at pH 3.5. However, for this chloramine, concentrations had to be corrected for thermal decay over the time course of the irradiation experiments. This was achieved by directly measuring the decay of the initial trichloramine solution with time using UV–vis spectrophotometry. The reaction rate constants (see Table 1) for the temperature-dependent reaction of the hydroxyl radical with NCl₃ was found to be well described by:

$$\ln {\mathrm{k}}_4=(21.26\pm 0.29)-\{(5770\pm 720)∕({8.314}^{\ast }\mathrm{T}∕\mathrm{K})\}$$ (16)

corresponding to an activation energy of E_a4 = 5.77 ± 0.72 kJ mol⁻¹ over the temperature range 7–39 °C. The room-temperature rate constant for this reaction, k₄ = 1.67 × 10⁸ M⁻¹ s⁻¹ is again slower than for NH₂Cl and NHCl₂, again consistent with the lower reactivity of 1,1,1-trichloroethane (k < 5 × 10⁶ M⁻¹ s⁻¹, Milosavljevic et al., 2005) and for trichloronitromethane ((4.8 ± 0.4) × 10⁷ M⁻¹ s⁻¹, Cole et al., 2007).

Overall, the second-order rate constant for the reaction between chloramines and the hydroxyl radical follows the order mono- > di- > trichloramine, with NH₂Cl being approximately a factor of two more reactive than NHCl₂ and NCl₃. In addition, there is a clear grouping of activation energies and Arrhenius pre-factors that is dependent upon the degree of chlorine substitution. We attribute these differences to the change in •OH reaction mechanism, from mainly H-atom abstraction for monochloramine to only Cl-atom abstraction for trichloramine. This is also reflected in the calculated Erying values, where for di- and trichloramine ΔH ≈ 3.5 kJ mol⁻¹ and ΔS ≈ −73 J mol⁻¹K⁻¹, in contrast to monochloramine where ΔH ≈ 6 kJ mol⁻¹ and ΔS ≈ −56 J mol⁻¹ K⁻¹.

3.5. Applicability to AOP treatment conditions

It is important to quantitatively account for the presence of all three chloramines within an AOP, especially as their presence will significantly interfere with the remediation chemistry that occurs, specifically through reaction of formed hydroxyl radicals with these disinfectants rather than removal of the target chemical contaminants of concern. Although development of a full kinetic model is beyond the scope of this work, as this would require quantitative evaluation of all the intermediates formed in these oxidations, as well as their radical reaction rate constants, the kinetic data from this study provides a useful foundation for the initial evaluation of the AOP efficiencies in removing trace levels of chemical contaminants. For example, on rare occasion 1,4-dioxane has been detected in OCWD AOP feedwater at a concentration of 1–3 μg/L. This amount was subsequently reduced to non-detect (<1 μg/L) in the UV/H₂O₂ AOP. This is a compound of concern (CDPH, 2010), and since 1,4-dioxane is not readily photolyzed, it requires •OH oxidation for its degradation (Patton et al., 2017). Assuming a concentration of 2 μg/L entering the AOP, Table 2 summarizes the most important •OH radical reactions for 1,4-dioxane in wastewater that has undergone primary, secondary, MF, and RO treatments prior to the final AOP with UV/H₂O₂. Based on these typically measured concentrations of these treated wastewater standard constituents, the photolytically produced •OH radicals will partition according to the relative rates of the individual reactions:

$$\frac{{\mathrm{R}}_{\mathrm{X}}}{{\sum }_i{R}_i}=\frac{{\mathrm{k}}_{\mathrm{X}}[X]}{{\sum }_i{k_i}^{\ast }[i]}\times 100\%$$ (17)

Table 2.

Summary of •OH reactivity with wastewater constituents in OCWD AOP system.

Reaction	k @ 25 °Cb (M−1 s−1	[Species] (mol dm−3)	•OH partitioning w/o chloramines (%)	•OH partitioning w/chloramines (%)
•OH + 1,4-dioxane → products	2.50 × 109	2.27 × 10−8	4.7 × 10−4	1.9 × 10−4
•OH + H2O2 → H2O + HO2•	9.00 × 107	8.82 × 10−5	65.5	26.2
$•\mathrm{OH}+{\mathrm{HCO}}_3^{-}\to {\mathrm{OH}}^{-}+{{\mathrm{CO}}_3}^{•-}$	8.50 × 106	3.74 × 10−4	22.9	9.2
•OH + DOM → H2O + DOM•	2.25 × 108	6.25 × 10−6	11.6	4.6
•OH + NH2Cl → H2O + NHCl•	6.06 × 108 a	2.38 × 10−5	–	47.6
•OH + NHCl2 → H2O + NCl2•/HOCl + NHCl•	2.57 × 108 a	1.43 × 10−5	–	12.1
•OH + NCl3 → HOCl + NCl2•	1.67 × 108 a	4.15 × 10−7	–	0.23

^aCalculated using measured Arrhenius behaviour.

^bBuxton et al., 1988 or this study.

Although the fraction of •OH radicals reacting with 1,4-dioxane is small under both circumstances, by including the scavenging reactions of chloramines into this calculation, it is seen that approximately 60% of the •OH radicals are scavenged by these three chloramine species, with a concomitant decrease of •OH reactivity with 1,4-dioxane by ~150%.

4. Conclusions

Absolute second-order temperature-dependent rate constants for hydroxyl radical reaction with mono-, di-, and trichloramine in aqueous solution have been determined. At 25 °C, these rate constants were calculated to be 6.06 × 10⁸, 2.57 × 10⁸, and 1.67 × 10⁸ M⁻¹ s⁻¹, respectively with corresponding activation energies of 8.57, 6.11, and 5.77 kJ mol⁻¹. The •OH radical reaction with monochloramine gives •NHCl as the radical product. For trichloramine, HOCl and the •NCl₂ radical are the proposed products for •OH reaction. Based on the transient absorbance intensity of the •NCl₂ radical formed for the hydroxyl radical reaction with NHCl₂, there were two sets of products; H₂O/•NCl₂ (20.3%) and HOCl/•NHCl (79.7%). A relative rates analysis showed that the presence of the three chloramine species in an advanced oxidation process treatment of wastewaters will significantly decrease the overall efficiency of hydroxyl radical reaction with chemical contaminants of concern, with up to 60% of produced •OH reacting with these species instead of contaminants of concern.

HIGHLIGHTS.

• Absolute hydroxyl radical rate constants for three chloramine species in water.

• Temperature-dependent oxidation kinetics and identified product species.

• Impacts on advanced oxidation process efficiency due to chloramines presence.