Impact of Microcystis aeruginosa Exudate on the Formation and

Subscriber access provided by University of Florida | Smathers Libraries

Article

Impact of Microcystis aeruginosa exudate on the formation and reactivity of iron oxide particles following Fe(II) and Fe(III) addition Shikha Garg, Kai Wang, and T. David Waite Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.7b00660 • Publication Date (Web): 15 Apr 2017 Downloaded from http://pubs.acs.org on April 16, 2017

Just Accepted “Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted online prior to technical editing, formatting for publication and author proofing. The American Chemical Society provides “Just Accepted” as a free service to the research community to expedite the dissemination of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully peer reviewed, but should not be considered the official version of record. They are accessible to all readers and citable by the Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore, the “Just Accepted” Web site may not include all articles that will be published in the journal. After a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web site and published as an ASAP article. Note that technical editing may introduce minor changes to the manuscript text and/or graphics which could affect content, and all legal disclaimers and ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or consequences arising from the use of information contained in these “Just Accepted” manuscripts.

Environmental Science & Technology is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

Page 1 of 38

Environmental Science & Technology

1

Impact of Microcystis aeruginosa exudate on the formation and reactivity of iron oxide

2

particles following Fe(II) and Fe(III) addition

3

Shikha Garg, Kai Wang and T. David Waite*

4

5

6

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

7

Sydney, NSW 2052, Australia

8

9

10

11

Revised

12

Environmental Science and Technology

13

April 2017

14

15

16

17

18

*

Corresponding author: Tel. +61-2-9385 5060; Email [email protected]

1 Environment ACS Paragon Plus


19

ABSTRACT

20

Impact of the organic exudate secreted by a toxic strain of Microcystis aeruginosa on the

21

formation, aggregation and reactivity of iron oxides that are formed on addition of Fe(II) and

22

Fe(III) salts to a solution of the exudate is investigated in this study. The exudate has a

23

stabilizing effect on the particles formed with decreased aggregation rate and increased

24

critical coagulant concentration required for diffusion-limited aggregation to occur. These

25

results suggest that the presence of algal exudates from Microcystis aeruginosa may

26

significantly influence particle aggregation both in natural water bodies where Fe(II)

27

oxidation results in oxide formation and in water treatment where Fe(III) salts are commonly

28

added to aid particle growth and contaminant capture. The exudate also affects the reactivity

29

of iron oxide particles formed with exudate-coated particles undergoing faster dissolution

30

than bare iron oxide particles. This has implications to iron availability, especially where

31

algae procure iron via dissolution of iron oxide particles as a result of either reaction with

32

reducing moieties, light-mediated ligand to metal charge transfer and/or reaction with

33

siderophores. The increased reactivity of exudate-coated particles is attributed, for the most

34

part, to the smaller size of these particles, higher surface area and increased accessibility of

35

surface sites.

36


Page 2 of 38

Page 3 of 38


37

1. INTRODUCTION

38

Amorphous iron oxides (AFO) play an important role in aquatic systems via a range of

39

processes including (i) adsorption of other inorganic and organic compounds, (ii) redox

40

reactions, and (iii) photocatalysis.1 Furthermore, the dissolution of AFO has the potential to

41

provide biologically significant levels of iron for uptake, particularly in conditions where

42

organisms are able to promote dissolution via redox processes and/or production of strong Fe

43

binding ligands.2-4 The formation of AFO in natural waters mostly occurs via oxidation of

44

ferrous iron (Fe(II)) that may enter via groundwater seepage or fluxes from anoxic bottom

45

sediments or may result from photo-reductive processes in sunlit surface waters. The

46

presence of natural organic matter (NOM) is recognized to exert a significant influence on the

47

formation and nature of AFO particles formed. 5-7 The presence of Fe complexing ligands in

48

NOM prevents a portion of the incoming inorganic Fe(II) (i.e. Fe(II)' ) from transforming

49

directly to AFO by binding Fe(III) formed on oxidation of Fe(II)' .8 The relative proportions

50

of the products (i.e. Fe(III)-NOM complex and AFO) resulting from Fe(II) oxidation depends

51

on the NOM concentration and the Fe binding characteristics of the NOM.8 The nature of the

52

AFO particles formed in the presence of NOM differs from that formed in the absence of

53

NOM. For example, Mylon and co-workers5 observed a significant difference in the colloid

54

stability of NOM-coated hematite particles compared to bare hematite particles in estuarine

55

environments. The presence of NOM in groundwater is also known to stabilize the AFO

56

formed on Fe(II) oxidation under these conditions.9 In water treatment systems, where

57

formation of iron oxides is induced via addition of Fe(III) salts to facilitate particle growth

58

and contaminant (including NOM) capture, the presence of NOM will cause formation of

59

AFO-NOM assemblages, the nature of which (in terms of size and reactivity) may differ from



Page 4 of 38

60

that of the bare AFO particles. In water treatment systems, the formation of AFO-NOM

61

assemblages is indeed favored for NOM removal.

62

There have been numerous studies of the means by which NOM influences the nature of the

63

iron oxide particles formed.5,

64

NOM molecules adsorb onto the surface of iron oxide particles 12 thereby changing important

65

factors that impact their aggregation, such as surface charge and/or hydrodynamic radius.

66

13, 14

67

concentrations stabilize iron oxide particles (towards aggregation) due to electrostatic

68

repulsion between negatively charged surfaces15 while steric effects result in the stabilization

69

of particles in the presence of excess high molecular weight humic substances.16 In contrast,

70

colloidal organic carbon, especially when present as chain like assemblages, increases

71

aggregation of iron oxides through formation of bridges.17 The presence of NOM not only

72

influences the kinetics of aggregate formation but also the structure of the aggregates formed.

73

Amal and co-workers showed that the fractal dimension of hematite aggregates that are

74

partially coated with fulvic acid molecules is higher than the fractal dimension of

75

assemblages formed in the absence of fulvic acid due to restructuring within the aggregates.10

76

However, at high fulvic acid concentration, bridging flocculation occurs resulting in more

77

open structures and hence lower fractal dimension of the aggregates.10 As a consequence of

78

changes in iron oxide stability and aggregate structure in the presence of NOM, the fate and

79

transport of these particles and associated contaminants in natural systems will also be

80

impacted. Furthermore, the ability of these modified AFO particles to supply bioavailable

81

iron may differ in comparison to pure oxides.

82

While there have been many studies of the impact of terrigenous NOM, including humic and

83

fulvic acids, on the formation, transformation and reactivity of iron oxide particles, there has

84

been very limited work investigating the effect of organic matter excreted by algae (algogenic

10-17

In cases where iron oxide particles are already present,

11,

Earlier studies have shown that the presence of fulvic acids at environmentally relevant


Page 5 of 38


85

NOM) on iron oxide formation and growth. Also, the currently available body of work

86

mainly describes the impact of NOM on iron oxide aggregation kinetics with limited research

87

on impact of NOM on the reactivity of these particles. The impact of algogenic NOM on iron

88

oxide particle characteristics will be of potential significance to iron bioavailability to algae

89

in natural waters and may also be important to iron oxide aggregation behaviour when iron

90

salts are added as coagulants in water treatment. Thus, in this study, we study the impact of

91

organic matter secreted by the freshwater alga Microcystis aeruginosa on the formation and

92

nature of the AFO particles formed. As implied above, the nature of the particles formed in

93

the presence of organic exudate will not only control the levels of labile iron in the milieu

94

surrounding these algae but may also impact the removal of particles in the water treatment

95

systems receiving waters from sources impacted by algal blooms. We employ dynamic light

96

scattering technique to study the aggregation kinetics and aggregate structure of exudate–

97

coated iron oxide and bare iron oxide particles. We also measure the reactivity of bare and

98

exudate-coated iron oxide using dissolution techniques since this process is the main

99

determinant of iron bioavailability and is commonly used as a measure of reactivity of iron

100

oxides. Since iron oxides in natural and engineered systems may be formed from either

101

Fe(II) or Fe(III) sources, we investigate the impact of algal exudate on the aggregation and

102

reactivity of iron oxides formed from both oxidation states of Fe. All experiments were

103

performed at pH 8 since this is representative of the conditions in the growth medium and,

104

more generally, reasonably representative of natural waters. Based on our experimental

105

results, we provide insight into the mechanism by which the exudate affects the formation,

106

aggregation and reactivity of iron oxides. We also measured the impact of Suwannee River

107

Fulvic acid (SRFA), a widely used model NOM, on iron oxide aggregation and reactivity and

108

compare these findings with the results obtained in the presence of algal exudate.



109

2. EXPERIMENTAL METHODS

110

2.1 Reagents

111

All reagent solutions were prepared using 18 MΩ.cm resistivity deionized water (DW,

112

Millipore) unless stated otherwise. All glassware and plasticware were pre-soaked in 5% HCl

113

for at least 24 h and rinsed at least three times using DW prior to use. All experiments were

114

performed at a controlled room temperature of 22±0.5°C at pH 8. All pH adjustments were

115

made using 1 M HCl (high purity 30% w/v, Sigma) and 1 M NaOH (Sigma) solution. All

116

stock solutions were stored at 4°C when not in use unless stated otherwise.

117

A primary stock solution of 4 mM Fe(II) was prepared in 0.2 M HCl and a 16 µM Fe(II)

118

working stock in 0.8 mM HCl was prepared daily by 250-fold dilution of the primary stock

119

solution by DW. A 0.5 mM Fe(III) working stock in 2 mM HCl was prepared on dilution of a

120

2.0 mM Fe(III) stock solution in 0.2 M HCl daily before experiments. Stock solutions of 0.1

121

M ferrozine (3-(2-pyridyl)-5,6-diphenyl-1,2,4-triazine-4′,4″-disulfonic acid sodium salt;

122

abbreviated as FZ; Sigma) and 0.1 M ascorbate (Sigma) were prepared in DW and adjusted

123

to pH 8. A 2 g.L-1 SRFA (International Humic Substance Society) stock solution was

124

prepared in DW.

125

2.2 Collection of algal exudate

126

Batch cultures of Microcystis aeruginosa were grown in sterilized modified Fraquil medium

127

(Fraquil*; total iron concentration 1 µM; pH 8)18 at 27°C under 90 µmol photons m-2.s-1 of

128

light supplied by cool-white fluorescent tubes on 14 h : 10 h light : dark cycles. During the

129

exponential growth phase, algal cells were collected by centrifuging the batch cultures for 10

130

min at 3500 rpm then washed twice using DW to remove loosely attached adsorbed iron and

131

organic matter. Washed cells were subsequently resuspended in a fresh iron and organic free

132

Fraquil* medium (referred here as “starved medium”; recipe shown in Table S1) to a final


Page 6 of 38

Page 7 of 38


133

density of 7.5×106 - 9×106 cells.mL-1. Iron and organic free medium was used in order to

134

preclude the impact of any pre-added iron and organics on the subsequent formation of iron

135

oxides. The starved culture was grown in the incubator under the conditions described above

136

for 48 h with this time period sufficient for release of a significant amount of exudate without

137

substantial cell rupture. After 48 h, algal cells were removed by filtration through a 0.22 µm

138

PVDF Millipore filter and the filtrate (referred to as “exudate” from here-on) used for further

139

experiments. This procedure of algal exudate preparation avoids introducing any changes to

140

the organic structure that may occur due to preparation techniques such as freeze drying

141

however results in production of only a limited quantity of exudate. The pH of the exudate

142

was adjusted to 8.0 ± 0.03 by addition of 1.5 mM NaHCO3 (Sigma) followed by 1 M HCl to

143

trim to the target value (if necessary) immediately before the AFO formation studies. The

144

total dissolved organic carbon (DOC) concentration of the exudate obtained was always in

145

the range 2.0 - 3.0 mg L-1 with similar cell status and cell density used at all times. No Fe was

146

detected in the exudate solution.

147 148

2.2 Characterization of exudate

149

The exudate characterisation was performed using a liquid chromatography-organic carbon

150

and nitrogen detector (LC-OCND) system. In LC-OCND, the fractionation is based on steric

151

interaction over a wide range of molecular weights with the various fractions eluted from the

152

column detected using three inline detectors, an organic carbon detector (OCD), UV

153

absorbance detector (UVD) and organic nitrogen detector (OND). LC-OCND separates DOC

154

into five different fractions, including biopolymers (Mw> 20000 g.mol−1), humic substances

155

(Mw ~ 500 to 1000 g.mol−1), building blocks (Mw ~ 300-500 g.mol−1), low molecular weight

156

(LMW) acids (Mw < 350 g.mol−1) and LMW neutrals (Mw < 350 g.mol−1). The sum of these

157

five fractions is termed chromatographable DOC. The hydrophobic DOC content was also



158

calculated based on the difference between measured DOC and chromotographable DOC.

159

More detailed description of the method used and analysis of the results obtained is provided

160

in Supporting Information (section S2).

161 162

2.3 Formation of AFO

163

AFO was formed as either exudate-coated AFO (denoted here as AFO-L) or AFO by adding

164

Fe (1-20 µM) to algal exudate and starved Fraquil* medium respectively (starved Fraquil*

165

medium has the same background composition as the algal exudate). Furthermore, both

166

coated and bare AFO particles were formed by addition of Fe(II) or Fe(III) and referred to

167

here as AFO(II) and AFO(III) respectively. While, the formation of AFO(III) occurs

168

instantaneously, the formation of AFO(II) occurs much more slowly due to the requirement

169

for Fe(II) oxidation prior to precipitation of Fe(III). Since the half-life of Fe(II) is expected to

170

be on the order of 1–2 min, 19, 20 >99% of Fe(II) will be oxidised in 10 minutes. As such, the

171

aggregation and dissolution kinetics of AFO(II) were measured 10 minutes after addition of

172

Fe(II) to ensure complete transformation of Fe(II) to AFO. Mixing is an important factor in

173

controlling the polymerization and aggregation behavior of AFO, hence mixing regimes were

174

closely replicated throughout each experiment. A detailed description of the procedure used

175

for preparation and mixing of AFO for dissolution rate and aggregation rate measurements is

176

provided in sections 2.4 and 2.5 respectively. The extent of variability between experiments

177

was assessed by performing experiments in triplicate. The concentration of Fe added was

178

high enough to result in precipitation of at least 95% of the Fe present with the extent of

179

precipitation confirmed by Fe speciation measurement using filtration and acidification

180

methods (see Table S2) as described in detail in our recent study.21


Page 8 of 38

Page 9 of 38


181

All experiments (i.e. aggregation kinetics and dissolution rate measurement) were performed

182

in air–saturated solutions with air saturation assured by maintaining sufficient head-space in

183

the reaction vessels.

184

2.4 Measurement of dissolution rates

185

For measurement of dissolution rates, AFO was formed by adding an appropriate volume of

186

Fe stock to 50 mL of algal exudate or exudate-free medium to achieve a final Fe

187

concentration of 1 µM while stirring rapidly at 1250 rpm using a 1.5 cm stirring bar. After 10

188

s, stirring was stopped and the AFO suspensions allowed to age for a certain time prior to

189

measurement of the reactivity of these oxides. The reactivity of bare and coated AFO was

190

assessed by measuring the dissolution rates of these particles using the reductive dissolution

191

(RD) method described in detail by Bligh and Waite.22 Briefly, ascorbate reduces Fe(III) at

192

the oxide surface23 to form Fe(II) which is trapped by added FZ. Thus, following addition of

193

ascorbate and FZ to the samples, the formation rate of Fe(FZ)3 was monitored by measuring

194

the absorbance at 562 nm using the Ocean Optics Spectrophotometry system described

195

elsewhere.22 The measured absorbance was due to formation of Fe(FZ)3 only with no

196

significant light scattering occurring due to the low concentration of iron oxide particles

197

present. A molar absorption coefficient of 30,000 M-1cm-1 for Fe(FZ)3 was obtained which is

198

consistent with the value reported earlier.24 Note that the reductive-dissolution of AFO by

199

exudate is not important at the time scales investigated here since no formation of Fe(FZ)3

200

was observed in the absence of ascorbate (data not shown). This is consistent with the

201

measured Fe(III) reduction kinetics in the presence of algal exudate at pH 8 in our recent

202

study21 which showed that Fe(II) generation on Fe(III) reduction by exudate occurs very

203

slowly (on the time-scales of hours). Furthermore, the role of exudate as electron shuttles to

204

facilitate AFO reduction by ascorbate also appears unlikely given that the rate of electron

205

transfer by ascorbate to AFO is limited by the rate of adsorption of ascorbate to the iron oxide



Page 10 of 38

206

surface with this process occurring at a similar rate as that of absorption of exudate to the

207

AFO surface. Thus, we conclude that any impact of exudate on AFO dissolution rate

208

observed here is due to changes in the aggregate structure rather than exudate-promoted

209

dissolution of AFO.

210

The rate law for dissolution of AFO 22, 25 is given as

211

 [AFO]  d [AFO] = −kd   [AFO0 ] dt  [AFO0 ] 

212

which, on integration, yields:

213

 [AFO]     [AFO0 ] 

214

where kd is the apparent dissolution rate constant for a given solution composition, [AFO 0 ]

215

is the initial concentration of AFO and γ is an exponent that captures the change in surface

216

area to mass ratio as dissolution proceeds and is expected to be dependent on both primary

217

particle size and size distribution of the particle assemblage.

218

Following addition of ascorbate and FZ to the AFO solution, the concentration of AFO

219

remaining (defined as [AFO 0 ] -[Fe(FZ)3]t) was measured over 5 minutes and eq.(2) fitted to

220

the data to determine the values of γ and kd . Note that no data could be collected for the

221

initial ~60 s due to the time taken for the sample to reach the flow cell used for Fe(FZ)3

222

absorbance measurement; as such, the data point at 60 s was assigned as the time zero of

223

dissolution and the AFO concentration at 60 s was assumed to be equal to [AFO 0 ] (i.e. 1

224

µM) since the resultant value of kd was insensitive to the small changes in the concentration

225

of AFO that occurred during the lag period (i.e. the initial 60 s of reaction). In comparison,

226

the kd value is more prone to error associated with extrapolation of absorbance to t=0 (see

γ

(1)

1−γ

= 1 − kd t (1 − γ )

(2)


Page 11 of 38


227

Supporting Information S4 for more details on the impact of extrapolation on the kd

228

calculation) and hence this was not performed. The same methodology was used to analyse

229

the AFO particles prepared from Fe(II) and Fe(III) sources as well as coated and bare

230

particles and hence any small discrepancy introduced due to the analysis method used here

231

will be similar in all cases. As such, it is reasonable to conclude that the differences in the kd

232

values obtained for the various suspension conditions used are not due to the analysis

233

method.

234

2.5 Dynamic light scattering

235

Dynamic light scattering (DLS) (Malvern Zetasizer Nano S) was used to monitor AFO

236

aggregation kinetics in the presence and absence of exudate. Incident light was produced by a

237

4 mW He−Ne laser with a wavelength (λ) of 663 nm. Scattered light was detected with an

238

avalanche photodiode detector at a scattering angle (θ) of 173°. A digital correlator was used

239

to develop an autocorrelation function that was analyzed using the method of cumulants

240

resulting in a z-averaged diameter dz (apparent intensity weighted mean hydrodynamic

241

diameter). For measurement of aggregation kinetics, an appropriate volume of 2 mM Fe

242

stock was pipetted into a cuvette in order to achieve the required Fe concentration in 3 mL of

243

solution. Subsequently, 3 mL of the starved medium or the algal exudate was added

244

vigorously to the cuvette, mixed rapidly and the particle size measured every 3 minutes. The

245

mixing procedure was closely replicated throughout each experiment. Triplicates of each

246

experimental condition were undertaken in order to account for any variation in the mixing

247

procedure. All measurements were undertaken at 22°C. The same instrument and procedure

248

was used to measure electrophoretic mobility which was subsequently transformed to zeta

249

potential using Smoluchowski’s approximation.

250



Page 12 of 38

251

3 RESULTS and DISCUSSION

252

3.1 Formation and aggregation of iron oxides in the absence and presence of organic

253

exudate secreted by Microcystis aeruginosa at pH 8

254

The rapid formation of particles and initial particle aggregation were examined by monitoring

255

the increase in hydrodynamic diameter (dz) of particles formed over time following addition

256

of Fe to both exudate-free medium (to form bare AFO) and to algal exudate (to form exudate

257

coated AFO) (Figure 1). As shown in Figure 1, the aggregation rate of bare AFO was much

258

higher than the exudate coated AFO particles formed following either Fe(III) or Fe(II)

259

addition. The initial particle aggregation rate is given by26

260

 dd z    ∝~ kN 0  dt t →0

(3)

261

where k is the aggregation rate constant and N0 is the initial number concentration of iron

262

oxide particles. The inverse stability ratio,

263

for a given experimental condition (k) by the rate of change for rapid (diffusion limited)

264

aggregation (kfast) occurring at high ionic strength; i.e.

265

1 , which is defined as the rate of change in dz W

 dd z   dt  1 k  t → 0 = = W k fast  dd z     dt t →0: fast

(4)

266

was calculated by conducting a linear least squares regression analysis of the initial change in

267

dz with time in our medium solution (Figure 1) and medium solution containing excess NaCl.

268

Note that the initial particle aggregation rate for AFO(II) (both bare and exudate coated) was

269

calculated by analysing the particle size data for t ≥10 min when complete formation of AFO

270

has occurred. The data for particle aggregation kinetics at high ionic strength is shown in

271

Supporting Information (Figures S1 and S2 for AFO(III) and AFO(II) respectively). No

272

significant (p>0.5 using single-tailed student’s t-test) impact of NaCl addition was observed


Page 13 of 38


273

on Fe(II) oxidation kinetics (Figure S3) supporting the conclusion that impact of NaCl

274

observed is due to changes in the AFO(II) aggregation behaviour. The minimum ionic

275

strength required for diffusion limited aggregation, i.e. the critical coagulation ionic strength

276

(CCIS), was determined by non-linear least squares fitting of the empirical relationship

277

between the inverse stability ratio and the relevant ionic strength (IS) (eq. 5) as described by

278

Mylon and co-workers;5 i.e.

279

1 1 = ' W 1 + (CCIS/[IS]) β

(5)

280

Note that since the experimental matrix contains both monovalent (Na+) and divalent (Ca2+)

281

cations, ionic strength rather than ionic concentration is used to determine colloidal stability.

282

As shown in Figure 2, a typical stability profile characterized by two regimes, reaction-

283

limited aggregation (RLA) and diffusion-limited aggregation (DLA), consistent with DLVO

284

colloid theory for iron oxide particles, is observed. Within the reaction-limited aggregation

285

regime at low salt concentrations (1/W < 1), the stability (towards aggregation) of the

286

particles decreased with increasing ionic strength leading to increasing rates of particle

287

aggregation. When the ionic strength exceeds CCIS at 1/W > 1, sufficient electrolyte was

288

present to eliminate the repulsive energy barrier between the particles, resulting in rapid

289

aggregation with the rate of change of the hydrodynamic diameter almost independent of the

290

ionic strength. Note that no data for particle aggregation kinetics could be obtained for IS
3 is obtained for exudate coated AFO(III) (data not shown) which

345

cannot be quantified as a mass fractal dimension since 1 ≤ dF ≤ 3. In order to explain this

While the dF for AFO(II) is

with the formation of particularly compact



346

result, we need to consider the assumptions underpinning the calculations. The main

347

assumption is that diffusion limited aggregation occurs with sticking occurring at every

348

collision however this may not be the case. Steric hindrance due to the presence of large

349

exudate molecules may prevent aggregation of particles, even at the high ionic strengths used

350

here. It is possible that diffusion-limited aggregation is also not occurring for exudate coated

351

AFO(II) with the result that, in reality, the mass fractal dimension of exudate coated AFO(II)

352

≤ 2.44. This value of dF is higher than that calculated for bare AFO(II) particles (1.72) and is

353

consistent with the results of an earlier study in which hematite particles partially coated with

354

fulvic acid exhibited a higher fractal dimension than the bare hematite particles due to

355

restructuring within the aggregate 10 however further work is required to accurately determine

356

the fractal dimension of exudate coated AFO and to definitively conclude that the presence of

357

algal exudate affects the aggregate structure of AFO.

358

3.3 Dissolution rates of bare and exudate coated iron oxides at pH 8

359

As shown in the previous sections, the presence of organic exudate has an impact on particle

360

aggregation kinetics as well as the structure of the aggregates formed. This will also impact

361

the accessibility of the surface sites and Fe centres and would be expected to affect the

362

reactivity of these particles. This is supported by our experimental observations (Figure 3)

363

which show that the dissolution rate, kd , of the exudate-coated iron oxide particles, formed

364

from Fe(III) and/or Fe(II) sources, is higher than that of the bare iron oxide particles with the

365

difference in dissolution rate much higher for the freshly formed particles compared to the

366

aged particles (plots showing the dissolution of AFO(II) and AFO(III) (including both bare

367

and exudate coated) from which values of kd were calculated are presented in Figures S7-

368

S10). As mentioned earlier, no reductive or ligand–promoted dissolution of iron oxide was

369

observed to induced by the exudate, at least on the time scales investigated here (data not

370

shown), thereby supporting the conclusion that the difference in the reactivity of bare and


Page 16 of 38

Page 17 of 38


371

coated iron oxide assemblages is due to differences in their physical properties such as

372

aggregate size and/or aggregate structure. The dissolution rates measured here for bare

373

AFO(III) and bare AFO(II) are similar to those previously determined for iron oxide

374

assemblages formed in seawater.22, 30

375

The hypothesis that the higher dissolution rate of the exudate-coated AFO compared to bare

376

iron oxide assemblages is due to the smaller size of the coated iron oxides might be

377

considered to account generally for the results obtained. As reported earlier, the rate of iron

378

oxide dissolution is proportional to the surface site concentration since ascorbate exchanges

379

electrons with Fe(III) atoms following its inner-sphere coordination to the oxide surface.31

380

Thus, the smaller size of exudate-coated iron oxides (Figure 1) may be envisaged to facilitate

381

faster dissolution of the iron oxides as a result of the larger surface area available and the

382

associated higher concentration of accessible Fe centres for formation of Fe-ascorbate surface

383

complexes. While changes in particle size might account for the difference in dissolution

384

rates of bare and exudate-coated iron oxides, particle size differences do not account for the

385

relative reactivity of AFO(II) and AFO(III) since AFO(II) is of smaller size than AFO(III)

386

with this result suggesting that other factors must be important. The lower dissolution rates of

387

AFO(II) compared to AFO(III) (including both bare and coated particles) measured here is

388

consistent with earlier findings and has been attributed to both smaller primary particle size

389

and less ordered molecular structure of AFO(III) with iron oxides formed via Fe(II) oxidation

390

typically exhibiting a more crystalline lepidocrocite structure compared to the ferrihydrite

391

structure of iron oxides formed directly via Fe(III) hydrolysis.22, 27

392

The decrease in the dissolution rate of both bare and coated iron oxide particles on aging

393

possibly occurs due to (i) rapid early aggregation that decreases ascorbate accessibility to

394

surface sites and/or (ii) changes in fractal structure with time.22 The possibility that change in

395

the crystal structure of the iron oxide over time is responsible for changes in their dissolution 17 Environment ACS Paragon Plus


396

rate is not considered here since the results of earlier work showed that the crystal structure of

397

iron oxides is established shortly following nucleation and hence cannot explain rapid early

398

aging of these particles.22 Higher fractal dimension, which represents a less open structure, is

399

expected to result in decreased accessibility to Fe centres and hence lower dissolution rates.

400

However, this is not consistent with our observation that AFO(III), which has a higher fractal

401

dimension than AFO(II), has higher dissolution rates than AFO(II) with this result supporting

402

the conclusion that the fractal structure of these assemblages has little influence on particle

403

reactivity. This is consistent with results of an earlier report which suggested that the impact

404

of differences in aggregate structure on particle reactivity is limited if participation of interior

405

sites in reactions is minimal due to diffusion limitations.32 As we have used FZ as the Fe(II)

406

trapping agent in the studies described here, this may only occur at the particle surface where

407

three molecules of FZ are able to coordinate Fe(II).

408

aggregation is most likely responsible for the observed decrease in dissolution rates with

409

time. Note that changes in the aggregate size is responsible for rapid early aging of the iron

410

oxide particles however there are other factors (such as crystal structure, primary particle

411

size) that are likely to control the overall dissolution rate of iron oxide particles.

412

The effects of aggregation on the reactivity of iron oxides may be modelled by considering

413

that physical loss of accessible surface area is the dominant factor responsible for the

414

decreasing dissolution rates of these particles over time. The surface area of the aggregates

415

over time was quantified using the particle aggregation data for 1 µM Fe (the concentration at

416

which dissolution rates are measured) shown in Figure S11. As described in our earlier

417

work,22 four different assumptions may be used to develop a quantitative measure of the

418

surface area of the fractal aggregates at any time during aggregation; i.e.

419

Model 1: Aggregates are solid smooth spheres with the surface area equivalent to the surface

420

area of the spheres. Hence,

It thus appears that rapid early


Page 18 of 38

Page 19 of 38


421

SATm r1 = SAT 1 rm

422

where SATm and SAT1 represent the surface area at any time during aggregation and initial

423

surface area respectively. r1 and rm represent the initial particle size and particle size at any

424

time during aggregation respectively.

425

Model 2: Aggregates are assumed to be solid spheres but with a fractal surface described by

426

the surface fractal dimension dS ~2.5. Hence,

427

SATm  r1  =  SAT 1  rm 

428

where SATm , SAT1 represent the surface area at any time during aggregation and initial surface

429

area respectively.

430

Model 3: Aggregates are assumed to grow as mass fractals with the mass contained in the

431

outer shell of thickness ∆r assumed to be accessible.

(7)

(3−d s )

(8)

dF

 r − ∆r  1−  m rm   = dF  r1 − ∆r  1−    r1 

432

VTm∆r VT 1∆r

433

where VTm∆r and VT 1∆r represents the volumes contained in the outer shell of thickness ∆r at

434

any time during aggregation and initially respectively.

435

Model 4: Aggregates are assumed to grow as mass fractals with the area of the enclosed

436

sphere assumed to be accessible. Hence,

437

E r  SATm = m  E SAT 1  r1 

(9)

2 −d F

(10)



438

E represent the initial enclosed surface area and enclosed surface area at where SATE1 and SATm

439

any time during aggregation respectively. Detailed description of these models was provided

440

by Bligh and Waite 22 and is also presented in Supporting Information (S4).

441

As shown in Figure 4, the decrease in Fe accessibility predicted by all the aforementioned

442

models displays the same pattern with rapid initial decline followed by a slower decline in a

443

manner similar to the measured profiles for kd as a function of time for both AFO(III) and

444

AFO(II). While the models used here are based on a simple representation of the aggregate

445

structure, the results obtained support the hypothesis that kd is a function of accessible

446

surface area. As mentioned earlier, since coordination of Fe(II) by FZ is restricted by steric

447

factors to exterior sites,33 the dependence of dissolution rates on surface site accessibility

448

seems reasonable though it should be recognised that these models may be deficient if

449

interior sites are involved as well. The modelling results further show that the solid particle

450

model captures the behaviour of the system with the model predicted decrease in dissolution

451

rates over time relatively similar to that observed experimentally. This is consistent with an

452

earlier study which also found that the solid particle model provided the best fit to the

453

formation of iron oxyhydroxide particles on dissociation of Fe(III)-organic complexes.34

454

While the decrease in accessible surface area of the particles due to aggregation describes the

455

dissolution rates of bare iron oxide particles very well, models based on this hypothesis are

456

unsuccessful in accounting for the observed decrease in dissolution rates of exudate-coated

457

AFO (both AFO(III) and AFO(II)) over time (see Figure 4). As shown in Figure S9, the

458

particles do not aggregate noticeably in the presence of exudate, especially at the low Fe

459

concentrations employed here, and hence the surface area of the particles does not change

460

significantly. This observation suggests that other factors play a role in controlling the

461

dissolution rates of exudate-coated iron oxide particles. It is possible that the organic exudate


Page 20 of 38

Page 21 of 38


462

coating rearranges over time providing more uniform surface coverage and, in so doing,

463

decreases the accessibility of FZ and ascorbate to surface Fe due to increased steric

464

hindrance, thereby resulting in a decrease in reactivity of these particles.

465

3.4 Impact of SRFA on iron oxide aggregation and reactivity

466

To compare the behaviour of algal exudate with terrigenous NOM such as SRFA (a material

467

that has been widely used as well-defined representative NOM) we measured the impact of

468

SRFA on iron oxide formation and reactivity under the experimental conditions investigated

469

here. SRFA coated AFO (both AFO(III) and AFO(II)) were prepared by addition of Fe (in the

470

same manner as undertaken for preparation of exudate-coated AFO) to starved Fraquil*

471

medium containing 5 mg.L-1 SRFA. A SRFA concentration of 5 mg.L-1 was used to maintain

472

equivalence with the algal exudate case. As observed in the presence of algal exudate (Figure

473

1), the aggregation rate of iron oxide particles was also reduced significantly in the presence

474

of SRFA (Figure 5) with increase in the stability (towards aggregation) of the iron oxide

475

particles, as indicated by a shift of CCIS to higher values in the presence of SRFA (see Table

476

1) compared to that observed for bare iron oxide. As shown, the CCIS value shifts from 0.7

477

mM to 12.1 mM for AFO(III) while it increases from 5.9 mM to 202.1 mM for AFO(II) in

478

the presence of SRFA. Furthermore, as found for the algal exudate, the reactivity of SRFA

479

coated AFO particles is higher than that of the bare AFO particles (Figure 6).

480

While the overall impact of SRFA on AFO aggregation and reactivity is consistent with that

481

observed in the presence of algal exudate, there are three notable differences. Firstly, particle

482

stabilization by SRFA is affected by the Fe to SRFA concentration ratio with the impact on

483

particle aggregation rate much more prominent at lower Fe concentrations (Figure 5). The

484

impact of [Fe]:[SRFA] on AFO stability is further supported by the shift of CCIS to lower

485

values with increase in Fe concentration (Figure S12). As shown, the CCIS value shifts from



486

12.1±0.5 mM to 6.7±0.7 mM with increase in Fe(III) concentration from 10 µM to 20 µM for

487

AFO(III) while it decreases from 202.0±18.9 mM to 26.3±11.3 mM with increase in Fe(II)

488

concentration from 15 µM to 25 µM for AFO(II) in the presence of SRFA. The influence of

489

SRFA concentration on AFO aggregation is consistent with an earlier report which suggested

490

that SRFA stabilizes iron oxides via charge stabilization15 with increased AFO coverage by

491

SRFA resulting in more charge stabilization and hence higher AFO stability. In contrast, the

492

exudate-coated AFO appears to be stabilized via steric hindrance as is reported to be the case

493

for large molecular weight humic acids16 with [Fe]:[exudate] having little impact on the iron

494

oxide aggregation kinetics. The high average molecular weight of exudate (13,500 g.mol-1)

495

compared to SRFA (567 g.mol-1) and the presence of large molecules such as polysaccharides

496

and proteins in the exudate can be considered to account for the differences in the particle

497

stabilization behaviour of the two organics.

498

The second prominent difference in the behaviour of SRFA and algal exudate is that the

499

reactivity of SRFA-coated AFO changes only during the initial 10 minutes with no

500

subsequent change in the particle reactivity over the next 24 h (Figure 6) while the reactivity

501

of exudate-coated AFO decreases continuously during 24 h (Figure 3). This result, combined

502

with the size data, suggests that SRFA stabilizes AFO (at least at the low [Fe]:[SRFA] ratio

503

employed here) for long periods with the lack of change in AFO particle size resulting in

504

limited loss of accessible surface area with limited resultant change in AFO reactivity. In

505

contrast, even though the exudate stabilizes iron oxide particles, increasing adsorption of

506

exudate and ongoing restructuring over time appears to result in a decrease in Fe accessibility

507

with consequent decrease in reactivity of these particles.

508

Thirdly, the reactivity of SRFA coated AFO(II) is higher than the SRFA coated AFO(III)

509

particles contrary to the behaviour observed for exudate coated and bare AFO particles. The

510

higher reactivity of bare AFO(III) compared to bare AFO(II) particles is usually attributed to 22 Environment ACS Paragon Plus

Page 22 of 38

Page 23 of 38


511

the less ordered molecular structure of AFO(III) with iron oxide particles formed via Fe(II)

512

oxidation typically exhibiting a more crystalline lepidocrocite structure compared to the

513

ferrihydrite structure of iron oxides formed directly via Fe(III) hydrolysis.22,

514

reactivity of SRFA coated AFO(II) is higher than the SRFA coated AFO(III), it appears that

515

the crystal structure of both iron oxide (SRFA-coated AFO(II) and SRFA-coated AFO(III))

516

particles is similar and/or does not play a role in controlling the dissolution rate of these

517

particles.

518

transformation of Fe(III) minerals to more crystalline, thermodynamically stable forms35, 36

519

and hence the presence of SRFA may possibly maintain the particles (both AFO(III) and

520

AFO(II)) in more reactive forms (such as ferrihydrite) with the reactivity of the particles

521

mainly controlled by the accessible surface area which is higher for SRFA-coated AFO(II) as

522

a result of its smaller particle size.

523

The difference in the behaviour of algal exudate and SRFA in controlling iron oxide

524

transformation and reactivity suggests that care should be taken in extrapolating the

525

observations made in the presence of fulvic acids to prediction of the transformation and fate

526

of iron oxide particles in waters impacted by algal blooms.

527

ENVIRONMENTAL IMPLICATIONS

528

Our results show that the presence of organic exudate secreted by Microcystis aeruginosa

529

impacts the formation, aggregation and reactivity of iron oxides formed at pH 8 with the

530

particles formed in the presence of organic exudate exhibiting higher stability (towards

531

aggregation) and reactivity (towards dissolution) than those formed in the absence of organic

532

exudate. The exudate-coated AFO particles undergo aging with decrease in their reactivity

533

over time in a manner similar to bare AFO and to SRFA coated AFO particles. Despite these

534

similarities, the aging mechanism appears to differ between these particle types. The aging of

535

bare AFO and SRFA-coated AFO can be explained purely in terms of a decrease in the

27

Since the

It has been reported previously that the presence of NOM inhibits the



536

accessible surface sites during aggregation but this mechanism appears to be minor with

537

regard to the aging of exudate coated AFO (AFO-L). Rather, increasing adsorption and

538

restructuring appear to be the principal mechanisms involved in aging of AFO-L.

539

The stabilizing effect of the organic exudate on iron oxide formation and growth may well

540

exert a negative impact on the effectiveness of water treatment using ferric iron coagulants as

541

these chemicals are added for the purpose, in part at least, of inducing the formation of large

542

particles that either sediment quickly or are captured effectively on passage through deep bed

543

sand filters. The increased reactivity of these particles may also result in some dissolution

544

within the sand filter where localised concentrations of natural organic matter (including algal

545

exudate) may be relatively high.

546

Particles resulting from the addition of Fe(II) (as may well be the case in a water body where

547

Fe(II) is released from anoxic sediments) are smaller in size and more reactive when formed

548

in the presence compared to the absence of algal exudate though these particles are not as

549

reactive as those formed from the addition of ferric salts, possibly because of the more

550

crystalline nature of the iron oxides formed. The fact that iron oxide particles formed in the

551

presence of algal exudate are significantly more reactive than those formed in the absence of

552

exudate suggests that they will more readily undergo reductive dissolution as a result of both

553

thermal (dark) and photochemical processes with implications to the supply of both major

554

and minor nutrients (such as phosphorus and iron respectively) to the resident biotic

555

population.

556

Since the impact of exudate on AFO aggregation and dissolution rate is similar for the range

557

of [Fe]:[exudate] ratios (2.2-44.8 w/w%) investigated here, the results presented here can be

558

generalized to natural water conditions where [Fe]:[DOC] ratios range between 0.0002 – 15

559

w/w%.37 It is important to note that for conditions where the [Fe]:[DOC] ratio is very low, the


Page 24 of 38

Page 25 of 38


560

exudate will interact with Fe resulting in formation of organically-complexed Fe21 rather than

561

AFO, thus these conditions are not investigated here.

562

Supporting Information

563

Details on addition experimental data on particle aggregation and dissolution kinetics are

564

provided in Supplementary Information which may be accessed free of charge at the ACS

565

website for this journal.

566

Corresponding Author

567

Professor T. David Waite; Email: [email protected]; Phone: +61-2-9385-5060

568

Acknowledgement

569

We acknowledge the assistance of Dr. Mark Bligh in dynamic light scattering experiments

570

and in interpretation of results from particle reactivity studies. We would also like to

571

acknowledge the help provided by Dr. Yuan Wang and Keng Han Tng for their help with

572

LC-OCND analysis. Financial support provided by the Australian Research Council through

573

Discovery Project DP120103234 and DECRA Award DE12102967 is gratefully

574

acknowledged.

575

References

576 577 578 579 580 581 582 583 584 585 586 587 588 589

1. Brown, G. E.; Henrich, V. E.; Casey, W. H.; Clark, D. L.; Eggleston, C.; Felmy, A.; Goodman, D. W.; Gratzel, M.; Maciel, G.; McCarthy, M. I.; Nealson, K. H.; Sverjensky, D. A.; Toney, M. F.; Zachara, J. M., Metal oxide surfaces and their interactions with aqueous solutions and microbial organisms Chem. Rev. 1999, 99, 77-174. 2. Kuma, K.; Matsunaga, K., Availability of colloidal ferric oxides to coastal marinephytoplankton. Mar Biol 1995, 122, 1-11. 3. 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. 4. Kraemer, S. M. In Iron oxide dissolution and solubility in the presence of siderophores, Symposium on Biogeochemical Controls on the Mobility and Bioavailability of, 2004; 2004. 5. Mylon, S. E.; Chen, K. L.; Elimelech, M., Influence of Natural Organic Matter and Ionic composition on the kinetics and structure of hematite colloid aggregation: Implications to iron depletion in estuaries. Langmuir 2004, 20, 9000-9006.



590 591 592 593 594 595 596 597 598 599 600 601 602 603 604 605 606 607 608 609 610 611 612 613 614 615 616 617 618 619 620 621 622 623 624 625 626 627 628 629 630 631 632 633 634 635 636 637 638 639

6. Elimelech, M.; Gregory, J.; Jia, X.; Williams, R. J., Particle deposition & aggregation: Measurement, Modelling and Simulation. Oxford: Butterworth-Heinemann, 1995. 7. Mayer, L. M., Aggregation of colloidal iron during estuarine mixing: Kinetics, mechanism, and seasonality. Geochim Cosmochim Acta 1982, 46, 2527-2535. 8. Bligh, M. W.; Waite, T. D., Role of Heterogeneous Precipitation in Determining the Nature of Products Formed on Oxidation of Fe(II) in Seawater Containing Natural Organic Matter. Environ. Sci. Technol 2010, 44, 6667-6673. 9. Liang, L.; MCCARTHY, J. F.; JOLLEY, L. W.; MCNABB, J. A.; MEHLHORN, T. L., Iron dynamics: Transformation of Fe( II) / Fe( III) during injection of natural organic matter in a sandy aquifer. Geochim Cosmochim Acta 1993, 57, 1987-1999 10. Amal, R.; Raper, J. A.; Waite, T. D., Effect of Fulvic Acid Adsorption on the Aggregation Kinetics and Structure of Hematite Particles. J. Colloid Interface Sci. 1992, 151, 244-257. 11. Amirbahaman, A.; Olson, T. M., Transport of humic matter-coated hematite in packed beds. Environ. Sci. Technol 1993, 27, 2807-2813. 12. Tipping, E., The adsorption of aquatic humic substances by iron oxides. Geochim Cosmochim Acta 1981, 45, 191-199. 13. Franchi, A.; O’Melia, C., Effects of Natural Organic Matter and Solution Chemistry on the Deposition and Reentrainment of Colloids in Porous Media. Environ. Sci. Technol 2003, 37, 1122-1129. 14. Tiller, C. L.; O’Melia, C., Natural organic matter and colloidal stability: Models and measurements. Colloids Surf. A 1993, 73, 89-102. 15. Liang, L.; Morgan, J. J., Chemical modelling of aqueous systems II. In ACS Symp. Ser., American Chemical Society: Washington DC, 1990. 16. Tipping, E.; Higgins, D. C., The effect of adsorbed humic substances on the colloid stability of haematite particles. Colloids Surfaces 1982, 5, 85-92. 17. Wilkinson, K. J.; Negre, J.-C.; Buffle, J., Coagulation of colloidal material in surface waters: the role of natural organic matter. J Contam Hydrol 1997, 26, 229-243. 18. Fujii, M.; Rose, A. L.; Omura, T.; Waite, T. D., Effect of Fe(II) and Fe(III) Transformation Kinetics on Iron Acquisition by a Toxic Strain of Microcystis aeruginosa. Environmental Science & Technology 2010, 44, (6), 1980-1986. 19. 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. 20. Santana-Casiano, J. M.; González-Dávila, M.; Millero, F. J., Oxidation of nanomolar levels of Fe(II) with oxygen in natural waters. Environ. Sci. Technol. 2005, 39, 2073-2079. 21. 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 In press. 22. 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. 23. Sulzberger, B.; Suter, D.; Siffert, C.; Banwart, S.; Stumm, W., Dissolution of Fe(III)(hydr)oxides in natural-waters; laboratory assessment on the kinetics controlled by surface coordination Mar Chem 1989, 28, 127-144. 24. Stookey, L. L., Ferrozine: a new spectrophotometric reagent for iron. Anal Chem 1970, 42, 779-781. 25. Larsen, O.; Postma, D., Kinetics of reductive bulk dissolution of lepidocrocite, ferrihydrite and goethite. Geochim Cosmochim Acta 2001, 65, 1367-1379.


Page 26 of 38

Page 27 of 38

640 641 642 643 644 645 646 647 648 649 650 651 652 653 654 655 656 657 658 659 660 661 662 663 664 665 666 667 668 669 670 671 672


26. Schudel, M.; Behrens, S. H.; Holthoff, H.; Kretzschmar, R.; Borkovec, M. J., Absolute Aggregation Rate Constants of Hematite Particles in Aqueous Suspensions: A Comparison of Two Different Surface Morphologies. Jounal of Colloid Interface Science 1997, 196, 241-253. 27. Boland, D. D.; Collins, R. N.; Miller, C. J.; Glover, C. J.; Waite, T. D., Factors controlling the Fe(II)-accelerated transformation of ferrihydrite to lepidocrocite and goethite. Environ. Sci. Technol 2014, 48, 5477 – 5485. 28. Meakin, P.; Vicsek, T.; Family, F., Dynamic cluster-size distribution in cluster-cluster aggregation -effects of cluster diffusivity Phy. Rev. B. 1985, 31, 564-569. 29. Amal, R.; Raper, J. A.; Waite, T. D., Fractal structure of hematite aggregates J. Colloid Interface Sci. 1990, 140, 158-168. 30. Rose, A. L.; Waite, T. D., Kinetics of hydrolysis and precipitation of ferric iron in seawater. Environ. Sci. Technol. 2003, 37, 3897-3903. 31. Suter, D.; Banwart, S.; Stumm, W., Dissolution of hydrous iron(III) oxides by reductive mechanisms. Langmuir 1991, 7, 809-813. 32. Liu, J.; Aruguete, D. M.; Murayama, M.; Hochella, M. F., Influence of size and aggregation on the reactivity of an environmentally and industrially relevant nanomaterial (PbS). Environ. Sci. Technol 2009, 43, 8178-8183. 33. Thompsen, J. C.; Mottola, H. A., Kinetics of the complexation of iron(II) with ferrozine Anal Chem 1984, 56, 755-757. 34. 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. 35. Jones, A. M.; Collins, R. N.; Rose, J.; Waite, T. D., The effect of silica and natural organic matter on the Fe(II)-catalysed transformation and reactivity of Fe(III) minerals. Geochim Cosmochim Acta 2009, 73, 4409–4422. 36. Schwertmann, U.; Thalmann, H., The influence of [Fe(II)], [Si] and pH on the formation of lepidocrocite and ferrihydrite during the oxidation of aqueous FeCl2 solutions. Clay Miner. 1976, 11, 189–200. 37. Kikuchi, T.; Fujii, M.; Terao, K.; Jiwei, R.; Lee, Y. P.; Yoshimura, C., Correlations between aromaticity of dissolved organic matter and trace metal concentrations in natural and effluent waters: A case study in the Sagami River Basin , Japan. Sci. Total Environ. 2017, 576, 36-45.

673 674



Page 28 of 38

675

Table 1: CCIS for the bare and coated iron oxide (shown as vertical lines in Figure 2 and

676

S10)

677

a

Fe concentration

CCIS for

CCIS for

(µM)

AFO(III) (mM)

AFO(II) (mM)

Bare iron oxide

10

0.7 ± 0.6

5.9 ± 2.1

Exudate coated iron oxide

10

14.0 ± 0.2

32.4 ± 4.1

SRFA coated iron oxide

10

12.1± 0.5

202.1± 4.1a

total Fe concentration is 15 µM

678

679


Page 29 of 38

680


Table 2: Fractal dimension of the bare and exudate coated iron oxide assemblages Fe concentration AFO(III)

681

(µM)

682

683

AFO(II)

Bare iron oxide

1

2.75 ± 0.22 1.93± 0.12

5

2.97 ± 0.15 1.67± 0.22

10

-

1.58± 0.32

-

-

-

2.37± 0.52

10

-

2.51± 0.24

5

2.64± 0.22

15

2.85± 0.32

1.14± 0.20

25

-

1.33± 0.15

Exudate coated 1 Iron oxide 5

SRFA coated iron oxide

684

685



686

Figure Captions

687

Figure 1: Size development of iron oxide particles formed in the absence (panel a) and

688

presence (panel b) of Microcystis aeruginosa exudate formed on addition of 10 µM (circles),

689

15 µM (squares) and 20 µM (triangles) Fe(III). Size development of iron oxide particles

690

formed in the absence (panel c) and presence (panel d) of Microcystis aeruginosa exudate

691

formed on addition of 10 µM (circles), 15 µM (squares) and 20 µM (triangles) Fe(II). Points

692

represent average of triplicate measurements; error bars represent the standard deviation of

693

triplicate measurements.

694

Figure 2: Representative particle stability plot for iron oxide particles formed in the absence

695

(triangles) and presence (squares) of organic exudate secreted by Microcystis aeruginosa

696

formed on addition of 10 µM Fe(III) (panel a) and 10 µM Fe(II) (panel b). The vertical

697

dashed lines indicate the CCIS for both AFO and AFO-L. Note that the x-axis scale was

698

selected to clearly show the empirical fits obtained using eq.(5) to the particle stability plot.

699

Figure 3: Dissolution rate constant of 1 µM (a) AFO(III)) and (b) AFO(II) following their

700

formation in the absence (triangles) and presence (squares) of organic exudate secreted by

701

Microcystis aeruginosa. Dissolution rate constants are determined from the rate of formation

702

of the Fe(FZ)3 complex following addition of 1 mM ascorbate and 1 mM FZ to AFO samples

703

at various times between 1 min (for AFO(III) or 15 min(AFO(II)) and 24 h.

704

Figure 4: Normalized decrease in dissolution rate constant and loss of surface area due to

705

aggregation. Symbols represent experimentally measured dissolution rate constant while lines

706

represent normalized surface area using various assumption. For mass fractals, a fractal

707

dimension Df = 2.2 and DS= 2.5 was used

708


Page 30 of 38

Page 31 of 38


709

Figure 5: (a) Size development of iron oxide particles formed in the presence of 5 mg.L-1

710

SRFA formed on addition of 10 µM (circles), 15 µM (squares) and 20 µM (triangles) Fe(III).

711

(b) Size development of iron oxide particles formed in the presence of 5 mg.L-1 SRFA

712

formed on addition of 15 µM (circles), 20 µM (squares) and 25 µM (triangles) Fe(II). Dashed

713

lines represent the size of bare iron oxide particles (from Figure 1) for comparison (note that

714

the size of bare AFO does not change with Fe concentration at least in the concentration

715

range investigated here). Points represent average of triplicate measurements; error bars

716

represent the standard deviation of triplicate measurements.

717

Figure 6: Dissolution rate constant of 1 µM AFO(III) (panel a) and AFO(II) (panel b)

718

following their formation in the presence of 5 mg.L-1 SRFA. Dissolution rate constants are

719

determined from the rate of formation of the Fe(FZ)3 complex following addition of 1 mM

720

ascorbate and 1 mM FZ to AFO samples at various times between 1 min (for AFO(III) or 15

721

min(AFO(II)) and 24 h.

722



723

Figure 1

724

725

726

727

728

729

730

731


Page 32 of 38

Page 33 of 38

732


Figure 2

733

734

735

736

737

738

739

740

741

742

743

744

745



746

Figure 3

747

748

749

750

751

752

753

754

755

756

757

758

759


Page 34 of 38

Page 35 of 38

760


Figure 4

761

762

763

764

765

766

767

768



769

Figure 5

770

771


Page 36 of 38

Page 37 of 38

772


Figure 6

773

774



775

Table of Contents graphic

776


Page 38 of 38

Impact of Microcystis aeruginosa Exudate on the Formation and

Recommend Documents