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Synergistic Effect of Reductive and Ligand Promoted Dissolution of Goethite Zimeng Wang, Walter D. C. Schenkeveld, Stephan M. Kraemer, and Daniel E. Giammar Environ. Sci. Technol., Just Accepted Manuscript • Publication Date (Web): 12 May 2015 Downloaded from http://pubs.acs.org on May 12, 2015

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Environmental Science & Technology

Synergistic Effect of Reductive and Ligand Promoted Dissolution of Goethite Zimeng Wang†*, Walter D.C. Schenkeveld‡, Stephan M. Kraemer‡, and Daniel E. Giammar§ † Department of Civil and Environmental Engineering, Stanford University, Stanford, CA, United States ‡ Department of Environmental Geosciences, University of Vienna, Vienna, Austria §Department of Energy, Environmental and Chemical Engineering, Washington University in St. Louis, St. Louis, MO, United States

*Corresponding author Jerry Yang and Akiko Yamazaki Environment & Energy Building 473 Via Ortega Stanford, California 94305, United States Email: [email protected] Web: www.stanford.edu/~wangzm

Manuscript Submitted to Environmental Science & Technology May 2015

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ACS Paragon Plus Environment


1

Abstract

2

Ligand promoted dissolution and reductive dissolution of iron (hydr)oxide minerals control

3

the bioavailability of iron in many environmental systems and have been recognized as

4

biological iron acquisition strategies. This study investigated the potential synergism between

5

ligands (desferrioxamine B (DFOB) or N,N'-Di(2-hydroxybenzyl)ethylenediamine-N,N'-diacetic

6

acid (HBED)) and a reductant (ascorbate) in goethite dissolution. Batch experiments were

7

performed at pH 6 with ligand or reductant alone and in combination and under both oxic and

8

anoxic conditions. Goethite dissolution in the presence of reductant or ligand alone followed

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classic surface controlled dissolution kinetics. Ascorbate alone does not promote goethite

10

dissolution under oxic conditions due to rapid reoxidation of Fe(II). The rate coefficients for

11

goethite dissolution by ligands are closely correlated with the stability constants of the aqueous

12

Fe(III)-ligand complexes. A synergistic effect of DFOB and ascorbate on the rate of goethite

13

dissolution was observed (total rates greater than the sum of the individual rates), and this effect

14

was most pronounced under oxic conditions. For HBED, macroscopically the synergistic effect

15

was hidden due to the inhibitory effect of ascorbate on HBED adsorption. After accounting for

16

the concentrations of adsorbed ascorbate and HBED, a synergistic effect could still be identified.

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The potential synergism between ligand and reductant for iron (hydr)oxide dissolution may have

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important implications for iron bioavailability in soil environments.

19 20

Introduction

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Iron is an essential micronutrient for plants and microorganisms. In calcareous soils and

22

marine waters, the solubility of ferric (hydr)oxides can only provide iron concentrations that are

23

far below the needs of most plants, fungi, and microorganisms. Organisms have developed

24

strategies to cope with conditions of low iron availability and to enhance iron solubility. Given

25

that the proton-promoted dissolution of iron (hydr)oxides is very slow at circumneutral pH,1, 2

26

mobility of iron in aquatic environments can be enhanced by (1) reduction of Fe(III) to the more

27

soluble form of Fe(II)3-8 and (2) formation of soluble Fe(III)-complexes with chelating ligands.

28

In oxic environments at circumneutral pH, Fe(III) (hydr)oxides are the thermodynamically

29

stable forms of iron. Fe(II), even if generated by a surface-acting reductant, cannot persist in the

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oxic solution9, 10 unless associated with certain specific Fe(II) binding ligands11 Iron is more

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effectively mobilized through complexation with high affinity organic ligands. In agriculture,

32

corrective measures against iron limitation involve the application of iron chelated with synthetic

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ligands such as EDDHA (ethylenediamine-N,N'-bis2-hydroxyphenylacetic acid).12,

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13

Such


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ligands can mobilize Fe and other metals from soil,14, 15 but the properties of environmental

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surfaces can strongly affect the kinetics of their reactions with metals.16 Similarly, in natural

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environments, an efficient biological iron acquisition strategy of bacteria, fungi, and

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graminaceous plants involves the exudation of siderophores.17-20 Siderophores (e.g.,

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desferrioxamine B, DFOB) are biogenic iron-binding ligands that are released under iron-

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deficient conditions. For typical hydroxamate siderophore concentrations in soil (10 to 100

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nM),21, 22 iron oxide solubility increases over a wide pH range. Above pH 4, ligand-promoted

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dissolution is faster than proton-promoted dissolution.23-26 Previous studies have provided

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spectroscopic evidence of siderophore surface complexes on iron oxides27,

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application of the ligand promoted dissolution mechanism to describe siderophore-iron oxide

44

interaction. Because an Fe(III)-siderophore surface complex is a precursor of the intermediate

45

species in the rate-determining step in ligand controlled dissolution, the dissolution rate is a

46

linear function of the adsorbed ligand concentration29 and of the solution saturation state30,

47

∆

( )

= ∆ = 1 − exp = !1 − "$ ' * #

%&

28

, allowing the

(1)

48

where RL[nmol g−1h−1] is the ligand controlled dissolution rate, kL [h−1] is the first order rate

49

coefficient, [L]ads [nmol g−1] is the adsorbed ligand concentration, ∆G is the Gibbs free energy of

50

reaction, R is the gas constant, T is the absolute temperature, Q is the activity quotient, Keq is the 3


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equilibrium constant of the overall reaction, and σ is Temkin’s coefficient31 (usually 1 to 3 for

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metal oxide dissolution) to account for the non-elementary nature of a reaction (for an

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elementary reaction the value is 1).31, 32

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In addition to having low concentrations in the environment, hydroxamate siderophores also

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do not adsorb strongly to iron oxide surfaces.23 Therefore, it has been proposed that siderophores

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act in conjunction with other molecules to increase iron bioavailability.33-35 Low molecular

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weight organic acids (many of which bind Fe(III) relatively weakly) are ubiquitous in soil, and

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can have synergistic effects on siderophore-promoted dissolution of iron oxides, where the net

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dissolution rates in the presence of both ligands are beyond the sum of the net dissolution rates

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observed in single ligand systems.36 This synergistic effect has been documented for various

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organic acids, different types of siderophores, and several iron(III) minerals,24, 26, 34, 37-40 with

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several possible mechanisms proposed. First, the strong affinity of the siderophore for Fe(III)

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increases the Fe(III) solubility and shifts the solution saturation state to maintain far from

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equilibrium conditions (i.e., f(∆G) →1).34, 38 Second, organic ligands can form kinetically labile

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iron complexes on the mineral surface, which can be rapidly extracted by the siderophore at

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higher rates than pristine Fe(III) sites.24, 37 For experiments with oxalate as a complexing ligand,

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the presence of a labile pool of Fe(III)-oxalato ternary surface complex was validated by isotope

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dilution experiments24, kinetic modeling40 and spectroscopy28, 41. Third, enhanced adsorption of

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certain positively charged siderophores, like DFOB, may result from a surface charge shift

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induced by adsorption of negatively charged compounds. This mechanism was used to interpret

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the synergistic effect of fulvic acid39 and anionic surfactant35 on DFOB-promoted dissolution of

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goethite.

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It has been suggested that extracellular reductants also contribute to iron bioavailability.33, 42

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Under iron-limited conditions, microorganisms and plant roots are known to release compounds

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with reducing capacity, even when the surrounding environments are oxic19, 42, 43. Exogenous

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reductants include pyridine-2,6-bis(monothiocarboxylate) (PDTC),44 ascorbate33 and various

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phenolic compounds (e.g., caffeic acid).19 However, the geochemical role of reductants in

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mediating iron bioavailability in oxic environments is not clear. In the presence of O2 at

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circumneutral pH, any reduced Fe(II) on the iron oxide surface is rapidly reoxidized.1,

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Nonetheless, several microbiological studies demonstrated a substantial role of a chemical

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reductant in iron acquisition by aerobic microorganisms.33,

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bioavailability based on cell growth, and their interpretation of the role of a reductive pathway

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was largely based on cell membrane and intracellular processes such as (1) the direct uptake of

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transient Fe(II) from the iron oxide surface by cells in very close proximity or direct contact47-49

45, 46

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9, 10

These studies assessed iron

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and (2) ascorbate acting as a reductant to “assist iron passage through the membrane

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transporter”33.

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The interplay of reductant and ligand on iron oxide dissolution is a classical research topic

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that began with pioneering studies by Stumm and colleagues. They suggested that ascorbate and

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oxalate/malonate in combination could dissolve iron oxides at rates beyond the sum of the

90

individual rates.3,

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together with oxalate and postulated a catalytic role of Fe(II)-oxalate complexes in acting as a

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surface-acting reductant for Fe(III).4 This pathway was observed to be faster than non-reductive

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oxalate-promoted iron oxide dissolution. As opposed to soil relevant oxic and neutral conditions,

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these investigations were performed under anoxic conditions at pH 3 or lower, where the

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reductive dissolution pathway itself can generate substantial soluble Fe. Dhungana et al. studied

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ferrihydrite dissolution at pH 7.5 by dipicolinic acid (a ligand) and H2S (a reductant), proposing

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a catalytic role of the in situ generated Fe(II).50 However, it is still not clear whether oxygen

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would compromise the reductive pathway and how it would affect the dissolution kinetics. A

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synergistic effect between photoreductive and ligand promoted dissolution of iron (hydr)oxides

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was reported in the context of oxic marine systems.51-53 Under UV irradiation, adsorbed oxalate

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can act as the electron donor to form surface Fe(II), and the siderophore then acts as a shuttle to

7

They later observed increased dissolution rates when Fe(II) was added

6



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transfer surface Fe(II) into solution. This synergistic effect has been demonstrated in irradiated

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systems but not in experiments performed in the dark.

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Building on the existing knowledge of iron oxide dissolution mechanisms, the objective of

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the present study was to systematically evaluate the interplay of reductive and ligand-promoted

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dissolution of goethite under pH conditions relevant to soils. This interplay can have implications

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for iron bioavailability to bacteria, fungi and plants. The central hypothesis was that a synergistic

108

effect can enhance iron mobility, especially under oxic conditions. To test the hypothesis,

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dissolution experiments were performed under anoxic and oxic conditions with varied reductant

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and ligand concentrations.

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Materials and Methods

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Materials. Goethite (α-FeOOH) was prepared according to the standard synthesis procedure

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in Cornell and Schwertmann.2 Briefly, KOH was added to Fe(NO3)3 solution with vigorous

115

stirring. The initial suspension of red-brown ferrihydrite was heated at 70 °C for 60 h to produce

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goethite. Excess reactants were removed by repeated dialysis against deionized water until the

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conductivity of the wash water was below the detection limit of 1 µS/cm. After the final dialysis

118

step, the product was centrifuged and stored as a concentrated slurry at 4 °C. Polyethylene

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containers were used for preparation and storage. X-ray powder diffraction (Rigaku

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GeigerflexD-MAX/A) confirmed the identity of the material as goethite. The specific surface

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area was determined to be 38 m2/g by a multipoint N2-BET adsorption method (Autosorb-1-

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C,Quantachrome). The L-ascorbic acid (C6H8O6, Sigma Aldrich, >95%), desferrioxamine B

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mesylate salt (C25H48N6O8·CH4O3S, DFOB, Sigma Aldrich, >92.5%) and N,N'-Di(2-

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hydroxybenzyl)ethylenediamine-N,N'-diacetic

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(C20H24N2O6•HCl•H2O, HBED, Strem Chemical Inc. >98%) were used as received without

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further purification. Ascorbate is a common biogenic reductant in soils, and ascorbate is a very

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weak ligand for iron that cannot mobilize Fe through ligand promoted dissolution (Figure S1b of

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the Supporting Information). DFOB is a microbial hydroxamate siderophore, and HBED is used

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in synthetic iron fertilizers for calcareous agricultural soils. Given the concern of possible

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degradation of these compounds in solutions, all the chemicals were kept in dry powder forms

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(as received) until the goethite dissolution experiments were performed. Ultrapure water

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(resistivity >18.2 MΩ·cm, Milli-Q, Millipore) was used in all experiments.

acid

monohydrochloride

hydrate

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Goethite Dissolution Experiments. The experimental conditions are summarized in Table 1.

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All experiments were performed at a constant temperature (22 °C) with reactors shielded with

135

aluminum foil to avoid potential photochemical redox reactions. Dissolution experiments were

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conducted in completely mixed batch reactors (100 mL) using magnetic stirring. A fixed ionic

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strength (10 mM) was provided by sodium chloride. A pH of 6.0 that is relevant to soil

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environments was set for all experiments. The pH was buffered by 1 mM 2-(N-

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morpholino)ethanesulfonic acid (MES, as sodium salt, pKa = 6.15), a widely used pH buffer

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verified to be applicable for goethite dissolution studies.39, 54 The solution pH was largely stable

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during the experiments (6.0 ± 0.05), and the pH was adjusted to this range using concentrated

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HCl or NaOH. Oxic experiments were performed using ultrapure water sparged with air (PO2 =

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0.21 bar) and with the headspaces of the reactors connected to the atmosphere. Anoxic

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experiments were performed in an anaerobic chamber (Coy Laboratory MI), where the gas

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composition was controlled to less than 1 ppm for O2 and between 2% and 5% (v/v) for H2. The

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ultrapure water used in these experiments was sparged by H2 with a Pd catalyst in the water. The

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reactors and solutions were equilibrated with the anoxic atmosphere in the chamber overnight or

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longer prior to the goethite dissolution experiments under anoxic conditions. All the experiments

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started with 1 g/L goethite added to freshly prepared solutions containing aliquots of ligand

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and/or reductant and with preadjusted and buffered pH. Samples were periodically collected,

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filtered with 0.1 µm poly(ether sulfone) (PES) syringe filters (Millipore), and preserved in 2%

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HNO3 for dissolved Fe analysis.

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Adsorption experiments. Adsorption experiments quantified the adsorption of reductant and

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ligands to the goethite surface. All adsorption experiments were conducted under the ambient

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atmosphere using batch reactors (50 mL) that were completely mixed on a shaker. The goethite

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concentration, ionic strength, pH, and buffer were selected to be consistent with those of the

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dissolution experiments. Wide ranges of ascorbate, DFOB and HBED loadings were equilibrated

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with 1 g/L goethite for 2 hours, which was found to be sufficient for reaching adsorption

159

equilibrium while having minimal goethite dissolution. Filtered (0.1 µm) samples were analyzed

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within 1 hour after filtration for the dissolved reductant and ligand concentrations.

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Aqueous Analysis. The pH of solutions and suspensions was measured with a glass pH

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electrode (Accumet). Dissolved Fe concentration was measured by inductively coupled plasma

163

mass spectrometry (ICP-MS, PerkinElmer ELAN DRC II). The detection limit for Fe was 10 µg

164

L-1 (0.18 µM). It should be noted that ICP-MS is not able to differentiate between the oxidation

165

states of dissolved Fe.

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In adsorption experiments, the dissolved concentrations of ascorbate, HBED and DFOB were

167

analyzed using a UV-Vis spectrophotometer (PerkinElmer-Lambda XLS). Absorbances of Fe-

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DFOB and Fe-HBED complexes were measured at 439 and 493 nm, respectively. These peaks

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were confirmed to be free of interference with the ascorbate peak. For ascorbate, the original

10



170

filtrate samples were measured for absorbance at 243 nm.3,

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DFOB or HBED in the filtrate was quantified by measuring the absorbance of Fe-DFOB or Fe-

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HBED complexes upon addition of a fixed concentration of Fe(III) that is in excess of the

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ligand.55 This method was successfully used in previous studies on ligand promoted goethite

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dissolution34,

175

described in the Supporting Information). The modified method was verified to be effective; the

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measured concentration correlated remarkably well (R2> 0.999) with the values obtained using

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the classical method (Figure S2 of the Supporting Information).

37, 56

4

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The concentration of dissolved

. We modified this method to overcome the ascorbate interference (details

178

The MES/NaCl/pH 6 solution was used as the blank and its spectrum referenced to DI water

179

does not have peaks overlapping those of interest. The absorbance of the original filtrate samples

180

without added Fe(III) was also measured at 439 or 493 nm to account for the DFOB in the

181

solution that was already complexed to Fe(III) released during the two-hour adsorption

182

experiments. Due to the short contact time for the adsorption experiments, the absorbance for the

183

original filtrate was verified to be negligible relative to that after the addition of excess Fe(III).

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The adsorbed ascorbate concentration, DFOB and HBED were calculated by dividing their

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concentration loss by the goethite concentration.

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Rate Calculation and Modeling. The dissolution rates of goethite were estimated by linear

187

regression of the dissolved Fe concentration versus time for the first three samples unless

188

otherwise noted. The regression was not forced through the origin, and a positive intercept was

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due to the fast initial dissolution, which is commonly observed in batch experiments of iron

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(hydr)oxide dissolution, possibly from small amounts of labile Fe species.24, 51, 57 After the rapid

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initial dissolution, the regression of dissolved iron versus time is linear as was anticipated for

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dissolution reactions far from equilibrium.58, 59 Caution was taken in the data selection for rate

193

calculations to make sure that the dissolution rate was captured without being affected by

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significant decrease of the free ligand concentration; this data selection allows the solution

195

saturation term in Eq. 1 to be ignored in the subsequent analysis. The fraction of the initial

196

goethite that was dissolved during the experiment was so small that the mass normalized

197

dissolution rate was calculated using the initial goethite concentration. Rate coefficients were

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obtained by linear correlation of dissolution rates versus adsorbed reductant or ligand

199

concentrations. Chemical equilibrium modeling was performed using MINEQL+ in which

200

activity coefficients were calculated by the Davies equation.60

201 202

Results and Discussion 12



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Adsorption Experiments. The adsorption isotherm of ascorbate suggests that a 1 mM total

204

ascorbate concentration already approached surface saturation (Figure S3a of the Supporting

205

Information). The addition of up to 80 µM DFOB or HBED could displace at most 26% of the

206

total adsorbed ascorbate (Figure S3b,c of the Supporting Information). All the ligand

207

concentrations used in this study (up to 20 µM for DFOB, and up to 10 µM for HBED) fell

208

within the initial steep range of their adsorption isotherms (Figure 1). Based on the initial slopes

209

of the isotherm data, adsorption of DFOB and HBED appears less favorable than adsorption of

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ascorbate.

211

Ascorbate and ligands were expected to compete for surface adsorption sites, and they can

212

also impact the adsorption equilibrium of each other by altering the surface charge of goethite.

213

The addition of 1 mM ascorbate decreased the adsorption of HBED to goethite, while it did not

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cause a substantial change in the DFOB adsorption. The difference between the effect of

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ascorbate on DFOB and HBED adsorption may be caused by differences in the charges of their

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surface species.

217

(H4DFOB+), while the dominant HBED species is negatively charged (H3HBED−) (Figure S4). If

218

the adsorbed forms of DFOB and HBED follow a similar trend, with the DFOB surface

219

complexes being either neutral or positively charged, then adsorption of ascorbate on goethite at

At pH 6, the dominant dissolved DFOB species is positively charged

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this pH (where the goethite surface has a positive surface charge in the absence of any adsorbing

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species), would decrease the positive charge of the goethite surface, making adsorption of the

222

negatively charged HBED less favorable and adsorption of the positively charged DFOB more

223

favorable.

224

Reductant or Ligand in Isolation: Control Experiments. The results from the control

225

experiments (Exp. 1-12) in which either a reductant or a ligand was added are consistent with

226

classical metal (hydr)oxide dissolution theories. As expected from solubility calculations, in the

227

absence of reductant and ligand, goethite dissolution (Exp. 1, 4) was too slow to be detectable,

228

regardless of whether O2 was present. Ascorbate could promote goethite dissolution under

229

anoxic condition (Exp. 2, 3). The established reductive dissolution mechanism, in which the

230

detachment of Fe(II) is the rate determining step, describes the dissolution rate as a linear

231

function of adsorbed reductant.4 The dissolution rates are linearly correlated with the surface

232

concentration of adsorbed ascorbate (Figure 2a, kascorbate = 5.7×10−4 h−1). However no net

233

dissolution of goethite by ascorbate was observed under oxic conditions since the dissolved iron

234

concentrations never exceeded the the detection limit (Exp. 5, 6). When dissolved O2 is available,

235

the surface Fe(II) generated by ascorbate does not enter the solution since the heterogeneous

236

oxidation of Fe(II) is very fast.

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Goethite dissolution by DFOB and HBED was not impacted by the presence or absence of

238

oxygen in the absence of a reductant (Exp. 7 and 9, 11 and 12). The lack of an effect of O2 on the

239

goethite dissolution rate in the presence of DFOB or HBED indicates that the process did not

240

involve any substantial release of Fe(II) as in a reductive dissolution mechanism. Some previous

241

research has seen evidence for hydrolysis and oxidation of siderophores on the goethite surface28,

242

50

243

enough extents to influence the measured Fe release. The dissolution rates were linearly

244

proportional to the adsorbed DFOB concentration (Figure 2b), which was consistent with the

245

ligand-promoted dissolution mechanism. HBED is expected to follow the same mechanisms,

246

although we only measured dissolution at one HBED loading. The rate coefficient (as in Eq. 1)

247

for DFOB (0.0089 h−1) was smaller than that for HBED (0.027 h−1) by a factor of three. At pH 6,

248

HBED is a stronger ligand for Fe(III) than DFOB (Figure S5 of the Supporting Information).

; however, if such reactions occurred in the present experiments, then they were not to great

249

The rate coefficients for goethite dissolution promoted by the ligands studied in this work

250

and in previous studies that were carried out under similar conditions (pH 6, etc.) show a

251

remarkable correlation with the stability constants of the 1:1 fully deprotonated Fe(III)-ligand

252

complex (Figure 3). It has been previously reported that dissolution rates increase with

253

increasing complex stability constants.61, 62 Our present analysis compiled the rate coefficients

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254

instead of the rates as the other studies did, which allowed us to exclude the variability in ligand

255

concentration and the different adsorption affinities of the ligands and possible thermodynamic

256

constraints (See the terms in Eq. 1). The correlation of rate coefficients with stability constants

257

suggests that the binding strength of the ligand to Fe in solution is related to the rate-limiting step

258

in ligand-promoted dissolution. Deeper interpretation of this correlation is hampered by the

259

unknown surface speciation of the adsorbed ligand in general and the coordinative environment

260

of the precursor of the rate determining step in particular.

261

Goethite Dissolution: Ascorbate + DFOB. Comparing the macroscopic dissolution rates of

262

goethite in the presence of ligand and reductant together with the rates in the presence of ligand

263

or reductant individually, a synergistic effect of ascorbate and DFOB on goethite dissolution was

264

observed. This effect was supported by goethite dissolution rates obtained from various

265

combinations of ascorbate and DFOB concentrations, both under oxic and anoxic conditions

266

(Figures 4 and 5). With the adsorption data and kinetic model, the synergistic effect could not be

267

explained merely by the change of the concentrations of adsorbed ligand or reductant (Table 1).

268

Under anoxic condition, iron was released as both Fe(III) (complexed by DFOB) and Fe(II)

269

(either as a free ion or complexed by DFOB with weaker affinity than to Fe(III)63). While

270

previous studies indicated that it is possible for Fe(II)-DFOB complexes to decompose with the

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DFOB’s hydroxamate group reduced to amide group and Fe(II) oxidized to Fe(III),64, 65 any such

272

auto-decomposition either did not occur or was negligible in the time scale of the present

273

experiments.

274

The synergistic effect was particularly strong under oxic condition (Figure 5). Ascorbate

275

alone did not lead to any goethite dissolution, but its addition together with DFOB resulted in

276

higher dissolution rates than with DFOB alone. The rates under oxic conditions were still lower

277

than those under anoxic conditions. During the experiment, ascorbate was consumed by

278

oxidation at the iron oxide surface and possibly by a reaction with oxygen. However, considering

279

the high excess of ascorbate in experiments relative to observed net iron reduction under anoxic

280

conditions, this may not significantly affect surface concentrations of ascorbate, particularly at 1

281

mM total ascorbate concentrations where adsorption densities appear close to surface saturation

282

(Figure S3a of the Supporting Information). Indeed, near constant dissolution rates throughout

283

the experiments indicate that this effect is small enough to not influence dissolution (Figure 5a

284

and b). Non-linear dissolution was only observed at the lowest DFOB concentration (5 µM).

285

Here we see that the dissolution occurred so fast that it approached equilibrium before the end of

286

the experiment. Theoretically 5 µM DFOB could solubilize 5 µM Fe(III) assuming that there is

287

no adsorption of Fe-DFOB complexes, but the dissolution rate decreased drastically as the

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system approached equilibrium and the concentration of uncomplexed free DFOB in solution

289

decreased (Eq. 1). Therefore, only the results from the first two samples were used for rate

290

calculations.

291

Experiments with varying concentrations of ascorbate and DFOB had two noteworthy trends.

292

First, for a fixed DFOB concentration, the synergistic effect in Fe mobilization increased with

293

increasing ascorbate concentration. Second, initial goethite dissolution rates in the presence of

294

both ascorbate and DFOB were independent of DFOB concentrations when the ascorbate

295

concentration was at a fixed value (Figure 5a to c). The second phenomenon was intriguing as it

296

suggests that the synergy may occur in the concentration ranges relevant in soils, where

297

reductants and siderophore are in submilli-molar and submicro-molar ranges, respectively.22, 66-68

298

Goethite Dissolution: Ascorbate + HBED. The interplay between ascorbate and HBED in

299

goethite dissolution was controlled by their interactive adsorption behavior. Under anoxic

300

conditions, the extent of goethite dissolution in the presence of both 1 mM ascorbate and 10 µM

301

HBED was marginally higher than the sum of the dissolution from the experiments with

302

reductant and ligand in isolation (Figure 6a). Taking into account the competitive adsorption

303

effect (i.e., ascorbate substantially displacing adsorbed HBED; Figure 1b) and using the

304

established surface-controlled dissolution rate model, the measured dissolution rate was a factor

18



Page 20 of 39

305

of two higher than the predicted one. This suggests an effect that is more than simply addition of

306

the two independent surface-controlled dissolution mechanisms. It may indicate that there is a

307

synergistic effect of ascorbate and HBED on the goethite dissolution rate coefficient, even

308

though the macroscopic effect is obscured by the influence of the competitive adsorption of

309

ascorbate and HBED on the net dissolution rate. Under oxic conditions, the addition of 1 mM

310

ascorbate even decreased the goethite dissolution rate compared with the rate for 10 µM HBED

311

only (Figure 6b). The adsorption data (Figure 1b) provide a reasonable explanation as the

312

reductant-ligand-surface interaction was dominated by the strong inhibitory effect of ascorbate

313

on HBED adsorption.

314

While the specific mechanisms of how adsorption of a reductant inhibits that of a ligand (or

315

vice versa) remain to be elucidated, the results from the ascorbate-HBED-goethite system

316

highlighted the possibility of intrinsic synergistic effects being obscured on a macroscopic scale

317

by adsorption interactions. In a slightly different but also environmentally relevant system

318

(oxalate-siderophore), Akafia et al. found that dissolution rates could increase, decrease, or

319

remain unchanged as compared to the rates in single-ligand systems, depending on the pH,

320

siderophore present, and the identity of the metal oxyhydroxide.62 Although no adsorption results

19


Page 21 of 39


321

were presented in that study, the authors suggested a possible role of decreased ligand adsorption

322

in accounting for the diminishing synergistic effect under certain conditions.

323

Possible Mechanisms for the Synergistic Effect. Excluding the effect of enhanced

324

adsorption of the ligands or reductants, there are at least three possible mechanisms that may

325

contribute to the observed synergistic effect, although our results based on macroscopic

326

measurements of dissolution rates and adsorption equilibrium are inconclusive regarding the

327

specific mechanism. The first potential mechanism is that the generation of surface Fe(II)

328

labilizes the goethite structure. Although ascorbate cannot lead to goethite dissolution by itself

329

when O2 is present, it can still facilitate redox reactions generating surface Fe(II) (probably short

330

lived). Knowledge of surface Fe(II) and goethite interactions has recently been advanced by the

331

discovery of electron transfer and atom exchange (ETAE) between aqueous Fe(II) and the Fe(III)

332

(hydr)oxide surface via a mechanism involving simultaneous Fe(II) oxidative adsorption coupled

333

to Fe(III) reductive dissolution.69, 70 The active recrystallization process may kinetically labilize

334

the goethite structure for faster ligand-promoted dissolution. Such an increase of the dynamics of

335

the goethite lattice would be consistent with previous observations of increased rates of iron

336

isotope exchange in the presence of Fe(II).70 The second possible mechanism is that DFOB may

337

detach a fraction of the short lived surface Fe(II) at a rate that outcompetes the reoxidation of

20



Page 22 of 39

338

surface Fe(II) back to Fe(III). Borer et al. presented a similar mechanism for the synergy

339

between a siderophore and a photoreductant for goethite dissolution.51 Although the Fe(II) that

340

was detached from the surface into solution by complexation with DFOB would be rapidly

341

oxidized by O2 to form an Fe(III)-DFOB complex, reoxidation would have been delayed and

342

relocated from the surface to the aqueous phase. In the presence of excess and highly accessible

343

O2, there might only be small amounts of surface Fe(II) due to its short lifetime. Our

344

experimental results indicated that the surface Fe(II) sites might be saturated by DFOB at

345

concentrations as low as 5 µM. When provided with 0.1 or 1 mM ascorbate, increasing the

346

DFOB concentration from 5 to 20 µM only marginally increased the dissolution rates (Figure S6

347

of the Supporting Information). The third possible mechanism is related to the radical scavenging

348

capability of DFOB71. As a one-electron-transfer process, reduction of surface Fe(III) is

349

accompanied by the oxidation of ascorbate to a radical species, which can re-oxidize surface

350

Fe(II). DFOB is a known radical scavenger that has been demonstrated to inhibit the reoxidation

351

of Fe(II).52, 72 It is not clear if HBED and other ligands could affect the redox reaction between

352

ascorbate and iron surface sites by a similar mechanism. Overall, our current results are not able

353

to attribute the observed synergistic effect to any specific mechanism. Labeled stable isotope

354

experiments and molecular dynamic simulations may provide more insights into the process.

21


Page 23 of 39


355

Environmental Implications. Knowledge of iron (hydr)oxide dissolution in the presence of

356

biogenic ligands is the foundation of understanding the natural strategies of iron acquisition by

357

microorganisms and plants as well as engineered interventions to improve bioavailability. Co-

358

exudation of reductants and ligands under iron limitation has been observed for both plants and

359

microorganisms. Our results demonstrate that reductants and ligands can act synergistically to

360

promote iron (hydr)oxide dissolution at soil pH values and concentrations that are relevant to

361

soils. This synergy may serve as an efficient biological iron acquisition strategy under oxic

362

conditions. Certain types of plants exude phenolics (e.g., caffeic acid) to mobilize Fe, many of

363

which are both ligands and reductants for Fe(III) with moderate potency.8, 73-75 The reductant-

364

ligand synergy may shed light on refined kinetic models for iron (hydr)oxide dissolution by

365

phenolics. Previous studies focusing on microbial growth proposed that iron reduction generated

366

a pool of iron that is readily bioavailable to aerobic microorganisms.33, 42, 45 The present results

367

further confirmed that iron reduction works in conjunction with ligands to increase the rate of

368

mobilization of Fe from poorly soluble iron sources.

369

While the reductant-ligand synergistic effect is supported by convincing evidence in abiotic

370

goethite model systems, translation of the results to real soil systems has to consider a number of

371

soil processes. Our HBED-ascorbate-goethite results illustrate that the synergistic effect may be

22



Page 24 of 39

372

masked by competitive adsorption (e.g., by phosphate, sulfate, humic and fulvic acids), which is

373

modulated

374

iron(hydr)oxides commonly contain several mol% substituted aluminum.2 An inhibitory effect of

375

Al was observed for the reductive dissolution pathway77-79 but not for ligand promoted

376

dissolution26 by DFOB. The complexation of aluminum with ligands (e.g., siderophores) may

377

help Al detach from iron (hydr)oxide surfaces.80 Therefore, the ligand-reductant synergy may be

378

more complex when considering the impurities in natural iron oxides. Moreover, further research

379

is required to evaluate the role of the ligand-reductant synergy in biological iron uptake in the

380

presence of microorganisms and plants, whose releases of biogenic exudates are regulated by the

381

surrounding iron bioavailability.

by

solution

pH

and

electrostatic

interactions.39,

76

Naturally occurring

382 383

Acknowledgements 384

The authors thank Jeffrey Catalano for his valuable inputs throughout this study. Chao Pan,

385

Manvitha Marni, Vrajesh Mehta provided laboratory assistance. This research was supported by

386

the International Center for Advanced Renewable Energy and Sustainability (ICARES) at

387

Washington University in St. Louis. Comments and suggestions of three anonymous reviewers

388

and Associate Editor Timm Strathmann helped us improve an earlier version of this paper.

23


Page 25 of 39


389 390

Supporting Information Available Additional information (including 2 tables, 6 figures and additional discussion about the

391 392

analytical methods) is available free of charge via the Internet at http://pubs.acs.org.

393 394

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586 587

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Table 1. Summary of the batch experiments for goethite dissolution at pH 6. (Part a) Reductant/Ligand alone Exp.IDa

Ascorbate (mM)

HBED (µM)

DFOB (µM)

PO2 (bar)

Dissolution rate (nmol·g-1h-1)

End time of regression (h)b

1 2 3 4 5 6 7 8 9 10 11 12

0 1 0.1 0 1 0.1 0 0 0 0 0 0

0 0 0 0 0 0 0 0 0 0 10 10

0 0 0 0 0 0 10 20 10 5 0 0

Synergistic Effect of Reductive and Ligand-Promoted Dissolution of

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