Degradation of polyamide membranes exposed to chlorine: an

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Degradation of polyamide membranes exposed to chlorine: an impedance spectroscopy study Mikhail Stolov, and Viatcheslav Freger Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.8b04790 • Publication Date (Web): 01 Feb 2019 Downloaded from http://pubs.acs.org on February 4, 2019

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Environmental Science & Technology

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Degradation of polyamide membranes exposed to chlorine: an impedance spectroscopy study

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Mikhail Stolov,1 Viatcheslav Freger1,2,3*

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1Wolfson

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e-mail: [email protected]

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Department of Chemical Engineering, Technion – IIT, Haifa 32000, Israel Water Research Institute, Technion – IIT, Haifa 32000, Israel 3Grand Technion Energy Program, Technion – IIT, Haifa 32000, Israel 2Grand

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ABSTRACT

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Polyamide is the key material in modern membrane desalination, however, its well-known and

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incompletely understood drawback is low tolerance of chlorine, the most efficient inline

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disinfectant. Here we report a first investigation of the mechanism and kinetics of chlorine attack

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using electrochemical impedance spectroscopy (EIS) that directly probes changes in ion

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permeation upon chlorination at different pH, focusing on its early stages and low chlorine

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concentrations (15 to 197 ppm). EIS results partly conform to established two-stage mechanism

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that proceeds as N-chlorination followed by either C-chlorination in acidic conditions or amide

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bond scission in alkaline conditions. However, early-time kinetics in acidic conditions shows

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inconsistencies with this model, explained by possible effects of direct ring chlorination and finite

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polymer relaxation rates. The findings indicate that (a) N-chlorination reduces membrane polarity

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and ion permeability, while C-chlorination has an opposite effect; (b) chlorination in acidic

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conditions must involve other reactions, such as direct ring chlorination, in addition to N-

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chlorination and Orton rearrangement; and (c) the ultimate chemical transformations (C-

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chlorination or amide bond scission) result in an irreversible increase in membrane polarity and

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loss of ion rejection. The results highlight the potential of EIS as a powerful and sensitive tool for

ACS Paragon Plus Environment

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studying chemical degradation of ion-selective materials that may assist in developing new

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chlorine-resistant membranes.

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INTRODUCTION

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Polyamide thin-film composite membranes have been increasingly used for water desalination and

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purification since 1980s.1 Their key component is an aromatic or semi-aromatic polyamide layer,

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50-200 nm thick, responsible for the high salt rejection and good water permeability.2 However,

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the well-known drawback of polyamide is its low tolerance to oxidizers, especially, chlorine or

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hypochlorite, which precludes the direct in-line use of these highly efficient disinfectants within

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membrane modules in desalination processes. Exposure to chlorine, even at a few ppm levels

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present in tap water, results in chemical degradation, which decreases performance, shortens

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membrane lifetime and increases operational costs.3

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Despite the large number of studies focusing on polyamide membranes damage by chlorine,

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contradictory reports on its effect on solute rejection, permeability, and surface properties may be

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found in the literature. Do et al. summarized the impact of pH and chlorine concentration on water

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permeability, surface hydrophilicity, NaCl and boron rejection of polyamide membranes.4 It is

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generally agreed that the decline in membrane performance due to chlorination is initiated by

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changes in the polymer chemistry.5-7 Proposed failure mechanisms include decreasing the number

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of cross-links, changes in the amount of charged groups and hydrogen bonding sites, subsequently

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modifying its structure and morphology as well as the rigidity, hydrophobicity, and free volume.8

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The chemical transformations that take place in polyamide upon chlorine attack were discussed in

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several recent publications.9-12 The commonly accepted general scheme is presented in Fig. 1.

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Oxidation of aromatic polyamide by chlorine species apparently proceeds through reversible N-

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chlorination of the amide nitrogen, followed by either hydrolysis of C-N bond or so-called Orton

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rearrangement that moves the Cl group to the aromatic ring.13, 14 Direct ring chlorination at para-

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position of the aromatic ring connected to the NH group is another route.15, 16 Exact chlorination


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scenario strongly depends on pH, as both N-chlorination and Orton rearrangement processes are

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promoted in acidic conditions17, whereas dechlorination of nitrogen atom as well as C-N bond

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cleavage occur in alkaline media.16, 18 Another factor is the speciation of free chlorine in aqueous

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solutions, which depends on pH. The relevant species are molecular chlorine Cl2, hypochlorous

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acid HClO, and hypochlorite anion OCl-, which dominate at pH ca. from 0 to 2, from 4 to 6, and

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above 10, respectively, with mixed regimes in between.11, 19 O

O N H

O

N H

Cl N H

n

O

O

O

O N H

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N H

IV

N H n

V

O

O

N Cl

II

III

N H

N H n

O

O

I

n OH H + N Cl

n N H

N H

O

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Figure 1. Mechanism of chlorine attack on polyamide membranes, where route I – N-chlorination; II –

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dechlorination; III – direct ring chlorination; IV – Orton rearrangement; V – chlorination-promoted

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hydrolysis.

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Notably, most studies of polyamide chlorination mechanism were based on surface analysis

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methods, such as X-ray photoelectron spectroscopy (XPS),15,

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Fourier transformed infrared spectroscopy (ATR-FTIR),22-24 and, recently, Rutherford

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backscattering spectrometry (RBS).11 The obvious limitation of these techniques is that they reflect

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average composition of the film, probe chemistry rather than transport characteristics, and are

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performed ex situ, as a post mortem analysis.

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attenuated total reflectance


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As a different and more direct approach, here we employ electrochemical impedance spectroscopy

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(EIS) to examine the effect and kinetics of polyamide layer chlorination. EIS proved to be a

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powerful tool for studying ion transport and selectivity in desalination membranes.25-28 Unlike

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surface-analytical and microscopic methods, it probes the trans-membrane conductivity, a

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transport property closely related to salt rejection capability of the polymer. It also reflects ion

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transport characteristics in the actual barrier part of the film, rather than average characteristics of

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the entire film. Conductivity is controlled by the most permeable ion of the salt (typically, the

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counter-ions) rather than least permeable (co-ions), as in the case of salt permeability observed in

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filtration. Thereby, conductivity, monitored by EIS in real time, more directly reflects changes in

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the ion exclusion due varying membrane charge and/or membrane polarity in the course of

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chlorination. These features are used here to monitor the chlorination kinetics and clarify the

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mechanism and impact of pH and chlorine concentration on the ion transport and exclusion in the

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polyamide layer.

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EXPERIMENTAL

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Chemicals and Materials. Sodium hypochlorite solution of concentration 4-4.99% was obtained

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from Sigma Aldrich. Exact concentration of free chlorine (interchangeably referred to below as

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concentration of chlorine, HClO or NaClO) was determined by iodometric titration with sodium

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thiosulfate.29 All other used chemicals were of analytical grade with purity over 99.5%. Ultrapure

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water (UPW, 18.2 MOhm∙cm) was used for preparing each solution. SWC5 reverse osmosis

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membrane was obtained from Hydranautics.

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Attaching polyamide film to the electrode surface. The method of isolating the polyamide layer

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from thin-film composite membranes was described in detail elsewhere.27, 30 Briefly, non-woven


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support layer was peeled off the membrane, then remaining part of membrane was attached, with

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the polyamide side facing down, to a surface of working PEEK-shrouded glassy carbon (GC)

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rotating disc electrode (Pine, GC diameter 5 mm) with a drop of iso-propanol. Polysulfone layer

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was then carefully washed away with N,N-dimethylformamide (DMF), dichloromethane, and

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again with DMF, and the remaining polyamide layer was dried to ensure it sticks firmly to the

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electrode surface. The procedure was then repeated twice to attach two more polyamide layers on

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top of each other. Triple layers were found to be more stable in chlorination experiments and had

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a larger and more reproducible resistance then they were preferred over single layers. Electrode

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covered with polyamide layer was stabilized by immersing in 1 M NaCl solution for at least 3

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hours before experiment.

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EIS measurements. The polyamide-covered working electrode (WE), a Pt foil (counter electrode,

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CE), and Ag/AgCl/KCl reference electrode (RE) were immersed in 1 M NaCl (Fig. 2a). This high

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NaCl concentration allowed good differentiation between membrane and solution resistance.

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Chlorination was initiated by adding sodium hypochlorite solution. When necessary, HCl and/or

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NaOH were added to adjust pH in amounts negligible compared to NaCl. The solution was

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continuously slowly purged with N2 to eliminate dissolved CO2 and was protected from direct light

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with an aluminum foil wrap. At the end of experiments, about one hour after chlorine addition, the

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residual chlorine was measured using the DPD (N,N-Diethyl-1,4-phenylenediamine) test31 and

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was found to exceed 80 % of initial value.

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EIS spectra were recorded on SP-300 potentiostat (BioLogic) every 1-1.5 min to monitor changes

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before and after adding chorine (hypochlorite). Each point in the EIS spectra was the average of

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three measurements. Using built-in EC-Lab software, the spectra were fitted to an equivalent

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circuit (see Fig. 2b),25 which yielded the values of membrane resistance Rm and trans-membrane


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conductivity Gm = 1/Rm. Due to inherent variability of small polyamide films, Gm could vary

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significantly (average 0.03 S, standard deviation 0.025 S), reflecting differences in effective

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thickness and charge of polyamide films. To account for these variations, the reported conductivity

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was normalized to the initial value stabilized in 1 M NaCl. Each point in the reported EIS spectra

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was the average of three measurements, and each experiment was duplicated, which showed a

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good reproducibility of normalized spectra and their variation with time. An example of

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consecutive EIS spectra recorded after chlorine addition and fitted values of resistance is shown

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in Fig. 2c.

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Figure 2. (a) Experimental setup, (b) diagram of the equivalent circuit and (c) EIS spectra of polyamide coated electrode in 1 M NaCl solution recorded at different times following addition of 81 ppm chlorine.

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RESULTS AND DISCUSSION

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Differentiating effects of chlorination and pH

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Figure 3a shows the change of membrane conductivity deduced from EIS spectra of polyamide

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film before and after adding NaClO to a NaCl solution at pH 7. It is seen that the membrane

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conductivity was stable before addition, but after exposure to NaClO it grew progressively, i.e.,

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membrane resistance dropped until fitting could no more differentiate it from the background

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resistance. Increasing conductivity indicates a damage by chlorine and loss of ion rejection. It must

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be noted that, since HClO is a weak acid (pKa = 9), addition of NaClO also increased pH to 9.6.

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Since, it is known that increased dissociation of the carboxylic groups in polyamide layer at

136

elevated pH also increases the membrane charge and swelling hence conductance,32,

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important to differentiate the direct effect of pH change from that of chlorination.

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For this purpose, two additional experiments were conducted. First, after stabilizing the membrane

139

in the NaCl solution, pH was adjusted to 9.6 using a NaOH solution, which indeed increased the

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membrane conductivity by ~30% to a value that stabilized after 10-15 min (Fig. 3b). This increase

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in conductivity could be attributed to increased pH only and the stabilization time should represent

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swelling and relaxation of the polymer in response to deprotonation and ionization of carboxylic

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charges, most of which has a pKa ~8-9.34 Note that this time is unlikely to be related to much faster

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dissociation kinetics or ion diffusion within the film. Indeed, for representative values of ion

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diffusivity in the membrane D ~ 10-11 m2/s and membrane thickness  ~ 100 nm, the diffusion time

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should be of the order 2/D ~ 1 ms.

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Finally, in a concluding experiment, the pH of NaCl solution was first adjusted to 9.6 by adding

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NaOH and membrane conductivity was let stabilize; thereafter the solution was replaced with a

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NaCl solution with added NaClO having the same pH.9.6 (Fig. 3c). Since after adding the last

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it was


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NaClO solution the pH did not change, subsequent increase of conductivity could be attributed

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solely to the chemical and physical impact of chlorination.

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153 154 155 156 157 158 159 160 161 162

Figure 3. Variations of membrane conductivity in alkaline (a-c) and acidic (d-e) conditions: (a) direct injection of NaClO with pH 9.6 obtained after injection; (b) pH adjustment to 9.6. with NaOH and stabilization; (c) pH adjustment to 9.6 with NaOH and stabilization, followed by replacement with NaClO solution pre-adjusted to pH 9.6; (d) pH adjustment to 2.4 with HCl and stabilization, followed by injection of NaClO and pH shift to 4.8; (e) pH adjustment to 2.4 with HCl and stabilization, followed by pH adjustment to 4.8 with NaOH and stabilization; (f) pH adjustment to 4.8 with HCl and stabilization, followed by replacement with NaClO solution pre-adjusted to pH 4.8. All solutions contained 1 M NaCl as background electrolyte, added chlorine concentration was 197 ppm. Results were normalized to the initial membrane conductivity G0 stabilized in 1 M NaCl solutions.

163 164

A similar approach was followed in acidic conditions. Figure 3d shows time dependence of

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membrane conductivity in NaCl solution after consecutive addition of HCl (to change pH to acidic)


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and NaClO. Addition of acid resulted in pH dropping to 2.4, which caused a significant drop in

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membrane conductivity due to protonation of carboxylic groups and decrease in membrane charge.

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This process was relatively slow, as expected of carboxylic group protonation and subsequent

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polymer relaxation to a less swollen state. After adding NaClO, pH rose to 4.8 and membrane

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conductivity rapidly dropped, though not as much as when pH was simply adjusted from 2.4 to

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4.8 with NaOH, without adding NaClO (Fig. 3e). This drop in conductivity after pH increasing

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from 2.4 and 4.8 could not be related to a change in membrane charge, since increased pH is

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supposed to increase dissociation of carboxylic charge and produce an opposite effect (cf. Fig. 3b).

174

However, RBS data on neutralization behavior of RO membranes indicate that in this pH range,

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well below pKa, minor pH changes were not supposed to significantly alter the membrane charge.

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We then presume that, in this pH range (2.4 to 4.8), the nearly instant drop in conductivity could

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partly be related to >100-fold drop in proton concentration in solution and therefore in the

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membrane as well, rather than to a change in membrane charge. Rapid removal of protons, known

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to have a very high affinity to polyamide,35 as well as high mobility, could result in such a fast

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change of conductivity in this pH range, in which carboxylic charges remain mostly non-

181

dissociated.

182

However, the drop in conductivity after simple adjustment to pH 4.8 with NaOH in Fig. 3e was

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much smaller that after adding NaClO, which caused a rapid and steady increase of the

184

conductivity up to the point where it could no more be determined (Fig 3d). Therefore, in the last

185

experiment displayed in Fig. 3f, pH was then kept constant throughout, by first adding HCl to the

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NaCl solution to adjust pH to 4.8 and, after membrane conductivity was stabilized, the solution

187

was replaced with an HCl solution with added NaClO of the same pH 4.8. The subsequent change


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of conductivity, i.e., a nearly instant small drop and subsequent rapid increase, could only reflect

189

the effect of chlorination and not that of pH.

190

The above experiments indicate that, while pH change could produce some moderate initial shift

191

in membrane conductivity, the subsequent change leading to a steady increase up to undetectably

192

large Gm values (i.e., small Rm) must be attributed to the effect of chlorination. This large increase

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in membrane conductivity as a result of chlorination is the manifestation of either a larger increase

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in membrane charge and/or weaker ion exclusion, consistent with reported reduced salt rejection

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of chlorine-damaged membranes.

196 197

Figure 4. Variation of normalized membrane conductivity with time after adding 197 ppm chlorine

198

for different pH. The quoted pH was measured shortly after adding hypochlorite. The inset shows

199

dependence of membrane conductivity measured at the initial point of each curve after adding

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hypochlorite. For all curves the conductivity was normalized to the stabilized conductivity in

201

neutral 1 M NaCl solution prior to changing pH and adding hypochlorite.

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Effect of pH on kinetics and mechanism of chlorination.

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In order to see the effect of pH in more detail, conductivity changes were examined for several pH

206

values in the range 4.8 to 9.6 for the same added hypochlorite concentration 197 ppm. In these

207

experiments, the simplified procedure displayed in Figs 3a and 3d was used. The membranes were

208

first preconditioned in neutral 1 M NaCl, then stabilized in chlorine-free 1 M NaCl adjusted to

209

desired pH, and, finally, 197 ppm hypochlorite was added. The pH quoted in Fig. 4 is the one

210

measured shortly after adding hypochlorite, while the shown conductivity values are normalized

211

to the one measured for the same membrane in the first neutral 1 M NaCl solution, as the common

212

basis for all pH conditions. The starting point of each curve then represents the complex primary

213

effect of pH, which is explicitly shown in the inset of Fig. 4. It is seen that normalized conductivity

214

at the starting point of each curve monotonically increases with pH, as expected from variations

215

of the membrane charge with pH.

216

However, comparison with the results in Fig. 3 d-f shows that for acidic pH the drop is significantly

217

larger than expected from direct effect of pH, i.e., protonation and subsequent change in membrane

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charge and swelling.35, 36 This strongly suggests that, at least in acidic conditions, the direct effect

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of pH is superimposed on some very fast primary effect of chlorination. This may also explain the

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different pattern of conductivity variations at low and high pH, whereby for all pH > 7 conductivity

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steadily increased with time roughly at a constant rate, while in acidic pH the initial slope is smaller

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and increases over time. The markedly different behavior at low and high pH suggests differences

223

in the chlorination mechanism, as further clarified below by analyzing the dependence on

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hypochlorite concentration.


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Figure 5. Chlorination of polyamide membrane at fixed pH: (a) variation of the normalized membrane

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conductivity after exposure to 15 ppm chlorine at pH 4.8 and 9.6; (b) conductivity variation after exposure

228

to different chlorine concentrations at pH 4.8; (c) Best fit of the data in (b) to eqs. 3 and 4; (d) conductivity

229

variation after exposure to different chlorine concentrations at pH 9.6. In (a) G0 was the conductivity

230

stabilized in 1 M NaCl solution before changing pH and adding chlorine, in (b) and (d) G0 was the stabilized

231

conductivity at specific pH before adding chlorine.

232 233

Effect of chlorine concentration on chlorination kinetics. The effect of initial chlorine

234

concentration was analyzed for two representative pH values, 4.8 and 9.6 in the range 15 to 197

235

ppm of initial chlorine. As in the previous section, the protocol included conditioning in neutral 1

236

M NaCl, followed by conditioning in 1 M NaCl adjusted to appropriate pH, but, at the final stage,

237

the entire solution was replaced with a hypochlorite solution adjusted to the same pH to rule out


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any direct effect of pH. Figure 6a displays the curves obtained for 15 ppm chlorine at the two pH,

239

4.8 and 9.6, normalized to the initial film conductivity in neutral 1 M NaCl, as in Fig. 4. It

240

highlights more than an order of magnitude difference between the conductivities measured at the

241

two pH. Therefore, in the subsequent Figures 5b and c the conductivity is normalized to the value

242

measured after stabilization at relevant pH shortly before adding hypochlorite. Such normalization

243

helps better resolve the kinetic behavior observed at early times and supplies the most consistent

244

common basis for comparing different samples, since the initial polyamide film conductivity could

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vary between samples.

246

The results in Fig. 5b obtained at pH 4.8 highlight the differences in chlorination kinetics for small

247

and large free chlorine concentrations [Cl]. At low [Cl], two distinct regimes are clearly observed,

248

whereby the conductivity first decreases, passes through a minimum, and thereafter increases with

249

time. The minimum shifts to earlier times and becomes shallower at higher [Cl]. Above 40 ppm

250

chlorine, the exact location of the minimum could no more be resolved, since it coincided with the

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very first point measured after hypochlorite addition. Also, variation of conductivity with time for

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higher chlorine concentrations, when minimum is no more resolved, became independent of [Cl].

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The transition from the initially decreasing conductivity to increasing one in acidic conditions (Fig.

254

5b) seems to conform to the accepted mechanism that involves two consecutive reactions, namely,

255

N-chlorination, i.e., irreversible reaction of free chlorine with amide nitrogen, followed by C-

256

chlorination via Orton rearrangement in acidic conditions or hydrolysis and subsequent amide link

257

cleavage in alkaline media.4, 11, 21 In acidic conditions, when de-chlorination is suppressed,18 this

258

corresponds to the following kinetic scheme

259

HClO, k1 k2 [N-H]   [N-Cl]   [C-Cl],

(1)


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where [N-H], [N-Cl] and [C-Cl] are concentrations of unreacted amide groups, chlorinated amide

261

groups and monochlorinated aromatic rings, respectively, and k1, and k2 are corresponding rate

262

constants. Assuming for N-chlorination a first order with respect to both [N-H] and free chlorine

263

[Cl] and for the Orton reaction first order with respect to [N-Cl], this yields the following kinetic

264

model

265

d [N-Cl]  k1[Cl]  [N-H]0  [N-Cl]  k2 [N-Cl], dt

266

whose solution is

267

[N-Cl]=  N-H 0

268

where [N-H]0 is initial amount of amide groups available for primary N-chlorination and

269

k  k1  Cl  k2 . For simplicity, eq. 2 implies that N-Cl groups return to be reactive N-H sites after

270

formation of C-Cl bonds, as well as after backward reaction. In fact, the XPS and RBS analysis

271

shows that at very high chlorine doses (104 ppm∙h) the ratio N:Cl saturates at about 1:1,11 which

272

suggests C-chlorination inactivates adjacent amide nitrogen. However, the approximation may be

273

reasonable for the very lower chlorine levels and doses used here, far from saturation.

274

Although EIS does not directly measure [N-Cl] and [C-Cl], the conductivity is assumed to

275

approximately obey the following relation (see discussion in the next section)

276

ln

277

where a and b are positive parameters and G0 is the conductivity just before adding hypochlorite.

278

The minus sign before a and plus before b reflect the fact that N-Cl groups reduce conductivity,

279

while C-Cl increase it (see next section).

k1  Cl k

1  e  ,  kt

[C-Cl]=  N-H 0

d [C-Cl]  k2 [N-Cl], dt

(2)

k2 k1  Cl  1  t  1  e  kt  ,  k  k 

G  a[N-Cl]  b[C-Cl], G0





(3)

(4)


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280

It follows from eqs. 3 and 4 that at short times (kt > 1), [N-Cl] will saturate and [C-Cl] will increase linearly with time thus the

284

impedance will asymptotically approach a linear dependence ln  G G0   At  B, with a slope

285

A  b  N-H 0

286

N-chlorination is the faster reaction, i.e., k  k1  Cl , the slope and intercept may simplify to

287

k   A  b  N-H 0 k2 and B    N-H 0  a  b 2  . k  

288

Unfortunately, the trends observed in Fig. 5b only partly agree with these relations. Namely, k

289

must increase with [Cl] for any specific values of constants k1, k2, a and b, therefore

k2 k1  Cl k

t  tmin  a bk2 . Conversely, at

k  k1  Cl  and intercept B    N-H 0  a  b 2  . If, as is apparently the case, k  k 

290

(a) At long times, when [Cl] increases, the slope A is expected to become approximately

291

constant and independent of [Cl], while the intercept B should become less negative and

292

closer to zero, as indeed observed in Fig 5b.

293

(b) The conductivity minimum is expected at about the same tmin, but should grow deeper for

294

larger [Cl]. However, Fig. 5b shows that the observed minimum shifts to earlier times,

295

grows shallower and tends to disappear with increasing [Cl], opposite to the model. Fig. 5c

296

demonstrates that a fit of the present data to eqs. 3 and 4 totally fails to reproduce the trend

297

of Gmin with increasing [Cl] observed in Fig. 5b.

298

This failure is even more puzzling, if we further observe that B (long times) and Gmin (short times)

299

are supposedly determined by the same ratio of N- and C-chlorination rates k1  Cl k2  k k2 , yet


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only for Gmin the effect of [Cl] is inconsistent with the model. This suggests that other mechanisms

301

might be involved at early times and reverse the effect of [Cl] on N- and C-chlorination rates in a

302

way that makes N-chlorination rate (or just its effect on conductivity) independent of [Cl] and, at

303

the same time, make C-chlorination rate increase with [Cl], opposite of eq. 3.

304

We speculate that this scenario is indeed likely and might result from (a) direct C-chlorination of

305

aromatic rings and (b) a finite rate of polymer relaxation that modifies the kinetics at short times.

306

Direct ring chlorination would add to C-chlorination rate in eq. 2 a term proportional to [Cl] with

307

yet another rate constant k3, as follows

308

d [C-Cl]  k2 [N-Cl]  k3 [Cl][C-H] , dt

309

Here, [C-H] may include sites not subject to the Orton rearrangement thus [C-H] may be unrelated

310

to [N-H]0. This would not contradict the Orton mechanism, which assumes that Cl2 required for

311

C-chlorination forms in the reaction of N-Cl group with Cl- ions.17 Note that Cl2 that may directly

312

chlorinate C-H sites is already present in NaCl + HClO solution at pH 4, according to the

313

equilibrium

314

HClO + H+ + Cl- ↔ Cl2 + H2O.

315

Availability of this Cl2 then opens a C-chlorination route parallel rather than consecutive to N-

316

chlorination. In present experiments, [H+] and [Cl-] were essentially fixed, then [Cl2] was

317

proportional to [Cl], and the equilibrium factor [Cl2]/[Cl] is part of k3 in eq. 5. At small times,

318

direct C-chlorination may outcompete the N-Cl-mediated reaction before [N-Cl] becomes

319

substantial and Orton mechanism takes over. This is formally equivalent to replacing k2 in eq. 3

320

with an effective value k2*  k2 1 

321

dependent at short times but, as [N-Cl] increases (eq. 3), approach a constant k2 at longer times.

 

(5)

(6)

 k3 [C-H] [Cl]  , which must be much larger than k2 and [Cl]k2 [N-Cl] 


17


Page 18 of 25

322

Furthermore, Figure 3 shows that polyamide conductivity may take up to 0.5 h to relax in response

323

to a pH change and a similar relaxation rates may be expected in response to chlorination. This

324

will more strongly obscure the effect of N-chlorination, the fastest process, and make it difficult

325

to observe the minimum and reduce its apparent depth, when tmin is short compared to the polymer

326

relaxation time. This means that, in order to describe conductivity variation at short times, the [Cl]-

327

dependent e

328

relaxation rate constant. The dotted line in Fig. 5b, which is a single exponential with a rate

329

constant k* = 20 h-1 (i.e., relaxation time 0.05 h) indeed envelops reasonably well the initial parts

330

of all curves. Ultimately, the finite relaxation rate and direct C-chlorination at short times would

331

modify the expressions for tmin and Gmin to t  tmin  a bk2* and ln  Gmin G0    k *  N-H 0 a 2 2bk2* .

332

These effective short-time values k2* and k * that replace k2 and k and show a different [Cl]-

333

dependence would make variation of tmin and Gmin with [Cl] consistent with present observations.

334

Yet, this modification would not affect the relations for asymptotic long-time slope A and intercept

335

B, still determined by parameters k2 and k. Obviously, future research will be required to support

336

this mechanism.

337

The situation seems to be more straightforward in alkaline conditions, where N-chlorination (Fig.

338

5d) is followed by amide bond hydrolysis. The process can be schematically presented as

339

k2   [N-Cl]  [N-H]   [R-COO  ] + [R -HN-Cl] , 

340

where R’-NH-Cl and R-COO- are the final products of amide bond hydrolysis in alkaline pH. In

341

contrast to the results for pH 4.8, the chlorination kinetics at pH 9.6 shown in Fig. 5d does not

342

reveal any significant dependence on chlorine concentration, even for the lowest ones used. Nor a

343

minimum of conductivity after addition of chlorine to alkaline solution of NaCl was observed.

 kt

term in eq. 3, may have to be replaced with a e

NaClO , k1

NaOH , k1

 k t

, where k* is a [Cl]-independent

(7)


18

Page 19 of 25


344

Likely, N-chlorination in alkaline conditions was much faster than in acidic conditions. It could

345

then be undetectable in the present experimental setup and chlorine concentrations involved.

346

However, the observed steady rise of conductivity, independent of chlorine concentration is

347

consistent with the two-stage mechanism, whereby the extremely fast first reaction with

348

hypochlorite presents no kinetic limitation whatsoever for all measured chlorine concentrations.

349

As a result, hypochlorite mainly acts as a promoter of the amide bond hydrolysis and formation of

350

free carboxylic charges. Curiously, the rate of conductivity increase is similar to what is observed

351

at pH 4.8 and high chlorine levels that might suggest a similar chlorination mechanism. However,

352

this similarity is apparently coincidental, since different mechanisms in acidic and alkaline

353

conditions are well established.

354

Discussion: physical-chemical mechanism behind conductivity variations. The previous

355

section presents the formal kinetic analysis that, in most aspects, agree with the kinetics of

356

chlorination obtained by other methods11,

357

enabled to extend the analysis to shorter times, smaller hypochlorite concentrations, and

358

chlorination exposures as low as a few tens ppm∙h. The steady increase of membrane conductivity

359

in alkaline pH is consistent with hydrolytic scission of amide bonds and formation of fixed charges.

360

However, the more complex effect of chlorination on conductivity in acidic conditions warrants

361

discussion. At acidic pH, neither N- nor C-chlorination should result in formation of new fixed

362

charges, therefore the main question is why they modify conductivity and why in the opposite way.

363

The steric effect of insertion of chlorine atoms into dense polyamide structure must be ruled out,

364

since N-Cl and C-Cl groups would produce comparable effects, while they are actually opposite.

365

Still, a drop in conductivity suggests that N-chlorinated membrane contains less ions, which,

366

without new charges formed, means stronger ion exclusion. This assumption also agrees well with

12.

Here, the sensitivity and time resolution of EIS


19


Page 20 of 25

367

reported increase in salt rejection upon minor chlorination in acidic conditions.37 In absence of

368

alternative mechanisms, this suggests that N-chlorination lowers the polarity and dielectric

369

constant of polyamide. This well agrees with the reported drop in water permeability observed

370

upon mild exposures to chlorine in acidic conditions.4, 7

371

Following the same kind of argument, it may be concluded that C-chlorination produces an

372

opposite effect, i.e., increases the polarity of polyamide, which weakens ion exclusion. Again, this

373

is consistent with increased conductivity observed here and reported drop in salt rejection and

374

increase of water permeability upon prolonged exposure to chlorine in acidic conditions.

375

The opposite effects of N- and C-chlorination on polyamide polarity may seem puzzling, but we

376

note that they may possibly be related to the polarity of the newly formed N-Cl and C-Cl bonds

377

relative to the pre-existing N-H and C-H. Since the electronegativity of N and Cl is relatively close

378

(3.04 and 3.16) and significantly larger than that of C and H atoms (2.55 and 2.20), N-H bond is

379

more polarized than N-Cl and N-chlorination is indeed supposed to reduce the polyamide polarity.

380

On the other hand, C-H bond should be less polar than C-Cl and C-chlorination is supposed to

381

increase polarity. For instance, the dielectric constant of chlorobenzene (5.52) is substantially

382

higher than that of benzene (2.27), similar to the effect of mono-chlorination on other organics.

383

Conversely, in the case of N-chlorination, analogous comparison of ammonia and chloramine

384

shows that the latter is much less basic (pKb 14 vs 4.5) and more acidic (pKb 15 vs 32), i.e., much

385

less readily binds and more readily releases proton, indicating a smaller polarization of N-Cl bond

386

and smaller negative atomic charge of N atom.

387

In addition to the changes in polarizability of polyamide, N- and C-chlorination may also modify

388

its propensity to H-bonding and thus vary water content in the polyamide layer. Hydrophobicity

389

of aromatic matrix makes H-bonding to amide sites pivotal to polyamide hydration, thereby


20

Page 21 of 25


390

removal of N-H sites by N-chlorination must have a strong negative impact on polymer swelling

391

and thus on water and ion permeation. In contrast, C-chlorination of aromatic matrix should

392

promote swelling and water permeation and thus also weaken ion exclusion. A similar effect

393

should be produced by hydrolytic cleavage of amide bonds, leading to formation of more polar

394

groups and higher ion content than those present in the original amide.

395

The above arguments explain and justify the reason for choosing the particular form of eq. 4 to

396

correlate conductivity with the changes in the content of N-Cl and C-Cl groups, in particular, the

397

minus sign before a and the plus sign before b. In addition, the logarithmic dependence was used

398

as a generic relation to match the non-linear effects of [N-Cl] and [C-Cl] on conductivity.

399

However, this choice may also reflect the more profound fact that logarithmic function relates the

400

ion concentrations in the membrane to the ion self-energy within polyamide, which is directly

401

related to the membrane dielectric properties, i.e., polarity.38 If formation of [N-Cl] and [C-Cl]

402

affects the ion-self-energy in polyamide in an approximately linear manner, the logarithmic

403

function becomes the most appropriate choice to describe their effect on conductivity.

404

Implications. EIS was employed for studying the mechanism of chlorination-induced polyamide

405

membrane degradation by monitoring early changes in their ion-conducting properties. Particular

406

advantageous is the fact that the method may examine the initial stages of chlorination and small

407

exposures and differentiate the effects of pH from chlorination-induced changes. EIS experiments

408

performed for 15 to 197 ppm chlorine concentrations and pH range 4.8 to 9.6 reveals a two-stage

409

mechanism, well observed in acidic conditions. Comparison with previous studies suggests that

410

the two stages could be identified as N-chlorination followed by either C-chlorination through

411

Orton rearrangement in acidic conditions or amide bond scission and formation of carboxylic

412

charges in alkaline conditions. However, the effect of chlorine on short-time kinetics in acidic


21


Page 22 of 25

413

conditions and, in particular, on observed conductivity minimum disagrees with the model. This

414

discrepancy was explained by possible direct aromatic ring chlorination, which would allow a C-

415

chlorination route that bypasses N-chlorination and Orton mechanism at short times, and polymer

416

relaxation that obscures fast initial kinetics at larger chlorine concentration and its effect on

417

observed conductivity.

418

Overall, the results indicate that (a) N-chlorination results in a reduced membrane polarity and

419

smaller ion permeability, while C-chlorination has an opposite effect; (b) the chlorination

420

mechanism in acidic conditions and its effect on conductivity cannot be fully explained by N-

421

chlorination followed by Orton rearrangement and may involve other processes, such as direct ring

422

chlorination, and (c) regardless of the specific conditions and short-time effects, the chemical

423

transformations (C-chlorination or amide bond scission) ultimately result in an irreversible

424

increase in membrane polarity and loss of ion rejection. The results highlight the potential of EIS

425

as a powerful and sensitive tool for studying chemical degradation of ion-selective materials that

426

may assist in developing new chorine-resistant membranes.

427 428

Acknowledgements

429

We thank Ms. Rhea Verbeke for fruitful discussions and suggestions. MS acknowledges the

430

support by the Center for Absorption in Science of the Israel Ministry of Immigrant Absorption.

431

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432 433 434 435 436 437 438

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Degradation of polyamide membranes exposed to chlorine: an

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