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Ferrous Iron Oxidation under Varying pO2 Levels: the Effect of Fe(III)/Al(III) Oxide Minerals and Organic Matter Chunmei Chen, and Aaron Thompson Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.7b05102 • Publication Date (Web): 01 Dec 2017 Downloaded from http://pubs.acs.org on December 1, 2017

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Environmental Science & Technology

To Environ. Sci. & Technol.

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Ferrous Iron Oxidation under Varying pO2 Levels: the Effect of Fe(III)/Al(III) Oxide

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Minerals and Organic Matter

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Chunmei Chen and Aaron Thompson*

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Depart of Crop and Soil Sciences

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The University of Georgia, Athens, GA, USA 30602

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Corresponding author *

Aaron Thompson, [email protected], (01) 706-410-1293

1 ACS Paragon Plus Environment


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Abstract

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Abiotic Fe(II) oxidation by O2 commonly occurs in the presence of mineral sorbents and

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organic matter (OM) in soils and sediments, however this tertiary system is rarely studied.

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Therefore, we examined the impacts of mineral surfaces (goethite and γ-Al2O3) and organic

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matter (Suwannee River Fulvic Acid, SRFA) on Fe(II) oxidation rates and the resulting Fe(III)

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(oxyhydr)oxides under 21% and 1% pO2 at pH 6. We tracked Fe dynamics by adding 57Fe(II) to

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Fe-labeled goethite and γ-Al2O3, and characterized the resulting solids using 57Fe Mössbauer

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spectroscopy. We found Fe(II) oxidation was slower at low pO2 and resulted in higher

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crystallinity Fe(III) phases. Relative to oxidation of Fe(II)(aq) alone, both goethite and γ-Al2O3

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surfaces increased Fe(II) oxidation rates regardless of pO2 levels, with goethite the stronger

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catalyst. Goethite surfaces promoted well-crystalline goethite formation, while γ-Al2O3 favored

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goethite and some lepidocrocite, and oxidation of Fe(II)aq alone favored lepidocrocite. SRFA

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reduced oxidation rates in all treatments except mineral-free systems at 21% pO2 and decreased

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Fe(III) phase crystallinity, facilitating low-crystalline ferrihydrite in the absence of mineral

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sorbents, low-crystalline lepidocrocite in the presence of γ-Al2O3, but either well-crystalline

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goethite or ferrihydrite when goethite was present. This work highlights that oxidation rate, the

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types of mineral surfaces, and OM control Fe(III precipitate composition.

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Introduction Iron (Fe) plays a major role in biogeochemical processes in soils and sediments.1 The

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transition between Fe(II) and Fe(III) is of particular interest in dynamic environments that

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experience alternating reducing and oxidizing conditions. The oxidation of Fe(II) and resultant

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precipitation of Fe(III) oxhydroxides (hereafter termed as Fe(III) oxides) with varying


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crystallinity occurs in redox dynamic environments, is a key process determining the fate of

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nutrients and contaminants.1,2 This is due to the inherent physicochemical properties of freshly

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formed oxides, which have the capacity to adsorb and/or incorporate a variety of contaminants

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and nutrients 3, 4. Fe(III) polymerization and precipitation by oxidation of Fe(II) may also trap

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large amounts of organic matter (OM) within the Fe(III) oxide structure, stabilizing the OM

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against microbial mineralization.5-9 Such iron/organic carbon associations are often invoked to

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explain the long-term stabilization of OM.10

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Fe(II) oxidation can occur through abiotic (chemical) and biotic (microbial) processes10,

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which are strongly influenced by pH and O2 concentration12-14. At circumneutral pH, abiotic

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oxidation dominates at high O2 levels, while abiotic and biotic oxidation occurs at comparable

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rates at low O2 levels.15,16 The abiotic oxidation of Fe(II) can proceed along two parallel

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pathways, one via homogeneous oxidation in solution and secondly via heterogeneous oxidation

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in association with mineral surfaces. In natural waters, OM and dissolved Fe(II) often co-exist,

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which may alter abiotic homogeneous Fe(II) oxidation kinetics by complexing both Fe(II) and

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Fe(III) species17-19. The impact of OM on O2-mediated abiotic Fe(II) oxidation is well studied

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and OM has been variably found to increase, decrease, or have no effect on the oxidation rate of

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Fe(II) in solution17-26. Previous studies of Fe(II) oxidation in the presence of OM have generally

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used high organic ligand concentrations in order to prevent the precipitation of Fe(III) oxides

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(focusing on how OM influences the rate of Fe(II) oxidation as opposed to the coupled oxidation

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and precipitation).17-26 In fact, a very limited number of studies have evaluated the impact of OM

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on the products of Fe(II) oxidation27,28. OM can interfere with the polymerization of Fe(III)

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phases, promote the formation of amorphous or less crystalline Fe(III) oxides, and substantially

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retard the Fe(II)-facilitated transformation into crystalline Fe(III) oxides27-33. Therefore,



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oxidation of Fe(II) in the presence of OM is likely to promote Fe(III) phases that are less-

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crystalline (more-reactive) than would occur in the absence of OM. This knowledge gap

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becomes even larger when we consider that neither Fe(II) oxidation rates nor the oxidation

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products have been evaluated in the common co-occurrence of Fe(II), OM, and oxide mineral

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surfaces in natural environments.

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Mineral surfaces are well known to catalyze the oxidation of Fe(II) (e.g. heterogeneous

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oxidation34-36) and most studies have used contaminants as the oxidant under O2-free (anoxic)

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conditions. For instance, the oxidation kinetics of Fe(II) on Fe/Al oxides by inorganic

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contaminant such as CrVI, AsV, HgII 37,38, or organic contaminants such as nitrobenzene and

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pertechnetate35,36,39-41. In these heterogeneous, contaminant reduction experiments, the repeated

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sorption and oxidation of Fe(II) on Fe/Al oxide surface results in the accumulation of new Fe(III)

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phases that often resemble the underlying oxide mineral surface. Larese-Casanova et al.35 for

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instance found de novo goethite forms via a templating effect on the initial parent goethite

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mineral surface and others have found hematite41and goethite36 form on α-Al2O3 surfaces. In

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natural environments, oxidant levels often fluctuate in response to rainfall, irrigation or as the

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ground water table rises and falls: in these situations, the primary oxidant is either NO3 or

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O211,42,43. Fe(II) oxidation by abiotic reaction with O2 or via Fe-oxidizing bacteria transferring

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electrons to O2 or NO3- produces a variety of secondary Fe(III) minerals, ranging from poorly-

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crystalline minerals such as ferrihydrite (Fh)42,44, to crystalline minerals including lepidocrocite

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(Lp)45-49, goethite (Gt)14,45,47,50, magnetite51 and green rust51.

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O2 concentrations are expected to control homogenous Fe(II) oxidation rates13, and

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influence the mineral composition of the resulting Fe(III) oxides14,52,53. However, it remains

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unclear how variations in pO2 levels will alter heterogeneous Fe(II) oxidation rates and the


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nature of the resulting Fe(III) oxides. There have been a few studies investigating O2-mediated

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abiotic heterogeneous Fe(II) oxidation kinetics in the presence of Fe(III) oxides at fully

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oxygenated conditions34,54,55, but these may not be applicable to lower O2-conditions. Although

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biotic Fe(II) oxidation becomes much more significant at lower O2 levels16, even in those

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systems the accumulation of biogenic Fe(III) oxides results in substantial abiotic oxidation via

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surface-catalyzed heterogeneous reactions at the periphery of the biotic reaction centers56,57. The

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nature of these newly precipitated Fe(III) oxides from heterogeneous oxidation is rarely

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characterized34, and no data appears to be available for O2-mediated Fe(II) oxidation on Al

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oxides, which likely involves different reaction kinetics than on Fe oxide surfaces as Al oxides

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do not participate in electron transfer reactions with the adsorbed Fe(II)36.

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Accordingly, our goal was to investigate Fe(II) oxidation kinetics and the resulting de

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novo Fe(III) solids in the presence of Gt, γ-Al2O3, and OM (using SRFA as a model compound)

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under varying pO2 levels, with homogenous oxidation of mineral-free Fe(II) solution as a control.

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We hypothesized that mineral surfaces would enhance Fe(II) oxidation rates at all pO2 levels and

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mediate more-crystalline Fe(III) oxide formation via templating effects, whereas the addition of

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SRFA would decrease the oxidation rates and result in less-crystalline Fe(III) oxide precipitation.

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We identify the de novo Fe(III) phases by taking advantage of the isotope specificity of 57Fe-

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Mössbauer spectroscopy to track the mineralogical evolution of the isotopically-labeled 57Fe(II)

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oxidation products on 56Fe-Gt and γ-Al2O3 phases.

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Methods

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Materials

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All treatments were prepared with O2-free ultrapure (18.2 MΩ) water prepared by

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purging the water with N2 and then storing it inside an anoxic glove box (95% N2, 5% H2) for at 5 ACS Paragon Plus Environment


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least 3 days prior to preparation of reagents. Acidic solutions of 57Fe(II) and 56Fe(II) were made

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by dissolving powders of 57Fe(0) (96% pure, Chemgas) and 56Fe(0) (99.9% pure, Isoflex),

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respectively, in anoxic 2 M HCl solutions and filtering the solutions prior to use58. Reference

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SRFA was obtained from the International Humic Substance Society and dissolved in anoxic,

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ultrapure water and stored in the glove box. 13C NMR estimates of carbon distribution of SRFA

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were included in SI Table S1.

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Fe-Gt was synthesized according to a modified method of Fh conversion59. The 56Fe(II)

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solution was removed from the anoxic glovebox and oxidized to Fe(III) by addition of hydrogen

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peroxide while stirring within a polypropylene co-polymer (PPCO) bottle. The solution was then

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filtered and the precipitates diluted with ultrapure water and the pH adjusted to 12.0 by adding

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5M KOH. The mineral suspension was then capped and aged at 70 °C for 20 d. The precipitate

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was freeze-dried following rinsing with anoxic ultrapure water and centrifuging 3 times at 20

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000 g for 10 min. The resulting precipitate was confirmed to be pure Gt by XRD. The specific

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surface area and total Fe content of the synthesized Gt is 18 m2 g-1 and 685 mg g-1, respectively.

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The synthesized 56Fe-Gt had negligible Mössbauer signal (i.e., negligible 57Fe isotope). The γ-

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alumina (Al2O3, Degussa) was purchased from Sigma-Aldrich and exhibited a SSA of 135 m2 g-1

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and contained no-detectable Fe based on acid-digestion and subsequent ICP analysis60. The γ-

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Al2O3 and synthesized 56Fe-Gt were stored in the anoxic glove box prior to use to remove any

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sorbed O2.

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Fe(II) oxidation experiments

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All Fe(II) oxidation experiments were done in triplicate and in the dark to avoid possible

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photochemical oxidation of Fe(II). Prior to the oxidation event, all the reagents and reaction

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bottles were prepared and stored in the glove-box. Experiments were conducted in 100 ml brown


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serum bottles containing 20 ml of 50 mM MES (2-(N-morpholino)-ethanesulfonic acid) buffer

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adjusted to pH 6 using anoxic NaOH or HCl solutions. Aliquots of the 57Fe(II) solution were

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added to achieve a 1.55 mM Fe(II) concentration. This concentration of 1.55 mM Fe(II) was

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chosen in order to form sufficient Fe(III) oxidation products to be detectable by 57Fe Mössbauer

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spectroscopy. This concentration is at the upper end of Fe(II) concentrations observed in many

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redox fluctuating systems such wetlands and paddy soils and is therefore justified6, 61. The solid

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phases were added at 30 mg 56Fe-Gt or γ- Al2O3 per sample for a solids loading of 1.5 g L-1.

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Anoxic SRFA solution was added to achieve a C/Fe molar ratio of 1.6 in the SRFA-containing

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reactors. This C/Fe molar ratio was chosen to facilitate Fe(III) precipitation as Fe(III)

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(oxyhydr)oxides, and it is within the range found across various natural environments6, 10, 61.

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Experiments were initiated by placing the reaction bottles on a rotary shaker in the glove box for

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0.5 h to equilibrate the added 57Fe(II) across the aqueous and solid phases under anoxic

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conditions prior to exposure to O2. The reactors were then sampled and Fe(II) oxidation was

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initiated by placing the reactors on a rotary shaker in custom-built, sealed-atmospheric chambers

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(fully contained within the anoxic glove-box) supplied with a continuous flow of either

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laboratory air (~21% O2) or a gas mixture containing 1% O2 balanced with 99% N2. Dissolved

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O2 (DO) measured by a Hach (USA) DO meter reached a maximum within 15 min exposure to

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the O2 source and remained at ~8.1 and ~0.4 mg/L for the 21% and 1% pO2 treatments

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respectively. Samples were taken from the reactors (under anoxic conditions) at various time

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intervals until Fe(II) was fully oxidized.

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Sampling and Analysis

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Aqueous samples were extracted in the glove-box using sterile syringes with 0.22-µm filters and analyzed for Fe(II) concentrations using a modified ferrozine method62. Total Fe(II)



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concentration was measured by extracting an unfiltered sample with 0.5M HCl for 2 hours63 and

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analyzing as described above. Sorbed Fe(II) was calculated based on the difference between

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HCl-extractable Fe(II) and dissolved Fe(II). Solid-phase samples were collected at the end of

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anoxic reaction, and at the end of Fe(II) oxidation event and analyzed for total C content and 57Fe

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mineral composition via Mössbauer spectroscopy. Samples for solid phase C analysis were

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freeze-dried following rinsing 3 times with anoxic ultrapure water and centrifuging at 20 000 g

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for 10 min. C content was measured with a vario Micro cube CHNS analyzer (Elementar). Solid

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samples for 57Fe Mössbauer analysis were collected after centrifuging at 20 000 g for 10 min,

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preserved between two layers of O2-impermeable Kapton tape (this step took 46 T at 5 K for Gt and Fh, and < 46 T for Lp) and QS values (near zero

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for Fh and Lp, while ~ -0.12 mm/s for Gt).

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In the OM-free systems, addition of γ-Al2O3 promoted Gt formation (CS ≈ 0.48 mm/s,

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QS = -0.08 to -0.13 mm/s, and H ≈ 47 T) in addition to Lp (CS ≈ 0.42 mm/s, QS ≈ 0.05 mm/s,

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and H ≈ 43 T), with a higher proportion of Gt at high O2 levels (Table 2; Figure 3E & 3F). In

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contrast, adding both γ-Al2O3 and SRFA facilitated the formation of Lp by inhibiting Gt

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formation relative to OM-free γ-Al2O3 system (Table 2; Figure 3G & 3H).

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In the absence of OM, Gt was the only product formed following 57Fe(II) oxidation in the

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presence of the parent 56Fe-Gt at both O2 levels (Table 2; Figure 3I & 3J). The overall parameters

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of the newly formed 57Fe-Gt (CS = 0.48 mm/s, QS ≈ -0.12 mm/s, and H ≈ 50 T) are similar to 11 ACS Paragon Plus Environment


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naturally abundant Gt standard (SI Table S2). When 56Fe-Gt and SRFA co-exists in the system, 57

Fe-Gt (CS = 0.48 mm/s, QS = -0.11 mm/s, and H = 49.7 T) was formed at low O2 levels

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(Figure 3M and Table 2), whereas high O2 levels led to the formation of 42% Fh with a QS close

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to 0 and of ~49 T and the balance Gt (Table 2; Figure 3L).

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Fe(III) oxides order magnetically and thereby transform from a doublet into a sextet in

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the MB spectrum at a characteristic measurement temperature, based on their relative

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crystallinity. Fe(III) oxides of lower crystallinity require lower measurement temperatures to

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magnetically order (and hence form a sextet) than Fe(III) oxides of higher crystallinity. In

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general, for all systems, 1% O2 led to a greater proportion of Fe(III) oxides magnetically ordered

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at higher temperature than 21% O2 oxidation, indicating that 1%-O2 oxidation led to more

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crystalline oxide formation, relative to 21% O2 (Figure 4).

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In the OM-free systems, γ-Al2O3 promoted the formation of a greater proportion of Fe(III)

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oxides ordered at or above 77K (largely attributed to Gt) (Figure 4; SI Figure S5 & S6; Table S4

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& S5) than in the mineral-free controls Figure 4; SI Figure S7 & S8; Table S6 & S7). Thus, γ-

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Al2O3 mediates the formation of more crystalline Gt relative to homogenous oxidation. The 56Fe-

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Gt enhanced this effect further as de novo 57Fe-Gt was almost completely ordered at 140K—

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similar to the parent Gt standard—and displayed the highest crystallinity of all treatments

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(Figure 4; SI Figure S9 & S10; Table S8 & S9).

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In all treatments, SRFA led to the formation of Fe(III) oxides with lower-ordering

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temperatures in the MB spectra, indicating lower crystallinity. In the system containing Fe(II)

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solution only or containing both Fe(II) and γ-Al2O3, all the resulting Fe(III) oxides in the

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presence of OM—identified as either Fh or Lp—ordered at or below 35K (Figure 4; SI Figure

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S11, S12, S13, S14; Table S10, S11, S12, S13): such solids are of extremely low crystallinity68.


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In the presence of Gt, SRFA resulted in less-crystalline oxide formation than SRFA-free reactors

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(Figure 4; SI Figure S15 & S16; Table S14 & S15).

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Discussion

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Slow Fe(II) oxidation at low pO2 results in more-crystalline Fe(III) oxide formation A number of studies have demonstrated the relative abundance of Gt and Lp produced

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from the homogenous oxidation of aqueous Fe(II) by O2 at circumneutral pH (without significant

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foreign ions) is manipulated by Fe(II) oxidation rates.52,53,69,70 Our study showed oxidizing Fe(II)

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with low pO2 was slower than with high pO2 and resulted in more crystalline solids in both the

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absence and presence of minerals. The longer exposure to sorbed Fe(II) in slow-oxidizing

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incubations (at low O2 levels) may have enhanced the recrystallization of Fe(III), relative to

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faster oxidations reactions. It is well known that, in the absence of foreign ions or OM, sorbed

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Fe(II) drives Fe mineral transformation to more crystalline Fe phases like Gt30-32,71 via the

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electron transfer between aqueous Fe(II) and Fe(III) solid phase 72-74. We found that even in the

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presence of SRFA, the Fh formed during oxidation of Fe(II) in 1% O2 treatment had a higher

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crystallinity than the Fh formed in the 21% O2 treatment. This may imply that Fe(II)-catalyzed

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changes in Fh crystallinity can occur without altering the mineral identity. Alternatively, the

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higher crystallinity resulting from the 1% O2 treatment may directly reflect faster rates of crystal

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nuclei formation, which are predicted to generate less crystalline structures 75.

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Mineral solid surfaces enhance Fe(II) oxidation rates, but drive the formation of more

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crystalline Fe(III) oxides

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Our results demonstrated that mineral surfaces strongly increased rates of O2-mediated

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Fe(II) oxidation relative to mineral-free homogenous conditions, regardless of pO2 levels. Gt was

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much more effective in catalyzing Fe(II) oxidation than γ-Al2O3, despite having a lower specific 13 ACS Paragon Plus Environment


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surface area. This is similar to O2-mediated Mn(II) oxidation in the presence of mineral surfaces,

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where Lan et al.76 found Fe(III) oxides had a greater catalytic activity than Al oxides; and during

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Cr(VI) reduction by Fe(II), where Buerge and Hug37 found Gt accelerated the reaction more than

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Al2O3, although they attributed this to higher Fe(II) sorption on Gt. In our study, oxidation rate

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constants are up to 7-fold greater for Gt than γ-Al2O3 (Table 1), whereas Gt sorbed only slightly

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more Fe(II) (~4% of total Fe(II)) than γ-Al2O3 (time = 0.5 hr, SI Figure S2). Thus we propose

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Gt’s semiconductor capabilities77, 78 facilitate electron transfer reactions that increase Fe(II)

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oxidation rates relative to those on γ-Al2O3, which has electrical insulator properties 79. In

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addition, adsorbed Fe(II) is valence stable at the γ-Al2O3 interface, whereas Fe(III) in Gt is an

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electron acceptor for sorbed Fe(II). Therefore, the Fe(III) in Gt may be able to transport electrons

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from Fe(II) to O2 when Fe(II)-Fe(III)-O2 complexes form at the surface, similar to electron

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transport in As(III)-Fe(II, III)-O2 complexes 80.

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Mineral surfaces increased Fe(II) oxidation rates relative to homogeneous rates, which

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should favor the formation of lower crystallinity Fe(III) oxides 75. However, the presence of γ-

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Al2O3 resulted in the formation of higher crystallinity Gt (e.g., higher MB ordering temperature)

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at the expense of Lp, while the presence of 56Fe-Gt promoted an even higher crystallinity 57Fe-Gt,

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which resembled the naturally abundant Gt standard. Templating effects likely play an important

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role here, though unfortunately they have only been investigated for adsorbed Fe(II) on Fe(III)

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oxides under anoxic conditions35. Our work suggests the formation of higher crystallinity Gt may

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be caused by repeated sorption and oxidation of Fe(II) on parent Gt surfaces, which likely

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provide a template for growth of a new phase.

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The potential for a templating effect is perhaps best evaluated in the presence of γ-Al2O3, which promotes Gt formation, in contrast to Lp formation during homogenous oxidation.


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Peretyazhko et al.41 found α-Al2O3 provided a template for the formation of α-Fe2O3 (hematite)

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during the oxidation of adsorbed Fe(II) by either Tc(VII) or O2. They attributed the formation of

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α-Fe2O3 to the fact that α-Al2O3 and hematite (α-Fe2O3) are isostructural and have similar

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interatomic distances for surficial oxygen atoms. If this is true, γ-Al2O3 should provide a

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template for the formation of isostructural γ-Fe2O3 (maghemite), yet this phase is typically

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formed via dehydration mechanisms or under hydrothermal conditions52. Instead, γ-Al2O3

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evidently favors linked corner-sharing rows of Fe(III)-hydroxide octahedral, which are

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precursors to Gt crystals. Subsequent sorption of Fe(II) on the precursor Gt nuclei may favor

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further Gt growth rather than γ-Fe2O3. A similar observation was made by Larese-Casanova et

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al.35, during oxidation of Fe(II) by nitrobenzene on α-Fe2O3 (hematite), which resulted in the

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formation of α-FeOOH (Gt) as opposed to α-Fe2O3. This Gt-favoring effect was also observed

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during heterogeneous Fe(II) oxidation on α-Al2O3 by nitrobenzene36. Taking together, it suggests

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Gt may be a common product of heterogeneous oxidation.

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OM decreases Fe(II) oxidation rates, but promotes less-crystalline Fe(III) oxide formation

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Previous studies have observed that the addition of SRFA at high C/Fe ratios (> 20)

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increased Fe(II)aq oxidation rates relative to organic-free Fe(II)aq oxidation rates.17, 21, 81 At these

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high C/Fe ratios significant Fe(II) complexation via SRFA occurs and the dissolved Fe(II)-SRFA

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complexes oxidize more rapidly than the free Fe(II)aq aquo-ion.17 Consistent with these previous

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studies, we found SRFA accelerated Fe(II) oxidation when added at a high C/Fe ratio of 15 (SI

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Figure S17), and this high C/Fe ratio also completely prevented the Fe(III) solid precipitation

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that occurs in the SRFA-free control.

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However, when we conducted the same experiments at lower C/Fe ratios that allowed

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Fe(III) precipitation in both treatments, we found quite different results. At the low C/Fe ratio

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(1.6) used for our primary experiments, we found that SRFA had a minor impact on Fe(II) 15 ACS Paragon Plus Environment


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oxidation rates at 21% O2, but substantially retarded Fe(II) oxidation at 1% O2 (Figure 1). At this

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low C/Fe ratio (1.6), Fe(II)-SRFAaq complexes likely represented only 8% of total aqueous Fe(II)

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prior to the oxidation (Visual Minteq, SI Table S16), thus Fe(III) solids were precipitated during

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oxidation in both the SRFA-free as well as the SRFA-addition treatment (Figure 3a-d and Table

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2), which should facilitate autocatalytic Fe(II) oxidation via surface catalysis13, 34 in both

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treatments. Evidently, surface adsorbed or coprecipitated SRFA on Fe(III) (oxyhydr)oxides (SI

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Table S17) decreased Fe(II) oxidation kinetics relative to SRFA-free Fe(III) (oxyhydr)oxides (SI

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Figure S18). We propose that carboxyl group of SRFA (SI Table S1) can bind to Fe(III) solids

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via ligand exchange 8, 82, prevent the inner-sphere complex formation between Fe(II) and

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bridging oxo-groups on oxide surface83 and hence reduce Fe(II) oxidation, since the oxidation

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rate of oxide surface-complexed Fe(II) species like FeOFe+ and ≡FeOFeOH0 are expected to be

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several orders of magnitude faster than that of dissolved Fe(II)34, 83, 84. This is consistent with our

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observation in the 1% O2 treatment that Fe(II) was oxidized much more rapidly in the absence of

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SRFA than in the SRFA addition treatment. Attenuation of the Fe(II) oxidation rate by the

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coprecipitated/adsorbed SFRA is still evident in the 21% O2 treatment, but the impact is very

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small (SI Figure S18A), likely because the reaction rates are overall much higher.

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When mineral surfaces were present initially, SRFA largely decreased Fe(II) oxidation

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rates at both pO2 levels. SRFA could lower heterogeneous Fe(II) oxidation rates in three ways: (1)

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SRFA could block part of the reactive mineral surface, which is supported by our findings that

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OM readily adsorbed to the Fe/Al oxides during the 0.5 h anoxic period (SI Table S17); (2),

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SRFA could coprecipitate with or adsorb on the newly formed Fe(III) solids (SI table S17) and in

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turn reduce the auto-catalytic Fe(II) oxidation; and (3) a portion of the sorbed Fe(II) could be

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preferentially complexed by carboxyl moieties in SRFA31,32, reducing the extent of Fe(II)


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adsorption on the mineral surfaces and for Gt, the amount of Fe(II)-Fe(III) electron transfer and

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therefore decreasing Fe(II) oxidation via this route. This scenario is supported by the appearance

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of a Fe(II) doublet in Mössbauer spectra of Fe(II)-reacted Gt-SFRA sample prior to oxidation

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(Figure 2b), which have QS values consistent with Fe(II) sorbed to OM67.

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Differentiating the types of OM-Fe co-precipitates formed across a matrix of low and

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high pO2, and homogeneous vs. heterogeneous oxidation fills a key knowledge gap of carbon

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behavior in environmental systems. Most studies synthesize Fe-OM co-precipitates via

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hydrolysis of Fe(III)-OM solutions5,7,8. Yet, in the environment there is ample evidence

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suggesting the main route of Fe-OM formation involves oxidation of Fe(II) and often in the

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presence of reactive Fe and Al surfaces6,9,85,86. Thus far, our conceptual understanding of Fe-OM

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phases is largely based on synthetic techniques that poorly approximate the likely natural

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synthesis routes. We find that despite reducing the oxidation rates—which should favor more

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crystalline Fe phases 75—OM resulted in the formation of less-crystalline Fh and Lp phases in all

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treatments, except for Fe(II) oxidation at 1% pO2 with 56Fe-Gt.

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In the mineral-free case, SRFA (initial C/Fe ratio= 1.6) leads to the co-precipitation of

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OM and Fh with a C/Fe ratio of ~1.2 in the coprecipitates (SI Table S17). This OM-Fh co-

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precipitate had a particularly low crystallinity and did not transform to a more stable Fe phase

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during the experiment. This agrees with earlier studies that OM can completely inhibit Fh

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conversion to more crystalline phases when the C/Fe ratio is >131, 32. Previous studies also

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observed the formation of Fh following oxidation of ferrous perchlorate in the presence of citrate,

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a low molecular weight organic acid at high citrate/Fe ratios.27, 87 These low-crystallinity OM-Fh

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co-precipitates can potentially stabilize OM and provide a positive feedback that increases

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carbon storage in natural environments.



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In the presence of γ-Al2O3, a large proportion of the SRFA was adsorbed onto γ-Al2O3

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surface prior to oxidation (SI Table S17). This may have eliminated some of the OM from the

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co-precipitation process as the resulting newly precipitated Fe(III) (oxyhydr)oxides had a low

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OM content (C/Fe ratio ~0.3, SI Table S17). Prior work has shown that low C/Fe ratios favor the

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formation of Lp over Fh and Gt32. Using the MB ordering temperature as a measurement of the

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Fe phase crystallinity67, we find the OM-Lp formed in presence of γ-Al2O3 at 21% O2 had a

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similarly low crystallinity to OM-Fh co-precipitates formed in homogeneous oxidation (Figure

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4). Traditionally Lp is considered a more thermodynamic stable phase (i.e., higher crystallinity)

404

than Fh, but it may be that OM can impart similar low crystallinity across a variety of phases,

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which might be similarly reactive—toward sorption, electron transfer reactions—as Fh.

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In the presence of Gt, pO2 levels had a strong effect on the resulting OM-Fe precipitates,

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with 21% O2 resulting in the formation of Fh and some Gt, while at 1% O2, only Gt was formed

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(C/Fe ~0.4) (Table 2 and SI Table S17). As the low-surface area Gt sorbs little SRFA prior to

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oxidation (SI Table S17), a large amount of the OM in the co-precipitates formed at 21% O2 is

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likely being associated with the 42% of the Fe(III) that precipitated as Fh (C/Fe ratio ~0.6, SI

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Table S17; Table 2). High C/Fe ratios in the OM-Fh co-precipitate coupled with a short exposure

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to Fe(II) (

The Effect of Fe(III) - ACS Publications - American Chemical Society

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