Journal of Membrane Science 258 (2005) 64–70 Enhanced oxidation of polyamide membranes using monochloramine and ferrous iron Christopher J. Gabelich a,∗ , John C. Frankin b , Fredrick W. Gerringer a , Kenneth P. Ishida c , I.H. (Mel) Suffet d a Metropolitan Water District of Southern California, 700 Moreno Avenue, La Verne, CA 91750, USA b United States Bureau of Reclamation, Yuma, AZ, USA c Orange County Water District, Fountain Valley, CA, USA d University of California, Los Angeles, CA, USA Received 4 January 2005; received in revised form 19 February 2005; accepted 21 February 2005 Available online 12 April 2005 Abstract Reverse osmosis testing using monochloramine (NH2 Cl) and free chlorine (HOCl) in the presence of ferrous iron [Fe(II)] resulted in accelerated chlorination of polyamide membranes. Similar effects were not observed when using ferric iron. Membrane damage was detected through irreversible increases in salt passage and the presence of chloride on the membrane surface using energy-dispersive spectroscopy. A mechanistic study suggested that the formation of an amidogen radical (• NH2 ) during NH2 Cl decomposition with Fe(II) led to the reduction of the activation energy for the chlorination reaction to proceed using NH2 Cl. Enhanced oxidation using HOCl and Fe(II) may have resulted from the formation of • OH radicals, which also lowered the chlorination-activation energy. Both pH suppression (from 8.0 to 6.0) and dechlorination (7.0 mg/L sodium sulfite) successfully halted the chlorination reaction. © 2005 Elsevier B.V. All rights reserved. Keywords: Reverse osmosis; Polyamide; Chlorine; Chloramine; Iron 1. Introduction The Metropolitan Water District of Southern California (MWDSC) is evaluating strategies to lower the salinity of Colorado River water. One option is through treatment with reverse osmosis (RO) using thin-film composite polyamide membranes. In an effort to reduce the costs of RO treatment, the use of existing conventional or direct-filtration treatment plants is desirable. Research at MWDSC using pretreatment with ferric chloride (FeCl3 ) coagulation and monochloramine (NH2 Cl) has shown both the presence [1] and absence [2,3] of polyamide membrane oxidation during RO testing. Although previous research has shown that severe membrane damage can occur when polyamide membranes are exposed to aqueous chlorine [4] and that ferrous iron [Fe(II)] can cat∗ Corresponding author. Tel.: +1 909 392 5113; fax: +1 909 392 5166. E-mail address: cgabelich@mwdh2o.com (C.J. Gabelich). 0376-7388/$ – see front matter © 2005 Elsevier B.V. All rights reserved. doi:10.1016/j.memsci.2005.02.034 alyze the reaction between aqueous chlorine and cellulose acetate membranes [5–7], no published research has shown that NH2 Cl in the presence of Fe(III) leads to impaired polyamide membrane performance. Additionally, other authors have reported stable membrane performance in the presence of Fe(III) without chlorine [8]. Polyamide membranes have free chlorine (HOCl) and chloramine tolerances of approximately 1000 and 300,000 ppm-h, respectively [9,10]. The solubility of Fe(III) at pH 8.0 is less than 1 × 10−8 mol/L (1 ␮g/L) as Fe(OH)3(s) [11], which typically leads to nondetectable levels of iron in MWDSC’s treatment plant effluents. This paper details pilot-scale experiments used to determine the potential pathway for an iron-mediated oxidation reaction of polyamide membranes in the presence of NH2 Cl. Polyamide membranes were exposed to HOCl and NH2 Cl alone and in combination with either Fe(II) or Fe(III). This research also evaluated the ability of citric acid, excess ammonia (NH3 ), and sodium sulfite (Na2 SO3 ) to inhibit mem- C.J. Gabelich et al. / Journal of Membrane Science 258 (2005) 64–70 65 (4) Fig. 1. Basic polyamide polymer unit. brane deterioration through iron sequestration, shifting chlorine equilibrium, and dechlorination, respectively. 2. Background Fig. 1 shows a basic polyamide–polymer unit. Chlorination of polyamides has been demonstrated through reaction of the free chlorine with the nitrogen of the secondarysubstituted amide group, which then resulted in ring chlorination [12]. When chlorine is introduced into water, either as chlorine (Cl2 ) gas or hypochlorite (OCl− ), it rapidly undergoes hydrolysis. The important reaction product is hypochlorous acid (HOCl), which is the dominant germicide used for drinking water disinfection [13]. As a “weak” acid, HOCl disassociates per Eq. (1). HOCl + H2 O ⇔ H3 O+ + OCl− , pKa at 20 ◦ C = 7.6 (1) The active germicidal and oxidation properties of chlorine are attributed to the +1 valence state of the chlorine atom (Cl1+ ) for hypochlorous acid [H1+ (O2− Cl1+ )− ], hypochlorite [(O2− Cl1+ )− ], and monochloramine [NH2 1− Cl1+ ] [13]. It is thought that the formation of Cl+ is the primary driver in polyamide membrane chlorination reactions. Halogenation reactions with aromatic rings are not thermodynamically favored unless a Lewis acid (i.e., an electron acceptor) is present [14]. In this case, the HOCl acts as a Lewis acid, thereby leading to the formation of Cl+ [see Eq. (2)], which then attacks the aromatic ring to form an amidesubstituted arenium ion [see Eq. (3)]. The energy of activation to form the arenium ion, which alters the aromaticity of the ring structure, has been shown to be much greater than the energy of activation leading to the final product [see Eq. (4)]. This phenomenon has been referred to as the “Orton rearrangement” [6]. OCl− Chlorination reactions with polyamide membranes have been shown to be pH dependent, with greater chlorine sensitivity observed at lower pH (e.g., pH 4–4.5) for linearchained polyamide membranes, but not for crosslinked polymeric membranes [15,16]. Soice and co-workers using acetanilide—an aromatic polyamide surrogate [17] and polyamide-coated drops [18] showed that ring chlorination occurred at pHs 4 and 7, with no ring chlorination at pH 10. The amount of HOCl is a function of pH, with 23% of chlorine present as HOCl at pH 8 (20 ◦ C), and 100 percent of chlorine present as HOCl at pH 4 (20 ◦ C) [13]. At pH 10, HOCl would account for less than 1% of the chlorine species in solution. Therefore, it may be assumed that HOCl is the active ion in polyamide membrane chlorination. Chlorine substitution should occur on the C2 or C4 carbon of the N H bonded ring, as N H bonds are ortho-paradirecting [14]. Chlorination of the carbonyl-linked aromatic ring has been observed to be much slower than the N H linked ring [19]. This finding may be caused by the complete electron set of the C O portion of the amide functional group, which results in a physical separation of the production of Cl+ ion via the amide nitrogen and the polymer ring. Direct chlorination of the aromatic ring has also been suggested [6,19], though experiments with tertiary-substituted aromatic amides showed no reaction with aqueous chlorine [12,15]. Kawaguchi and Tamara [12], Lowell et al. [15], and Soice et al. [17] observed that irreversible chlorination occurred only with secondary-substituted amides, leading to the conclusion that the presence of the N H bond plays a critical role in polyamide chlorination and that direct chlorination of the aromatic ring does not occur. By adding either a methyl or benzene group, a localized positive charge is created on the N H group, which disrupts the ability of the arenium ion [see Eq. (3)] to form, resulting in a chlorine-tolerant membrane. The formation of monochloramine has been shown to be a rapid reaction (k1 = 5.1 × 106 L/mol s) [20], with the reverse reaction, decomposition of monochloramine to NH3 and HOCl, being considerably slower (k2 = 3 × 10−5 s−1 ) [21] [see Eq. (5)]. The resulting equilibrium rate contact (Keq ) for monochloramine formation is 1.7 × 1011 [21]. ⊕ HO Cl + HO Cl ⇋ H2 O Cl ⇋ H2 O + Cl+ (2) (3) 66 C.J. Gabelich et al. / Journal of Membrane Science 258 (2005) 64–70 HOCl + NH3 (aq)k1 ⇔k2 NH2 Cl + H2 O (5) Keq = k1 = 1.7 × 1011 k2 (6) Keq = [NH2 Cl][H2 O] [HOCl][NH3 ] (7) The chlorine tolerance of polyamide membranes is considerably higher for NH2 Cl than for HOCl [9,10]. Nonaqueous FeCl3 has been used in halogenation reactions for benzene and other aromatic compounds [13]. In an aqueous environment with natural alkalinity, FeCl3 hydrolyzes based on the following general equation [22]: 2FeCl3 + 3Ca(HCO3 )2 ⇒ 2Fe(OH)3 ↓ + 3CaCl2 + 6CO2 (8) In addition to Fe(OH)3(s) , other hydrated iron species are also present, depending on pH. At pH values less than 8.0, the principle soluble species are Fe3+ , Fe(OH)2 + , and FeOH2+ , though the concentrations of these species are orders of magnitude lower than that of amorphous Fe(OH)3(s) [11]. At pH 8.0, the solubility of iron is less than 1 × 10−8 mol/L (1 ␮g/L) as Fe(OH)3(s) . However, based on sedimentation and filter efficiency, higher concentrations of particulate Fe(III) may be present. In the aqueous system, Fe(OH)3(s) may act similarly to nonaqueous FeCl3 in creating the Cl+ ion to react directly with the polymer ring [see Eq. (2)]. Formation of the arenium intermediary is unchanged [see Eq. (3)]. The resulting Fe(OH)4 − is then regenerated to Fe(OH)3(s) through hydrolysis with water. In this fashion, Fe(OH)3(s) may act as a catalyst, and it is not consumed in the overall reaction. Both Fe(III) and Fe(II) act as effective catalysts in both biological (e.g., cytochrome P450) and inorganic (e.g., Fenton’s reagent) oxidation reactions. However, these reactions utilize • OH radicals, either directly or through hydrogen peroxide addition. Ferrous iron has been shown to catalyze the reduction of NH2 Cl to chloride (Cl− ) between pH 5.8 and pH 7.9 [23]. The following reaction pathways have been proposed by Vikesland and Valentine [23,24]: Fe(II)soln + NH2 Cl ⇒ Fe(III) + • NH2 + Cl− (9) Fe(II)soln + • NH2 + H+ ⇒ Fe(III) + NH3 (10) Fe(II)soln + > Fe(OH)3(s) ⇔ Fe(II)surf (11) Fe(II)surf + NH2 Cl ⇒ • NH 2 + Cl − + Fe(III) (12) Fe(II)surf + • NH2 + H+ ⇒ +Fe(III) + NH3 (13) Dissolved ferrous iron [Fe(II)soln ] and NH2 Cl readily reacted to form a radical intermediary, amidogen (• NH2 ), Fe(III) solid, and chloride (Cl− ) [see Eq. (9)]. The • NH2 then rapidly decomposes through reaction with Fe(II)soln to form ferric oxide [Fe(OH)3 ] and NH3 [see Eq. (10)]. When Fe(II)soln is oxidized by NH2 Cl, Fe(III) is formed. The presence of Fe(OH)3 surfaces [>Fe(OH)3(s) ] has been shown to accelerate the decomposition of NH2 Cl, most probably through the formation of an Fe(II) surface complex [Fe(II)surf ] [Eqs. (11)–(13)]. In natural water systems, pH may be a contributing factor in NH2 Cl decomposition, in that Fe(II) oxidation is favored at higher pH [25], thereby further accelerating NH2 Cl reduction by increasing Fe(II)surf [see Eq. (12)]. Other scavenging reactions for • NH2 were proposed, but are not relevant to this discussion. Although the reduction of NH2 Cl by Fe(II) was demonstrated, the final end product was hypothesized to be Cl− . This is inconsistent with polyamide membrane chlorination theory, which is predicated on the presence of Cl+ . 3. Experimental methods Source water from the Colorado River was processed through a 0.2 ␮m nominal pore size, hollow-fiber microfiltration (MF) unit (Model 12M10C, U.S. Filter/Memcor, Timonium, Md.). A 2.5 mg/L chloramine residual was maintained in the MF feed, using a 4:1 w/w Cl2 to NH3 ratio. A free chlorine concentration of 2.0 mg/L was achieved via breakpoint chlorination, and dechlorination was achieved by quenching with 7.0 mg/L Na2 SO3 . Chemical feeds prior to RO treatment included 7.9 × 10−7 M ferrous sulfate (0.10 mg/L technical grade, Spectrum Quality Products, Gardena, CA) and ferric chloride (0.13 mg/L NSF-certified, Kemiron Pacific, Fontana, CA). Retention time between the Fe(II)/Fe(III) addition point and membrane exposure was less than 10 s. Two identical, pilot-scale RO units were used for membrane testing. Each RO test unit had two parallel vessels, each containing two 2.5-in. (diameter) by 40 in. (long), spiral-wound elements (ESPA3, Hydranautics, San Diego, CA). While each vessel was operated with a common feed (15–16 L/min per vessel at 900 kPa net driving pressure), chemical additions were added independently to each vessel. After each experiment, the RO units were cleaned with a citric acid solution (pH 2.5), reoperated with chloraminated MF filtrate to reestablish clean-membrane performance, and the RO elements were then autopsied. Salt transport coefficients were normalized to 25 ◦ C [26]. Osmotic pressure and salt transport coefficients were calculated based on conductivity (␮S/cm). Conductivity-to-total-dissolved-solids (mg/L TDS) conversion factors were empirically derived from influent, concentrate, and permeate samples. All water quality constituents were analyzed according to Standard Methods for the Examination of Water and Wastewater [27] except the trace metals (Al, Fe, Mn, Ba, and Sr), which were analyzed according to EPA Method 200.8 using inductively coupled plasma mass spectrometry (ELAN 6000 ICP-MS, PerkinElmer Life and Analytical Sciences, Boston, MA). Chlorine residuals were measured using method 4500Cl G [27]. Membrane surface analyses included scanning electron microscopy (SEM), energy-dispersive spectroscopy (EDS), and attenuated total reflectance/Fourier-transform infrared (ATR/FT-IR) spectroscopy. SEM analyses were conducted using a field-emission gun (XL30-FEG, Philips Ad- C.J. Gabelich et al. / Journal of Membrane Science 258 (2005) 64–70 Table 1 Mean microfiltered water quality data Parameter Value Total dissolved solids (mg/L) Total hardness (mg/L as CaCO3 ) Total alkalinity (mg/L as CaCO3 ) Total organic carbon (mg/L) Dissolved oxygen (mg/L) Aluminum (␮g/L) Iron (␮g/L) pH (units) Temperature (◦ C) Silt density index Turbidity (NTU) 918 (162) 368 (75) 169 (34) 2.4 (0.4) 8.9 (0.1) 8.8 (13) <20 MDL 7.9 (0.3) 25 (7) 1.6 (0.5) 0.10 (0.02) Data in parenthesis indicate standard deviation. vanced Metrology Systems, Natick, MA) with EDS (EDAX, Mahwah, NJ) capabilities per Goldstein et al. [28]. The EDS data were given in percent weights of detected elements greater than 12 atomic mass units. Membrane oxidation was detected through ATR/FT-IR spectroscopy (Magna 550, Thermo Nicolet, Madison, WI) [29]. 4. Results Table 1 lists selected MF effluent water quality data. Water quality was consistent during the 21 weeks of testing. Prior to RO chemical addition, no iron (method detection limit [MDL], 20 ␮g/L) was detected in the MF effluent at any time. The median effluent turbidity and silt density index were 0.10 NTU and 1.9, respectively. The RO feed water was taken directly after the MF. The mean ambient conditions for Colorado River water were 973 mg/L total dissolved solids, pH 8.0, 20 ◦ C, and 8.9 mg/L dissolved oxygen. Fig. 2 shows the salt (B value) transport coefficients of membranes exposed to equal molar concentrations of Fe(II) and Fe(III), using either free or combined chlorine. Free chlorine exposure was shown to have a negative impact on membrane properties, more so with the presence of Fe(II). Ferrous addition led to enhanced salt passage (nonrecoverable after citric acid cleaning) using both HOCl and NH2 Cl 67 Table 2 Energy-dispersive spectroscopy data for polyamide membrane surfaces exposed to free chlorine and chloramines Test Element O S Cl Fe HOCl HOCl + Fe(III) HOCl + Fe(II) NH2 Cl NH2 Cl + Fe(III) NH2 Cl + Fe(II) 71 66 59 60 64 65 27 31 35 40 36 32 2.2 2.5 5.4 – – 2.3 – – – – – 1.1 Data on a percent w/w basis. All membranes operated at pH 8.0. (Fig. 2). Membranes exposed to Fe(III) exhibited increases in salt transport similar to those exhibited by membranes exposed to Fe(II); however, upon acid cleaning, the salt passage recovered relative to the non-Fe(III)-exposed control. This latter finding may be attributable to enhanced concentration polarization resulting from excess cake formation at the membrane surface caused by the ferric oxide colloids [30,31]. Hoek et al. [30,31] demonstrated that dense-colloidal packings hindered back diffusion of solutes away from the membrane surface, resulting in enhanced osmotic pressure and increased salt passage. An alternative hypothesis would be Fe(III) facilitating a reversible N–Cl chlorination reaction, as reported in Kawaguchi and Tamura [12], that was removed after acid cleaning. This line of research was not pursued further during this study. As the B value for the Fe(III)-exposed membranes recovered and the B value for Fe(II) did not, it may be concluded the Fe(II) in both HOCl and NH2 Cl conditions permanently altered the membrane chemistry, whereas Fe(III) addition did not. EDS data (Table 2) detected chlorine addition for all HOCl-exposed membranes, more so for the HOCl and Fe(II) experimental conditions [5.4% Cl for HOCl and Fe(II) versus 2.2% Cl for HOCl and 2.5% for HOCl and Fe(III)]. No chlorine addition to the membrane structure was observed for the NH2 Cl and the NH2 Cl and Fe(III)-exposed membranes. However, the NH2 Cl and Fe(II)-exposed membrane showed chlorine content equivalent to the HOCl-exposed membrane (2.3% versus 2.2% Cl, respectively), indicating Fig. 2. Salt transport coefficients of polyamide membranes exposed to NH2 Cl and HOCl, with and without Fe(III) (left column) or Fe(II) (right column). Star symbols indicate acid cleanings. All data taken at pH 8.0 and normalized to 25 ◦ C. 68 C.J. Gabelich et al. / Journal of Membrane Science 258 (2005) 64–70 Fig. 3. ATR/FT-IR spectra of polyamide membranes. Membrane oxidation was detected by observing declines in intensities of (a) the C C ring vibrations near 1610 and 1448 cm−1 ; (b) the out-of-plane C H bend near 780 cm−1 ; (c) the amide II N H bend near 1540 cm−1 , which also shifted to a lower wavenumber. that the presence of Fe(II) accelerated the chlorination reaction of polyamide membranes in both HOCl and NH2 Cl environments. ATR/FT-IR spectroscopy (Fig. 3) confirmed membrane oxidation/chlorination only for polyamide membranes exposed to HOCl, HOCl and Fe(II), and HOCl and Fe(III), but not for membranes exposed to NH2 Cl, NH2 Cl and Fe(II), and NH2 Cl and Fe(III). Changes in the amide II N H bend near 1540 cm−1 , the C C ring vibrations near 1610 and 1448 cm−1 , and the out-of-plane C H bend near 780 cm−1 were visible in the ATR/FT-IR spectra of chlorine-damaged membranes when compared to virgin membrane material. The C C and C H out-of-plane bending vibrational bands decreased in intensity, whereas the amide II band intensity not only decreased but also shifted to a lower wavenumber. These changes probably result from chlorine addition to the aromatic ring [32]. Figs. 4 and 5 show experimental results using various potential mitigation measures to halt the polyamide membrane chlorination reaction. As with the aqueous chemistry Fig. 5. Effect of sodium sulfide (7.0 mg/L Na2 SO3 ), citric acid (1.0 mg/L) and ammonia (20 mg/L NH3 ) addition to control for Fe(II)/NH2 Cl oxidation of polyamide membranes. Star symbols indicate acid cleanings. All data taken at pH 8.0 and normalized to 25 ◦ C. for HOCl, Fe(II), Fe(III), and NH3 is governed by pH, two pH ranges were tested: pH 6.0 and 8.0 (Fig. 4). When the pH was decreased from 8.0 to 6.0, it was observed that (1) the rapid decline in water flux (not shown) stopped and (2) the increase in salt passage (B value) after membrane cleaning also stopped. Fig. 5 shows that neither the presence of citric acid (CA) nor that of excess NH3 could arrest the membrane degradation reactions. However, the quenching of chlorine by Na2 SO3 did halt the decline in membrane performance. 5. Discussion Fig. 4. Effect of pH on Fe(II)/NH2 Cl oxidation of polyamide membranes. Star symbols indicate acid cleanings. All data normalized to 25 ◦ C. As ammonium sulfate and sodium hypochlorite were added at 4.4 × 10−5 M as NH3 (0.75 mg/L) and 3.5 × 10−5 M as Cl2 (2.5 mg/L), respectively, the theoretical free chlorine concentration was approximately 2.6 × 10−10 M [see Eq. (7)]. Therefore, the level of free chlorine present was approximately five orders of magnitude lower when using combined C.J. Gabelich et al. / Journal of Membrane Science 258 (2005) 64–70 chlorine—far below the MDL for free chlorine. This low level of free chlorine alone would probably not result in degradation of the polyamide membrane surface. However, in a chloraminated environment, the presence of Fe(II) resulted in chlorination of the polyamide membrane. These data suggest that previous polyamide membrane oxidation using Fe(III) coagulation [1] may have been caused by the Fe(III) oxidizing substandard iron fittings or parts, leading to the liberation of Fe(II) ions just prior to RO treatment. The authors hypothesize that • NH2 radicals produced during the Fe(II)/NH2 Cl redox reaction [see Eq. (9)] [24] attack the N H bond on the membrane amide group, thereby lowering the activation energy for chlorination of the mphenylenediamine ring. The • NH2 lowered the chlorination reaction activation energy sufficiently such that the NH2 Cl reacted directly with the polyamide membrane. The source of Cl+ was not the Fe(II)/NH2 Cl redox reaction that results in Cl− formation, but rather the NH2 Cl itself, which was present in excess. Further proof of this hypothesis can be found in Fig. 5, where excess ammonia was added in the presence of Fe(II) and NH2 Cl. By increasing NH3 from 0.6 to 20 mg/L, the free chlorine concentration would decrease from 2.6 × 10−10 to 8.6 × 10−11 M HOCl, yet no change in reaction rate (as indicated by increasing salt passage [B value]) was observed. Therefore, the conclusion may be that HOClderived Cl+ does not play a significant role in polyamide membrane oxidation in NH2 Cl environments. Similarly to Fe(II)/NH2 Cl redox chemistry, the reduction of HOCl in the presence of Fe(II) may result in • OH radical formation [see Eq. (14)]: 69 ence of Fe(III) solids enhances the formation of • NH2 (Eqs. (11)–(13)) such that at higher pH, the conversion of Fe(II) to Fe(III) accelerates the • NH2 mediated chlorination pathway. The determination of the exact pathway was beyond the scope of this paper. Dechlorination with Na2 SO3 successfully halted the chlorination reaction (Fig. 5) by reducing the chlorine state in NH2 Cl to Cl− , thus alleviating the potential adverse affects of Fe(II) on the polyamide membrane surface. The addition of citric acid did little to halt the chlorination reaction, as the affinity of citric acid for dissolved iron [Fe(II)] is significantly less than that for colloidal iron [Fe(III)]; therefore, Eq. (9) might still have proceeded. 6. Conclusions and recommendations This study provided evidence for a Fe(II)-mediated pathway for NH2 Cl attack on the polyamide membrane. This paper clearly shows that the practice of chloramination for biological fouling control may be detrimental to long-term membrane performance. Over time, NH2 Cl may react with the membrane surface. Whether chlorination under these conditions will result in the rate-limiting step for membrane life has yet to be determined. Longer-term testing is needed to establish whether membrane chlorination is an important byproduct of biogrowth control through chloramination. Acknowledgments Fe(II) + HOCl ⇒ Fe(III) + • OH + Cl− (14) It is also hypothesized that the • OH radical may then act similarly to • NH2 , attacking the N H bond on the membrane amide group. This hypothesis can be demonstrated by the increased salt passage of the HOCl and Fe(II)-exposed membrane relative to HOCl alone (Fig. 3) and the increased presence of chlorine on the polyamide surface of the HOCl and Fe(II)-exposed membrane (Table 2). Oxidation of Fe(II) by oxygen has been shown to be a complex function of solution pH [33]. Mean dissolved oxygen during the present study was 8.9 mg/L. This complex pH dependence is hypothesized to occur as a result of the parallel oxygenation of Fe(II); its hydroxo-complexes [Fe(OH)+ , Fe(OH)2 , Fe(OH)3 − ]; and carbonate complexes [FeCO3 , Fe(CO3 )2 2− , Fe(CO3 )(OH)− ] [34]. These hydroxyl and carbonyl ligands donate electron density to the Fe(II) atom, thus making the metal a better reductant [23,24]. Consequently, as pH increases from 6.0 to 8.0, increased ligand binding would occur as well as, possibly, increased reactivity of Fe(II) to reduce NH2 Cl. Fig. 4 shows sequential experiments using Fe(II) and NH2 Cl at pH 6.0 and 8.0. These data suggest that after a short period of time at lower pH, the chlorination reaction does not progress—potentially through suppression of Fe(II) reactivity. An alternative hypothesis is that the pres- This project was sponsored by the Desalination Research and Innovation Partnership, with funding graciously provided by the California Energy Commission’s Public Interest Energy Research Program. All RO testing was conducted at the U.S. Bureau of Reclamation’s (USBR’s) Water Quality Improvement Center in Yuma, Arizona. Burns and Roe (B&R) Service Corporation conducted the operation and maintenance of the research platforms. Special thanks are extended to Michael Norris (USBR), Angela Adams (USBR), Brent Corbett (B&R), Ed Lahota (USBR), and all the B&R operators. Kudos to Dr. Richard Valentine at the University of Iowa for helping the authors understand the importance of ammonia radical chemistry. Special thanks, also, to Dr. Krassimir N. Bozhilov at the Central Facility for Advanced Microscopy and Microanalysis, University of California, Riverside, for the excellent spectroscopic analysis. Additional thanks to James Lozier (CH2M Hill) for his advice early in the project. References [1] C.J. Gabelich, T.I. Yun, B.M. Coffey, I.H. Suffet, Effects of aluminum sulfate and ferric chloride coagulant residuals on polyamide membrane performance, Desalination 150 (2002) 15–30. 70 C.J. Gabelich et al. / Journal of Membrane Science 258 (2005) 64–70 [2] C.J. Gabelich, T.I. Yun, B.M. Coffey, I.H. 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