Effects of Second Metal Oxides on Surface-Mediated Reduction of

Feb 28, 2019 - We believe that the conduction bands of goethite and TiO2 were used as conduits for electron transfer from Fe(II) through TiO2 to goeth...
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Effects of Second Metal Oxides on Surface-Mediated Reduction of Contaminants by Fe(II) with Iron Oxide Jianzhi Huang,† Yifan Dai,‡ Chung-Chiun Liu,‡ and Huichun Zhang*,† †

Department of Civil Engineering and ‡Department of Chemical Engineering, Case Western Reserve University, Cleveland, Ohio 44106, United States

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S Supporting Information *

ABSTRACT: This work examined the effects of two second metal oxides (SiO2 and TiO2) on the reductive reactivity of Fe(II)/goethite, an important natural reductant. SiO2 significantly inhibited the reductive reactivity, as quantified by the reduction kinetics of p-cyanonitrobenzene (pCNB) as a probe compound, while TiO2 greatly enhanced the reactivity. Silicate showed comparable inhibitory effects as SiO2 particles, indicating that the inhibition effect of SiO2 was dominated by its dissolution. Pseudo-first-order rate constants (k) of Fe(II)/goethite + TiO2 mixtures were higher than the sum of the k values of the respective single oxide. For the mixtures of Fe(II)/goethite + TiO2, the k values followed rutile > TiO2−P25 > anatase, despite a different trend in the adsorbed Fe(II) amount. This reactivity trend agreed well with their conduction band potentials. Higher loadings of TiO2 also led to higher reactivity. When TiO2 was physically separated from goethite by confining it in a dialysis bag, k of Fe(II)/goethite + TiO2 was comparable to the sum of the k values of the respective single oxide. We believe that the conduction bands of goethite and TiO2 were used as conduits for electron transfer from Fe(II) through TiO2 to goethite and eventually to reduce pCNB. This type of interparticle electron transfer was for the first time discovered for dark conditions, which might play a previously unrecognized yet important role in contaminant reduction and Fe(II)/Fe(III) redox cycling in the environment. KEYWORDS: Fe(II)−iron oxides, surface-mediated reduction, interparticle electron transfer (IPET), second metal oxides, adsorption, conduction band ical properties of iron oxides,3,11−13 solution pH,3,9,10,14,15 Fe(II) surface speciation,5,7,16−18 metal ions,19,20 and natural organic matter (NOM),21,22 have been found to affect the reduction rates of various contaminants. For example, the Fe(II) monohydroxo surface complex (FeIIIOFeIIOH0) was believed to be an effective reductant toward U(VI) and nitroaromatic compounds,7,16 and the reduction rate was proportional to its concentration. Almost all of the previous studies on the reductive reactivity of surface-adsorbed Fe(II) in model systems only investigated

1. INTRODUCTION Electron transfer reactions are important in affecting the transformation and remediation of contaminants in the environment. They are generally slow in homogeneous solutions but can be significantly enhanced in the presence of minerals.1 For example, pyrite can promote the rapid reaction between thiosulfate and molecular oxygen.2 In addition, Fe(II) is a weak electron donor in anoxic environments; however, surface-adsorbed Fe(II) on iron oxides has been established to be a much better electron donor. The redox reaction of Fe(II)/Fe(III) in the environment plays an important role in controlling biogeochemical processes. Previous studies have shown that surface-complexed Fe(II) on iron oxides is capable of reducing various pollutants in natural and engineered anoxic environments, such as organic compounds,3−5 nitrite,6 and uranium(VI).7,8 Many factors, including the dissolved Fe(II) concentration,9,10 physiochem© XXXX American Chemical Society

Special Issue: Iron Redox Chemistry and Its Environmental Impact Received: Revised: Accepted: Published: A

December 20, 2018 February 7, 2019 February 28, 2019 February 28, 2019 DOI: 10.1021/acsearthspacechem.8b00210 ACS Earth Space Chem. XXXX, XXX, XXX−XXX

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ACS Earth and Space Chemistry

acetonitrile, FeSO4·7H2O, ferrozine, HCl (TraceMetal-grade concentrated), 2-(N-morpholino)ethanesulfonic acid (MES buffer), and NaOH, were obtained from Fisher Scientific or Sigma-Aldrich. The stock solution of pCNB was prepared in ultrapure (UP) water containing 1% acetonitrile. Goethite (Bayferrox 910) was purchased from Lanxess. Anatase and rutile were purchased from Nanostructured and Amorphous Material, Inc. TiO2−P25 was purchased from Nippon Aerosil Co., Ltd, and SiO2 was purchased from Evonik Industries. The properties of different metal oxides are listed in Table 1.

single well-defined iron oxides. However, in natural aquatic environments, bulk and surface properties of mixed solids can behave substantially differently from the respective single solid.23 Specifically, iron oxides in natural environments commonly exist in mixtures with other metal oxides, such as SiO223,24 and TiO2,25−27 in which the interactions can largely affect trace element adsorption,28 redox reaction,29−31 and photocatalytic reaction.32 For example, it has been shown that the adsorption of Cd(II) and Ca(II) on SiO2 was enhanced by Al(OH)3 as a result of the increase in the surface site concentration and change in the surface potential.33 It has also been reported that TiO2 inhibited the oxidative reactivity of MnO2 toward triclosan but only when a small amount of triclosan was present.31 In addition, nanosized metal oxides, such as TiO2 and SiO2, could be released into natural and engineered environments,34 raising concerns about their fate and impact in the environment. There are numerous studies about the transport and fate of nanomaterials, such as aggregation35,36 and interactions with other nanoparticles.37 However, very limited work has quantitatively examined the effects of second metal oxides on the reductive reactivity of surface-adsorbed Fe(II). Further, it remains unknown if the addition/presence of environmentally benign and abundant oxides, such as SiO2 and TiO2,38,39 might enhance the reactivity of surface Fe(II) to facilitate site remediation efforts. Semiconductor minerals, such as goethite and TiO2, play an important role in the redox reaction in natural environments.40−42 Most existing studies on semiconductor minerals focused on photocatalytic reactions, but there has been relatively little research to explore the role semiconductor minerals may play in non-photochemical processes.2,43 Some researchers have proposed and validated that semiconductor minerals, such as pyrite, can act as a conduit between dissolved oxygen and thiosulfate under dark conditions.43,44 Therefore, it would be interesting to examine the interactions between two common semiconductor minerals (e.g., goethite and TiO2) and how the interactions affect the reductive reactivity of adsorbed Fe(II) under dark conditions. The aim of this study was to elucidate the impact of second metal oxides on Fe(II)−goethite reductive reactivity. Goethite was chosen as the representative iron oxide because it is one of the most thermodynamically stable iron oxides in soil and sediments. p-Cyanonitrobenzene (pCNB), a widely used probe compound, was selected as a result of its facile reduction and low adsorption to mineral surfaces.45−48 pCNB was examined for its reduction kinetics as an indicator for the reactivity of surface-adsorbed Fe(II). The effects of two metal oxides (SiO2 and TiO2) on the reactivity of Fe(II)/goethite were investigated. The impact of silicate ions on the reductive reactivity was studied to evaluate the relative contribution of soluble ions versus SiO2 oxide. In addition, we proposed and validated the occurrence of interparticle electron transfer (IPET) from TiO2 to goethite by examining the impact of different types and amounts of TiO2 and pH. The results of this work offered a new view to the importance of semiconductor minerals on non-photocatalytic redox reactions. The previously unrecognized IPET might significantly affect the fate and transformation of contaminants and Fe(II)− Fe(III) redox cycling in many dark environments.

Table 1. Properties of the Metal Oxide Particlesa

a

name

BET (m2/g)

particle size (nm)

pHzpc

FeOOH TiO2 (P25) anatase rutile rutile with SiO2 SiO2

17.5 35−65 200−220 20−40 160 200 ± 25

length 300−1500, diameter 100 25 10−30 60 10−40 12

8.0 6.3 5.9 5.4 5.4 2.4

The properties were obtained from refs 30, 31, and 49.

2.2. Reactor Setup. All solutions in the experiments were degassed by boiling under vacuum before transferred into an anaerobic chamber with an atmosphere of 95−98% N2 and 2− 5% H2 (Coy Laboratory Products, Inc.). All experiments were conducted in 50 mL amber glass serum bottles under constant stirring. NaCl (0.1 M) was added to maintain ionic strength. Reactors were prepared by adding goethite and/or second metal oxides to UP water containing buffer and NaCl. The reactors were then equilibrated overnight to ensure complete hydration of the oxide surfaces. After that, a certain amount of Fe(II) was added to the reactors, which were then equilibrated for 24 h at 25 ± 2 °C. Reactions were initiated by adding a certain amount of pCNB to the reactors. Aliquots of suspensions were collected and filtered using 0.2 μm nylon filters. Typical reaction conditions were 1 g/L goethite, 180 μM Fe(II), 0−1 g/L SiO2 or TiO2, 3 μM pCNB, 0.1 M NaCl, and pH 6.3. In addition, to better illustrate the IPET effect, a dialysis bag (Spectra/Por 7) with a small molecular weight cutoff (MWCO = 3500) was used to eliminate the direct contact between TiO2 and goethite particles (Figure S1 of the Supporting Information). During the reaction, TiO2 particles were confined within the dialysis bag, while goethite particles, Fe(II), and pCNB were added to the reactor but outside of the bag. All other conditions were the same as the typical conditions. 2.3. Analytical Methods. The concentrations of pCNB in the supernatants were analyzed by an Agilent 1200 reversedphase high-performance liquid chromatography (HPLC) system with a diode array detector and a Zorbax XDB-C18 column (4.6 × 250 mm, 5 μm) at a flow rate of 1 mL/min. For pCNB, the mobile phase was UP water and methanol (60:40). Gradient elution was run to separate pCNB and its reaction products. Aqueous Fe(II) and total Fe were analyzed by a modified ferrozine method.50,51 A total of 100 μL of reagent A (10 mM ferrozine with 0.1 M ammonia acetate) was added to 1 mL samples. The absorbance was measured by an ultraviolet−visible (UV−vis) spectrophotometer (Agilent 8670) at 562 nm for Fe(II) concentrations. Dissolved metal ion concentrations were measured by inductively coupled plasma atomic emission spectroscopy (ICP−OES, iCAP, Thermo Scientific) after filtering the suspensions using 0.2

2. EXPERIMENTAL SECTION 2.1. Chemicals and Reagents. pCNB was purchased from Sigma-Aldrich at >97%. Other chemicals, including B

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Figure 1. Effects of varying (a) SiO2 and (b) TiO2 loading on the reduction kinetics of pCNB by Fe(II)/goethite. Reaction conditions: 3 μM pCNB, 1 g/L goethite (G), 180 μM Fe(II), 0.2−1 g/L second metal oxides (SiO2 or TiO2−P25), 0.1 M NaCl, and pH 6.3.

μm nylon filters followed by acidifying the supernatants. Cyclic voltammetry (CV) measurements were conducted using a three-electrode configuration cell. The reference electrode was an Ag/AgCl electrode (CH Instruments). The working electrode was a glassy carbon electrode (3 mm diameter from CH Instruments), while the counter electrode was a platinum wire. The scan rate was 80 mV/s. 2.4. Effect of Silicate Ions on Reductive Reactivity. To evaluate the contribution of silicate ions to the overall effect of SiO2 oxide on the reductive reactivity, the effects of SiO2 and silicate ions were compared. First, after the dissolution of different loadings of SiO2 either alone or with goethite (1 g/L), the silicate concentrations in the supernatants were measured by ICP−OES. On the basis of the concentrations measured in the supernatants and the obtained adsorption isotherm of silicate ions on goethite (1 g/L), a range of initial concentrations of silicate, prepared from stock solution of Na2SiO3, was used to study the effect of silicate ions on the reductive reactivity. Ti4+ was not studied because the solubility of TiO2 is very low.31,52 2.5. Sedimentation. Previous work examined equilibrium aggregation states of oxide mixtures by conducting sedimentation experiments.30 Sedimentation of goethite and its mixtures with second oxides was studied by monitoring the optical absorbance at 508 nm of the suspension as a function of time using the UV−vis spectrophotometer.30 Here, it was measuring scattering of light rather than absorption of light. Normalized absorbance measurements (A/A0) were reported as a function of time. Similar to the previous research,30 0.2 g/L goethite was used in the single oxide experiment. In the binary metal oxide systems, 0.1 g/L of each metal oxide was used to maintain a total metal oxide concentration of 0.2 g/L. Similar to the kinetic experiments, goethite and its mixtures were stirred overnight on a magnetic stir plate, the pH of the reactors was adjusted to 6.3, and the mixtures were then equilibrated for 24 h before the experiments.

oxides examined in this study are listed in Table 1. Rate constants for the reduction of pCNB (k) were calculated on the basis of the pseudo-first-order kinetics. In the absence of either goethite or Fe(II), the pCNB concentration remained constant during the entire time course (data not shown). In addition, the adsorption of pCNB on metal oxides was found to be negligible (Table S1 of the Supporting Information). In the presence of both Fe(II) and goethite [or Fe(II) and TiO2], the pCNB concentration decreased over time as a result of reduction (Figure S2 of the Supporting Information), similar to previous research.3 However, in the presence of both Fe(II) and SiO2, the k values were very small, ranging from 0.014 to 0.019 h−1 when the SiO2 loading was 0.2−1 g/L. For different metal oxides, the k values decreased in the following order: goethite ≫ TiO2 ≫ SiO2. As shown in Figure 1a, the k values of pCNB reduction decreased as the loading of SiO2 increased in the mixtures of goethite and SiO2. The k value decreased from 1.176 to 0.104 h−1 when SiO2 loading increased from 0 to 0.2 g/L. As SiO2 loading increased to 1 g/L, k decreased further to 0.011 h−1. Therefore, SiO2 had a strong inhibiting effect. To our surprise, unlike SiO2, TiO2 increased the k value as its loading increased (red circles in Figure 1b). In addition, the k values of the goethite + TiO2 mixtures (black squares in Figure 1b) were higher than the sum of the k values of the respective single oxide (blue triangles in Figure 1b).10 It is interesting to see the two metal oxides behaved drastically differently on the reductive reactivity of Fe(II)/goethite, as will be explained below. 3.2. Inhibition Mechanism of SiO2. The inhibition effect of SiO2 was examined first. Previous studies have shown that SiO2 can affect the oxidation of MnO2 by silicate anions adsorbed onto the metal oxide surface.31 On the basis of the calculation using MINEQL+ 4.6, the solubility of SiO2 is about 1.86 mM at pH 6.3 (Figure S3 of the Supporting Information). This results in a significant amount of silicate released into the aqueous phase, which might play an important role in the inhibition effect. To evaluate the relative contribution of silicate versus SiO2 oxide particles to the overall inhibitory effect of SiO2 on the reductive reactivity of Fe(II)/goethite, we examined the effect of silicate on the degradation of pCNB. The initial concentrations of silicate in SiO2 suspensions were determined on the basis of the dissolution experiments of

3. RESULTS AND DISCUSSION 3.1. pCNB Degradation in the Presence of Second Metal Oxides. To quantitatively examine the reductive reactivity of Fe(II)/goethite mixed with different second metal oxides, experiments were designed to monitor the reduction kinetics of pCNB by Fe(II)/goethite under various second metal oxide loadings. Selected properties of the metal C

DOI: 10.1021/acsearthspacechem.8b00210 ACS Earth Space Chem. XXXX, XXX, XXX−XXX

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ACS Earth and Space Chemistry different loadings of SiO2 (Text S1 and Table S2 of the Supporting Information) and the adsorption isotherm of silicate on 1 g/L goethite (inset in Figure 2). For example, in

Figure 3. Adsorption of Fe(II) onto the mixtures of goethite and SiO2 or goethite and silicate (G, 1 g/L goethite; G−SiO2-0.2, G−SiO2-0.5, and G−SiO2-1, 1 g/L goethite + 0.2−1 g/L SiO2; Si-100, Si-500, and Si-1000, 1 g/L goethite + 100−1000 μM Na2SiO3; and SiO2-1, 1 g/L SiO2). Other reaction conditions: 0.1 M NaCl, 180 μM Fe(II), and pH 6.3.

Figure 2. Effect of the silicate ion concentration on the reduction kinetics of Fe(II)/goethite mixed with either SiO2 (0.2−1 g/L) (G + SiO2) or silicate (50−1000 μM) (G + soluble Si). (Inset) Adsorption isotherm of silicate on 1 g/L goethite in the presence of 0.1 M NaCl at pH 6.3.

the mixture of 1 g/L goethite + 0.2 g/L SiO2, the silicate concentration in the supernatant was 413 μM (Table S2 of the Supporting Information). On the basis of the adsorption isotherm (Figure 2), the amount of silicate adsorbed qe was 67 μmol/g or 67 μM on 1 g/L goethite. Therefore, the total concentration of silicate in the mixture was 480 μM. Figure 2 shows comparable rate constants of Fe(II)/goethite with either SiO2 suspension (koxide) or silicate (ksol) under the same silicate concentrations. In the presence of silicate, ksol decreased as the metasilicate loading increased from 50 to 1000 μM. Under the same silicate concentrations, koxide was only slightly smaller than ksol. The results demonstrate that silicate played a dominant role in inhibiting the redox reactivity. The inhibition effect by silicate might result from two reasons. First, as demonstrated by previous research,3,10 the amount of Fe(II) adsorbed can significantly affect the reductive reactivity. As the silicate loading increased from 0 to 1000 μM in the mixtures of goethite and silicate, the amount of Fe(II) adsorbed decreased from 32.1 to 22.1% (Figure 3), hence leading to lower reactivity. Comparable decreases in the amount of Fe(II) adsorbed also occurred in the mixtures of goethite with 0−1 g/L SiO2 oxide. Similarly, silicate adsorption onto metal oxides has been reported to inhibit the adsorption of arsenate and humic substances by iron hydroxides.53,54 Second, part of adsorbed Fe(II) likely complexed with the oxygen groups of the adsorbed silicate rather than with goethite, inhibiting Fe(II)-to-Fe(III) electron transfer.55 Before adding Fe(II), goethite and SiO2 had been preequilibrated overnight; thus, the goethite surface had been covered with silicate ions. This was proven by the lower amounts of silicate dissolved in goethite + SiO2 than in SiO2 alone (Table S2 of the Supporting Information) and the change in the goethite aggregation patterns, as indicated by the much slower sedimentation kinetics (and, hence, formation of smaller aggregates) when SiO2 was added (Figure 4). In the binary mixtures of goethite and SiO2, the adsorption of silicate onto goethite would change the surface charge from positive to negative because SiO2 surfaces were negatively charged (pHpzc 2.4; Table 1), which inhibited goethite aggregation. In

Figure 4. Sedimentation time courses of goethite mixing with different second metal oxides in the presence of 0.1 M NaCl at pH 6.3 (G, 0.2 g/L goethite; G + TiO2, 0.1 g/L goethite + 0.1 g/L TiO2; TiO2, 0.2 g/L TiO2; and G + SiO2, 0.1 g/L goethite + 0.1 g/L SiO2).

addition, the reduction potentials of Fe(II)/goethite in the presence of 0−1 g/L SiO2 were measured, as shown in Table S3 of the Supporting Information. Increasing SiO2 loading led to increasingly less negative reduction potentials for Fe(II)− goethite, suggesting lower reductive reactivity. Nevertheless, more details of electron transfer in ternary mixtures of Fe(II)− silicate−goethite should be studied in future research. Besides the dominant effect of silicate, part of Fe(II) was adsorbed on SiO2 whose reductive reactivity was very low (Figure 1), contributing to the minor difference in the reactivity between Fe(II)/goethite + silicate and Fe(II)/ goethite + SiO2 (Figure 2). The Fe(II) adsorption onto SiO2 was proven by the direct adsorption onto 1 g/L SiO2 and the observed differences in the Fe(II) adsorption by goethite + silicate versus by goethite + SiO2 (Figure 3). For example, the initial silicate concentration was similar between goethite + 0.2 g/L SiO2 and goethite + 500 μM silicate; however, the Fe(II) D

DOI: 10.1021/acsearthspacechem.8b00210 ACS Earth Space Chem. XXXX, XXX, XXX−XXX

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Figure 5. (a) Effect of varying TiO2 types on the reduction kinetics of pCNB by Fe(II)/goethite (bars) and the adsorption percentages of Fe(II) onto goethite + TiO2 (line). The differences between different data were statistically significant based on t tests (p < 0.05). (b) Reductive reactivity of Fe(II)−goethite−TiO2 in batch (G + TiO2) versus in the dialysis bag reactors. “G + TiO2(DB)” means goethite and TiO2 were separated by the dialysis bag; “TiO2(DB)” means that TiO2 was in the dialysis bag without adding goethite in the reactor; and “G(DB)” means that goethite was in the reactor without TiO2 in the dialysis bag. Reaction conditions unless otherwise specified: 3 μM pCNB, 1 g/L goethite, 1 g/L TiO2−P25, 180 μM Fe(II), 0.1 M NaCl, and pH 6.3.

or slightly above the redox potential of the reductant can catalyze the reduction, and the closer the redox potentials of the reductant and the conduction band, the faster the reaction.43,57 For example, Xu et al. found that, among many metal-doped sphalerites, Ni-doped sphalerites had the highest activity in thiosulfate oxidation because its conduction band potential was the closest to the redox potential of the S2O32−/ S4O62− couple.43 To examine the validity of the proposed mechanism, the effects of different types of TiO2 were investigated. As shown in Figure 5a, Fe(II)/goethite + rutile coated with less than 5% SiO2 (G−rutile−SiO2) had the lowest k value, although it adsorbed the highest amount of Fe(II). This is mainly due to the inhibition effect of dissolved silicate ions, as discussed in section 3.2. The k value of Fe(II)/goethite + anatase (G− anatase) was smaller than that of Fe(II)/goethite + TiO2−P25 (G−P25) or Fe(II)/goethite + rutile (G−rutile), despite the higher amount of Fe(II) adsorbed. This result can be explained by the semiconductor property of TiO2, as shown below. Figure 6 depicts the flat band potentials of the valence and conduction bands of goethite and TiO2 at pH 6.3. As shown, the flat band potential of anatase is ∼0.1 eV more negative than that of rutile, so that the conduction band of anatase is 0.1 eV above that of rutile. The conduction band of rutile is hence closer than that of anatase to the redox potential of the reductants, such as surface-adsorbed Fe(II) and aqueous species Fe(OH)2+/Fe(OH)20. Therefore, according to previous studies,2,43 electron transfer from the reductants through the conduction band of rutile to that of goethite is faster than through the conduction band of anatase, resulting in the larger k value of G−rutile than that of G−anatase. The k value of G− P25 is between that of G−rutile and G−anatase because TiO2−P25 is a mixture of anatase and rutile. Therefore, the reductive reactivity by Fe(II)/goethite + TiO2 correlates well with the band energies of different TiO2 types, validating the mechanism proposed above. This IPET has been commonly studied for its role in photocatalytic reactions.44,58,59 Here, we believe IPET is applicable in our dark system in improving the reductive reactivity of Fe(II)/goethite in the presence of TiO2. In addition, a dialysis bag was employed to eliminate the direct contact between TiO2 and goethite to better illustrate the IPET. As shown in Figure 5b, when TiO2 and goethite

adsorption in the former was higher than in the later, indicating that additional Fe(II) was adsorbed by SiO2. 3.3. Enhancement Mechanism of TiO2. Unlike SiO2, TiO2 did not affect the goethite aggregation pattern (Figure 4). This is because the surface charge of TiO2 was about zero at pH 6.3 (Table 1) to allow for significant heteroaggregation between TiO2 and goethite. The heteroaggregation was also demonstrated in the transmission electron microscopy (TEM) images of goethite, TiO2, and their mixture, as shown in Figure S5 of the Supporting Information. In addition, previous work already reported that TiO2 was poorly soluble in water;52 thus, there was a negligible impact of Ti4+ ions. It was believed that adsorbed Fe(II) on metal oxides could directly reduce aqueous contaminants.7,16 If so, the value of kG+P25 for the mixture of goethite and TiO2−P25 should not be greater than kG + kP25 (i.e., the sum of the k values for single oxides) because the amount of Fe(II) adsorbed by the mixture cannot be greater than the sum by goethite and TiO2 separately. However, our results showed that kG+P25 ≫ kG + kP25 (Figure 1b), suggesting that TiO2 did not simply provide additional surfaces to enhance the reactivity of Fe(II)/goethite. Therefore, we propose that the conduction bands of goethite and TiO2 can act as conduits for electron transfer from Fe(II) to pCNB. This mechanism resembles the reported conceptual model by Gorski and Scherer for the electron transfer mechanism in Fe(II)−iron oxides.56 In the goethite and TiO2 mixtures, there were two types of adsorbed Fe(II). One was adsorbed onto goethite [i.e., ads.Fe(II)G], and the other was adsorbed onto TiO2 [i.e., ads.Fe(II)T]. Ads.Fe(II)G can reduce the contaminant through the bulk conduction band of goethite, as suggested in the revised conceptual model.56 Ads.Fe(II)T can play two roles in the reduction of pCNB. First, part of ads.Fe(II)T can directly reduce pCNB through the conduction band of TiO2, as indicated by the control experiments of Fe(II)/TiO2 (Figure 1b). However, the reactivity of ads.Fe(II)T was rather low compared to that of ads.Fe(II)G. Second, part of ads.Fe(II)T could donate electrons through the conduction band of TiO2 to that of goethite, which could then reduce pCNB on goethite surfaces. If the above proposed mechanism is correct, then only those semiconductors whose conduction band potentials are below E

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Figure 7. Linear correlation of the experimental initial rate constant (k) for degradation of pCNB versus adsorbed Fe(II) (μM) [Fe(II), 180 μM; NaCl, 0.1 M; pH, 6.3; goethite, 1 g/L; and TiO2−P25, 0.2− 4 g/L].

Figure 6. Diagram depicting the redox potentials of the conduction and valence bands and the band gap energies for TiO2 and FeOOH at pH 6.3. The potentials of the conduction and valence bands of goethite were obtained from ref 60; the potentials of the conduction and valence bands of anatase and rutile were obtained from ref 40; and the reduction potentials of Fe(II)/goethite, Fe(II)/TiO2, and Fe(OH)2+/Fe(OH)20 were obtain from refs 17, 55, and 61.

increased; thus, more electrons could participate in IPET, resulting in a higher reductive reactivity. Finally, the effect of pH on the reductive reactivity of binary metal oxides was studied. Here, we used two types of metal oxide systems. One contained 2 g/L goethite and the other contained 1 g/L goethite plus 1 g/L TiO2−P25. As discussed in section 3.1, the k value of Fe(II)/goethite was much higher than that of Fe(II)/TiO2; however, when the two metal oxides were mixed together, the k values in these two systems were comparable (Figure 8), validating the promoting effect of

were physically separated, the reductive reactivity of the ternary system G + TiO2(DB) was comparable to the sum of that of Fe2+−G [G(DB)] and Fe2+−TiO2 [TiO2(DB)]. Note that the much lower reactivity of TiO2 + Fe2+ was because of the mass transfer limitation imposed by the dialysis bag on Fe2+ and pCNB. This is very different from the results in Figure 1b, where the reductive reactivity of the ternary system was much higher than the sum of the k values of the respective single oxide. This is because the electrons of ads.Fe(II)T in the dialysis bag were not able to transfer to goethite to reduce pCNB, verifying that IPET was important in enhancing the reactivity. It should be noted that Fe(II) here was the ultimate electron donor; however, the fate of the electrons in the reaction systems is not yet clear. Gorski and Scherer have proposed that the electrons provided by Fe(II) in Fe(II)/iron oxide might have three likely fates: (1) immobilized in a trapping site of iron oxide, (2) transfer through a conduction band and then react with contaminants, and (3) move near the surface and then cause surface-bound Fe3+ to undergo reductive dissolution.56 As shown in Figure 5a, the amount of adsorbed Fe(II) in goethite + anatase was higher than that in goethite + rutile. However, the reductive reactivity of the former was lower, which might result from the fact that part of the electrons from Fe(II) were immobilized in the trapping sites of the two semiconductor minerals. Nevertheless, the fate of electrons in the mixed metal oxides should be further studied in the future. In addition, TiO2−P25 was used to study how adsorption of Fe(II) by TiO2 affected the reactivity in binary metal oxides. One might reasonably expect that the addition of TiO2 would increase Fe(II) adsorption because the mineral loading was higher, as indeed observed in Figure S6 of the Supporting Information. Previous studies have believed that Fe(II) adsorption and speciation in heterogeneous systems were key factors to influence the redox reactivity of metal oxides.3,7,16,61,62 In our experiments, the k value was linearly correlated with the amount of Fe(II) adsorbed when TiO2 loading was increased from 0 to 4 g/L, while both goethite and Fe(II) loadings were fixed (Figure 7). This is because, as the loading of TiO2−P25 increased, the amount of ads.Fe(II)T also

Figure 8. Effect of pH on the reductive reactivity of Fe(II)−goethite (2 g/L) and Fe(II)−goethite (1 g/L) + TiO2 (1 g/L). Reaction conditions: 3 μM pCNB, 0.1 M NaCl, and 180 μM Fe(II).

TiO2. The same promoting effect also occurred under different pH values (Figure 8). This again indicates that IPET may occur between goethite and TiO2, even under different pH conditions.

4. CONCLUSION In summary, this work systematically examined the reduction behavior of surface-adsorbed Fe(II) in binary oxide mixtures. The properties of the second metal oxides significantly affected the surface-mediated reductive reactivity of Fe(II)/goethite. SiO2 strongly inhibited the reductive reactivity of adsorbed Fe(II) by dissolved silicate ions. In contrast, TiO2 substantially F

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ACS Earth and Space Chemistry

nitroaromatic compounds under iron-reducing subsurface conditions. Environ. Sci. Technol. 1999, 33 (9), 1479−1487. (5) Klupinski, T. P.; Chin, Y.-P.; Traina, S. J. Abiotic degradation of pentachloronitrobenzene by Fe(II): Reactions on goethite and iron oxide nanoparticles. Environ. Sci. Technol. 2004, 38 (16), 4353−4360. (6) Tai, Y.-L.; Dempsey, B. A. Nitrite reduction with hydrous ferric oxide and Fe(II): Stoichiometry, rate, and mechanism. Water Res. 2009, 43 (2), 546−552. (7) Liger, E.; Charlet, L.; Van Cappellen, P. Surface catalysis of uranium(VI) reduction by iron(II). Geochim. Cosmochim. Acta 1999, 63 (19), 2939−2955. (8) Taylor, S. D.; Marcano, M. C.; Becker, U. A first principles investigation of electron transfer between Fe(II) and U(VI) on insulating Al- vs. semiconducting Fe-oxide surfaces via the proximity effect. Geochim. Cosmochim. Acta 2017, 197, 305−322. (9) Pecher, K.; Haderlein, S. B.; Schwarzenbach, R. P. Reduction of polyhalogenated methanes by surface-bound Fe(II) in aqueous suspensions of iron oxides. Environ. Sci. Technol. 2002, 36 (8), 1734−1741. (10) Amonette, J. E.; Workman, D. J.; Kennedy, D. W.; Fruchter, J. S.; Gorby, Y. A. Dechlorination of carbon tetrachloride by Fe(II) associated with goethite. Environ. Sci. Technol. 2000, 34 (21), 4606− 4613. (11) Strathmann, T. J.; Stone, A. T. Mineral surface catalysis of reactions between FeII and oxime carbamate pesticides. Geochim. Cosmochim. Acta 2003, 67 (15), 2775−2791. (12) Elsner, M.; Schwarzenbach, R. P.; Haderlein, S. B. Reactivity of Fe(II)-bearing minerals toward reductive transformation of organic contaminants. Environ. Sci. Technol. 2004, 38 (3), 799−807. (13) Vikesland, P. J.; Valentine, R. L. Iron oxide surface-catalyzed oxidation of ferrous iron by monochloramine: Implications of oxide type and carbonate on reactivity. Environ. Sci. Technol. 2002, 36 (3), 512−519. (14) Wang, S.; Arnold, W. A. Abiotic reduction of dinitroaniline herbicides. Water Res. 2003, 37 (17), 4191−4201. (15) Klupinski, T. P.; Chin, Y.-P. Abiotic degradation of trifluralin by Fe(II): Kinetics and transformation pathways. Environ. Sci. Technol. 2003, 37 (7), 1311−1318. (16) Charlet, L.; Silvester, E.; Liger, E. N-compound reduction and actinide immobilisation in surficial fluids by Fe(II): The surface FeIIIOFeIIOH° species as major reductant. Chem. Geol. 1998, 151 (1), 85−93. (17) Silvester, E.; Charlet, L.; Tournassat, C.; Géhin, A.; Grenèche, J.-M.; Liger, E. Redox potential measurements and Mössbauer spectrometry of FeII adsorbed onto FeIII (oxyhydr)oxides. Geochim. Cosmochim. Acta 2005, 69 (20), 4801−4815. (18) Amirbahman, A.; Kent, D. B.; Curtis, G. P.; Marvin-DiPasquale, M. C. Kinetics of homogeneous and surface-catalyzed mercury (II) reduction by iron (II). Environ. Sci. Technol. 2013, 47 (13), 7204− 7213. (19) Maithreepala, R.; Doong, R.-a. Synergistic effect of copper ion on the reductive dechlorination of carbon tetrachloride by surfacebound Fe(II) associated with goethite. Environ. Sci. Technol. 2004, 38 (1), 260−268. (20) Tao, L.; Zhu, Z.; Li, F. Fe(II)/Cu(II) interaction on αFeOOH, kaolin and TiO2 for interfacial reactions of 2-nitrophenol reductive transformation. Colloids Surf., A 2013, 425, 92−98. (21) Colón, D.; Weber, E. J.; Anderson, J. L. Effect of natural organic matter on the reduction of nitroaromatics by Fe(II) species. Environ. Sci. Technol. 2008, 42 (17), 6538−6543. (22) Vindedahl, A. M.; Stemig, M. S.; Arnold, W. A.; Penn, R. L. Character of humic substances as a predictor for goethite nanoparticle reactivity and aggregation. Environ. Sci. Technol. 2016, 50 (3), 1200− 1208. (23) Anderson, P. R.; Benjamin, M. M. Surface and bulk characteristics of binary oxide suspensions. Environ. Sci. Technol. 1990, 24 (5), 692−698. (24) Stumm, W.; Sulzberger, B. The cycling of iron in natural environments: Considerations based on laboratory studies of

enhanced the reactivity as a result of the IPET between TiO2 and goethite. Previously unrecognized IPET could significantly affect the fate and transformation of contaminants and Fe(II)− Fe(III) redox cycling in dark environments. The obtained results will contribute to the development of accurate predictive tools to estimate the fate and transformation of contaminants in the environment. In addition, the new findings point to a better way to design more efficient, environmentally benign catalysts.

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ASSOCIATED CONTENT

* Supporting Information S

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acsearthspacechem.8b00210.

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Dissolution of SiO2 (Text S1), concentrations of pCNB (μM) in the aqueous phase after mixing with different oxides for 50 min and 2 h (Table S1), soluble silicon concentrations in single and binary oxide suspensions at pH 6.3 (Table S2), half-wave potentials (E1/2) of Fe(II)/goethite in the presence of 0−1.0 g/L SiO2 (Table S3), experiment setup for separating TiO2 and goethite with a dialysis bag (Figure S1), effects of varying (a) SiO2 and (b) TiO2 loading on the reduction kinetics of pCNB by Fe(II)/goethite (Figure S2), solubility of 1 g/L SiO2 (log of the molar concentration of total soluble Si) as a function of pH at 0.1 M NaCl, calculated using MINEQL+ 4.6 (Figure S3), cyclic voltammograms (CV) of Fe(II)/goethite in the presence of SiO2 (0−1 g/L), from −1.0 to 1.0 V at a scan rate of 80 mV/s (Figure S4), TEM images of (a) goethite, (b) TiO2, and (c) goethite + TiO2 (Figure S5), and adsorption of Fe(II) onto different goethite and TiO2 mixtures (Figure S6) (PDF)

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

Yifan Dai: 0000-0002-1009-5790 Chung-Chiun Liu: 0000-0002-4313-8064 Huichun Zhang: 0000-0002-5683-5117 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This material is based on work supported by the National Science Foundation under Grants CBET-1762691 and CHE1762686.

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REFERENCES

(1) Schoonen, M.; Strongin, D. R. Catalysis of electron transfer reactions at mineral surfaces. Environ. Catal. 2005, 37−60. (2) Xu, Y.; Schoonen, M. A. The stability of thiosulfate in the presence of pyrite in low-temperature aqueous solutions. Geochim. Cosmochim. Acta 1995, 59 (22), 4605−4622. (3) Klausen, J.; Troeber, S. P.; Haderlein, S. B.; Schwarzenbach, R. P. Reduction of substituted nitrobenzenes by Fe(II) in aqueous mineral suspensions. Environ. Sci. Technol. 1995, 29 (9), 2396−2404. (4) Hofstetter, T. B.; Heijman, C. G.; Haderlein, S. B.; Holliger, C.; Schwarzenbach, R. P. Complete reduction of TNT and other (poly) G

DOI: 10.1021/acsearthspacechem.8b00210 ACS Earth Space Chem. XXXX, XXX, XXX−XXX

Article

ACS Earth and Space Chemistry heterogeneous redox processes. Geochim. Cosmochim. Acta 1992, 56 (8), 3233−3257. (25) Brown, L. S. The occurrence of Leucoxene in some of the Permian Mid-Continent sediments. Am. Mineral. 1928, 13, 233−235. (26) O’Neill, H. S. C.; Pownceby, M. I.; Wall, V. J. Ilmenite-rutileiron and ulvospinel-ilmenite-iron equilibria and the thermochemistry of ilmenite (FeTiO3) and ulvospinel (Fe2TiO4). Geochim. Cosmochim. Acta 1988, 52 (8), 2065−2072. (27) Pearce, C. I.; Liu, J.; Baer, D. R.; Qafoku, O.; Heald, S. M.; Arenholz, E.; Grosz, A. E.; McKinley, J. P.; Resch, C. T.; Bowden, M. E.; Engelhard, M. H.; Rosso, K. M. Characterization of natural titanomagnetites (Fe3−xTixO4) for studying heterogeneous electron transfer to Tc(VII) in the Hanford subsurface. Geochim. Cosmochim. Acta 2014, 128, 114−127. (28) Xu, Y.; Axe, L. Synthesis and characterization of iron oxidecoated silica and its effect on metal adsorption. J. Colloid Interface Sci. 2005, 282 (1), 11−19. (29) Taujale, S.; Baratta, L. R.; Huang, J.; Zhang, H. Interactions in ternary mixtures of MnO2, Al2O3, and natural organic matter (NOM) and the impact on MnO2 oxidative reactivity. Environ. Sci. Technol. 2016, 50 (5), 2345−2353. (30) Zhang, H.; Taujale, S.; Huang, J.; Lee, G.-J. Effects of NOM on oxidative reactivity of manganese dioxide in binary oxide mixtures with goethite or hematite. Langmuir 2015, 31 (9), 2790−2799. (31) Taujale, S.; Zhang, H. Impact of interactions between metal oxides to oxidative reactivity of manganese dioxide. Environ. Sci. Technol. 2012, 46 (5), 2764−2771. (32) Tada, H.; Jin, Q.; Nishijima, H.; Yamamoto, H.; Fujishima, M.; Okuoka, S. i.; Hattori, T.; Sumida, Y.; Kobayashi, H. Titanium (IV) dioxide surface-Modified with iron oxide as a visible light photocatalyst. Angew. Chem. 2011, 123 (15), 3563−3567. (33) Meng, X.; Letterman, R. D. Effect of component oxide interaction on the adsorption properties of mixed oxides. Environ. Sci. Technol. 1993, 27 (5), 970−975. (34) Gottschalk, F.; Sonderer, T.; Scholz, R. W.; Nowack, B. Modeled environmental concentrations of engineered nanomaterials (TiO2, ZnO, Ag, CNT, fullerenes) for different regions. Environ. Sci. Technol. 2009, 43 (24), 9216−9222. (35) Chowdhury, I.; Cwiertny, D. M.; Walker, S. L. Combined factors influencing the aggregation and deposition of nano-TiO2 in the presence of humic acid and bacteria. Environ. Sci. Technol. 2012, 46 (13), 6968−6976. (36) Godinez, I. G.; Darnault, C. J. Aggregation and transport of nano-TiO2 in saturated porous media: Effects of pH, surfactants and flow velocity. Water Res. 2011, 45 (2), 839−851. (37) Tong, T.; Fang, K.; Thomas, S. A.; Kelly, J. J.; Gray, K. A.; Gaillard, J.-F. o. Chemical interactions between nano-ZnO and nanoTiO2 in a natural aqueous medium. Environ. Sci. Technol. 2014, 48 (14), 7924−7932. (38) Bird, M. I.; Longstaffe, F. J.; Fyfe, W. S. Oxygen-isotope fractionation in titanium-oxide minerals at low temperature. Geochim. Cosmochim. Acta 1993, 57 (13), 3083−3091. (39) Boonfueng, T.; Axe, L.; Xu, Y. Properties and structure of manganese oxide-coated clay. J. Colloid Interface Sci. 2005, 281 (1), 80−92. (40) Schoonen, M. A.; Xu, Y.; Strongin, D. R. An introduction to geocatalysis. J. Geochem. Explor. 1998, 62 (1), 201−215. (41) Selli, E.; De Giorgi, A.; Bidoglio, G. Humic acid-sensitized photoreduction of Cr(VI) on ZnO particles. Environ. Sci. Technol. 1996, 30 (2), 598−604. (42) Brown, G. E.; Henrich, V. E.; Casey, W. H.; Clark, D. L.; Eggleston, C.; Felmy, A.; Goodman, D. W.; Grätzel, M.; Maciel, G.; McCarthy, M. I.; Nealson, K. H.; Sverjensky, D. A.; Toney, M. F.; Zachara, J. M. Metal oxide surfaces and their interactions with aqueous solutions and microbial organisms. Chem. Rev. 1999, 99 (1), 77−174. (43) Xu, Y.; Schoonen, M. A. A.; Strongin, D. R. Thiosulfate oxidation: Catalysis of synthetic sphalerite doped with transition metals. Geochim. Cosmochim. Acta 1996, 60 (23), 4701−4710.

(44) Serpone, N.; Maruthamuthu, P.; Pichat, P.; Pelizzetti, E.; Hidaka, H. Exploiting the interparticle electron transfer process in the photocatalysed oxidation of phenol, 2-chlorophenol and pentachlorophenol: Chemical evidence for electron and hole transfer between coupled semiconductors. J. Photochem. Photobiol., A 1995, 85 (3), 247−255. (45) Zhang, H.; Weber, E. J. Identifying indicators of reactivity for chemical reductants in sediments. Environ. Sci. Technol. 2013, 47 (13), 6959−6968. (46) Simon, R.; Colón, D.; Tebes-Stevens, C. L.; Weber, E. J. Effect of redox zonation on the reductive transformation of p-cyanonitrobenzene in a laboratory sediment column. Environ. Sci. Technol. 2000, 34 (17), 3617−3622. (47) Colón, D.; Weber, E. J.; Baughman, G. L. Sediment-associated reactions of aromatic amines. 2. QSAR development. Environ. Sci. Technol. 2002, 36 (11), 2443−2450. (48) Colón, D.; Weber, E. J.; Anderson, J. L.; Winget, P.; Suárez, L. A. Reduction of nitrosobenzenes and N-hydroxylanilines by Fe(II) species: Elucidation of the reaction mechanism. Environ. Sci. Technol. 2006, 40 (14), 4449−4454. (49) Kosmulski, M. The significance of the difference in the point of zero charge between rutile and anatase. Adv. Colloid Interface Sci. 2002, 99 (3), 255−264. (50) Chen, Y.; Zhang, H. Complexation facilitated reduction of aromatic N-oxides by aqueous FeII−tiron complex: Reaction kinetics and mechanisms. Environ. Sci. Technol. 2013, 47 (19), 11023−11031. (51) Viollier, E.; Inglett, P.; Hunter, K.; Roychoudhury, A.; Van Cappellen, P. The ferrozine method revisited: Fe(II)/Fe(III) determination in natural waters. Appl. Geochem. 2000, 15 (6), 785− 790. (52) Wiese, G.; Healy, T. Solubility effects in Al2O3 and TiO2 colloidal dispersions. J. Colloid Interface Sci. 1975, 52 (3), 452−457. (53) Davis, C. C.; Knocke, W. R.; Edwards, M. Implications of aqueous silica sorption to iron hydroxide: Mobilization of iron colloids and interference with sorption of arsenate and humic substances. Environ. Sci. Technol. 2001, 35 (15), 3158−3162. (54) Davis, C. C.; Chen, H.-W.; Edwards, M. Modeling silica sorption to iron hydroxide. Environ. Sci. Technol. 2002, 36 (4), 582− 587. (55) Jones, A. M.; Griffin, P. J.; Collins, R. N.; Waite, T. D. Ferrous iron oxidation under acidic conditionsThe effect of ferric oxide surfaces. Geochim. Cosmochim. Acta 2014, 145, 1−12. (56) Gorski, C. A.; Scherer, M. M. Fe2+ sorption at the Fe oxide− water interface: A revised conceptual framework. In Aquatic Redox Chemistry; Tratnyek, P. G., Grundl, T. J., Haderlein, S. B., Eds.; American Chemical Society (ACS): Washington, D.C., 2011; ACS Symposium Series, Vol. 1071, Chapter 15, pp 315−343, DOI: 10.1021/bk-2011-1071.ch015. (57) Morrison, S. R. The Chemical Physics of Surfaces; Springer Science & Business Media: Boston, MA, 2013. (58) Yang, J.; Pan, B.; Li, H.; Liao, S.; Zhang, D.; Wu, M.; Xing, B. Degradation of p-nitrophenol on biochars: Role of persistent free radicals. Environ. Sci. Technol. 2016, 50 (2), 694−700. (59) Lu, J. J.; Lu, Y. M.; Tasi, S. I.; Hsiung, T. L.; Wang, H. P.; Jang, L. Y. Conductivity enhancement and semiconductor−metal transition in Ti-doped ZnO films. Opt. Mater. 2007, 29 (11), 1548−1552. (60) Sherman, D. M. Electronic structures of iron (III) and manganese (IV)(hydr) oxide minerals: Thermodynamics of photochemical reductive dissolution in aquatic environments. Geochim. Cosmochim. Acta 2005, 69 (13), 3249−3255. (61) Li, F.-b.; Tao, L.; Feng, C.-h.; Li, X.-z.; Sun, K.-w. Electrochemical evidences for promoted interfacial reactions: The role of Fe(II) adsorbed onto γ-Al2O3 and TiO2 in reductive transformation of 2-nitrophenol. Environ. Sci. Technol. 2009, 43 (10), 3656−3661. (62) Nano, G. V.; Strathmann, T. J. Application of surface complexation modeling to the reactivity of iron(II) with nitroaromatic and oxime carbamate contaminants in aqueous TiO2 suspensions. J. Colloid Interface Sci. 2008, 321 (2), 350−359. H

DOI: 10.1021/acsearthspacechem.8b00210 ACS Earth Space Chem. XXXX, XXX, XXX−XXX