Impact of Divalent Cations on Dark Production of Hydroxyl Radicals

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Impact of Divalent Cations on Dark Production of Hydroxyl Radicals from Oxygenation of Reduced Humic Acids at Anoxic-Oxic Interfaces Peng Liao, Yuzhen Liang, and Zhenqing Shi ACS Earth Space Chem., Just Accepted Manuscript • DOI: 10.1021/ acsearthspacechem.8b00181 • Publication Date (Web): 16 Jan 2019 Downloaded from http://pubs.acs.org on January 20, 2019

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ACS Earth and Space Chemistry

Impact of Divalent Cations on Dark Production of Hydroxyl Radicals from Oxygenation of Reduced Humic Acids at AnoxicOxic Interfaces Peng Liao†, †School

Yuzhen Liang†, ‡,

and Zhenqing Shi†, ‡, *

of Environment and Energy, South China University of Technology,

Guangzhou, Guangdong 510006, People’s Republic of China

‡The

Key Lab of Pollution Control and Ecosystem Restoration in Industry Clusters,

Ministry of Education, South China University of Technology, Guangzhou, Guangdong 510006, People’s Republic of China

*Corresponding authors: Zhenqing Shi [email protected]

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ABSTRACT Humic acids (HAs) are redox-active and can serve as either electron acceptors or electron donors to participate in multiple redox reactions. In nature water, HA can be intimately associated with divalent cations, such as Fe2+, Ca2+, and Mg2+, through a series of reactions which may in turn affect the redox reactivity of HA. Recent advances have demonstrated that the oxygenation of reduced HA in the dark can produce •OH at anoxic-oxic interfaces. However, little is known about the roles of the divalent cations complexed with HA on the production of •OH from reduced HA. This study provides new knowledge regarding the impact of Fe2+, Ca2+, and Mg2+ ions on the dark production of •OH from oxygenation of reduced HA at anoxic-oxic interfaces over a wide range of environmentally relevant conditions. Results show that the rates and yields of •OH production increase with increasing Fe2+ concentration (0.18–0.89 mM). This is largely attributed to the formation of complexed Fe(II) with HA, which increases the number of Fe(II)/Fe(III) cycles and enhances the decomposition of formed H2O2, accelerating the rates of Fenton reactions under circumneutral conditions. However, the promotional effect of Fe2+ on •OH formation is greatly suppressed in the coexistence of high Ca2+/Mg2+ concentration (5–20 mM), likely due to the retarded formation of HA-Fe(II) complexes and competition of HA’s surface reactive sites by Ca2+/Mg2+ ions. Findings improve the current understanding of the dark production of •OH from reduced HA and provide valuable insights towards understanding of carbon cycling and contaminant fate at anoxic-oxic interfaces. 2 ACS Paragon Plus Environment

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KEYWORDS: reduced humic acids, oxygenation, dark hydroxyl radical production, divalent cations, mechanisms, anoxic-oxic interfaces

1. INTRODUCTION Humic acids (HAs) are a group of redox-active organic macromoleculars that occur in most environments and can participate in a wide range of redox reactions that affect the biogeochemical cycling of carbon and the redox transformation of metals and organic contaminants in both natural and engineered systems.1‒5 HA is capable of serving as either an electron acceptor and an electron donor, depending on prevailing redox conditions.6‒9 Quinones and phenolic moieties are considered to be the dominant electron-accepting and electron-donating groups for the redox activity of HA.3,6,9‒11 Under anoxic conditions, HA can be chemically or microbially reduced with quinones being the major reducible moieties.3,6‒8,10‒14 The reduced HA has a high reducing capacity and can reduce many metals (e.g., iron (Fe) and mercury) and organics (e.g., chlorinated aliphatic pollutants).15‒17 A recent study demonstrated that the oxidation of reduced HA by O2 can produce •OH in the dark.3 The formation of •OH has been attributed to the redox-active quinone moieties, which could transfer three electrons to O2 to form •OH.3 The yields of •OH production were shown to vary from 42 to 160 mmol per mole of electrons donated by the reduced HA.3 As the most potent reactive oxygen species in the environment, the produced •OH was expected to oxidize dissolved HA to yield low molecular weight 3 ACS Paragon Plus Environment


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organic compounds and CO2.3,18‒20 The produced •OH was also suggested to have a profound impact on the contaminant redox transformation and degradation.3,20 Although the dark oxidation of reduced HAs by O2 represents an important source of •OH formation at anoxic-oxic interfaces,3,18,19 the production of •OH from a given reduced HA may likely be influenced by the presence of cations such as Fe2+, Ca2+, and Mg2+, which can form surface complexes with HA. Ferrous Fe is a redox-active divalent cation and can form HA-Fe(II) complexes that play a pivotal role in the formation of •OH under circumneutral conditions.21‒24 The HA-Fe(II) complexes can be detected in many anoxic-oxic interfaces with dynamic redox environments such as groundwater, wetlands, and marine environments.25‒27 A large body of work demonstrated that the formation of Fe(II)-L complexes (e.g., L= humic or fulvic acids) may either accelerate or retard the rates of •OH formation from oxidation of Fe2+ by O2 and H2O2.22,24,28‒31 For example, a recent study showed that Fe(II)-L complexes enhanced the Fe(II)-mediated reduction of O2 to O2- and decomposition of H2O2, facilitating •OH formation.24 In contrast, Fe(II)-L complex can also significantly decrease the yield of •OH formation upon oxygenation of Fe(II) in seawater.29 Although great strides have been made in studying the properties of Fe(II)-L complexes,22‒24,28‒31 most of these previous studies focused primarily on the effect of oxidized HA (or fulvic acid) on the kinetic of Fe2+ oxidation and concurrent formation of •OH, and little has been made to understand the impact of Fe2+ on the formation of •OH from oxidation of reduced HA. While recent field studies supported that the 4 ACS Paragon Plus Environment

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oxidation of reduced dissolved organic matter (DOM) and Fe2+ by O2 produced abundant •OH at anoxic-oxic interfaces in soil and surface waters,18,19 the underlying mechanisms are not thoroughly appreciated or understood. Both Ca2+ and Mg2+ ions are abundant in natural waters and may have a significant impact on the properties of HA and HA-Fe(II) complexes.4,32,33 Ca2+ and Mg2+ can increase the solution ionic strength and decrease the negative charges of HA, both of which may result in aggregation of HA.32‒34 This may subsequently influence the redox reactivity of HA.35 In natural waters, the concentrations of Ca2+ and Mg2+ were found to be typically higher than that of Fe2+.36 Such a higher concentration of alkaline-earth metals may significantly compete with Fe2+ for the binding sites in HA,31,37,38 affecting the formation rates of HA-Fe(II) complexes. While there is a large number of studies pertaining to the interactions of HA with Ca2+/Mg2+,32‒34,37‒41 information regarding the influence of Ca2+ and Mg2+ on the dark formation of •OH from oxygenation of reduced HA is largely absent. The objective of this study was to fundamentally investigate the impact of Fe2+, Ca2+, and Mg2+, individually and in combination, on the dark production of •OH from oxygenation of reduced HA at anoxic-oxic interfaces over a wide range of environmentally relevant concentrations of Fe2+ (0−0.89 mM) and Ca2+/Mg2+ (0−20 mM) at pH 7.0. The rate and yield of •OH production affected by divalent cations were deciphered through a combination of well-controlled batch experiments and an array of complementary characterization techniques. Knowledge obtained in this study 5 ACS Paragon Plus Environment


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improves our current understanding of the dark production of •OH under more complex aquatic conditions and contributes to the prediction of the fate of carbon cycle and contaminant degradation at anoxic-oxic interfaces.

2. EXPERIMENTAL SECTION 2.1. HA Isolation and Characterization. Aldrich HA (AHA, Sigma Aldrich) and Suwannee River Humic Acid (SRHA, 2S101H, International Humic Substances Society (IHSS)) were selected as model HAs in this study. Both HAs isolates were chosen as they are well-characterized and have been extensively used for the study of the redox chemistry of HA in aquatic environments.3,6,10,33,40,42 Stock solutions of each HA isolate were prepared by dissolving 0.5 g of HA solid in 100 mL ultrapure water (resistivity > 18.2 MΩ·cm, Milli-Q, Millipore) adjusted to pH 10.5 using 1 M NaOH in the dark. The solution was then filtered using a 0.45 μm nitrocellulose filter (Millipore). These solutions were referred to as oxidized HAs (AHAox and SRHAox). To mimic the reduced state of HA present at anoxic conditions, aliquots of above oxidized HA solutions were chemically reduced following the procedure of our previous studies.33,40 Briefly, a solution of 80 mL of each of AHAox and SRHAox was reacted with 5% H2 in the presence of a Pd catalyst (0.5% wt on Al2O3 spheres, 1 g/L, Sigma-Aldrich) for 24 h rotation in the dark. The resulting solutions were subsequently purged with ultrapure N2 to remove excess H2 and filtered again using a 0.45 μm nitrocellulose filter to remove Pd catalyst. These solutions were referred to as reduced HAs (AHAred and 6 ACS Paragon Plus Environment

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SRHAred). Stock solutions of reduced HAs were stored in the dark in an anoxic glovebox (95% N2 and 5% H2, < 1 ppm O2, Coy Lab Products Inc., MI) prior to use. The total organic carbon (TOC) in the HA stock solutions was determined by a TOC analyzer (Multi N/C 3100, Analytik Jena, Germany). The specific ultraviolet absorbance (SUVA254), an indicator of HA aromaticity, for both HAs were determined by diluting HAs stock solutions (i.e., 1−10 mg C/L) and measured at 254 nm using a UV-vis spectrophotometer (Cary 60, Agilent). The redox capacity of both oxidized and reduced HA was measured by adding a aliquot of HA stock solution to a 5 mM Fe(III) citrate for 1 h under anoxic conditions as described by Lovley et al.12 followed by determination of formed Fe(II) by a modified 1,10-phenanthroline assay described below. The electron accepting capacity of HA was quantified as the difference between the reducing capacities of oxidized and reduced HA. Additional characteristics (i.e., elemental composition,

13C

NMR, and trace metals) of the HAs are summarized in

Table S1 in the Supporting Information (SI). 2.2. Oxidation of Reduced HA by O2. Duplicate batch experiments were performed to delineate the effect of Fe2+, Ca2+, and Mg2+ on the rate and yield of •OH production from the oxygenation of AHAred and SRHAred in the dark (see detailed experimental conditions in Table S2). All experiments were conducted in magnetically stirred polypropylene reactors (30 mL) at room temperature (22 ± 1 °C) that were shielded with aluminum foil to prevent any photochemical redox reactions of HAs that may generate reactive oxygen species. A pH of 7.0 that is relevant to aquatic 7 ACS Paragon Plus Environment


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environments was set for all experiments. To avoid the introduction of pH buffers that may complicate the reaction systems, the pH was controlled by an autotitrator operated in pH-stat mode (902 Titrando, Metrohm USA Inc.). Variations in solution pH were less than 0.3 units throughout the experiments, as shown in an exemplary pH-time profile (Figure S1). Prior to the oxidation event, all the solution preparations and mixing operations were conducted inside the glovebox. A suite of Fe2+ (0.18−0.89 mM), Ca2+ (5−20 mM), and Mg2+ (5−20 mM) was separately added from stock solutions to a vigorously stirred solution containing 50 mg C/L AHAred or SRHAred and 10 mM sodium benzoate (99.5%), an •OH probe, for 30 min to ensure well-mixing. The HA concentration (i.e., 50 mg C/L) and the concentration ranges of Fe2+, Ca2+, and Mg2+ were

selected

to

mimic

the

conditions

found

in

anoxic-oxic

interface

environments.3,4,31,33,36,40,41,43‒48 The dissolved oxygen (DO) concentration in the wellmixed solutions was below the detection limit (0.01 mg/L) of the DO probe (HACH, USA). No detectable oxidation of Fe(II) was observed in any sample during the time scale of 30 min under anoxic conditions (data not shown). After the short anoxic mixing, the suspension was immediately transferred out from the glovebox and aerated to represent a transition to oxic conditions. Oxidation experiments were initiated by uncapping the reactors and exposing the samples to air (PO2 = 0.21 bar) with magnetic stirring in the dark. For all experiments, the DO concentrations reached a maximum (~ 8.2 mg/L) within 30 min of exposure to the air as measured by the DO probe. Controls containing reduced HA and/or oxidized HA 8 ACS Paragon Plus Environment

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alone (without divalent cations addition) and oxidized HA in the presence of Fe2+ (0−0.89 mM) were set up in the same manner as that used for the scenarios described above. Additional experiments were performed to study the coexistence of Fe2+ (0.18 mM) and Ca2+/ Mg2+ (5−20 mM) and the effect of added H2O2 (10 μM) and catalase (20 mg/L, Sigma-Aldrich) on •OH production. Additional experiments were also performed in the absence of sodium benzoate in order to quantify the evolution of particle size and divalent cation concentrations. All other experimental procedures were identical to those used for the initial experiments as described above. For all experiments, samples were periodically collected and a portion of them were filtered immediately through a 0.22 μm syringe filter (PES, Millipore) for analysis of •OH, dissolved Ca, and dissolved Mg. The remainder of the unfiltered subsamples were used for total Fe(II) and particle size analysis. 2.3. Analytical Techniques. Production of •OH was quantified using a benzoic acid probe technique as described previously.36,49 Briefly, benzoic acid reacted with •OH to from a stable product of p-hydroxybenzoic acid (p-HBA). The concentration of p-HBA in the filtrate was measured using an Agilent HPLC equipped with a UV-vis detector and a Cosmosil C18-PAQ column (4.6 × 250 mm). The mobile phase was a mixture of acetonitrile and 1% TFA/water (65:35, v/v) at 1 mL/min and the detection wavelength was 255 nm. The adsorption of p-HBA on PES filter membrane was negligible. The instrument detection limit for p-HBA was 0.1 μM. The cumulative •OH concentration was calculated from the 5.87 times of the concentration of p-HBA.36 9 ACS Paragon Plus Environment


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Total Fe(II) concentration was measured by a modified 1,10-phenanthroline method at a wavelength of 510 nm using a UV-vis spectrophotometer.50 Samples were digested by 2 M HCl for 3 h in the dark. All manipulations were conducted in anoxic conditions to prevent any potential oxidation of Fe(II) by O2 in the presence of high Clconcentration.51 An predetermined amount of fluoride was added initially to the sample in an effort to effectively inhibit Fe(III) reduction by HA during acidification and to prevent the formation of colored Fe(III)-phenanthroline complexes (ε510 nm = 182 M−1 cm−1). The particle size distribution of HAs in the absence and presence of divalent cations was monitored through dynamic light scattering (DLS) using a Zetasizer Nano instrument (Malvern). All measurements were performed in 10 consecutive scans for each duplicate sample. The accuracy of the instrument was validated by using polystyrene latex standard particles (DTS1060, Malvern) with an average diameter of 63 ± 3 nm. A double-spherical aberration-corrected scanning transmission electron microscope (Cs-STEM, Thermo Fisher, Titan Themis G2 60-300) was launched for selected samples that were collected at the end of the oxidation experiments to probe the associations of divalent cations with HAs at nano scales. In addition, a control sample containing HA-Fe(II) complexes formed under anoxic conditions was also imaged for comparison. The instrument was operated at 300 kV and equipped with a high brightness X-FEG Schottky field emission gun and four Super-X silicon drift X-ray 10 ACS Paragon Plus Environment

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energy-dispersive spectroscopy (EDS) detectors. Imaging was conducted in STEM mode using the high angle annular dark field (HAADF) detector. STEM samples were prepared by pipetting ~ 20 μL of suspension onto a 20 nm thick window of Si3N4 membrane (SN100, SiMPore Inc, USA) following by immediate evaporation of the remaining water at room temperature under vacuum. 2.4. Data Analysis. We used a simple pseudo-first-order kinetics model (eq 1) to quantify the rates of •OH production. The assumption that the •OH production follows a pseudo-first-order rate law is in part supported by the substantially lower amounts of reduced moieties present in HA compared to O2.9,52 More evidence supporting this assumption is provided in the section S1 of SI. The pseudo-first-order kinetic rate constant (k) of •OH production was calculated by fitting the experimental data using a nonlinear least-squares algorithm. 𝐶𝑡 = 𝐶𝑒𝑞(1 ― 𝑒 ―𝑘𝑡)

(1)

where Ct and Ceq are the •OH concentration at time t and equilibrium, respectively.

3. RESULTS AND DISCUSSION 3.1. Validation of Dark •OH Production upon Oxygenation of Reduced HAs. Significant production of •OH was observed for the oxygenation of chemically reduced AHA and SRHA in the dark (Figure 1a). In contrast, no •OH production was observed under anoxic conditions (data not shown). These observations suggested that the transport of reduced HA from anoxic zones across interfaces to more oxic conditions 11 ACS Paragon Plus Environment


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could trigger the production of •OH, supporting the notion that the dark oxygenation of reduced HA represented a significant source of •OH at anoxic-oxic interfaces.3 For both AHAred and SRHAred, the cumulative concentrations of •OH increased rapidly within 3 h and then approached a steady-state for the remainder of the experiments (24 h total, Figure 1a), indicative of the fast kinetics of •OH formation (Table S2). Prior studies have reported that the oxidation of reduced HA by O2 occurred very quickly, with kinetics on the order of seconds to minutes.53 The rapid increase in •OH production during the first 3 h of O2 exposure likely suggested that some reduced moieties in HA persisted in solution and could further donate electrons to O2. This was supported by the findings from Bauer and Kappler that the exposure of reduced HA to O2 within short time frames (e.g., minutes) did not oxidize HA completely.53 In contrast to the reduced HAs, no detectable •OH production occurred from the oxygenation of AHAox and SRHAox (Figure 1a). This was due primarily to the fact that the oxidized HA had a low electron transfer capacity to donate electrons to produce •OH.14,52 While the oxidized or untreated HAs can be slowly reduced during storage in the glovebox and may contribute to dark •OH production,3 this was not occurred in this study, likely due to the short storage time (< 24 h) of AHAox and SRHAox in the glovebox prior to O2 exposure. Oxygenation of AHAred produced a greater amount of •OH than that of SRHAred (Figure 1a, Table S2). Differences in •OH production may be attributed to differences in redox-active components between HA of different origins and compositions. 12 ACS Paragon Plus Environment

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Measurements of the redox capacity of HA indicated that the electron accepting capacity (EAC) of AHAred was higher compared to that of SRHAred (3220 ± 158 versus 2465 ± 151 μeq/gC), which was similar to the values reported by Jiang and Kappler.14 Previous electrochemical measurements also showed a higher EAC of AHA relative to SRHA (923 ± 60 versus 671 ± 8 μmole-/gHA, Table S1).6,10 Further, the SUVA254 value of AHA was higher than that of SRHA (9.5 ± 0.4 versus 6.3 ± 0.3 L/mg·m, Table S1), which is similar to previous finding that the EAC of the humic substances increased with increasing SUVA254.9,54 These observations suggested that the amount of redoxactive functional groups in AHAred was higher than that in SRHAred. Thus, the substantially higher concentration of the redox-active functional groups in AHAred was likely responsible for the higher yield of •OH from AHAred. An increasing body of evidence has substantiated that quinones are the main redox-active functional moieties in HA.6,7,14,52,54 However, the non-quinone structure can also be accounted for 21–66% of the total electron-carrying capacity in Pd-H2 treated AHAred and SRHAred.42 Evidence from electrochemical and spectroscopic studies suggested that the reduced forms of nitrogen- and sulfur-containing moieties in HAs were redox-active.55 In addition to quinones, the noticeably higher amount of organically bound sulfur in AHA relative to SRHA (3.5 versus 0.54%, Table S1) might also be responsible for the higher yield of •OH formation for AHAred. Although the presence of redox-active trace metals such as Fe in HA can enhance the production of •OH, the ability of Fe to exert an influence on •OH production was 13 ACS Paragon Plus Environment


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limited in this study. First, the production of •OH from AHAred and SRHAred did not stoichiometrically correlate with their Fe contents. For example, the Fe amount in AHA was greater than 10-fold that of in SRHA (300.1 ± 3.6 versus 22.9 ± 0.5 μmol/gHA, Table S1), while the •OH concentration produced from AHAred was double that of produced from SRHAred (Figure 1a). Second, the steady-state concentrations of •OH increased linearly with increasing AHAred and SRHAred concentrations (R2 > 0.99, Figure 1b), suggesting that the formation of •OH was strongly correlated with the concentration of the redox-active functional groups and not caused by the intrinsic Fe. The negligible contribution of Fe to •OH formation was further supported by the inductively coupled plasma-mass spectrometry (ICP-MS) measurements, which showed very low concentrations of total Fe (< 5 μg/L) in working solutions (i.e., 50 mg C/L HAs in total) throughout the experiments. Such low concentrations of Fe have been found to have a low contribution (e.g., < 7%) to the measured electron transfer capacities of HA.8,10 H2O2 has been hypothesized to be a key reaction intermediate in the formation of •OH from oxygenation of reduced HA.3 To further explore this, we tested the formation of •OH in the presence of externally H2O2 and catalase (that can rapidly decompose H2O2). Results showed that the addition of 10 μM H2O2 significantly increased production of •OH relative to H2O2-free system and that the presence of catalase completely suppressed •OH formation (Figure 1c,d). These observations strongly support the conclusion of the intermediacy of H2O2 in the formation of •OH (eq 2). This 14 ACS Paragon Plus Environment

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can be further confirmed by a recent study, which provided qualitative evidence that the micromole concentration of H2O2 was generated from oxygenation of the reduced Leonardite HA.52 As expected from the literature, oxygenation of the reducible quinone moieties (e.g., semiquinones) produced H2O2, which can subsequently be converted to •OH through a metal-independent organic Fenton-like reaction.3 𝐻𝐴𝑟𝑒𝑑 + 𝑂2 → 𝐻𝐴𝑜𝑥 + 𝐻2𝑂2

(2)

3.2. Fe(II) Significantly Enhanced the Production of •OH. The addition of Fe2+ to both AHAred and SRHAred significantly increased the rates and yields of •OH production, with the greatest •OH generated at higher Fe2+ concentrations (Figure 2a,b, Table S2). This finding supports the recent field evidence that the dark formation of •OH from oxygenation of soil and lake water increased with increasing Fe(II) concentrations.18,56 Control experiments showed no detectable •OH production from oxygenation of free Fe2+ (0.18–0.89 mM) solutions (not shown). In addition to reduced HA, addition of Fe2+ to both AHAox and SRHAox also increased the yield of •OH production, although the rates and extent of •OH production from reduced HA were noticeable higher compared to oxidized HA (Figure 2c,d, Table S2). The cumulative •OH concentrations were linearly dependent on the initial Fe2+ concentrations (R2 > 0.98, Figure S2), emphasizing the importance of Fe2+ for •OH production from reduced HAs. The enhanced production of •OH may be related to (1) the formation of HA-Fe(II) complexes that affected the Fe cycling and redox of HA and (2) the enhanced efficiency of Fenton reactions under circumneutral conditions.22,24,28,52 15 ACS Paragon Plus Environment


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It has been found that the complexation with HA can alter the oxidation rate of Fe2+, a key player in the •OH formation based on the Haber–Weiss mechanism.22,24,28‒30,57 The formation of HA-Fe(II/III) complexes was supported by the ultrahigh resolution HAADF-STEM-EDS images (Figure 3), which showed that C and Fe had similar elemental distributions at nano scales. The lack of detectable truly soluble Fe (< 1–3 nm, filtered by10 000 Da cellulose ultrafiltration membranes (EMD, Millipore)) measured at the end of each experiment further confirms that all Fe was complexed with colloidal HAs (not shown). Previous studies demonstrated that Fe(II) complexed with the oxidized HAs can be oxidized faster than free Fe2+ ions over a broad range of pH.28,58 However, our results showed that Fe(II) complexed with the reduced HAs substantially lowered the pseudo first-order rates of Fe(II) oxidation (see detailed derivation of pseudo first-order equation in section S1 of SI) (Figure 4). For all Fe2+ concentrations tested (0.18–0.89 mM), total Fe(II) was completely oxidized within 1 h in the absence of HA (Figure 4). In contrast, Fe(II) could still be detected after 3 h in the presence of both reduced HAs (Figure 4). This was similar to the recent work of Daugherty et al.52 who reported > 50% of initial Fe(II) that was complexed with the reduced HA remained reduced after 4 h of O2 exposure. In addition to reduced HAs, oxidized HAs also retarded the oxidation of Fe(II) to a lesser extent (Figure 4). This result, which is opposite to prior observations,28,58,59 may be due to the fact that the oxidized HA contained certain electron-accepting moieties.3 Inspection of the rates of •OH formation versus Fe(II) oxidation showed that they 16 ACS Paragon Plus Environment

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exhibited a opposite relationship; the lower net rates of Fe(II) oxidation led to higher •OH production (Figures 2 and 4, Table S2). A plausible explanation for this observation was likely due to the fact that HA-Fe(II) complexes maintained the relatively high concentrations of Fe(II) under oxic conditions,52,60 promoting the generation of •OH through equations 3 and 4. 𝐻𝐴 ― 𝐹𝑒(𝐼𝐼) + 𝑂2 → 𝐻𝐴 ― 𝐹𝑒(𝐼𝐼𝐼) + 𝑂2―

(3)

𝐻𝐴 ― 𝐹𝑒(𝐼𝐼) + 𝐻2𝑂2 → 𝐻𝐴 ― 𝐹𝑒(𝐼𝐼𝐼) + 𝑂𝐻 ― + •𝑂𝐻

(4)

The cycle of Fe(II)/Fe(III) may be described by a net Fe(II) oxygenation kinetics, considering forward Fe(II) oxidation and feedback from Fe(III) reduction (eq 5).61 ―

(

𝑑[𝐹𝑒(𝐼𝐼)] 𝑑𝑡 =

)

=

𝑛𝑒𝑡

∑𝑘

∑𝑘

𝑜𝑥[𝐹𝑒(𝐼𝐼)

𝑛𝑒𝑡[𝐹𝑒(𝐼𝐼)

― 𝐿][𝑂2]

― 𝐿][𝑜𝑥𝑖𝑑𝑎𝑛𝑡𝑠] ―

∑𝑘

𝑟𝑒𝑑[𝐹𝑒(𝐼𝐼𝐼)

― 𝐿][𝑟𝑒𝑑𝑢𝑐𝑡𝑎𝑛𝑡𝑠]

(5)

where knet is the generic net rate constant of the oxygenation of Fe(II)-Lred species (L represents HA), kox and kred are the generic rate constants of the summed Fe(II)-Lred oxidation and Fe(III)-Lox reduction, respectively. The decrease in the net oxidation of Fe(II) by reduced HAs over the course of the first 3 h suggested that the back Fe(III) reduction provided a source of Fe(II) (eq 6). This explanation was supported by the results of Daugherty et al.52, who observed the reduction of HA-Fe(III) complexes by the reduced HA. Further evidence of this phenomenon was observed by Bauer and Kappler,53 who reported the reduction of Fe(III) oxides by the reduced HA, even under oxic conditions. Reduced quinone-like moieties that intrinsically presented in HAs, have been suggested to be responsible for 17 ACS Paragon Plus Environment


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the reduction of HA-Fe(III) complexes.52,62,63 Other reductants such as superoxide (O2•-) may also contribute to Fe(III) reduction (eq 7).52,64 𝐻𝐴𝑟𝑒𝑑 +𝐹𝑒(𝐼𝐼𝐼) ―𝐻𝐴 → 𝐻𝐴𝑜𝑥 +𝐹𝑒(𝐼𝐼) ―𝐻𝐴 𝐹𝑒(𝐼𝐼𝐼) ―𝐻𝐴 + 𝑂2― → 𝐹𝑒(𝐼𝐼) ―𝐻𝐴 + 𝑂2

(6) (7)

In addition to increasing the number of Fe(II)/Fe(III) cycles, HA-Fe(II) complex may be expected to expedite the efficiency of molecular oxygen activation, producing more •OH production through a sequence of electron transfer reactions (e.g., O2 → O2•→ H2O2 →•OH), based on the findings of Wang et al.65 Furthermore, the complexed Fe(II) may catalyze the decomposition of H2O2 produced from oxygenation of (uncomplexed) reduced HA (eq 2), which shifts the pathway of •OH formation from metal-independent organic Fenton-like reaction to Fe(II)-catalyzed Fenton reaction, accelerating •OH production.57 The validity of this mechanism is supported by the findings of Daugherty et al.52, who observed the fast suppression of H2O2 production in the presence of complexed Fe(II). The higher extent of the increase in •OH production from oxygenation of AHA compared to SRHA due to the presence of Fe2+ may be attributable to (1) the lower net oxidation rate of Fe(II) complexed AHA relative to Fe(II) complexed SRHA and (2) the difference in complexation capacity of Fe(II) with the (redox-active) functional groups of AHA and SRHA. 3.3. Ca2+ and Mg2+ Retarded the Production of •OH. In contrast to Fe2+, Ca2+ and Mg2+ lowered the rates and yields of •OH production from oxygenation of both AHAred and SRHAred in a concentration-dependent manner (Figure 5). For both AHAred 18 ACS Paragon Plus Environment

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and SRHAred, the •OH production rates decreased dramatically as the concentrations of Ca2+ increased from 0 to 5 mM, and then progressively approached a low value with a further increase in Ca2+ concentration to 20 mM (Figure 5a,b). The inhibitory effect of Mg2+ followed the similar qualitative trends as that of Ca2+, but the extent of the decrease in •OH production was significantly lower for Mg2+ than for Ca2+ (Figure 5c,d). For example, the presence of 5 mM Mg2+ only slightly (6–14%) decreased the rates of •OH production, and the •OH production rates in the presence of 20 mM Mg2+ was similar or slightly higher than the rates observed in the presence of 5 mM Ca2+ (Figure 5, Table S2). The decrease in the production of •OH caused by Ca2+/Mg2+ was likely due to multiple factors such as (1) blocking of the active site of reduced HA due to complexation and aggregation and (2) suppression of the redox reactivity of reduced HA.32,35,39,66,67 Ca2+ and Mg2+ can complex with the carboxyl groups of HAs, which subsequently results in aggregation of HAs.32,33,41 HAADF-STEM-EDS elemental mappings showed a similar distribution of C and Ca/Mg across the imaged regions (Figure 6), suggesting the association of HA with Ca2+ and Mg2+. DLS results showed that the Z-averaged hydrodynamic diameters of both reduced HA suspensions without Ca2+/Mg2+ were in the range of 110–250 nm throughout the experiments, while they were much higher (300–3000 nm) in the presence of 5–20 mM Ca2+/Mg2+ (Figure S3a-d). HADDF-STEM images supported DLS observations, showing that the presence of Ca2+ and Mg2+ increased the particle/aggregate size of HA (Figure S4). The rapid formation of larger 19 ACS Paragon Plus Environment


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HA aggregates may result in the blocking of the surface sites that may suppress the electron transfer available for •OH production. This hypothesis was possibly supported by the negative correlations between the steady-state •OH concentrations and the AHA aggregate sizes (Figure S5a,b). The higher extent of the inhibition of •OH production for Ca2+ compared to Mg2+ may be attributed to the stronger complexation and bridging effect of Ca2+ relative to Mg2+.33,41,68 DLS results supported this proposition, showing that Ca2+ had a more pronounced effects on aggregation of HA than Mg2+ (Figure S3ad). In addition, both Ca2+ and Mg2+ can increase the ionic strength of solution that neutralizes the negative surface charges of HAs, which could lower the inter- and intramolecular electrostatic repulsions, leading to the molecular shrinking of HAs.35 This molecule compression has been found to decrease the electron-transfer capacity of HA.35 For example, Lu et al. reported that the EAC of DOM extracted from sewage sludge decreased from ~ 600 to ~ 500 μmole/g C when the KCl concentrations were raised from 100 to 500 mM.35 Because divalent ions can increase the ionic strength more effectively than monovalent ions such as K+,32 Ca2+ and Mg2+ present at high concentrations may also be expected to lower the EAC of reduced HA and, consequently, •OH production. Furthermore, recent studies substantiated that the EAC of particulate HA was much lower than that of dissolved and colloidal HA,9,67 implying that aggregation caused by Ca2+ and Mg2+ may also lower the EAC of HA. 3.4. Effect of Coexistence of Fe2+ and Ca2+/Mg2+ on •OH Production. The 20 ACS Paragon Plus Environment

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mixtures of 0.18 mM Fe2+ and 5–20 mM Ca2+/Mg2+ produced •OH at rates 2–40% higher than the rates of •OH production from the reduced HA alone (Figure 7, Table S2). This suggested that, when in competition, the impact of Fe2+ dominated that of Ca2+/Mg2+. However, for a given Ca2+/Mg2+ concentration and a reduced HA, a mixture of Fe2+ and Ca2+/Mg2+ produced •OH at a rate 42–64% lower than the rate from the Fe2+ alone, and this effect was more pronounced in the case of Ca2+ (Figure 7, Table S2). This observation is important in anoxic-oxic interfaces (e.g., surface water-groundwater interaction zone, groundwater recharge, riverbank filtration, and flood events), where the presence of Ca2+/Mg2+ could result in a retarded production of •OH from oxygenation of reduced HAs and Fe2+. Possible explanations for the inhibitory effect of Ca2+ and Mg2+ on •OH production from reduced HAs and Fe2+ may be related to the role of Ca2+/Mg2+ as a competitor for the complexation of HA-Fe(II) and as a facilitator for the aggregation of HA-Fe complexes.31,33,37,38, 66 It has been proposed that Ca2+ and Mg2+ can compete with Fe2+ for the same binding sites on HA and thus decrease the affinity of HA for Fe(II).31,37,38 Our HAADF-STEMEDS images showed spatial associations among C, Fe, and Ca/Mg atoms on a nanoscale (Figure 8), supporting the possible formation of ternary HA-Fe-Ca/Mg complexation. This finding relates well to the work of Sowers et al.69 who reported the formation of Fe-Ca-NOM association in systems containing NOM, ferrihydrite, and Ca likely through Ca2+-bridging of carboxylic C to ferrihydrite. Additionally, the inspection of several comparable EDS images (data not shown) revealed that the atomic 21 ACS Paragon Plus Environment


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ratio of C/Ca or C/Mg in HA-Fe-Ca/Mg systems (i.e., C/Ca = 2.9–3.7; C/Mg = 8.9– 11.6) was higher than that in HA-Ca/Mg systems (i.e., C/Ca = 0.8–1.1; C/Mg = 2.3– 4.5), suggesting that the association of Fe was probably stronger than Ca/Mg to C. These observations are somewhat consistent with previous studies that the apparent formation constant of NOM-Fe(II) complexes (log K = 5.77) at pH 5 was much higher compared to NOM-Ca2+ complexes (log K = 2.92) and NOM-Mg2+ complexes (log K = 2.09).70 Further, in systems containing both Fe2+ and Ca2+/Mg2+, Ca2+ and Mg2+ increased the oxygenation rates of Fe(II), with most oxidation occurring at high Ca2+/Mg2+ concentrations (Figure 9). This was likely due to the decrease in (1) the formation of oxidation-resistant HA-Fe(II) complexes and (2) the effective HA-Fe(III) concentrations available for Fe(III) reduction.31,71 This result also supports a recent assertion that the presence of other divalent cations in HA-Fe(II) system may retard the Fe complexation and redox buffer capabilities of reduced HA.52 HAs were more susceptible to aggregation in the presence of both Fe2+ and Ca2+/Mg2+ versus in the presence of Ca2+/Mg2+ alone, as the former showed appreciably higher hydrodynamic diameters than the later across all Ca2+/Mg2+ concentration ranges (Figure S3). In contrast, in the absence of Ca2+/Mg2+, the HA-Fe complexes across all Fe2+ concentrations were stable against aggregation, as their hydrodynamic diameter varied in the range of 50–250 nm (Figure S6). The accelerated aggregation was further ascertained by HAADF-STEM images, which showed that, Ca2+/Mg2+ (particularly Ca2+) may bridge the HA-Fe primary particles and aggregates together (Figure S4). 22 ACS Paragon Plus Environment

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Atomic resolution images (the insert in Figure S4e,f) showed differences in morphology between HA-Fe-Ca complexes and HA-Fe-Mg complexes. As discussed earlier, blocking of the HA surface sites due to aggregation may lower the •OH production. In HA-Fe-Ca/Mg systems, the extensive aggregation of HA-Fe by Ca2+/Mg2+ may also result in the surface site blockage, decreasing the number of HA reactive sites occupied by Fe(II), and thus making less complexed Fe(II) available for the enhanced •OH production. The observations that the steady-state •OH concentrations were correlated with the HA aggregate sizes (Figure S5c,d) probably support the conclusion that the blockage of HA surface sites plays a negative role in the •OH production.

4. CONCLUSIONS This study, to our knowledge, is the first report discerning the influence of divalent cations, including Fe2+, Ca2+ and Mg2+, on the dark production of •OH from oxygenation of reduced HAs at anoxic-oxic interfaces. Through a combination of batch experiments and a powerful array of characterization techniques that include HAADFSTEM-EDS, and DLS, we have provided multiple lines of evidence that the rates and yields of •OH production from reduced HAs were substantially influenced by the presence of Fe2+, Ca2+ and Mg2+. Specifically, Fe2+ dramatically enhanced the production of •OH, due primarily to the formation of complexed Fe(II) which increased the number of Fe(II)/Fe(III) cycles and accelerated the decomposition of H2O2. In 23 ACS Paragon Plus Environment


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contrast, Ca2+ and Mg2+ greatly retarded the production of •OH, likely due to the blocking of the reactive sites of HAs and suppression of the EAC of HAs. Coexistence of Fe2+ and Ca2+/Mg2+ produced •OH at rates higher than the rates from reduced HA alone, suggesting that, when in competition, the effect of Fe2+ dominated that of Ca2+ and Mg2+. However, the promotional effect of Fe2+ on •OH production from the reduced HAs was greatly suppressed by the coexistence of high Ca2+ and Mg2+. Taken together, our findings extended the knowledge of the production of •OH from oxygenation of reduced HAs in simple solutions to more complex conditions, and provided critical insight into the understanding of the cycling of C and nutrient and the fate of contaminants, microbes, and viruses at anoxic-oxic interfaces. We realize that the natural environments are highly complicated than systems presented in this study, and ongoing research is being conducted to extend the framework developed here by evaluating the role of numerous variables (e.g., pH, DO dynamics, NOM compositions, and microorganisms) for •OH production with real world waters.

Supporting Information Additional information regarding HA characteristics (Table S1), summary of the batch experiments for •OH production (Table S2), kinetics model for Fe(II) oxidation (section SI), representative pH-time profile (Figure S1), correlations between the •OH concentrations and the Fe2+ concentrations for HAs (Figure S2), effects of cations on particle size of HAs (Figure S3), representative high-resolution HAADF-STEM images 24 ACS Paragon Plus Environment

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(Figure S4), correlations between the •OH concentrations and the particle size of HA (Figure S5), and the effect of Fe(II) on the particle size of HA (Figure S6) (PDF).

ACKNOWLEDGMENTS This work was supported by the China Postdoctoral Science Foundation (2017M610530), the Natural Science Foundation of China (No. 41703128), and the Basic Research Project of Shenzhen (JCYJ20170307110055182). HAADF-STEMEDS measurements were performed in the Pico Center at Southern University of Science and Technology, which support from the Presidential Fund and Development and Reform Commission of Shenzhen Municipality. Peng Liao thank the staff of the Environmental Biogeochemistry Laboratory (EBGL) at Southern University of Science and Technology for their assistance with DLS. Peng Liao also acknowledges financial support from the Program for Guangdong Introducing Innovative and Entrepreneurial Teams (2017ZT07Z479). We appreciate the comments of anonymous reviewers that helped us improve the presentation and interpretation of our study.

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Figure 1. (a) Production of •OH upon oxygenation of 50 mg C/L AHAred and SRHAred in the dark at pH 7.0. Please note that the dark production of •OH from oxygenation of 50 mg C/L AHAox and SRHAox was negligible. (b) Correlations between the steadystate •OH concentrations and the initial reduced HA concentrations (0−100 mg C/L). (c,d) Effect of externally H2O2 (10 μM) and catalase (20 mg/L) on the dark production of •OH from oxygenation of (c) AHAred and (d) SRHAred. Solid lines represent the pseudo-first-order model fits using equation 1. Error bars represent the standard deviations of at least duplicate measurements. Where error bars are not visible, they are smaller than the data symbols.



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Figure 2. (a-d) Effect of Fe2+ concentrations (0.18−0.89 mM) on the dark production of •OH from oxygenation of 50 mg C/L (a) AHAred, (b) SRHAred, (c) AHAox, and (d) SRHAox at pH 7.0. Solid lines in panels a-d represent the pseudo-first-order model fits using equation 1. Error bars represent the standard deviation of at least duplicate measurements. Where error bars are not visible, they are smaller than the data symbols.


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Figure 3. (a,c) Representative high-resolution HAADF-STEM images and the corresponding (b,d) EDS elemental mapping of (a,b) AHA-Fe(II) complexes and (c,d) AHA-Fe(III) complexes. The AHA-Fe(III) suspensions were collected at the termination of the 24-h oxic experiments, with initial AHAred and Fe(II) concentration of 50 mg C/L and 0.89 mM, respectively. The AHA-Fe(II) suspensions (control experiment) were captured by mixing 50 mg C/L AHAred with 0.89 mM Fe(II) under anoxic conditions for 2 h. All images were taken at a comparable magnification of scale bar of 200 nm.



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Figure 4. Oxygenation kinetics of added Fe2+ in (a-c) AHA-Fe and (d-f) SRHA-Fe systems. Because the total Fe(II) concentrations were undetectable after 6 h of reaction, we only report the results for the time scale of the first 6 h. Fitted curves were derived from first-order equation (Ct = C0•e-k’t, where C0 and Ct are the concentrations of Fe(II) at time t = 0 and t = t, respectively, and k’ is a pseudo first-order rate constant of Fe(II) oxygenation). Error bars represent the standard deviations of at least duplicate measurements. Where error bars are not visible, they are smaller than the data symbols.


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Figure 5. Effects of (a,b) Ca2+ and (c,d) Mg2+ concentrations on the dark production of •OH from oxygenation of 50 mg C/L (a,c) AHAred and (b,d) SRHAred at pH 7.0. Solid lines represent the pseudo-first-order model fits using equation 1. Error bars represent the standard deviations of at least duplicate measurements. Where error bars are not visible, they are smaller than the data symbols.



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Figure 6. (a,c) Representative high-resolution HAADF-STEM images and the corresponding (b,d) EDS elemental mapping of (a,b) AHA-Ca complexes and (c,d) AHA-Mg complexes collected at the termination of the 24-h oxic experiments. The initial concentrations of AHAred and Ca2+/Mg2+ were 50 mg C/L and 10 mM, respectively. All images were taken at a comparable magnification of scale bar of 200 nm.


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Figure 7. Effects of the (a,b) mixtures of Fe2+ and Ca2+ and (c,d) mixtures of Fe2+ and Mg2+ on the dark production of •OH from oxygenation of 50 mg C/L (a,c) AHAred and (b,d) SRHAred at pH 7.0. Solid lines represent the pseudo-first-order model fits using equation 1. Error bars represent the standard deviations of at least duplicate measurements. Where error bars are not visible, they are smaller than the data symbols.



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Figure 8. (a,c) Representative high-resolution HAADF-STEM images and the corresponding (b,d) EDS elemental mapping of (a,b) AHA-Fe-Ca complexes and (c,d) AHA-Fe-Mg complexes collected at the termination of the 24-h oxic experiments. The initial concentrations of AHA, Fe2+, and Ca2+/Mg2+ were 50 mg C/L, 0.89 mM, and 10 mM, respectively.


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Figure 9. Oxygenation kinetics of added Fe2+ (0.18 mM) in the systems of (a) AHAFe(II)-Ca2+, (b) SRHA-Fe(II)-Ca2+, (c) AHA-Fe(II)-Mg2+, and (d) SRHA-Fe(II)-Mg2+. Because the total Fe(II) concentrations were undetectable after 6 h of reaction, we only report the results for the time scale of the first 6 h. Fitted curves were derived from firstorder equation (Ct = C0•e-k’t, where C0 and Ct are the concentrations of Fe(II) at time t = 0 and t = t, respectively, and k’ is a pseudo first-order rate constant of Fe(II) oxygenation). Error bars represent the standard deviations of at least duplicate measurements. Where error bars are not visible, they are smaller than the data symbols.



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TOC


Impact of Divalent Cations on Dark Production of Hydroxyl Radicals

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