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Light-mediated reactive oxygen species generation and iron redox transformations in the presence of exudate from the cyanobacterium Microcystis aeruginosa Kai Wang, Shikha Garg, and T. David Waite Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.7b01441 • Publication Date (Web): 26 Jun 2017 Downloaded from http://pubs.acs.org on June 26, 2017

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Environmental Science & Technology

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Light-mediated reactive oxygen species generation and iron redox transformations in

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the presence of exudate from the cyanobacterium Microcystis aeruginosa

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Kai Wang, Shikha Garg, and T. David Waite*

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School of Civil and Environmental Engineering, The University of New South Wales, Sydney,

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NSW 2052, Australia

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Revised

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Environmental, Science and Technology

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June 2017

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1 Environment ACS Paragon Plus


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ToC Art

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ABSTRACT

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The photochemical properties of the organic exudate secreted by a toxic strain of Microcystis

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aeruginosa were studied by measuring reactive oxygen species (ROS) generation and redox

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transformations of iron in the presence of the organic exudate under acidic (pH 4) and

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alkaline (pH 8) conditions. Our results show that the organic exudate generates nanomolar

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concentrations of superoxide and hydrogen peroxide on irradiation with simulated sunlight in

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a manner consistent with that reported for terrigenous natural organic matter. The photo-

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generated superoxide plays an important role in Fe(III) reduction under alkaline conditions

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with nearly 45% of the observed Fe(II) generation on Fe(III) reduction occurring via Fe(III)

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reduction by superoxide while the rest of the Fe(III) reduction occurs via a ligand-to-metal

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charge transfer (LMCT) pathway. In contrast, under acidic conditions, 100% of the observed

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photochemical Fe(II) generation on Fe(III) reduction occurs via a LMCT pathway. These

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results suggest that steady-state dissolved Fe concentrations and hence Fe availability in

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natural waters will significantly increase in the presence of these algal exudates.

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Furthermore, significant diel variation in Fe(II) concentration is to be expected, even in acidic

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waters, since time scales of light-mediated Fe(III) reduction and thermal Fe(III) reduction

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differ markedly. A kinetic model is developed that adequately describes both the generation

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of ROS and the photochemical redox transformations of iron in the presence of M.

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aeruginosa exudate.



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1. INTRODUCTION

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Iron is a critical micronutrient in natural environments and is implicated in a range of

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important cellular functions including respiration and photosynthesis. The reduced ferrous

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form of iron (Fe(II)) is substantially more bioavailable than the thermodynamically favoured

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oxidized ferric form (Fe(III)) due to its higher solubility and weaker affinity for organic

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ligands.1,

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transformations of iron and suggested that these processes can strongly influence iron

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availability in natural waters.1,

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reduction of Fe(III); namely, reduction by superoxide ( O•− 2 )

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transfer (LMCT)

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The relative importance of these pathways is dependent particularly on the nature and

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concentration of the organic moieties present and the pH. For example, reduction by O•− 2 is

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likely to be more important in the presence of organic moieties capable of generating O•− 2

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and at high pH conditions where O•− 2 is relatively long-lived (with a half-life of a few

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minutes).15 Similarly, LMCT will be important in the presence of organic moieties which are

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capable of forming photo-labile complexes with Fe(III).

2

As a result, researchers have given particular attention to reductive

10, 11

3-7

Three major pathways are known to account for the 8, 9

, ligand to metal charge

and/or reduction by semiquinone or hydroquinone-like moieties.12-14

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The kinetics of Fe(II) oxidation is also impacted by the nature of the organic moieties

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present and the pH of the solution. Different organic groups exhibit distinct Fe(II) binding

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affinities with the strength of binding markedly affecting the Fe(II) oxidation rate.

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described by Pham and Waite20, the second order oxidation rate constant of organically-

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complexed Fe(II) is affected by the stability constant of the Fe(III)L and Fe(II)L complexes

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with an increase in the K Fe( III ) L / K Fe( II ) L value resulting in an increase in the Fe(II) oxidation

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rate constant. Furthermore, generation of organic entities such as semiquinone and/or peroxyl

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radicals and reactive oxygen species (ROS) such as O•− and hydrogen peroxide (H2O2) on 2


16-20

As

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irradiation of organic matter may further enhance the Fe(II) oxidation rate.21-23 Thus, the

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nature of organic compounds present is important in controlling the bioavailability of Fe

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given that the steady state concentration of the highly soluble Fe(II) form is recognized, in

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most aquatic systems, to dominate the rate and extent of Fe uptake.

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Natural organic matter (NOM) can be broadly classified into two categories,

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autochthonous NOM and terrigenous NOM. While the redox transformations of iron in the

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presence of terrigenous NOM (particularly through use of the reference NOM Suwannee

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River Fulvic acid (SRFA)) has been widely investigated,

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investigation of iron redox transformations in the presence of autochthonous NOM which is

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generated as a result of active exudation by algae and bacteria. These organic exudates are a

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complex mixture of poorly characterized organic compounds and may be present in high

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concentrations in surface waters, particularly when large numbers of algae are present.25 As

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such, there is a need to understand the impact of these organic exudates on the

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biogeochemical cycling of Fe both because of the implications to Fe supply to biota and

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because of the impact on the fate of any other entities associated with iron species

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(particularly iron oxides).

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tertiolecta can complex Cu(II)

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seawater.27 In a recent study, Santana Casiano and co-workers showed that algal exudates

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possess Fe(III) reducing phenolic groups and hence can favor the persistence of Fe(II) in

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coastal regions.28 We recently investigated the redox transformations of Fe in the presence of

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organic exudate secreted by Microcystis aeruginosa, a common alga in cyanobacterial

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blooms in temperate fresh waters.29 Our work showed that short-lived reduced organic

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moieties are naturally present in the algal exudate and are able to induce significant reduction

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of Fe(III) under both acidic and alkaline conditions. Furthermore, our results indicated that

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these organic exudates have strong Fe complexing ability which results in more rapid Fe(II)

4, 5, 21, 24

there have been limited

A previous study showed that the exudates of Dunaliella 26

and also impact the oxidation kinetics of Fe(II) in



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oxygenation than in the absence of the exudate due to changes in the redox potential of the

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Fe(III)/Fe(II) redox couple.

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While most of the earlier work with algal exudate is focused on redox transformations

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occurring in the absence of light, in this work we investigate the impact of light on the

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organic exudate secreted by M. aeruginosa with particular attention given to light-mediated

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ROS ( O•− 2 and H2O2) generation and Fe redox transformations under acidic (pH 4) and

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alkaline (pH 8) conditions. The main aims of this study were to investigate the mechanism of

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(i) O•− 2 and H2O2 generation and (ii) Fe redox transformation by the algal exudate under

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irradiated conditions in order to improve our understanding of the impact of these organic

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moieties on the solubility, bioavailability and transformation of iron in natural waters. The

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pH 8 studies used are representative of conditions in the growth medium (and, more

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generally, reasonably representative of natural waters) while the pH 4 condition has been

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chosen principally to avoid complications associated with the precipitation of iron oxides and

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fast Fe(II) oxygenation and thus assist in quantifying the rate and extent of dissolved Fe

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redox transformations mediated by the algal exudate. M. aeruginosa was chosen since it is

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one of the most common blue-green algal species found in freshwater environments. Based

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on our experimental results, we have developed a mathematical kinetic model that adequately

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describes the generation rate of O•− 2 and H2O2 and Fe redox transformations in irradiated

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solutions containing algal exudate.

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2. EXPERIMENTAL METHODS

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2.1 Reagents

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All reagent solutions were prepared using 18 MΩ.cm resistivity Milli-Q water (MQ) unless

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stated otherwise. All experiments were performed at a controlled room temperature of


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22.0±0.5 °C in the algal exudate at pH 4 or 8. The pH of the algal exudate was adjusted to

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either pH 4 by addition of 1 M HCl (high purity 30% w/v, Sigma) or pH 8 by addition of 1.5

130

mM NaHCO3 followed by 1 M HCl to trim to the target value (if necessary) immediately

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before the experiments. In order to mimic surface water conditions, where light-mediated

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ROS generation and Fe redox transformations are most prevalent, all experiments were

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performed in air–saturated solutions unless stated otherwise with air saturation assured by

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maintaining sufficient head-space in the reaction vessels.

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Primary and working stock solutions of Fe(II) and Fe(III) were prepared as described

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in our earlier work.12 Stock solutions of ferrozine (3-(2-pyridyl)-5,6-diphenyl-1,2,4-triazine-

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4′,4″-disulfonic acid sodium salt; abbreviated as FZ; Sigma), desferrioxamine B (DFB;

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Sigma), superoxide dismutase (SOD; Sigma) and Amplex Red (AR; Invitrogen) mixed with

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horseradish peroxidase (HRP; Sigma) were prepared and stored as reported in our earlier

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work.12 A 1 µM stock solution of methyl Cypridina luciferin analogue (MCLA; Sigma) was

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prepared in acetate buffer at pH 6 as described earlier.30

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diethylenetriaminepentaacetate (DTPA; Sigma) was prepared in MQ.

A 5 mM stock solution of

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•− For measurement of O•− 2 decay kinetics and calibration of O 2 generation results, the

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O•− stock was generated by photolysis of a 0.1 M borate buffer solution at pH 11.5 2

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containing 15 µM DTPA (to bind any trace metals present), 4% v/v ethanol (100% high

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purity; Fluka) and 0.4% v/v acetone (BDH) in a 1 cm quartz cuvette using a 15 W Hg-lamp.

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31

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2.2 Collection of algal exudate

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The algal exudate was prepared as described in our earlier work.29 The total dissolved

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organic carbon (DOC) content of the exudate was measured using a TOC analyser (TOC-

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5000A, Shimadzu; detection limit ~ 0.1 mg L-1) and was always in the range 2.0 - 3.0 mg L-1



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with cell status and cell density similar for all experiments. No change in DOC of the exudate

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occured during irradiation over the time scales investigated here.

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2.3 Experimental setup

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For measurement of O•− 2 and H2O2 generation under irradiated conditions at pH 4 and 8,

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3.5 mL of algal exudate was irradiated in a 1 cm quartz cuvette for 1, 2, 5 or 10 minutes and

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the concentration of O•− 2 and H2O2 formed after irradiation measured using the MCLA

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method and AR method respectively as described in the following section. All measurements

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of O•− 2 and H2O2 generation were performed in DTPA (1.5 µM) amended solution to prevent

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trace metal catalysed dismutation of O•− 2 .

161

For measurement of O•− 2 decay at pH 8 in non-irradiated and previously-irradiated

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exudate solutions, an appropriate volume of photochemically generated O•− 2 stock solution

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was added to 3.5 mL of the non-irradiated algal exudate or exudate irradiated for a certain

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•− time (2, 5 or 10 minutes) prior to O•− 2 addition, respectively. The concentration of O 2

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remaining and H2O2 formed was measured for 1 minute continuously using the MCLA and

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AR methods respectively as described in the following section. All measurements of O•− 2

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decay and consequent H2O2 generation were performed in solutions containing 1.5 µM

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DTPA to prevent trace metal catalysed O•− 2 dismutation. However, DTPA may not be able to

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bind trace metals incorporated in cellular biomolecules (for example, as Fe-SOD enzymes)

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and, if present, these entities may catalyse O•− 2 decay.

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For measurement of Fe redox transformations under irradiated conditions at pH 4, 3

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mL of algal exudate containing appropriate concentrations of Fe(III) or Fe(II) (in the

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concentration range 25-150 nM) was irradiated in a 1 cm quartz cuvette for 1, 2, 5 or 10


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minutes and the concentration of Fe(II) formed or remaining after irradiation was measured

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using the modified FZ method. 12

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For measurement of Fe(III) reduction rates under irradiated conditions at pH 8, an

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appropriate concentration of Fe(III) (in the concentration range 25-100 nM) was added to 3

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mL of algal exudate containing 1 mM FZ and was irradiated in a 1 cm quartz cuvette for 1,2

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5 or 10 minutes. The presence of FZ during irradiation traps any Fe(II) formed prior to its

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oxidation and hence the Fe(II) generation rate measured at pH 8 represents the Fe(III)

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reduction rate. Proper control experiments were performed to ensure that FZ-mediated Fe(III)

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reduction is negligible under the conditions investigated here. For measurement of Fe(II)

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oxidation kinetics under irradiated conditions at pH 8, an appropriate concentration of Fe(II)

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(in the concentration range 50-150 nM) was added to 3 mL of algal exudate and the mixture

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subsequently irradiated in a 1 cm quartz cuvette for various times (in the range 10-120 s) and

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the concentration of Fe(II) remaining measured using the FZ method. 32

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To determine the role of dioxygen in Fe(II) oxidation, we conducted one set of

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experiments in which the concentration of dioxygen was reduced by sparging the buffer

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solution with Ar in a sealed reactor for 4 hours prior to experiments as described in our

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earlier work.29 Detailed description of the experimental setup used for deoxygenation

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experiments is also provided in Supporting Information (S1).

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A ThermoOriel 150W Xe lamp equipped with AM0 and AM1 mass filters was used

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as the light source. The spectral irradiance of the lamp and absorbed spectral irradiance by the

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exudate as function of wavelength was reported in our earlier work12 and is also shown in

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Figure S1 in Supporting Information. The total absorbed spectral irradiance (or fluence rate)

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is 222.5 µEinstein.m-2.s-1 which corresponds to a dose of 133.5 mEinstein.m-2 during 10

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minutes of irradiation. All control experiments were performed in exudate-free medium that

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had the same composition, ionic strength (5 mM at pH 4, and 6.5 mM at pH 8) and pH as the



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algal exudate. No change in pH of solutions over the course of experiments was observed

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under the conditions investigated here.

201

2.4 Measurement of Fe(II)

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The concentration of Fe(II) at pH 4 and 8 was measured using the FZ method as described in

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our earlier work.12 The details of the method are also presented in Supporting Information

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(section S1).

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2.5 H2O2 determination

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The concentration of H2O2 generated on irradiation of exudate was quantified

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fluorometrically using a Cary Eclipse fluorescence spectrophotometer (Agilent Technologies)

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with 2 µM AR and 1 kU.L-1 HRP as the fluorescence regent. Since this method is only

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effective in the circumneutral pH range (i.e. 7.5-8.5) 33, for H2O2 measurement at pH 4, 1 mL

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of irradiated sample was mixed with 2 mL of 10 mM phosphate buffer (pH 7.5) to adjust the

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pH prior to measurement. H2O2 concentrations in all pH 8 samples were measured directly

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without pH adjustment. The equipment setting and calibration procedure used were described

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in our earlier study.

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fluorescence reagent (i.e. 2 µM AR and 1 kU.L-1 HRP) was added to the cuvette containing

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3.5 mL of the algal exudate (non-irradiated or previously irradiated) prior to addition of O•− 2

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stock in order to trap any H2O2 formed instantaneously since O•− 2 decay is rapid under these

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conditions. The concentration of H2O2 was measured continuously for 1 min. Control

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experiments were performed in the presence of catalase (an enzyme which causes decay of

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H2O2) to ensure that the fluorescence measured was due to generation of H2O2 only. The

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detection limit of the H2O2 measurement method is approximately 3 nM (defined as 3 times

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the standard deviation of the reagent blank).

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2.6 Superoxide measurement

30

For measurement of H2O2 concentration formed on O•− 2 decay, the


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Superoxide was measured using a FeLume chemiluminescence (CL) system

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(Waterville Analytical) with 1 µM MCLA (a chemiluminescent reagent that is highly

225

selective for superoxide) in 50 mM acetate buffer at pH 6.0.34 A peristaltic pump was used to

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separately deliver the experimental solution and MCLA at a flow rate of 1.5 mL.min-1 into

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the flow cell of a FeLume CL system where mixing yielded superoxide-specific CL that was

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detected by the instrument’s photomultiplier tube. The O•− 2 concentration in the experimental

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solution was monitored for 2 min however no data could be acquired during the initial 30 s

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after O•− 2 addition due to the time taken for sample to reach the FeLume flow cell.

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The system was calibrated by standard additions of O•− 2 stock solution that was

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generated photochemically immediately prior to use as described earlier (section 2.1). The

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procedures for standard additions and method for calculation of O•− 2 concentration has been

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described in detail previously8 and is also presented in Supporting Information (see section

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S1). The detection limit of the method is approximately 0.1 nM (defined as 3 times the

236

standard deviation of the reagent blank).

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2.7 Kinetic modeling

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Kintek Explorer was used for kinetic modelling of our experimental results.35 Kintek

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Explorer is a kinetic simulation program that allows multiple data sets to be fit

240

simultaneously to a single model based on numerical integration of the rate equations

241

describing the reaction mechanism.35 As described in our earlier work,29 the model is

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developed based on a proposed set of chemical reactions and their associated rate

243

equations. The rate constant for each reaction is generally either drawn from the

244

literature or, if the rate constant is dependent on experimental conditions, deduced from

245

studies of that particular reaction under the conditions of interest. Hence, agreement

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between model results and experimentally measured values over a range of conditions



247

indicates that the chosen reaction set is a reasonable description of the underlying

248

processes operating under the experimental conditions investigated.

249

RESULTS and DISCUSSION

250

3.1 Photochemical generation of superoxide and hydrogen peroxide by the M.

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aeruginosa exudate at pH 4 and 8

252

As shown in Figure 1, considerable amounts of O•− 2 and H2O2 are generated on irradiation

253

of organic exudate at pH 8 however only a small concentration of H2O2 is generated at pH 4.

254

•− No measurement of O•− 2 was performed at pH 4 due to the short lifetime of O 2 (a few

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seconds for nanomolar concentrations of [O•− 2 ]

256

8, on irradiation, O•− 2 concentration increases rapidly and then reaches steady-state within 10

257

minutes of irradiation. The H2O2 concentration increases linearly at both pH 4 and 8 on the

258

time-scales investigated here. No impact of aging (on storage in the dark for ~ 24h) of the

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exudate was observed on O•− 2 and H2O2 generation rates (data not shown) suggesting that the

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O•− 2 and H2O2 generating moieties are quite stable under the conditions investigated here. The

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presence of DTPA had no effect on the concentration of H2O2 concentration generated,

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confirming that production of these species was solely due to photolysis of exudate with no

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significant production due to either direct photolysis of DTPA or as a result of oxygenation of

264

Fe(II) or Cu(I) which may be present in trace amounts. This is further supported by the

265

observation that no impact of Fe(III) addition is observed on rates of H2O2 generation at pH 4

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and 8 (Figure S2) supporting the hypothesis that Fe cycling has no impact on H2O2

267

generation at least on the time scales investigated here.

15

) under acidic conditions. As shown, at pH

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Note that the generation of other ROS such as singlet oxygen (1O2) and hydroxyl

269

radicals ( OH • ) may also occur on irradiation of exudate however the steady-state

270

concentrations of these entities is expected to be very low due to their rapid decay kinetics 12 Environment ACS Paragon Plus

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(1O2 undergoes rapid relaxation due to its interaction with water36 while OH • is rapidly

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scavenged by various inorganic (i.e. Cl⁻, HCO3⁻ etc.) and organic groups present in our

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experimental matrix)37 and hence no attempt was made to measure these entities in our

274

experimental system.

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3.1.1 Mechanism of H2O2 decay and generation at pH 4 and 8

276

Increasing the initial H2O2 concentration had no effect on the H2O2 production rate at

277

pH 8 (Figure S3) indicating that H2O2 did not react measurably with other entities in this

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system. Thus, on the timescale of these experiments, H2O2 appears to be a stable end product

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of exudate photolysis at pH 8. In contrast, at pH 4, increasing the initial H2O2 concentration

280

resulted in a decrease in the H2O2 generation rate (Figure S3b) suggesting that H2O2

281

consuming entities are present in the exudate solution at pH 4. Using the data shown in

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Figure S3b (see section S1.6 for detailed description of this calculation), the H2O2 decay rate

283

constant is determined to be 4.0×10-4 s-1 which is similar to that measured when H2O2 was

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added to algal exudate in the dark (Figure S4). This result supports the conclusion that these

285

H2O2 consuming entities are not generated on irradiation but are present intrinsically in the

286

exudate.

287

As reported earlier, the generation of H2O2 occurs via disproportionation of

288

38 photochemically generated O•− However, the H2O2 generation rates measured here 2 at pH 8.

289

are much higher than those predicted from uncatalysed disproportionation of measured O•− 2

290

concentrations (Figure S5), based on a rate constant of 3.5×104 M-1.s-1 reported by Beilski

291

and co-workers at pH 8.15 Furthermore, no effect of SOD (30 kU.L-1) addition is observed on

292

H2O2 generation at pH 8 (Figure S6). Addition of SOD catalyses the disproportionation of

293

O•− 2 and hence, in the presence of SOD, disproportionation outcompetes other oxidative sinks

294

of O•− 2 not resulting in H2O2 generation. While both these observation may support the



295

conclusion that H2O2 is formed via a pathway not involving O•− 2 , this result may also be

296

explained by the presence of entities in the exudate that catalyse O•− 2 disproportionation as

297

reported to be the case for SRFA,39 thereby resulting in lower O•− and higher H2O2 2

298

concentrations than that expected if uncatalysed O•− This also 2 disproportionation occurs.

299

explains the absence of any effect of SOD addition on H2O2 generation since the rapid

300

exudate-catalysed disproportionation kinetics of O•− 2 may already exceed other oxidative

301

sinks of O•− We investigate the mechanism and kinetics 2 not resulting in H2O2 generation.

302

•− of O•− 2 decay in the following section and confirm that catalysed O 2 disproportionation

303

occurs in the presence of algal exudate.

304

3.1.2 Mechanism and kinetics of O•− 2 decay under non-irradiated conditions at pH 8

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As shown in Figure 2, when nanomolar concentrations of O•− 2 were added to solutions

306

containing algal exudate at pH 8 in the dark, the O•− 2 decays at a rate much higher than that

307

measured in the exudate-free medium. Similar O•− 2 decay rates are observed in the absence

308

and presence of DTPA indicating that trace metal catalysed O•− 2 decay is not important under

309

the conditions investigated here. These observations thus supports our conclusion that the

310

exudate catalyses O•− 2 decay in a manner similar to that observed for terrigenous natural

311

organic matter.39

312

Based on the observed decay of O•− 2 , a pseudo-first order rate constant of 0.09 ± 0.01

313

-1 s-1 for O•− 2 decay is calculated which is close to the value (0.12 s ) reported by Goldstone

314

and co-workers for the catalytic disproportionation of O•− 2 in the presence of SRFA at pH

315

9.5.39 No impact of aging (≥ 24 h) of the exudate on O•− 2 decay was observed supporting the

316

conclusion that the redox-active groups involved in O•− 2 decay are long-lived.


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As shown in Figure 2b, H2O2 is formed during the decay of O•− 2 with ~0.5 moles of

318

H2O2 formed for each mole of O•− 2 decayed supporting the conclusion that exudate-mediated

319

•− O•− as was 2 decay results in formation of H2O2 via catalytic disproportionation of O 2

320

reported to be the case for SRFA.38, 39 Furthermore, a pseudo-first order decay rate constant

321

of 0.13 ±0.04 s-1 was calculated based on the measured H2O2 concentration formed on O•− 2

322

decay (Figure 2b) using eq.(1) (see Supporting Information S1 for derivation of this equation)

323

which is statistically similar (p> 0.1 using single student’s t-test) to that determined from the

324

O•− 2 decay data:

325

1 - kddark t [H 2O2 ]= [O•− 2 ]0 1-e 2

326

dark where kd is the pseudo first order O•− decay rate constant in non-irradiated exudate 2

327

•− solution and [ O•− 2 ]0 represents the initial concentration of O 2 .

(

)

(1)

328

The catalytic disproportionation of O•− 2 in the presence of algal exudate explains the

329

lower O•− 2 concentration and higher H2O2 concentration measured on irradiation of the

330

exudate at pH 8 (Figure 1).

331

In the following section, we investigate the impact of irradiation of the exudate on the

332

•− O•− 2 decay kinetics to determine the mechanism of O 2 decay in irradiated exudate solution.

333

3.1.3 Mechanism and kinetics of O•− 2 decay under irradiated conditions at pH 8

334

As shown in Figure 3, when O•− 2 is added to previously-irradiated exudate solution in

335

the dark, it decays more rapidly as indicated by the lower O•− 2 concentration remaining after

336

30s in pre-irradiated exudate solution compared to that measured in non-irradiated solution.

337

Furthermore, the O•− 2 decay rate increases with increase in pre-irradiation time. These results

338

support the conclusion that organic moieties are generated on irradiation of the exudate that



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further catalyse the decay of O•− 2 . This result is consistent with observations that the lifetime

340

38 of O•− Although, the decay rate 2 is shorter in irradiated than non-irradiated SRFA solution.

341

of O•− 2 increases in previously-irradiated exudate solution, the concentration of H2O2 formed

342

on O•− decay in previously-irradiated solution is lower than that measured under dark 2

343

conditions suggesting that the photo-generated organic moieties oxidize O•− 2 to dioxygen

344

thereby diminishing the extent of H2O2 generation. The pseudo-first order O•− 2 decay rate

345

constant, calculated based on the measured O•− 2 concentration remaining after 30 s in exudate

346

solution previously irradiated for 10 min is 1.2-fold higher (p 0.05 s-1 produces the same

494

results.

495

3.4.2 Other sinks of superoxide and hydrogen peroxide

496

As discussed, H2O2 decays in the presence of exudate at pH 4. These H2O2 consuming

497

moieties are not generated on irradiation of the exudate but are present intrinsically in the

498

exudate solution since similar H2O2 decay rates were observed under non-irradiated and

499

irradiated conditions (see section S1.6 for details on the calculation of H2O2 decay rates in

500

irradiated and non-irradiated conditions). We have assumed that the concentration of the

501

H2O2 consuming moieties is in excess of the H2O2 concentration used here and hence

502

remained constant during the reaction, at least on the time scales investigated here. The

503

pseudo-first order rate constant of H2O2 decay was determined to be 4×10-4 s-1 (reaction 13,

504

Table 1) based on the measured rate of H2O2 decay in non-irradiated exudate solutions at pH

505

4.


Page 22 of 41

Page 23 of 41

506


In the presence of SRFA, oxidative decay of O•− 2 via interaction with photo-generated 38

507

organic moieties was shown to be important

and was found to be significant here as well

508

since the O•− 2 decay rate in irradiated exudate solution (Figure 3) was higher than that

509

measured in non-irradiated solution at pH 8 (Figure 2). Reactions 14 and 15 represent the

510

generation of oxidizing organic moieties and their reaction with O•− 2 respectively with the

511

rate constants for these reactions determined based on the measured rates of generation of

512

•− O•− 2 and H2O2 (Figure 1) and O 2 decay rate in irradiated exudate solution at pH 8 (Figure

513

3). Since no measurement of O•− 2 concentration was possible at pH 4 (in view of the rapid

514

rate of disproportionation), the rate constant for these reactions cannot be determined at pH 4.

515

3.4.3 Ligand-to-metal charge transfer mediated Fe(III) reduction at pH 4 and 8

516

As discussed earlier, LMCT is the main pathway for Fe(III) reduction at pH 4

517

(reaction 16, Table 1) with the pseudo-first order rate constant for this reaction determined

518

from the best fit to the measured Fe(II) concentrations generated on Fe(III) reduction at pH 4.

519

Since Fe(III) mostly exists in organically-complexed form at pH 4, the rate constant

520

determined here represents that for light-mediated LMCT of Fe(III)L.

521

As noted earlier, LMCT-mediated Fe(III) reduction plays some role at pH 8 with

522

exudate-coated iron oxyhydroxide particles undergoing reductive dissolution and any Fe(III)

523

' present as either Fe(III) or Fe(III)L also undergoing reduction. The pseudo-first order rate

524

constant for reduction of Fe(III) present as iron oxyhydroxide particles via LMCT at pH 8

525

(reaction 17, Table 1), as has been reported to occur in earlier studies,

526

based on the measured Fe(II) generation rate on Fe(III) reduction in the pH 8 exudate-free

527

medium (Figure S8) where Fe(III) exists almost exclusively as iron oxyhydroxide particles.

528

While it is possible that the reactivity of exudate-coated iron oxyhydroxide particles may be

529

quite different to the particles formed in the exudate-free medium, the reactivity is expected


42, 46

was determined


Page 24 of 41

530

to vary by two-three fold at most (as confirmed in a separate study) which does not impact

531

the model output significantly. The pseudo-first order rate constant for reduction of Fe(III)L

532

via LMCT at pH 8 was assumed to be the same as that determined at pH 4. The rate constant

533

for Fe(III)' reduction by LMCT (reaction 8, Table 1) was determined based on the measured

534

Fe(II) formation rate on Fe(III) reduction in exudate-free medium at pH 4 under irradiated

535

' conditions (Figure S8) where Fe(III) exists in dissolved form as Fe(III) .

536

3.4.4 Superoxide mediated Fe(III) reduction at 8

537

•− As discussed, Fe(III) reduction by O•− 2 is important at pH 8 with O 2 contributing to

538

reduction of Fe(III)L and Fe(III)' (reaction 19 and 20, Table 1) that is formed via dissolution

539

of iron oxyhydroxide (reaction 9, Table 1) as reported in an earlier study.41 The rate constant

540

' for Fe(III) reduction by O•− 2 was used as reported earlier at pH 8.

541

Fe(III)L reduction by O•− 2 was determined based on the best-fit to the measured Fe(II)

542

formation rate at pH 8.

543

Fe(III) reduction at pH 8 with the rate constant for this reaction determined from best fit to

544

the measured rate and extent of Fe(III) reduction under irradiated conditions and is ten-fold

545

higher than that determined under non-irradiated conditions in our earlier work.29 To explain

546

this discrepancy, we need to consider the time scales at which the AFO dissolution is

547

occurring in the two studies. While the Fe(III) reduction via AFO dissolution under irradiated

548

conditions is measured over 10 minutes, the time-scale for the measurement of Fe(III)

549

reduction under dark conditions was 4 hours during which the AFO undergoes

550

significant aging (i.e. decrease in reactivity). The dissolution rate constant used here is the

551

same as that reported for 1 min aged AFO47 while that determined under non-irradiated

552

conditions was representative of AFO aged for a few (3-6) hours. 47

5

5

The rate constant for

The dissolution of AFO (reaction 9, Table 1) controls the rate of


quite

Page 25 of 41

553


3.4.6 Superoxide-mediated Fe(II) oxidation at pH 8

554

Based on our experimental results (Figure S12), superoxide-mediated Fe(II) oxidation

555

occurs at pH 8 and this reaction is included in the kinetic model (reaction 21, Table 1) with

556

the rate constant for this reaction obtained based on the best fit to the measured rates of Fe(II)

557

oxidation (Figure 5) and H2O2 generation on Fe(II) addition (Figure S14) at pH 8. The fitted

558

value of the rate constant for Fe(II)L oxidation is the same as that reported earlier for

559

organically complexed Fe(II)

560

part, in organically complexed form.

23

in accord with the hypothesis that Fe(II) exists, for the most

561

Note that the interaction of short-lived Fe(III) reducing groups (Exred) with Fe(III)

562

(reaction 1, Table 1), as reported to occur in our earlier work 29 and shown here, do not play

563

a significant role in Fe(III) reduction in irradiated conditions since Fe(III) reduction by O•− 2

564

and LMCT occur at much faster rates.

565

As described in our earlier work,29 the rate constants for formation and dissociation of

566

Fe(III)L (reaction 7 and 10 respectively) and the Fe binding capacity of the exudate was

567

assumed to be similar to that reported for Fe(III)SRFA. Reactions 7 and 10 play an important

568

role in controlling the Fe(III) reduction rate at pH 8 and hence further work is required to

569

verify the rate constant for this reaction. We have included a range of values for the

570

formation and dissociation rate constants for Fe(III)L at pH 8 which are consistent with the

571

experimental results obtained here (see Table 1). As shown in Table 1, conditional stability

572

constant (i.e. Kcond; the ratio of the formation and dissociation rate constant) values > 107.5

573

and 108 at pH 4 and 8 respectively are consistent with our experimental results. Similarly, any

574

values of the rate constants ≥ 5×103 M-1s-1 and ≥ 1×105 M-1s-1 for Fe(II)L formation at pH

575

4 and 8 respectively are consistent with our observations. Note that the values of formation

576

rate constants and conditional stability constant obtained here are based on the Fe binding

577

capacity of 65.6 nmoles.g-1C of exudate as reported for SRFA


48

and suitable adjustments


578

need to be made to these constants for any changes in the Fe binding capacity. Our future

579

work will focus on measuring the Fe binding capacity of the exudate and the conditional

580

stability constant of exudate-complexed Fe to further constrain our model.

581

As shown in Figures 1, 4 and 5, the kinetic model provides an excellent description of

582

our experimental results at both pH 4 and 8. The small discrepancies in the model predicted

583

O•− 2 concentrations and experimentally-measured values (Figure 1) are possibly due to the

584

error introduced in the measured O•− concentrations as a result of the extrapolation 2

585

procedure (section S1.5) used for calculation of the concentration of this transient species

586

given that O•− 2 decays rapidly under the experimental conditions investigated here. The small

587

discrepancy between model predicted results and the experimentally measured values is

588

well within the variability in experimental data expected at such low Fe and ROS

589

concentrations. In all cases, model-predicted results are within the 99% confidence interval of

590

the experimentally measured values.

591

Using the mathematical model developed here, we also calculated the turnover

592

frequency (TOF) of Fe under irradiated conditions in the presence of exudate where the TOF

593

of Fe is defined as: TOF =

594

Fe(III) reduction rate [Fe]T

595

Using the measured Fe(III) reduction rate, Fe TOF of 15.6 h-1 is obtained at pH 4 under

596

irradiated conditions which is similar to the Fe TOF reported in irradiated SRFA solution

597

(12.7 h-1)

598

further supports the hypothesized strong Fe complexing ability of the exudate as reported to

599

be the case for SRFA.48, 50 The TOF value at pH 8 (1.3 h-1) is much lower than that obtained

600

at pH 4 mainly because of the relatively low reactivity of the dominant Fe(III) species (i.e.

601

iron oxyhydroxide) present at pH 8. The TOF of Fe is also expected to vary with the exudate

49

indicating similar Fe(III) reduction rates via LMCT in the two solutions. This


Page 26 of 41

Page 27 of 41


602

concentration as well as with time of day due to variation in the light intensity. The actual

603

turnover frequencies in the milieu surrounding the algal cells are possibly higher than those

604

predicted here due to the continuous generation of algal exudate.

605 606

ENVIRONMENTAL IMPLICATIONS

607

Our results show that the exudate excreted by M. aeruginosa induces generation of

608

O•− 2 on irradiation which rapidly reduces Fe(III) under alkaline conditions. Furthermore,

609

these organic exudates have strong Fe binding capacity, with the metal centre in these Fe(III)-

610

exudate complexes prone to reduction via a ligand-to-metal charge transfer (LMCT) pathway

611

on irradiation. These are important observations and suggest that the high Fe binding affinity

612

and O•− 2 generating ability of these organic exudates may enhance Fe solubility and Fe

613

availability in natural waters. These observations further suggest that the algae may produce

614

exudates to facilitate their iron acquisition since these organic moieties promote both

615

solubilisation and extracellular reduction of Fe.

616

The Fe(II)-Fe(III) turnover frequency in the presence of these organic exudates is

617

much higher than that measured in organic free waters (Figure 6) with significant diel

618

variations occurring in the TOF since the time scales of Fe(III) reduction differ markedly

619

under irradiated and non-irradiated conditions. While the solution conditions investigated

620

here are similar to freshwater systems, the results can be easily extended to marine systems.

621

Overall, our results show that the organic exudate secreted by M. aeruginosa could play a

622

significant role in the biogeochemical cycling of Fe. While the focus of the current study is

623

on the impact of the exudate on photochemical Fe redox transformations under the conditions

624

of the medium in which the algae have been grown, additional attention should be given in

625

future studies to the effects of pH and other major solution constituents on exudate-mediated

626

iron redox transformations.



627

Supporting Information

628

Details on additional experimental methods, results on generation of ROS by exudate and its

629

impact on Fe redox transformation under irradiated condition are illustrated in Supplementary

630

Information which may be accessed free of charge at the ACS website for this journal.

631 632

Corresponding Author

633

Professor T. David Waite; Address: School of Civil and Environmental Engineering, UNSW,

634

Kensington, Australia, 2052; Email: [email protected]; Phone: +61-2-9385-5060; Fax:

635

+61-2-9385-6139

636

Acknowledgement

637

We gratefully acknowledge funding provided by the Australian Research Council through

638

ARC Discovery Grant DP150102248 and ARC DECRA Award DE120102967.

639

References

640 641 642 643 644 645 646 647 648 649 650 651 652 653 654 655 656 657 658

1. Shaked, Y.; Kustka, A. B.; Morel, F. M. M., A general kinetic model for iron acquisition by eukaryotic phytoplankton. Limnol Oceanogr 2005, 50, 872-882. 2. Voelker, B. M.; Morel, F. M. M.; Sulzberger, B., Iron redox cycling in surface waters: effects of humic substances and light. Environ. Sci. Technol. 1997, 31, 1004-1011. 3. Rose, A. L.; Salmon, T. P.; Lukondeh, T.; Neilan, B. A.; Waite, T. D., Use of superoxide as an electron shuttle by the marine cyanobacterium Lyngbya majuscula. Environ. Sci. Technol. 2005, 39, 3708-3715. 4. Rose, A. L.; Waite, T. D., Role of superoxide in photochemical reduction of iron in seawater Geochim Cosmochim Acta 2006, 70, 3869-3882. 5. Fujii, M.; Rose, A. L.; Waite, T. D.; Omura, T., Oxygen and superoxide-mediated redox kinetics of iron complexed by humic substances in coastal seawater. Environ. Sci. Technol 2010, 44, 9337-9342. 6. Öztürk, M.; Croot, P. L. C.; Bertisson, S. B.; Abrahamssond, K.; Karlsone, B.; David, R. D.; Franssong, A. F.; Sakshauga, E., Iron enrichment and photoreduction of iron under UV and PAR in the presence of hydroxycarboxylic acid: implications for phytoplankton growth in the Southern Ocean. Deep Sea Research II 2004, 51, 2841-2856. 7. Salmon, T. P.; Rose, A. L.; Neilan, B. A.; Waite, T. D., The FeL model of iron acquisition: Non-dissociative reduction of ferric complexes in the marine environment. Limnol Oceanogr 2006, 51, 1744-1754

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8. Rose, A. L.; Waite, T. D., Reduction of organically complexed ferric iron by superoxide in a simulated natural water. Environ. Sci. Technol. 2005, 39, 2645-2650. 9. Garg, S.; Rose, A. L.; Waite, T. D., Superoxide Mediated Reduction of Organically Complexed Iron(III): Comparison of Non-Dissociative and Dissociative Reduction Pathways. Environ. Sci. Technol. 2007, 41, (9), 3205-3212. 10. Sima, J.; Makánová, J., Photochemistry of iron(III) complexes. Coord. Chem. Rev. 1997, 160, 161-189. 11. Barbeau, K.; Moffett, J. W., Laboratory and field studies of colloidal iron oxide dissolution as mediated by phagotrophy and photolysis. Limnol Oceanogr 2000, 45, 827-835. 12. Garg, S.; Ito, H.; Rose, A. L.; Waite, T. D., Mechanism and Kinetics of Dark Iron Redox Transformations in Acidic Previously Photolyzed Acidic Natural Organic Matter Solutions Environ. Sci. Technol 2013, 47, 1861-1867. 13. Santana-Casiano, J. M.; González-Dávila, M.; González, A. G.; Millero, F. J., Fe (III) reduction in the presence of catechol in seawater. Aquatic geochemistry 2010, 16, 467-482. 14. Jiang, C.; Garg, S.; Waite, T. D., Hydroquinone-Mediated Redox Cycling of Iron and Concomitant Oxidation of Hydroquinone in Oxic Waters under Acidic Conditions: Comparison with Iron−Natural Organic Matter Interactions. Environ. Sci. Technol 2015, 49, 14076−14084. 15. Bielski, B. H. J.; Cabelli, D. E.; Arudi, R. L.; Ross, A. B., Reactivity of HO2/O2radicals in aqueous solution. J Phys Chem Ref Data 1985, 14, (4), 1041-1100. 16. Rose, A. L.; Waite, T. D., Kinetic model for Fe(II) oxidation in seawater in the absence and presence of natural organic matter. Environ. Sci. Technol. 2002, 36, 433-444. 17. Santana-Casiano, J. M.; González-Dávila, M.; Millero, F. J., The oxidation of Fe(II) in NaCl–HCO3- and seawater solutions in the presence of phthalate and salicylate ions: a kinetic model. Mar Chem 2004, 85, 27-40. 18. Santana-Casiano, J. M.; González-Dávila, M.; Rodríguez, M. J.; Millero, F. J., The effect of organic compounds in the oxidation kinetics of Fe(II). Mar Chem 2000, 70, 211222. 19. Miller, C. J.; Rose, A. L.; Waite, T. D., Impact of natural organic matter on H2O2mediated oxidation of Fe(II) in a simulated freshwater system. Geochim Cosmochim Acta 2009, 73, 2758-2768. 20. Pham, A. N.; Waite, T. D., Modeling the Kinetics of Fe(II) Oxidation in the Presence of Citrate and Salicylate in Aqueous Solutions at pH 6.0−8.0 and 25 °C. J Phys Chem A 2008, 112, 5395-5405. 21. Garg, S.; Jiang, C.; Waite, T. D., Mechanistic insights into iron redox transformations in the presence of natural organic matter: Impact of pH and light Geochim Cosmochim Acta 2015, 165, 14 - 34. 22. Miller, C. J.; Lee, S. M. V.; Rose, A. L.; Waite, T. D., Impact of Natural Organic Matter on H2O2-Mediated Oxidation of Fe(II) in Coastal Seawaters. Environ. Sci. Technol 2012, 46, 11078-11085. 23. Rush, J. D.; Bielski, B. H. J., Pulse radiolytic studies of the reactions of HO2/O2- with Fe(II)/Fe(III) ions. The reactivity of HO2/O2- with ferric ions and its implication on the occurrence of the Haber-Weiss reaction. J Phys Chem 1985, 89, 5062-5066. 24. Garg, S.; Rose, A. L.; Waite, T. D., Pathways Contributing to the Formation and Decay of Ferrous Iron in Sunlit Natural Waters. In Aquatic Redox Chemistry, ACS symposium series: 2011; Vol. 1071. 25. Hassler, C. S.; Alasonati, E.; MancusoNichols, C. A.; Slaveykova, V. I., Exopolysaccharides produced by bacteria isolated from the pelagic Sothern Ocean - role in Fe binding, chemical reactivity, and bioavailability. Mar Chem 2011, 123, 88-98.



709 710 711 712 713 714 715 716 717 718 719 720 721 722 723 724 725 726 727 728 729 730 731 732 733 734 735 736 737 738 739 740 741 742 743 744 745 746 747 748 749 750 751 752 753 754 755 756 757 758

26. González-Dávila, A. G.; Santana-Casiano, J. M.; Pérez-Pena, J.; Millero, F. J., Binding of Cu(II) to the Surface and Exudates of the Alga Dunaliella tertiolecta in Seawater. Environ Sci Technol 1995, 29, 289-301. 27. González, A. G.; Santana-Casiano, J. M.; González-Dávila, M.; Pérez-Almeida, N.; Suárez de Tangil, M., Effect of Dunaliella tertiolecta organic exudates on the Fe (II) oxidation kinetics in seawater. Environ Sci Technol 2014, 48, 7933-7941. 28. Santana-Casiano, J. M.; González-Dávila, M.; González, A. G.; Rico, M.; López, A.; Martel, A., Characterization of phenolic exudates from Phaeodactylum tricornutum and their effects on the chemistry of Fe (II)–Fe (III). Mar Chem 2014, 158, 10-16. 29. Wang, K.; Garg, S.; Waite, T. D., Redox transformation of iron in the presence of exudate from the cyanobacterium Microcystis aeruginosa under conditions typical of natural waters. Environ. Sci. Technol 2017, 51, 3287-3297. 30. Garg, S.; Rose, A. L.; Waite, T. D., Production of Reactive Oxygen Species on Photolysis of Dilute Aqueous Solutions of Quinones. Photochem. Photobiol. 2007, 83, 904913. 31. McDowell, M. S.; Bakac, A.; Espenson, J. H., A convenient route to superoxide ion in aqueous solution. Inorg. Chem. 1982, 22, 847-848. 32. Stookey, L. L., Ferrozine: a new spectrophotometric reagent for iron. Anal Chem 1970, 42, 779-781. 33. Zhou, M.; Diwu, Z.; Panchuk-Voloshina, N.; Haugland, R. P., A stable nonfluorescent derivative of resorufin for the fluorometric determination of trace hydrogen peroxide: Applications in detecting the activity of phagocyte NADPH oxidase and other oxidases. Anal. Biochem. 1997, 253, 162-168. 34. Boland, N. E. Quantification of Nanomolar Superoxide in Aqueous Solution: Flow Injection Analysis Using the Chemiluminescent Reagent MCLA. Honours thesis, Colby College, Waterville, Maine, 2001. 35. Johnson, K. A.; Simpson, Z. B.; Blom, T., Global Kinetic Explorer: A new computer program for dynamic simulation and fitting of kinetic data. . Anal. Biochem. 2009, 387, 2029. 36. Dalrymple, R. M., Correlations between Dissolved Organic Matter Optical Properties and Quantum Yields of Singlet Oxygen and Hydrogen Peroxide. Environ Sci Technol 2010, 44, 5824-5829. 37. Goldstone, J. V.; Pullin, M. J.; Bertilsson, S.; Voelker, B. M., Reactions of hydroxyl radical with humic substances: bleaching, mineralization, and production of bioavailable carbon substrates. Environ. Sci. Technol. 2002, 36, 362-372. 38. Garg, S.; Rose, A. L.; Waite, T. D., Photochemical Production of Superoxide and Hydrogen Peroxide from Natural Organic Matter Geochim Cosmochim Acta 2011, 74, 43104320. 39. Goldstone, J. V.; Voelker, B. M., Chemistry of superoxide radical in seawater: CDOM associated sink of superoxide in coastal waters. Environ. Sci. Technol. 2000, 34, 1043-1048. 40. Bligh, M. W.; Waite, T. D., Formation,reactivity, and aging of ferric oxide particles formed from Fe(II) and Fe(III) sources: Implications for iron bioavailability in the marine environment. Geochim Cosmochim Acta 2011, 75, 7741-7758. 41. Fujii, M.; Rose, A. L.; Waite, T. D.; Omura, T., Superoxide-Mediated Dissolution of Amorphous Ferric Oxyhydroxide in Seawater Environ. Sci. Technol. 2006, 40, 880-887. 42. Waite, T. D.; Morel, F. M. M., Photoreductive dissolution of colloidal iron oxides in natural waters. Environ. Sci. Technol. 1984, 18, 860-868. 43. Faust, B. C., A review of the photochemical redox reactions of iron(III) species in atmospheric, oceanic, and surface waters: influences on geochemical cycles and oxidant


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formation. In Aquatic and Surface Photochemistry, Helz, G. R.; Zepp, R. G.; Crosby, D. G., Eds. CRC Press: Boca Raton, Florida, 1994; pp 3-37. 44. Khaiken, G. I.; Alfassi, Z. B.; Huie, R. E.; Neta, P., Oxidation of Ferrous and Ferrocyanide Ions by Peroxyl Radicals. J Phys Chem 1996, 100, 7072-7077. 45. Chevallier, E.; Joliboisa, R. D.; Meuniera, N.; Carliera, P.; Monodb, A., Fenton-like’’ reactions ofmethylhydroperoxide and ethylhydroperoxide with Fe2+ in liquid aerosols under tropospheric conditions. Atmos. Environ. 2004, 38, 921-933. 46. Borer, P.; Kraemer, S. M.; Sulzberger, B.; Hug, S. J.; Kretzschmar, R., Photodissolution of lepidocrocite (γ-FeOOH) in the presence of desferrioxamine B and aerobactin. Geochim Cosmochim Acta 2009, 73, 4673-4687. 47. Rose, A. L.; Waite, T. D., Kinetics of hydrolysis and precipitation of ferric iron in seawater. Environ. Sci. Technol. 2003, 37, 3897-3903. 48. Fujii, M.; Imaoka, A.; Yoshimura, C.; Waite, T. D., Effects of Molecular Composition of Natural Organic Matter on Ferric Iron Complexation at Circumneutral pH. Environ. Sci. Technol 2014, 48, 4414-4424. 49. Garg, S.; Jiang, C.; Miller, C. J.; Rose, A. L.; Waite, T. D., Iron Redox Transformations in Continuously Photolyzed Acidic Solutions Containing Natural Organic Matter: Kinetic and Mechanistic Insights. Environ. Sci. Technol 2013, 47, 9190-9197. 50. Rose, A. L.; Waite, T. D., Kinetics of iron complexation by dissolved natural organic matter in coastal waters. Mar Chem 2003, 84, 85-103. 51. Pham, A. N.; Rose, A. L.; Feitz, A. J.; Waite, T. D., Kinetics of Fe(III) precipitation in aqueous solutions at pH 6.0-9.5 and 25°C. Geochim Cosmochim Acta 2006, 70, (3), 640650. 52. Bligh, M. W.; Waite, T. D., Formation, aggregation and reactivity of amorphous ferric oxyhydroxides on dissociation of Fe(III)-organic complexes in dilute aqueous suspensions. Geochim Cosmochim Acta 2010, 74, 5746-5762.

785 786



787

Figure Captions

788

Figure 1: Photochemical generation of superoxide (squares) and hydrogen peroxide (circles)

789

by algal exudate at (a) pH 4 and (b) pH 8. Superoxide measurement was not performed at pH

790

4 due to its short lifetime at pH 4. Symbols represent experimental data (average of triplicate

791

measurements) while lines represent model values. Error bars represent standard deviation of

792

triplicate measurements.

793

Figure 2: (a) Decrease in superoxide concentration on addition of 450 nM O•− 2 to pH 8

794

exudate-free medium (triangles) and algal exudate (squares) in dark. Circles represent

795

•− decrease in O•− 2 concentration on addition of 200 nM O 2 to algal exudate in dark. Symbols

796

represent measured data while lines represent extrapolated values based on the measured data

797

points. (b) Concentration of H2O2 generated on O•− 2 decay when added to algal exudate at

798

pH 8 in dark for various initial O•− 2 concentration as indicated on the plot.

799

Figure 3: Superoxide concentration remaining after 30s (solid bars) on addition of 475 nM

800

•− O•− 2 to pH 8 algal exudate irradiated for a certain time prior to addition of O 2 in dark. Final

801

concentration of H2O2 generated on decay of 475 nM O•− 2 (open bars) when added to pH 8

802

algal exudate irradiated for a certain time prior to addition of O•− 2 in dark.

803

Figure 4: Photochemical generation of Fe(II) on addition of 100 nM (squares), 50 nM

804

(circles) and 25 nM (triangles) Fe(III) to fresh algal exudate at pH 4 (panel a) and pH 8 (panel

805

b). Diamonds in panel b represent photochemical generation of Fe(II) on addition of 100 nM

806

Fe(III) to fresh algal exudate containing 30 kU.L-1 SOD (diamonds) at pH 8. Symbols

807

represent experimental data (average of triplicate measurements); lines represent model

808

values. Error bars represent 95% confidence interval of triplicate measurements.


Page 32 of 41

Page 33 of 41


809

Figure 5: (a) Decrease in Fe(II) concentration on addition of 150 nM (squares), 100 nM

810

(circles), 50 nM (triangles) Fe(II) to fresh exudate at pH 8 under irradiated conditions.(b)

811

Decrease in 100 nM Fe(II) concentration in the presence of organic exudate containing 0

812

(squares), 50 nM (circles), 100 nM (triangles) and 150 nM (diamonds) H2O2 under irradiated

813

condition at pH 8. Symbols represent experimental data (average of duplicate measurements);

814

lines represent model values. Error bars represent 95% confidence interval of triplicate

815

measurements.

816

Figure 6: Diurnal cycle of Fe cycling rate in the presence and absence of exudate at pH 8.

817

Normalized values of the light intensity with respect to the peak intensity is shown on the plot

818

assuming sinusoidal variation in light intensity.

819 820

821

822 823



824

Figure 1

825 826 827 828 829 830 831 832 833


Page 34 of 41

Page 35 of 41

834


Figure 2

835 836 837 838 839 840 841 842 843 844



845

Figure 3

846

847 848 849 850 851 852 853 854 855 856 857 858 859 860 861 862 863


Page 36 of 41

Page 37 of 41

864


Figure 4

865

866 867 868 869 870 871 872 873 874



875

Figure 5

876 877 878 879 880 881 882 883 884


Page 38 of 41

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885


Figure 6

886 887 888 889 890 891 892 893 894 895 896 897 898 899 900 901 902



Page 40 of 41

903 904

Table 1. Kinetic model to explain iron redox transformation and ROS generation in the

905

presence of exudate released by M. Aeruginosa

Rate No.

a

Reaction

Constant

Published valuea,b

Rate Constanta

Published valuea,b pH 8

pH 4 Dark Fe transformation 1

Fe(III)L+Ex red  → Fe(II)L+Ex ox

1.2×103

4×103 12

1.2×103

_

2

aging Ex red  → Ex ox

1×10-4 s-1

_

1×10-4 s-1

_

3

Fe(II)L + O2  → Fe(III)L + O•− 2

0.5

_

60

93 19

4

Fe(II)L + H 2O2  → Fe(III)L + OH• + OH −

_

_

1.3×105

1.6×105 19

k

_

_

1.6×107

1.6×107 51

k

_

_

4×106

4×106 52

5

p _ hom Fe(III)' + Fe(III)'  → 2AFO-L

6

p _ het Fe(III)' + AFO-L  → 2AFO-L

7

Fe(III)' + L  → Fe(III)L

≥ 1×105c

_

0.1-5×106c

2.5×106 48

8

Fe(II)' + L  → Fe(II)L

≥ 5×103c

_

≥ 1×105c

12 g-1.L.s-1 50

9

d AFO-L  k → Fe(III) '

_

_

2.3×10-4 s-1

2.3×10-4 s-1 47 d

10

Fe(III)L  → Fe(III) ' + L

≤ 3×10-3

_

≤ 1×10-3

9×10-5 s-1 50

Light -mediated O•− 2 and H2O2 generation and consumption 9×10-11 11

12

hυ

*

O2

•− 2

_

4 ×10-10

_

Q → Q → O

1 1 Ex O•− → H 2O 2 + O 2 2  2 2

M.s-1 0.12

M.s-1 1.1×107 15


0.12

0.12 s-1 39

Page 41 of 41


Ex H 2O2  → O2 /H 2O

13

4×10-4 s-1e

_

_

_

_

_

1.5×10-10

1.7×10-9

M.s-1

M.s-1f38

1×105

1×105 38

hυ Ex  → R•

14

R • + O•− → R − + O2 2 

15

_

_

Light-mediated Fe(III) reduction hυ Fe(III)L  → Fe(II)' +Lox

16

3×10-3s-1

7.5×10-3

_ 3×10-3s-1

s-149 17

hυ AFO-L  → Fe(II)' +OH •

1.2×10-4 s-1g

_

18

hυ Fe(III)'  → Fe(II)'

5×10-4s-1h

5×10-4s-1h

19

Fe(III)L + O•− → Fe(II)L + O2 2 

(3-100)×104

(2.5-5.6)×104 15

20

Fe(III)' + O•− → Fe(II)' + O 2 2 

1.5×108

1.5×108 23

2×106

0.43-2.3×106 5

Light-mediated Fe(II) oxidation

Fe(II)L + O•− → Fe(III)L + H 2O2 2 

21 906 907 908 909 910 911 912 913 914 915 916 917 918 919 920 921 922 923 924

a

the rate constant listed are in units of M-1s-1 unless stated otherwise published value of the rate constant shown here are the values reported in SRFA system unless stated otherwise. c based on L concentration of 65.6 nmoles.mg-1 C of exudate as reported to be the case for SRFA48 d reported for 1 min aged iron oxyhydroxide e calculated based on the measured H2O2 decay rate in exudate solution at pH 4 f calculated using the reported pseudo-first order rate constant the bulk organic concentration involved in R• generation in SRFA solution g calculated based on the measured Fe(II) formation rate on Fe(III) reduction in exudate-free medium at pH 8 h calculated based on the measured Fe(II) formation rate on Fe(III) reduction in exudate-free medium at pH 4 AFO-L- exudate-coated iron oxyhdroxide; Fe(III)L and Fe(II)L- exudate-complexed Fe(III) and Fe(II) respectively; Fe(III)' and Fe(II)' - inorganic Fe(III) and Fe(II) respectively ; L- Fe binding ligand in exudate ; Ex red - Fe(III) reducing moieties intrinsically present in the b

• • exudate ; Ex - bulk organic moieties in the exudate involved in O•− 2 and R generation; R photo-generated oxidizing organic moieties


Light-Mediated Reactive Oxygen Species ... - ACS Publications

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