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Oxidative dissolution of silver nanoparticles by chlorine: Implications to silver nanoparticle fate and toxicity Shikha Garg, Hongyan Rong, Christopher J. Miller, and T. David Waite Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.6b00037 • Publication Date (Web): 17 Mar 2016 Downloaded from http://pubs.acs.org on March 21, 2016

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Environmental Science & Technology

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Oxidative dissolution of silver nanoparticles by chlorine: Implications to

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silver nanoparticle fate and toxicity

3 Shikha Garg, Hongyan Rong, Christopher J. Miller and T. David Waite*

4 5 6

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

7

NSW 2052, Australia

8 9 10

Revised

11

Environmental Science and Technology

12

March 2016

13 14 15 16 17

*

18

[email protected]

Corresponding author: Tel. +61-2-9385 5060; FAX +61-2-9385 6139; Email

1 Environment ACS Paragon Plus


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ABSTRACT

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The kinetics of oxidative dissolution of silver nanoparticles (AgNPs) by chlorine is

21

investigated in this work with results showing that AgNPs are oxidized in the presence of

22

chlorine at a much faster rate than observed in the presence of dioxygen and/or hydrogen

23

peroxide. The oxidation of AgNPs by chlorine occurs in air-saturated solution in

24

stoichiometric amounts with two moles of AgNPs oxidized for each mole of chlorine added.

25

Dioxygen plays an important role in OCl − -mediated AgNPs oxidation, especially at lower

26

OCl − concentrations, with the mechanism shifting from stoichiometric oxidation of AgNPs

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by OCl− in the presence of dioxygen to catalytic removal of OCl− by AgNPs in the absence of

28

dioxygen. These results suggest that the presence of chlorine will mitigate AgNPs toxicity by

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forming less-reactive AgCl(s) following AgNPs oxidation though the disinfection efficiency

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of OCl − may not be significantly impacted by the presence of AgNPs since a chlorine-

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containing species is formed on OCl − decay that has significant oxidizing capacity. Our

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results further suggest that the antibacterial efficacy of nanosilver particles embedded on

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fabrics may be negated when treated with detergents containing strong oxidants such as

34

chlorine.

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1. INTRODUCTION

38

Silver nanoparticles (AgNPs) possess unique antimicrobial properties and hence are one of

39

the most widely used nanoparticles in a range of consumer goods such as food packaging,

40

clothing, medical devices and cleaning agents.1, 2 The worldwide production of AgNPs was

41

approximately ~320 tons/year in 2011

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incorporated into more and more consumer products. Due to their widespread use, AgNPs

43

will inevitably be released into the environment1 and may have an impact on human and

44

ecological health as these particles have been shown to be toxic to bacteria,4 algae,5 and other

45

aquatic organisms1,

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continued use of AgNPs will lead to the development of resistance among harmful bacteria to

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AgNPs and may also damage or alter beneficial microbial communities.8-10 Additionally,

48

there are concerns that oxidation of AgNPs could diminish their persistence and potency.

49

Although oxidative dissolution of AgNPs initially results in formation of even more toxic

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Ag(I) species,11 the presence of chloride and/or sulfide in waters may result in the

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precipitation of AgCl(s) and/or Ag2S(s) respectively thereby minimizing Ag(I) toxicity.12-14

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Even though the presence of excess chloride (as is the case in marine waters) may result in

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formation of dissolved Ag-Cl species ( AgCl 2 , AgCl 3 ) rather than AgCl(s), the toxicity of Ag

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would still be expected to decrease as reported earlier.15,16 Thus, to evaluate the

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environmental and human health risks associated with the use of these nanoparticles, there is

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a need to understand the redox transformation that these particles undergo under various

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environmental conditions. It is also important that AgNPs transformations under common

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aqueous conditions be understood to ensure lasting efficacy when these particles are

59

employed as bactericides in commercial products or biomedical applications.

5

3

with this value increasing as nano silver is

as well as to human cells.6,

−

7

Furthermore, there is a risk that the

2−



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Investigators have recently shown that treatment of fabric embedded with AgNPs with

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chlorine-containing solution results in more than 50% of the antibacterial AgNPs being

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transformed to less-reactive AgCl(s).17 Other workers have suggested that AgNPs are

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removed during chlorine disinfection via oxidative-dissolution of the nanoparticles however

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neither the mechanism nor timescale of removal are particularly clear.

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study, we investigate the kinetics and mechanism of reaction of AgNPs with chlorine, and,

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based on our experimental results, we highlight the key features of the mechanism of AgNPs

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oxidation by chlorine and discuss the implications of these results to AgNPs fate and toxicity

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under aqueous conditions typical of natural waters.

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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 unless stated

72

otherwise. All experiments were carried out in 2 mM NaHCO3 buffer at pH 8 unless

73

specifically stated. Solution pH was adjusted by addition of either 1 M HNO3 or 1 M NaOH.

74

A maximum pH variation of ± 0.1 units was considered acceptable during experiments. A 50

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mM stock solution of trisodium citrate, sodium borohydride (NaBH4) and Ag(I) (added as

76

AgNO3) was prepared in Milli-Q water and stored at 4°C when not in use. A stock solution

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containing ~ 0.2 mM citrate-stabilized AgNPs was prepared as described earlier.19 Briefly,

78

0.24 mL of 50 mM Ag(I) stock solution was added dropwise under vigorous stirring to 59.8

79

mL of solution containing 0.4 mM of NaBH4 and 0.6 mM of sodium citrate kept in an ice-

80

bath forming AgNPs. Following 3 h of additional stirring at room temperature, soluble

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byproducts were removed by centrifugal ultraﬁltration (Amicon Ultra-15 3K, Millipore, MA)

82

and Milli-Q water addition in two cycles, after which the ~0.2 mM AgNP stock suspensions

83

were stored at 4 °C for later use. The z-average diameter, dz(the intensity weighted mean


18

As such, in this

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hydrodynamic diameter) was determined to be 48.1 ± 3.7 nm using Zetasizer Nano S

85

(Malvern) with a 633 nm laser source and detection angle of 1730. A 6 mM N,N-diethyl-p-

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phenylenediamine (DPD; Fluka Analytical) stock solution was prepared in 50 mM H2SO4

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solution (pH ≈ 1) to prevent DPD auto-oxidation. An approximately 6 mM NaOCl stock

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solution was prepared by 100-fold dilution of concentrated sodium hypochlorite (Sigma) and

89

was standardized by measuring the UV absorbance in the range 280 - 300 nm.20 A stock

90

solution of 0.55 mM phthalhydrazide (Sigma) was prepared in a solution containing 2 mM

91

NaHCO3 with the final solution pH adjusted to 8.21 All experiments were performed in air-

92

saturated solution unless stated otherwise. All experiments were performed at controlled

93

room temperature of 22°C in plastic bottles covered with aluminium foil to avoid interaction

94

with the ambient light. Samples were stirred continuously during the duration of the

95

experiments. For investigating the role of dioxygen, experiments were performed in an

96

anaerobic chamber after leaving the solution in the chamber for 2-4 hours. No Ag(I)

97

reduction was observed in the anaerobic chamber.

98

Since, at pH 8, chlorine exists principally as OCl − ,20 we use OCl − to represent total chlorine

99

from here on in our discussion.

100 101

2.2 Chlorine determination

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Concentrations of OCl − were measured using the DPD method.22, 23 DPD reacts with OCl −

103

and produces the radical cation DPD•+ which exhibits an absorption peak at 551 nm.22-24 For

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measurement of OCl−, 300 µL of 50 mM phosphate buffer ([NaH2PO4]:[Na2HPO4] = 3:1)

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and 100 µL of 6 mM DPD stock solution was added to 2.6 mL of sample and the absorbance

106

was measured at 551 nm using a Cary 50 UV-Visible Spectrophotometer (Varian). The final

107

pH of the solution was 6.3 at about 24 °C, which is within the ideal pH range (6.2 to 6.5) for

108

the DPD colorimetric method.24 To account for DPD oxidation occurring due to the presence 5 Environment ACS Paragon Plus


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of other oxidants in our experimental matrix, parallel experiments were performed in which

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50 µM glycine was added prior to DPD addition in order to selectively remove OCl−. Since

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glycine converts OCl− instantaneously into chloroaminoacetic acid but has no effect on other

112

DPD oxidants,

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DPD•+ formed by all other oxidants except OCl−. Control experiments were performed (see

114

Supporting Information SI-1 for more details) to ensure that complete consumption of OCl−

115

occurs on glycine addition (Figure S1); however no interaction occurs with other DPD

116

oxidant like H2O2. The concentration of OCl− present at any time was calculated as the

117

difference in the absorbance measured in the absence and presence of glycine.22,

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described in the results, these additional DPD oxidants account for 25-100% of the DPD•+

119

absorbance observed in our experimental matrix. Calibration was performed by standard

120

addition of NaOCl (in the concentration range of 0 to 10.0 µM) either before or after the

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sample measurement using the experimental procedure described above. A molar extinction

122

coefficient 19,000±2000 M−1cm−1 was achieved which is close to the published value of

123

21,000 M−1cm−1.24

22, 25

the absorbance observed in samples containing glycine corresponds to

26

As

124 125

2.3 AgNP Characterization

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AgNPs were characterized by measuring their absorbance in the range 300-700 nm using a

127

Cary 50 spectrophotometer (Varian). The change in the concentration of AgNPs was

128

measured using the peak surface plasmon resonance (SPR) absorbance that was observed to

129

occur at 392 nm under all experimental conditions investigated here (see Figure S2). Since

130

there was no significant (p> 0.1 using single tailed student t-test) shift in the SPR peak during

131

the duration of our experiments (indicating that AgNP aggregation did not occur), the peak

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SPR absorbance is linearly correlated to the AgNP concentration. For measurement of

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AgNPs decay in the presence of OCl−, 2 mL of sample was withdrawn from the reactor at


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various times following OCl − addition and transferred to 1 cm quartz cuvette, shaken for 3-4

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s prior to absorbance measurement. As the measurement time was around 10 s, insignificant

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particle aggregation or settling was likely to have occurred during this period. The first

137

sample point was always measured at 30 ± 2 s. Particles were also characterized by high

138

resolution transmission electron microscopy (TEM) combined with electron diffraction

139

analysis. The TEM used in this work was an FEI Tecnai G2 instrument housed within the

140

Mark Wainwright Analytical Centre at the University of New South Wales.

141 142

2.4 Dissolved Ag(I) measurement

143

The final concentration of Ag(I), formed on oxidation of AgNPs by OCl − , was measured

144

using ICP-MS (Perkin Elmer) analysis after removing AgNPs by centrifuging the samples for

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45 min at 4000 rpm using Amicon centrifugal ultrafilters (Amicon Ultra-15 3K, Millipore,

146

MA) containing porous cellulose membranes with a nominal pore size of 1−2 nm. Samples

147

were diluted 3-5 times prior to measurement by ICP-MS. Samples for Ag(I) measurement

148

were always taken after 30 minutes of reaction time by which time the reaction had

149

essentially reached completion. It is to be noted here that any AgCl(s) and/or Ag2O(s), if

150

formed, will also be removed during centrifugation and hence the Ag concentration measured

151

using this method represents the total dissolved Ag(I) concentration.

152 153 154

3. RESULTS AND DISCUSSION

155

3.1 AgNPs oxidation by chlorine in air-saturated solution

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Under our experimental conditions, the decrease in AgNP concentration in the absence of

157

− OCl − was very small, however, AgNPs are rapidly removed upon addition of OCl (Figure 1)



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with concomitant increase in dissolved Ag(I) concentration (Figure 2), suggesting that OCl−

159

mediates the oxidative dissolution of AgNPs and that the rate of this reaction is much faster

160

than the AgNPs oxygenation rate. Note that the AgNP SPR peak does not significantly (p>

161

0.1 using single tailed student t-test) broaden or red shift during reaction with OCl− (Figure

162

S2) contrary to that observed by Li and co-workers

163

attributed to formation of Ag2O(s) at the AgNP surface. Additionally, high resolution TEM

164

analysis combined with electron diffraction analysis also revealed no evidence for Ag2O(s)

165

formation.

166

The oxidation rate of AgNPs in the presence of OCl − decreases over time, presumably due to

167

consumption of both rate controlling reactants, with the reaction reaching completion within

168

10 minutes of addition of OCl − . The oxidation rate of 5 µM AgNPs increases with increase

169

in the initial OCl − concentration, with complete removal of AgNPs occurring within 30

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seconds on addition of ≥10 µM OCl − . The final dissolved Ag(I) concentration, formed on

171

oxidation of 5 µM AgNPs varies with initial OCl − concentration between 0 to 2.5 µM,

172

however, at OCl− concentrations >2.5 µM nearly complete (> 90%) oxidation of AgNPs

173

occurs, supporting stoichiometric constraints on the reaction (Figure 2a).

174

concentration of dissolved Ag(I) formed on oxidation of AgNPs in the presence of 5 µM

175

OCl − also increases linearly with increase in the initial AgNPs concentration, however, it

176

reaches a limiting value for initial AgNPs concentrations >10 µM (Figure 2b), also reflecting

177

the stoichiometric constraints of the overall reaction. It is to be noted here that in these

178

experiments no Cl− was present and hence all the Ag(I) formed on oxidation of AgNPs will

179

exist as Ag+; however in the presence of sufficient Cl−, at least a portion of the Ag(I) formed

180

on oxidation of AgNPs will precipitate as AgCl(s) as demonstrated in earlier studies.17, 18

27

during AgNP oxygenation and


The final

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3.2 Chlorine decay in the presence of AgNPs in air-saturated solution

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As shown in Figure 1, the concentration of OCl − decreases in the presence of AgNPs with

183

the decay rate increasing with increase in AgNPs and OCl − concentration (see Figure S3).

184

The stoichiometry of Ag(I) formation relative to OCl − consumption is approximately 2:1 for

185

all the AgNP and OCl − concentrations investigated here. This observation indicates that 2

186

moles of electrons are accepted by each mole of OCl − , suggesting that Cl − is the final

187

product on reduction of OCl − in this system. However, if glycine, which rapidly consumes

188

OCl−, is added to a sample prior to adding DPD, formation of significant DPD•+ is still

189

observed (Figure S4a). Furthermore, in the absence of glycine, no decrease in DPD•+

190

absorbance over time is observed (Figure S4b), thereby suggesting that the product formed on

191

OCl − decay is capable of oxidizing DPD and hence cannot be Cl − . This observation further

192

supports the conclusion that all Ag(I) formed is present as Ag+ rather than AgCl(s) in our

193

experimental matrix due to the absence of Cl − . The possible formation of chlorite ( ClO 2− ) or

194

chlorate ( ClO 3− ) as the final product is also rejected as both of these species are not expected

195

to react with DPD at the concentration and pH investigated here.22 The formation of chlorine

196

dioxide ( ClO 2 ) which is known to react with DPD22 is also excluded based on the

197

observation that sparging with argon, which will remove ClO 2 if present, did not decrease the

198

observed DPD•+ absorbance in both the presence and absence of glycine (Figure S5). This

199

additional oxidant formed on OCl − decay is also neither H2O2 and/or HO• (data not shown;

200

see Supporting Information SI-1 for description of these experiments) since no formation of

201

these species was observed in our experimental matrix. The possibility that DPD reacts with

202

Ag+ is also rejected since no DPD•+ absorbance was measured on addition of DPD to a

203

solution containing Ag+. The possible formation of some oxidizing intermediate as a result of

204

reaction between citrate (which may be released to solution following oxidation of AgNPs)



205

and OCl − is also rejected since no reaction between these entities was observed in our

206

experimental matrix at the pH and concentration of citrate (≤ 100 µM) and OCl − (≤ 10 µM)

207

used here. Although the identity of this additional DPD oxidant(s) is not clear from our

208

work, it appears to be quite stable (lifetime > 24 h; see Figure S6) and possibly includes some

209

oxy- chlorine species. The concentration of this additional DPD oxidant formed on OCl −

210

decay (referred to as Cl DPD-ox from hereon) in air-saturated solutions under various

211

experimental conditions is shown in Figure S7. The concentration of Cl DPD-ox formed rapidly

212

increases and reaches a maximum value in the first few minutes and then decreases slightly

213

(< 10%) possibly as a result of very slow decay after its formation. The maximum

214

concentration of Cl DPD-ox formed is the same as the concentration of OCl − consumed for all

215

concentrations of AgNPs and OCl − investigated here supporting the conclusion that Cl DPD-ox

216

is formed as a result of OCl − decay. Given that the decay of Cl DPD-ox is a relatively minor

217

process at the time scale of our experiments, this process is not discussed further.

218

3.3 Role of dioxygen

219

As shown in Figure 3, removal of dioxygen inhibits OCl − - mediated oxidative dissolution of

220

AgNPs completely, with no formation of dissolved Ag(I) (measured after 30 minutes of

221

reaction time) when AgNPs concentration is in excess (≥ four-fold) of OCl − concentration;

222

however, little effect of dioxygen removal is observed on dissolved Ag(I) formation when

223

OCl − concentration is similar to AgNPs concentration (Figure 3a). This suggests that

224

dioxygen plays an important role in OCl − -mediated AgNPs oxidative dissolution, especially

225

at lower OCl − concentrations.

226

The decay of OCl − is also affected by removal of dioxygen, with the stoichiometry of

227

formation of additional DPD oxidant relative to the amount of OCl − consumed decreasing


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from 1:1 in air-saturated solution to approximately 0.5:1 in deoxygenated solution (Figure

229

3b). This observation suggests that under deoxygenated conditions a portion of OCl − decays

230

to form species that are not capable of oxidizing DPD (referred as Clno-DPD-ox), which may

231

possibly include Cl − , ClO −2 , and ClO 3− .

232

The stoichiometry of Ag(I) formation relative to OCl − consumption under deoxygenated

233

conditions is also completely different to that observed in air-saturated solution (2:1) and, as

234

shown in Figure 3c, is dependent upon the initial AgNPs: OCl − concentration ratio. At higher

235

AgNPs: OCl − concentration ratios, there is no oxidation of AgNPs but complete removal of

236

OCl − still occurs in deoxygenated solution, suggesting that AgNPs acts catalytically under

237

such conditions.

238

3.4 Mechanism of oxidative dissolution of AgNPs by chlorine

239

Based on our experimental observations and the discussion presented in previous sections,

240

the mechanism of OCl − -mediated AgNPs oxidation must incorporate the following features:

241

(i)

AgNPs oxidized for each mole of OCl − consumed in air-saturated solution.

242 243

(ii)

244 245

The product formed on OCl − decay in air-saturated solution is capable of oxidizing DPD.

(iii)

Dioxygen plays an important role in controlling the oxidation of AgNPs, especially at low OCl − concentration.

246 247

AgNPs oxidation by OCl − occurs in stoichiometric amounts with 2 moles of

(iv)

While the fate of dioxygen is not clear, it appears to be mostly reduced to H2O

248

since there is no evidence of formation of superoxide (no effect of SOD addition

249

on AgNPs oxidation was observed; data not shown), H2O2 or hydroxyl radicals

250

under the conditions examined.



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A simple overall reaction of the form shown in eq (1) can explain the formation of Ag(I) and

252

OCl − decay on oxidation of AgNPs by OCl − in air-saturated solution.

253

2Ag0n + OCl −  → 2Ag(I)+Cl DPD-ox

254

However, this overall reaction cannot explain AgNPs oxidation under deoxygenated

255

conditions and hence warrants more detailed analysis. One hypothesis to explain the role of

256

dioxygen in AgNPs oxidation is that the initial product of the AgNPs- OCl − reaction reacts

257

with dioxygen to yield Ag(I) in air-saturated solution but reforms AgNPs and/or reacts with

258

OCl − to form Ag(I) in the absence of dioxygen, as shown in eq (2). This type of reaction

259

mechanism is consistent with the mechanism proposed for oxidative dissolution of AgNPs by

260

H2O219,

261

intermediate which is further oxidized by H2O2 to yield Ag(I). In this case, the reactive

262

intermediate formed on reaction of AgNPs and OCl − appears to be oxidized by both O2 and

263

OCl − . At low OCl − concentration in deoxygenated solution, most of the reactive

264

intermediate decays to reform AgNPs; however, in the presence of excess OCl − , the reactive

265

intermediate is oxidized to yield Ag(I).

266

28

(1)

with the reaction between AgNPs and H2O2 initiating formation of a reactive

O

2   → Ag 0 ...OCl−  Ag 0n + OCl− ← → Ag(I)  n OCl−

(2)

267

A second alternative is that the initial reaction between OCl − and AgNPs result in formation

268

of Ag(I) and a reduced chlorine species which reacts with dioxygen in air-saturated solution,

269

but, in the absence of dioxygen, reduces Ag(I) to reform AgNPs and/or reacts with OCl − to

270

form another chlorine species which is capable of oxidizing DPD (eq. 3-4).

271

 → Ag + +Cl red Ag0n + OCl − ← 

(3)

272

O2 Cl red  → Cl DPD-ox OCl −

(4)

273

Both possibilities are consistent with Ag(I) formation in air-saturated and deoxygenated

274

solution. However, alternative 2 appears less likely based on our earlier work for two main


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reasons: (i) the presumably related AgNPs-H2O2 reaction is reported to proceed via formation

276

of a reactive intermediate19, 28 rather than complete one-electron transfer, and (ii) the Cl red

277

species shown in eq (3) is expected to be Cl• based on charge-balance; Cl• reacts with

278

dioxygen to form ClOO• (a distinct isomer differing from the aforementioned ClO2• which is

279

of the form OClO•) which is known to dissociate rapidly to reform Cl• and O 2 , 29 i.e., would

280

not react as required in reaction 4.

281

A reaction schematic which is capable of explaining the formation of Ag(I) on AgNPs

282

oxidation under air-saturated and deoxygenated conditions is shown in Figure 4. As shown,

283

the reaction of AgNPs and OCl − initially results in formation of a reactive intermediate

284

which further reacts with dioxygen in air-saturated solution to yield Ag(I). The reaction of the

285

reactive intermediate with OCl − and AgNPs in air-saturated solution is expected to be

286

unimportant under the conditions of our experiment given that dioxygen concentration is 25-

287

250 fold higher than the AgNPs and OCl − concentrations used here. Under deoxygenated

288

conditions however, the reactive intermediate either reacts with OCl − to form Ag(I) or

289

decays back to reform AgNPs. The reformation of AgNPs could occur either due to simple

290

dissociation of the reactive intermediate or due to reaction with AgNPs resulting in formation

291

of Ag(I) and charged AgNPs ( Ag0* n ), as hypothesized to occur in the AgNPs-H2O2

292

reaction,19 which react together to reform AgNPs.

293

A kinetic model has been developed based on the reaction schematic presented in Figure 4 to

294

explain our experimental results (Table 1; see Supporting Information SI-3 for more details

295

of the kinetic model). We have assumed that the stoichiometry of Ag(I) formation to OCl −

296

decay is possibly controlled by formation of a reactive intermediate in a 2:1 stoichiometric

297

reaction between AgNPs and OCl − . As shown in Figures 1-3, the kinetic model explains our

298

experimental data very well. The model also explains the effect of dioxygen removal on



299

AgNPs oxidation with varying degrees of AgNPs oxidation observed at different AgNPs:

300

OCl − concentration ratios. The model also predicts the concentration of the additional DPD

301

oxidant formed (Figure S7) in air-saturated and deoxygenated solution (Figure 3) very well.

302

However, we cannot reject the possibility that there may be other reactions which may play a

303

role in Ag(I) formation. Furthermore, the kinetic model presented here does not describe the

304

exact reactions controlling the fate of OCl − due to the unknown nature of the Cl DPD-ox

305

formed. Additionally, although we assumed that the stoichiometry of Ag(I) formation to

306

OCl − decay is controlled by formation of a reactive intermediate in a 2:1 stoichiometric

307

reaction between AgNPs and OCl − , there are equally valid alternative mechanisms that we

308

cannot exclude, such as the product of OCl − decay in the initial step (which results in Ag(I)

309

formation) may oxidize another mole of AgNPs via a similar mechanism or, alternatively,

310

may decay via bimolecular dismutation with reformation of OCl − thereby resulting in an

311

overall 2:1 stoichiometry of Ag(I) formed to OCl − consumed. Thus, further work is required

312

to identify the product(s) formed on OCl − decay and the reactions controlling OCl − decay to

313

explain the complete mechanism of AgNPs oxidation by OCl − .

314 315 316

4. ENVIRONMENTAL IMPLICATIONS

317

The oxidation rate of AgNPs in the presence of OCl − is much faster than both its

318

oxygenation rate

319

by OCl − in air-saturated solution occurs in stoichiometric amounts with two moles of Ag(I)

320

formed for each mole of OCl − that is decayed, with an unknown product also formed in this

321

reaction that is capable of oxidizing DPD. Dioxygen plays an important role in OCl − -

322

mediated AgNPs oxidation, especially at lower OCl − concentration, with the mechanism

Our experimental results show that OCl − rapidly oxidizes AgNPs in air-saturated solution.

27

and the rate of oxidation of AgNPs by H2O2.19 The oxidation of AgNPs


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shifting from stoichiometric oxidation of AgNPs by OCl− in the presence of dioxygen to a

324

catalytic removal of OCl− by AgNPs in the absence of dioxygen.

325

Based on the results presented here we can conclude that the presence of OCl − will control

326

the AgNPs concentration (if present) and also mitigate its toxicity. Since the species formed

327

on OCl − decay has significant oxidizing capacity (as evident from DPD oxidation observed

328

in our experimental matrix), it is possible that the disinfection efficiency of OCl − is not

329

significantly impacted by the presence of AgNPs. This reaction also has significant

330

implications to the fate of Ag(0) nanoparticles embedded in antimicrobial fabrics (such as

331

used, for example, in burn dressings). Such fabrics when treated with chlorine bleach may

332

result in oxidation of Ag(0) and lead to leaching of Ag(I) which will be scavenged by any

333

Cl − present to form less-reactive AgCl(s). This suggests that the efficacy of Ag as an

334

antimicrobial agent may be restricted after washing in solutions containing strong oxidizing

335

agents such as OCl − . Similarly, leaching of silver may occur in water disinfection systems

336

(such as the Swach water filter that is marketed for household water treatment in India) if the

337

AgNPs-loaded material comes in contact with water containing high concentrations of

338

chlorine, thereby resulting in decreased disinfection efficiency.

339

The findings reported here also have implications in organic contaminant degradation using

340

AgCl(s) as the photocatalyst. The photocatalytic activity of AgCl(s) has previously been

341

shown to decrease due to its reduction to Ag(0) by photo-generated electrons.30, 31 However,

342

in the presence of OCl − , the photocatalytic activity of AgCl(s) may not change as any Ag(0)

343

formed may be re-oxidized to Ag(I) with subsequent reformation of AgCl(s). Our recent

344

investigation using AgCl(s) as a photocatalyst show this is indeed the case with presence of

345

chlorine increasing the recycling rate of AgCl(s), and leading to enhanced degradation rates

346

of target species.32



347

Overall, the work presented here provides important new insights into the fate and toxicity of

348

AgNPs in natural and engineered aquatic environments.

349

ACKNOWLEDGEMENTS

350

Financial support from the Australian Research Council through Discovery Grants

351

DE120102967 and DP120103222 is gratefully acknowledged. We also acknowledge the

352

invaluable assistance provided by Dr Quadir Zakaria from the Mark Wainwright Analytical

353

Center at UNSW in undertaking electron diffraction analyses.

354 355

Supporting Information available

356

Details on experimental procedure, manipulative experiments to determine the additional

357

DPD oxidant formed on chlorine decay, and the kinetic model is provided. This material is

358

available free of charge via the Internet at http://pubs.acs.org.

359 360

REFERENCES

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1. Benn, T. M.; Westerhoff, P., Nanoparticle silver released into water from commercially available sock fabrics. Environ. Sci. Technol 2008, 42, 4133-4139. 2. Luoma, S. Silver Nanotechnologies and the Environment: Old Problems or New Challenges; Woodrow Wilson International Center for Scholars: 2008. 3. Nowack, B.; Krug, H. F.; Height, M., 120 years of nanosilver history:Implications for policy makers. Environ. Sci. Technol 2011, 44, 1177-1183. 4. Sondi, I.; Salopk-Sondi, B., Silver nanoparticles as antimicrobial agent: a case study on E. coli as a model for Gram negative bacteria. J. Colloid Interface Sci. 2004, 275, (1), 177-182. 5. He, D.; Dorantes-Aranda, J. J.; Waite, T. D., Silver NanoparticleAlgae Interactions: Oxidative Dissolution, Reactive Oxygen Species Generation and Synergistic Toxic Effects. Environ. Sci. Technol 2012, 46, 8731-8738. 6. AshaRani, P. V.; Mun, G. L. K.; Hande, M. P.; Valiyaveettil, S., Cytotoxicity and Genotoxocity of Silver nanoparticles in Human Cells. ACS Nano 2009, 3, (2), 279-290. 7. Kim, S.; Choi, J. E.; Choi, J.; Chung, K. H.; Park, K.; Yi, J.; Ryu, D. Y., Oxidative stress dependent toxocity of silver nanoparticles in human hepatoma cells. Toxicology In Vitro 2009, 23, (6), 1076-1084. 8. Panyala, N.; R.; Pena-Mendez, E. M.; Havel, J. J., Silver or silver nanoparticles: a hazardous threat to the environment and human health. Appl. Biomed. 2008, 6, 117–129.


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9. Fabrega, J.; Fawcett, S. R.; Renshaw, J. C.; Lead, J. R., Silver nanoparticle impact on bacterial growth: Effect of pH, concentration, and organic matter. Environ. Sci. Technol 43, (19), 7285–7290. 10. Vasileiadis, S.; Puglisi, E.; Trevisan, M.; Scheckel, K. G.; Langdon, K. A.; McLaughlin, M. J.; Lombi, E.; E., D., Changes in soil bacterial communities and diversity in response to long-term silver exposure. FEMS Microbiol Ecol 2015, 91, (10). 11. Lok, C. N.; Ho, C. M.; Chen, R.; He, Q. Y.; Yu, W. Y.; Sun, H. Z.; Tam, P. K. H.; Chiu, J. F.; Che, C. M., Proteomic analysis of the mode of antibacterial action of silver nanoparticles. J Proteome Res 2006, 5, 916-924. 12. Smetana, A. B.; Klabunde, K. J.; Marchin, G. R.; Sorensen, C. M., Biocidal activity of nanocrystalline silver powders and particles. Langmuir 2008, 24, 7457-7464. 13. Devi, G. P.; Ahmad, K. B. A.; Varsha, M. S.; Shrijha, B.; Lal, K. S.; Anbazhagan, V.; Thiagrajan, R., Sulfidation of silver nanoparticles reduces its toxicity in zebrafish. Aquat Toxicol (Amst) 2015, 158, 149-156. 14. Reinsch, B.; Levard, C.; Li, Z.; Ma, R.; Wise, A.; Gregory, K.; Brown, G. J.; Lowry, G., Sulfidation of silver nanoparticles decreases Escherchia Coli growth inhibition. Environ. Sci. Technol 2012, 46, 6992-7000. 15. Lee, D.-Y.; Fortin, C.; Campbell, P. G. C., Contrasting effects of chloride on the toxicity of silver to two green algae, Pseudokirchneriella subcapitata and Chlamydomonas reinhardtii. Aquat Toxicol (Amst) 2005, 75, 127-135. 16. Leblanc, G. A.; Mastone, J. D.; Paradice, A. P.; Wilson, B. F., The influence of speciation on the toxicity of silver to fathead minnow (pimephales promelas). Environ. Toxicol. Chem. 1984, 3, 37-46. 17. Impellitteri, C. A.; Tolaymat, T. M.; Scheckel, K. G., The Speciation of Silver Nanoparticles in Antimicrobial Fabric Before and After Exposure to a Hypochlorite/Detergent Solution. J Environ Qual 2009, 38, 1528-1530. 18. Yuan, Z.; Chen, Y.; Li, T.; Yu, C., Reaction of silver nanoparticles in the disinfection process. Chemosphere 2013, 93, 619-625. 19. He, D.; Garg, S.; Waite, T. D., H2O2-mediated Oxidation of Zero-valent Silver and Resultant Interactions between Silver Nanoparticles, Silver Ions and Reactive Oxygen Species. Langmuir 2012, 28, 10266-10275. 20. Morris, J. C., The Acid Ionization Constant of HOCl from 5 to 35°. The Journal of Physical Chemistry 1996, 70, 3798-3805. 21. Miller, C. J.; Rose, A. L.; Waite, T. D., Phthalhydrazide chemiluminescence method for determination of hydroxyl radical production: Modifications and adaptations for use in natural systems. Anal Chem 2011, 83, (1), 261-268. Standard Methods For the Examination of Water and Wastewater. American Public 22. Health Association, American Water Works Association, Water Environment Federation: The United States of America, 1998. 23. Palin, A. T., Current DPD Methods for Disinfectant Residual measurement. Journal of the Institution of Water Engineers and Scientists 1986, 40, 501-510. 24. Bader, H.; Sturzenegger, V.; Hoigne, J., Photometric Method for the Determination of Low Concentrations of Hydrogen Peroxide by the Peroxidase Catalyzed Oxidation of N,NDiethyl-p-Phenylenediamine(DPD). Water Research 1988, 22, 1109-1115. 25. Peskin, A. V.; Midwinter, R. G.; Harwood, D. T.; Winterbourn, C. C., Chlorine transfer between glycine, taurine, and histamine: Reaction rates and impact on cellular reactivity. Free Radical Biol. Med. 2004, 37, 1622-1630. 26. Yua, H.-W.; Oh, S.-G.; Kim, I. S.; Peppera, I.; Snyder, S.; Jang, A., Formation and speciation of haloacetic acids in seawater desalination using chlorine dioxide as disinfectant. Journal of Industrial and Engineering Chemistry 2015, 26, 193-201.



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27. Li, X.; Lenhart, J. J.; Walker, H. W., Dissolution-Accompanied Aggregation Kinetics of Silver Nanoparticles. Langmuir 2010, 26, 16690–16698. 28. He, D.; Miller, C. J.; Waite, T. D., Fenton-like zero-valent silver nanoparticlemediated hydroxyl radical production. J. Catal. 2014, 317, 198-205. 29. Dunn, R. C.; Simon, J. D., Excited-State Photoreactions of Chlorine Dioxide in Water Journal of American Chemical Society 1992, 114, 4856-4860. 30. Ma, T.; Garg, S.; Miller, C. J.; Waite, T. D., Contaminant degradation by irradiated semiconducting silver chloride particles: Kinetics and modelling. J. Colloid Interface Sci. 2014, 446, 351-357. 31. Tian, B.; Zhang, J., Morphology-Controlled synethesis and applications of silver halide photocatalytic materials. Catalysis surveys from Asia 2012, 16, 210-230. 32. Garg, S.; Rong, H.; Miller, C. J.; Waite, T. D., Chlorine-mediated Regeneration of Semiconducting AgCl(s) Following Light-induced Ag0 Formation: Implications to Contaminant Degradation. J Phys Chem C in press.


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Table 1: A kinetic model to explain oxidation of AgNPs by OCl −

a

Reaction no

Reaction

Rate constant used

1

2Ag 0 + OCl −  → 2Ag 0 ...OCl −

2.5×109 M-3s-1

2

2Ag 0 ...OCl− + O2  → 2Ag + + ClDPD-ox

≥1×106 M-1s-1

3

2Ag 0 ...OCl− + OCl−  → 2Ag + + ClDPD-ox + Clno-DPD-ox

1×107 M-1s-1

4

1 1 2Ag0 ...OCl− + Ag0  → 2Ag0* + Ag + + Clno-DPD-ox + ClDPD-ox 2 2

1×106 M-1s-1

5

Ag 0* + Ag +  → Ag 0 + Ag 0 a

≥1×106 M-1s-1

0*

Note that the reduction of dioxygen by charged AgNPs (i.e. Ag ) is not included since these charged AgNPs are formed in the absence of dioxygen only.



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Figure Captions

454

µM AgNPs with varying initial OCl − concentrations (indicated on the plot) in air-saturated

455

solutions at pH 8. Decrease in AgNPs (panel b) and OCl − concentration (panel d) on

456

reaction of 5 µM OCl − with varying initial AgNPs concentrations (indicated on the plot) in

457

air-saturated solutions at pH 8. Points represent the average of triplicate measurements; lines

458

represent model results. Error bars represents standard deviation of triplicate measurements.

Figure 1: Decrease in concentration of AgNPs (panel a) and OCl − (panel c) on reaction of 5

459 460

Figure 2: (a) Final dissolved Ag(I) concentration formed on oxidation of 5 µM AgNPs in the

461

presence of various OCl − concentrations in air-saturated solutions at pH 8.

462

dissolved Ag(I) concentration generated on oxidation of AgNPs by 5 µM OCl − in air-

463

saturated solution at pH 8. Points represent the average of triplicate measurements; lines

464

represent model results. Error bars represents standard deviation of triplicate measurements.

(b) Final

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Figure 3: (a) Fraction of AgNPs oxidized in the presence of OCl − in air-saturated and

467

deoxygenated solutions for various initial AgNPs and OCl − concentration ratios. (b)

468

Concentration of the additional DPD oxidant formed per mole of OCl − consumed (denoted as

469

∆[OCl−]) on reaction of AgNPs and OCl − in air-saturated and deoxygenated solutions for

470

various initial AgNPs and OCl − concentration ratios. (c) Concentration of Ag(I) formed per

471

mole of OCl − consumed on reaction of AgNPs and OCl − in air-saturated and deoxygenated

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solutions for various initial AgNPs and OCl − concentration ratios. Points represent the

473

average of triplicate measurements; lines represent model results.

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Figure 4: Proposed reaction schematic showing oxidation of AgNPs by OCl − .

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Figure 3

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Figure 4

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ToC Graphic

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Oxidative Dissolution of Silver Nanoparticles by ... - ACS Publications

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