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C: Surfaces, Interfaces, Porous Materials, and Catalysis

Dissolution, Adsorption, and Redox Reaction in Ternary Mixtures of Goethite, Aluminum Oxides, and Hydroquinone Huichun Judy Zhang, Kowsalya Devi Rasamani, Shifa Zhong, Saru Taujale, Laura R. Baratta, and Zijie Yang J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.8b12217 • Publication Date (Web): 28 Jan 2019 Downloaded from http://pubs.acs.org on February 3, 2019

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The Journal of Physical Chemistry

Dissolution, Adsorption, and Redox Reaction in Ternary Mixtures of Goethite, Aluminum Oxides, and Hydroquinone

Huichun Zhanga,b,*, Kowsalya Devi Rasamani b, Shifa Zhong a, Saru Taujale b, Laura R. Baratta b, and Zijie Yang a a

Department of Civil Engineering, Case Western Reserve University, 2104 Adelbert Road, Cleveland, OH 44106 b

Department of Civil and Environmental Engineering, Temple University 1947 North 12th Street, Philadelphia, PA 19122

*Corresponding Author, contact e-mail: [email protected], phone: (216)368-0689

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Abstract To better understand the oxidative reactivity of iron oxides in the fate of contaminants in acidic environments, we examined the reactivity of goethite in binary mixtures with Al 2 O 3 by carrying out oxidation experiments of hydroquinone (HQ) in the presence of goethite and/or Al 2 O 3 at pH 3. Kinetic results revealed inhibiting effects of 0.2 – 20 g/L of three different Al 2 O 3 on the oxidative reactivity of goethite. Surprising, soluble Al ions of 0.18 – 18 mM had negligible impact on the reactivity. It turned out the Fe3+ dissolved from goethite partly contributed to the observed HQ oxidation, and the Al 2 O 3 adsorbed the Fe3+ to lead to the slower HQ oxidation. The observed pseudo-first order rate constants in HQ oxidation had a strong linear correlation with Fe3+ concentration in various goethite and Al 2 O 3 mixtures. Separate experiments confirmed the reactivity of Fe3+ toward HQ and the linear correlation between [Fe3+] and HQ oxidation reactivity. Finally, sedimentation experiments showed negligible heteroaggregation between goethite and AluC-Al 2 O 3 or nAl 2 O 3 , but intensive heteroaggregation between goethite and Alu 65-Al 2 O 3 , which explained the observed the highest inhibition effect of Alu 65. Overall, oxide mixtures are very complex whose reactivity is determined by many factors such as oxide dissolution.


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Introduction Iron(III) (hydr)oxides (simplified as oxides hereafter), the most abundant Fe form in many sedimentary environments, affect a wide range of geological processes. 1-6 They also play an important role in oxidative transformation of a number of organic contaminants, such as phenols, hydroquinones, anilines, and fluoroquinolones.7-12 To accurately assess the risks of these contaminants in the environment, it is necessary that their redox transformation is properly accounted for regarding both reaction kinetics and products.13 The oxidation mechanism of organic compounds by iron oxides is known to be surface related, involving preequilibrium adsorption of the compounds by the oxide surface, followed by electron transfer from the compounds to the oxide and subsequent product formation.7, 8, 10 Factors that can influence the reaction kinetics include solution pH, the presence of co-solutes such as metal ions and dissolved organic matter (DOM), and the initial concentration and structural characteristics of the contaminant and the oxide.7, 8, 14 To more accurately estimate the fate and transformation of contaminants in soils and sediments, we should recognize that soils and sediments are complex mixtures of minerals and other constituents, including metal oxides, clay minerals, and natural organic matter. Most of the existing work on redox transformation in lab-prepared, simplified model systems, however, only focused on one iron oxide under changing solution conditions, e.g., pH, co-solutes. 8, 15, 16 Therefore, the observed reactivities in the model systems will likely significantly deviate from those in the environment. To decrease such deviations, our recent work examined the oxidative reactivity of MnO 2 in binary mixtures with another oxide, including Al 2 O 3 , Fe(III) oxides, SiO 2 ,



and TiO 2 , and found largely inhibited oxidative reactivity of MnO 2 with the inhibition effect following Al 2 O 3 > Fe(III) oxides >> SiO 2 > TiO 2 . 17-19 The main inhibition mechanisms are highly oxide dependent. 17-19 Given the similarities in the oxidation mechanisms of iron and Mn oxides,8, 20 it is natural to hypothesize that the above observed interaction mechanisms will also affect the oxidative reactivity of iron oxides upon mixing with another metal oxide. In other words, we might expect to see another metal oxide to inhibit the oxidative reactivity of iron oxides through heteroaggregation, surface complexation, and/or competitive adsorption. However, it remains unknown whether and to what extents these mechanisms are applicable to iron oxides. To address this question, we selected hydroquinone (HQ) as a probe compound for oxide reactivity because (1) HQ-like moieties have been widely believed to exist in natural organic matter as the redox-active moiety;21, 22 (2) the oxidation kinetics of HQ at pH 3.75 have been frequently used as a quantitative measure for the reactivity of various Fe(III) oxides;23-25 and (3) the oxidation of HQ by goethite has been systematically examined under changing pH and other solution conditions; for instance, it was found that HQ oxidation was highly pH dependent, with the rate decreased to below the detection limit when the solution pH increased from 4.5 to 6.0.15 Much faster oxidation of AQDS and other quinone-like compounds than HQ by different Fe(III) oxides was also reported.26-28 Because H+ is consumed in the reduction of Fe(III) oxides,29 the reduction potentials of Fe(III) oxides (and hence their oxidizing reactivity) should decrease with increasing pH. Together with the lower reduction potentials of iron oxides than those of Mn oxides and hence lower oxidative reactivity of the former,29 iron oxides are more redox-relevant in acidic environments, such as acid mine drainage sites or hydrothermal vents. Interactions in acidic 4 ACS Paragon Plus Environment

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environments, such as particle interactions, surface complexation and adsorption of contaminants, are expected to be largely different from those under neutral and alkaline conditions and hence deserves special attention. When using HQ as the probe compound, it is worth mentioning that the Fe3+ dissolved from iron oxides may also oxidize HQ, especially in the beginning of the reaction since the iron oxides have been typically equilibrated with solution overnight, although the contribution of the Fe3+ to the overall kinetics has not been explored. 15 Here, at low pH (< 3.5), there was typically a non-zero intercept in the amount of Fe2+ formed at time zero (up to 15 µM depending on the pH), implying fast reduction of the Fe3+ by HQ. Indeed, Pracht et al. reported very fast oxidation of HQ by dissolved Fe(III) with pH varying between pH 0.11 and 6.72.12 The positive intercepts in the amount of Fe(II) formed at time zero suggest fast initial oxidation of either HQ 24 or ascorbate 30 by the dissolved Fe(III). Rapid stoichiometric oxidation of 2-methoxyhydroquinone by Fe(III) was also reported to complete within 5-25 min at pH 4-6, with Fe(II) generation rate increasing with increasing pH because the monodissociated HQ was the most kinetically active species.31 The reaction equilibrium between HQ/Q and Fe(III)/Fe(II) species is highly dependent on thermodynamic constraints such as reduction potential under different pH, quinone structure, and iron speciation.32 However, although the fast reaction between Fe3+ and HQ has been recognized in the previous work,15, 30 its role on the observed oxide reactivity was not examined. Further, the contribution of soluble Fe3+ to the oxidation of HQ points to a potential effect of iron oxide dissolution to the overall oxide redox reactivity under acidic conditions. However, there is little information available regarding how dissolution of iron oxides affects the observed oxidative reactivity. Because oxide dissolution might be influenced by the 5 ACS Paragon Plus Environment


presence of another oxide, the oxidative reactivity of the iron oxides may be changed as a result. The goal of this work was to examine the interaction mechanisms in binary mixtures of iron and Al oxides and to elucidate the impact of the interactions on the oxidative reactivity of the iron oxide. Al oxides were selected because Al is the most abundant metal in the Earth’s crust. Goethite and Al 2 O 3 are widely examined as representative iron and Al oxides and were thus selected in this work to study their interaction mechanisms. First, the reactivity of goethite, as quantified by the oxidation kinetics of HQ, was examined under different conditions including three different types of Al 2 O 3 and different amounts of soluble Al ions, Fe3+, and Fe2+. Next, the dissolution of goethite and Al 2 O 3 as either single oxides or in binary oxide mixtures was monitored. The adsorption of soluble Al ions, Fe3+, and Fe2+ by goethite and Al 2 O 3 was also examined. Then, the extents of aggregation in binary oxide mixtures were quantified by conducting sedimentation experiments. Our results, together with our previous findings on MnO 2 , pointed to the complexity of oxide mixtures and that different systems should be studied individually to avoid misinterpretation of the interaction mechanisms.

Experimental Methods Chemicals and Oxide Preparation. nAl 2 O 3 - γ was purchased from nanostructured and amorphous materials Inc. (Table 1). Aeroxide® AluC Al 2 O 3 (66% γ and 33% δ) and Alu 65 Al 2 O 3 (with Θ and δ structure) were obtained from Evonik Industries. Goethite (Bayferrox 910) was obtained from Lanxess®. Hydroquinone (HQ, 99.5% purity), 1,4-benzoquinone (Q, 99% purity), 6 ACS Paragon Plus Environment

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NaCl, HCl, and NaOH were purchased from Fisher Scientific. All chemicals were used without further purification. The stock solution of hydroquinone was prepared in methanol at 10 mM and stored at 4oC. Table 1. Metal oxides and their properties.a

180

Average particle size (nm) b 25

8.9333

Experimentally determined particle size (nm)b 660

AluC Al 2 O 3 b

100 ± 15

2034

8.9333

846

AluC-Al 2 O 3

Alu 65 Al 2 O 3 b

65 ± 10

-

8.9333

666

Alu 65-

Metal Oxide nAl 2 O 3 -γ a

BET surface area (m2/g)

pH zpc

Referred to as nAl 2 O 3

Al 2 O 3 Goethite18 a

15.0

8.0

100 × 600

-

FeOOH

Properties provided by the manufacturers unless otherwise specified. Manufacturer: a)

Nanostructured and Amorphous Material Inc., b) Evonik Industries.

b

“Experimentally

determined particle size” was the particle size of hydrated particles in solution, different from the size of primary particles (i.e., “average particle size”) reported by the manufacturer. Reactor setup. All kinetic experiments were performed in duplicates using ≥ 18.0 mΩ nanopure water in 125 mL screw-cap amber bottles with Teflon caps. Most reactions were conducted with AluC-Al 2 O 3 unless otherwise specified. Typically, the reactors contained 2 g/L goethite, 1 g/L AluC-Al 2 O 3 , 0.1 M NaCl, and 100 µM HQ at pH 3.0. All reactions were conducted at pH 3 in air to allow the oxidation kinetics (especially in oxide mixtures) to be monitored within a reasonable time frame. As the preliminary data shown in Figure S1a, negligible HQ was oxidized



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by goethite (with or without Al 2 O 3 ) for 12+ days at pH 4 and 5 in the absence of oxygen. In addition, the oxidation of HQ by goethite in the presence of oxygen became increasingly more complicated at pH≥4 (Figure 1a vs. S1b), as O 2 has been demonstrated to be increasingly involved in the radical-based reactions.35 Additional rationale for the focus on pH 3 in this work is provided in Text S1 of the supporting information (SI). pH was adjusted by an Accumet AB 150 pH meter by adding HCl or NaOH. The total volumes of the reactors were 50 mL. In reactors containing Al 2 O 3 or Al ions, they were added in the order of goethite followed by Al 2 O 3 /AlCl 3 . The bottles were then equilibrated for 24 h by stirring on a magnetic stir plate at 22±2oC. This allows the oxide surfaces to be sufficiently hydrated. During the pre-equilibrium and the following reaction periods, the suspension pH was periodically checked and re-adjusted to 3.0 when needed. Reactions were then initiated by spiking an aliquot of HQ stock into the reactors to attain a concentration of 100 µM. At predetermined time intervals within about 200 h, 1 mL of the suspensions were filtered by 0.22 µm Nylon filters. The filtrates were then transferred to separate vials for HQ and Q analysis within the same day by an Agilent 1200 HPLC with a Zorbax XDB-C18 (4.6 × 250 mm, 5 µm) column. Pseudo-first order reaction rate constants k obs were calculated based on the degradation kinetics of HQ (eq. 1) and formation kinetics of Q (eq. 2): ln[HQ] = ln[HQ] 0 – k obs ×t

(1)

ln[HQ] = ln([HQ] 0 -[Q]) – k obs ×t

(2)

where [HQ] and [Q] are the concentrations of HQ and Q at the reaction time t, and [HQ] 0 is the initial HQ concentration. k obs was used to quantify the oxidative reactivity of the mixtures, with


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larger k obs values indicating higher reactivity. Separate batch experiments were also conducted to examine the oxidation kinetics of HQ by Fe3+ under similar conditions, as described above. Dissolution of Al 2 O 3 and goethite. Dissolution experiments of Al 2 O 3 and goethite either separately or in binary oxide mixtures were carried out at pH 3.0, both in the absence and presence of HQ. 1 g/L of Al 2 O 3 and 2 g/L of goethite were added to reactors containing 50 mL DI water and 0.1 M NaCl and adjusted to pH 3.0. The mixtures were then equilibrated for 24 – 96 h on a magnetic stir plate, during which the solution pH was repeatedly re-adjusted by HCl to 3.0. The pH adjustment was more frequent during the first a few hours, especially when there was a larger amount of Al 2 O 3 present. Therefore, the observed oxide dissolution might be somewhat enhanced by the pH adjustment process. At certain times, aliquots of the mixtures were centrifuged followed by filtration through 0.22 µm filters. The filtrates were acidified using 1 M HCl; Al concentrations were measured using an inductively coupled plasma optical emission spectrometry (ICP-OES); and Fe concentrations were measured using a modified Ferrozine method.36 Note that 20 nm syringe filters were also employed to filter the mixtures and similar Fe3+ concentrations were detected (Table S1), therefore, 0.22 µm filters were employed in all the following experiments. For each dissolution condition, additional sets of experiments were prepared and monitored under identical conditions but with the addition of (i) HQ to examine whether HQ affected the dissolution or (ii) 0 – 20 µM Fe3+ or 0 – 120 µM Fe2+ to examine the impact of these ions on goethite dissolution. Adsorption of soluble Al ions, Fe3+, and Fe2+ by oxides. Adsorption of soluble Al ions, Fe3+, and Fe2+ by goethite and Al 2 O 3 was carried out under the experimental conditions identical to the



kinetic experiments. The initial concentration (C i ) of the ions ranged from 0 – 18 mM, 0 – 20 µM, and 0 – 120 µM, respectively. After adding the ions to the reactors, the suspensions were stirred on a magnetic stir plate and equilibrated for up to 96 hours, with periodic pH readjustment to 3.0. The suspensions were then centrifuged and filtered through 0.22 µm filters. The filtrates were acidified by adding 2 M HCl to obtain a final pH of

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