Role of Secondary Particle Formation in the Persistence of Silver

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Role of Secondary Particle Formation in the Persistence of Silver Nanoparticles in Humic Acid Containing Water under Light Irradiation Tuoya Zhang, Dawei Lu, Lixi Zeng, Yongguang Yin, Yujian He, Qian Liu, and Guibin Jiang Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.7b04115 • Publication Date (Web): 22 Nov 2017 Downloaded from http://pubs.acs.org on November 22, 2017

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Environmental Science & Technology

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ARTICLE

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Role of Secondary Particle Formation in the Persistence of Silver

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Nanoparticles in Humic Acid Containing Water under Light

4

Irradiation

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Tuoya Zhang1,2, Dawei Lu1,2, Lixi Zeng3, Yongguang Yin1, Yujian He2, Qian Liu1,2,4,*, Guibin

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Jiang1,2 1

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State Key Laboratory of Environmental Chemistry and Ecotoxicology, Research Center for Eco-Environmental Sciences, Chinese Academy of Sciences, Beijing 100085, China

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2

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University of Chinese Academy of Sciences, Beijing 100049, China

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3

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Guangdong Key Laboratory of Environmental Pollution and Health, Jinan University, Guangzhou

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510632, China

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School of Environment, Guangzhou Key Laboratory of Environmental Exposure and Health, and

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Institute of Environment and Health, Jianghan University, Wuhan 430056, China

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*Corresponding author: Prof. Qian Liu. 18 Shuangqing Road, Haidian District, Beijing 100085,

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China. Tel: +86-10-62849124. Fax: +82-10-62849339. Email: [email protected].

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Keywords: silver nanoparticle; isotope fractionation; environmental process; dissolved organic

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matter; secondary particle

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ACS Paragon Plus Environment


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Abstract

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The wide use of silver nanoparticles (AgNPs) leads to the increasing release of AgNPs into the

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environment. Dissolved organic matter (DOM) is a key factor affecting the behaviors and fate of

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AgNPs in the aquatic environment. However, the mechanisms for the DOM-mediated

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transformations of AgNPs are still not fully understood. In this study, we investigated the

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persistence of AgNPs in aquatic environment in the presence of different concentrations of humic

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acid (HA) over periods of time up to 14 days. The Ag species were monitored and characterized by

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absorption spectrometry, transmission electron microscopy (TEM), inductively coupled plasma

29

mass spectrometry (ICP-MS), and multicollector ICP-MS (MC-ICP-MS). Results showed that the

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long-term persistence of AgNPs in HA-containing water was determined by two critical

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concentrations of HA. When the HA concentration exceeded a lower critical value, AgNPs could be

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persistent in the solution and a large number of AgNPs were formed secondarily from the

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HA-induced reduction of the Ag+ ions released from the primary AgNPs, causing a redistribution of

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the particle size. With the HA concentration above a higher critical value, AgNPs could persist in

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the solution without significant change in particle size. Notably, we used Ag isotope fractionation to

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investigate the transformation mechanism of AgNPs. The natural isotopic analysis by MC-ICP-MS

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revealed that the size redistribution of AgNPs caused significant Ag isotope fractionation, which

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gave additional evidence for the proposed mechanisms. This study provides new insights into the

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environmental fate of engineered AgNPs and highlights the usefulness of stable isotope

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fractionation in environmental nanotechnology.

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1. Introduction

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In recent years, silver nanoparticles (AgNPs) are widely applied in various physical, biological,

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and medicinal fields due to their excellent optical, electrical, catalytic, and especially

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broad-spectrum antibiotic properties.1-3 The rapidly growing production and usage amount of

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AgNPs leads to the inevitably increasing level of AgNPs released into the environment,4-6 which

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raises great concerns about their environmental toxicity and safety. A large number of studies have

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demonstrated that AgNPs can cause various types of toxic effects to organisms such as apoptosis,

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organic damage, and neurological hazards,7-10 and that the toxicity of AgNPs is mainly attributed to

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the release of Ag+ ions.11-13

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As Ag is a highly active element, AgNPs are easy to transform into other Ag species (e.g., Ag+,

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Ag2O, AgClx(x-1)-, and Ag2S) in the aquatic environment. Various environmental factors, such as pH,

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temperature, dissolved oxygen, ionic strength, and sunlight radiation, can influence the stability and

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transformations of AgNPs, which subsequently affect the fate and toxicity of AgNPs in the

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environment.14-17 Specifically, as an omnipresent part of natural water, dissolved organic matter

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(DOM) plays an important role in determining the environmental fate and toxicity of AgNPs. It has

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been reported that DOM can improve the stability of AgNPs in aquatic media and suppress the Ag+

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ion release and homoaggregation of AgNPs.18 Yang et al.19 exposed nematode to AgNP suspensions

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with DOM, and found that DOM could reduce the toxicity of AgNPs by decreasing the intracellular

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uptake of AgNPs as a result of the formation of DOM-AgNP composites. Recent studies reported

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that Ag+ can be photoreduced to form AgNPs in river water or synthetic natural water samples

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containing DOM.20-22 The reduction of Ag+ was supposed to follow two possible mechanisms: (1)

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direct charge transfer from the phenol groups in DOM to metal ions by complexation between Ag+

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and DOM;23,

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(2) mediation by superoxide ions (O2•-) produced from photoirradiation of the 3



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phenol groups in DOM.22 The possible mechanism can be described by the following chemical

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reactions:25

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HAred + O2 → HAox + O2 •-

(1)

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Ag+ + O2 •-→ Ag0 (AgNP) + O2

(2)

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where HAred and HAox represent the reduced and oxidized forms of humic acid (HA) as a model

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DOM, respectively. Yu et al.26 reported the transformations between AgNPs and Ag+ in DOM-rich

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water within 48 h. Furtado et al.27 investigated the fate of AgNPs in boreal lake mesocosms and

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found that AgNPs were relatively stable in the lake environment with a half-life of total Ag of ~20

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days, whereas the detailed processes or mechanisms were unclear. Overall, due to the highly

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dynamic properties of AgNPs28 and complex composition of DOM,29 the behaviors of AgNPs in

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natural water and the underlying mechanisms are still not entirely understood.

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Another cause for the poor understanding of environmental behaviors of AgNPs is the shortage

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of research tools. Although there have been multiple techniques applicable for the characterization

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of nanoparticles, such as ultraviolet-visible (UV-vis) absorption spectrometry, electron microscopy,

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and dynamic light scattering (DLS), the information gave by these techniques is not sufficient to

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fully reveal the complex environmental behaviors of AgNPs. In our previous work, we used

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multicollector inductively coupled plasma mass spectrometry (MC-ICP-MS) to measure the subtle

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variations in stable isotopic composition of Ag during different transformation processes of AgNPs

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in natural water.30 Ag has two naturally occurring stable isotopes (107Ag and

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abundances (51.8%:48.2%). Results illustrated that different transformation processes of AgNPs

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could cause different Ag isotope fractionation, and the Ag isotope fractionation could indicate the

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presence of concurrent processes and might even distinguish the origins of AgNPs in the

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environment. Thus, isotope fractionation can be used as a unique tracer for environmental processes, 4


109

Ag) with close

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which makes it different from other characterization techniques. This study preliminarily showed

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the promise of stable isotope fractionation as a new tool to study the nanoparticles in the

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environment, while the application potential of this tool is yet to be explored.

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In the present study, we investigated the persistence of AgNPs and the stabilization mechanism

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in a HA-containing water. We incubated polyvinylpyrrolidone (PVP)-coated AgNPs as a model type

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of engineered AgNPs in the presence of different concentrations of HA for up to 14 days. We

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monitored and characterized the Ag species by different techniques, including UV-vis absorption

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spectrometry, transmission electron microscopy (TEM), energy dispersive X-ray spectroscopy

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(EDS), inductively coupled plasma mass spectrometry (ICP-MS), and MC-ICP-MS. We show that

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the HA-mediated secondary formation of AgNPs is an important mechanism for the stabilization of

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AgNPs in HA-containing water. Notably, the Ag isotopic analysis by MC-ICP-MS provide strong

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evidence for the formation of secondary AgNPs. This study not only helps to better understand the

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behaviors and fate of AgNPs in the environment, but also further demonstrates the usefulness of

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stable isotope fractionation as a new tool for studying nanoparticles in complex systems.

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2. Materials and methods

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2.1 Materials

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PVP-AgNPs were purchased from Zhuiguang Technologies (Shanghai, China). The stock

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solution was stored at 4 °C in dark. The AgNP working solution was prepared after ultrafiltration

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using ultra-15 centrifugal filter units (10 KDa; Millipore, MA, USA) and washed with ultrapure

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water for four times to remove excess Ag+. HNO3 (65%) was purchased from Merck (Darmstadt,

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Germany). Suwannee River humic acid standard II (SRHA) was obtained from the International

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Humic Substances Society (IHSS, St. Paul, MN). The SRHA stock solutions were prepared in

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ultrapure water and vibrated in shaker overnight, and then filtered through a 0.45-µm filter.

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Ultrapure water (Merck Millipore, Darmstadt, Germany) was used to prepare all working solutions.

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Multi-element calibration standard for ICP-MS analysis was purchased from Agilent

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Technologies (8500–6940 for silver standard and 5188–6525 for internal standard; CA, USA). The

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Ag isotope standard SRM 978a for MC-ICP-MS measurement was purchased from the National

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Institute of Standards and Technology (NIST, Gaithersburg, MD). The Pd internal standard was

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purchased from the Delta Scientific Laboratory Products Ltd (Mississauga, ON, Canada). All

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standard solutions were used after dilution with 5% HNO3.

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2.2 Characterization of silver nanoparticles

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The AgNPs were characterized by a transmission electron microscope (H-7500, Hitachi, Japan)

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and a high-resolution transmission electron microscope equipped with an energy dispersive X-ray

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spectroscopy (TECANI G20, FEI, Hillsboro, OR). The sample solution was dropped onto the

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ultrathin carbon-coated copper grid (Zhongjingkeyi Technology Co., Beijing) and dried under an

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infrared lamp. The AgNP solution was also characterized by a UV-vis spectrometer (Shimadzu

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UV-3600, Kyoto, Japan). 6


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2.3 The Ag+ release experiment

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The Ag+ release experiments were conducted in a solar simulator (SN-500, Beifang Lihui

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Instrument Equipment Co., Beijing) equipped with three wind-refrigerated 2500 W Xe lamps as

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simulated sunlight sources. The light intensity was maintained at 550 W/m2 and the temperature

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was kept at 35 ± 2 °C. AgNP solutions (17.0 mg/L) were incubated with a series of concentration (0,

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0.1, 0.5, 1, 2, 4, 8, 10, 16, 20, and 50 mg/L) of SRHA in 500 mL beakers for up to 14 days. The

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dissolved oxygen concentration in the system varied within the range of 4.6-5.0 mg/L. The

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experiments were also carried out at a lower concentration (1.0 mg/L) of AgNPs. At each time

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interval, 10 mL of the samples was collected and centrifuged by an ultrafiltration device (Amicon

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ultra-15 centrifugal filter units, 10 kDa; Millipore, MA, USA) at 8000 g for 30 min to separate the

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dissolved Ag+ ions from AgNPs. The ultrafiltration process has been examined to ensure that no

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significant Ag+ loss or artificial Ag isotope fractionation was caused.30 The control experiments

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were performed in dark by wrapping the beakers with aluminum foils, and other conditions were the

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same as those in sunlight experiments. All experiments were performed in triplicate (n = 3). Note:

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the conclusions drawn from sunlight experiments were also validated under real sunlight.

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Considering that it is unable to control the real sunlight intensity, data were collected only under

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simulated sunlight.

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2.4 Sample preparation for ICP-MS analysis

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The concentration of total Ag and released Ag+ was measured by an Agilent Technologies 8800

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inductively coupled plasma mass spectrometer (CA, USA). The total Ag content was measured after

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digestion of sample with 65% HNO3 and H2O2 followed by dilution with 5% HNO3. The released

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Ag+ samples (containing only Ag+) were directly diluted with 5% HNO3 for ICP-MS analysis.

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2.5 Sample preparation for Ag isotope ratio measurement by MC-ICP-MS 7



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The Ag isotope ratio was measured by a Nu II Plasma multicollector inductively coupled

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plasma mass spectrometer (Wrexham, UK) equipped with 16 Faraday cups and working in

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low-resolution mode. Each 2 mL of the residual AgNP solution and dissolved Ag+ samples were

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digested in a CEM Mars 5 Microwave Reaction System (Matthews, NC) after adding 5 mL of 65%

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HNO3 and 1 mL of H2O2. The resulting solution was diluted to a Ag concentration at about 100

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µg/L with 5% HNO3 for MC-ICP-MS analysis. All liquid samples were introduced by a Desolvation

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Nebulizer System (DSN-100) using self-aspiration mode at a flow rate of 0.1 L/min. The blank

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signals obtained from 5% HNO3 were subtracted from all sample and standard signals. Three

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parallel measurements were made for all samples (n = 3).

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A δ value relative to a Ag isotope standard solution (NIST SRM 978a) was used to describe the Ag isotope ratio (109Ag/107Ag) in the sample: δ109/107 Ag =

Sample

RT

RStandard T

Sample

- 1 × 1000‰

(3)

and RStandard are mass bias-corrected (or ‘‘true’’) Ag isotope ratios in the sample T

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where RT

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and standard, respectively. The Ag isotope fractionation during a process of AgNPs was described

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by a ∆109/107Ag value as follows:

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∆109/107Agproduct = δ109/107Agproduct − δ109/107Aginitial

(4)

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where ∆109/107Agproduct represents the Ag isotope fractionation in the transformation product

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(released Ag+ ions or residual AgNPs) compared to the initial AgNPs. δ109/107Agproduct and

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δ109/107Aginitial represent the Ag isotopic compositions of the product and initial AgNPs, respectively.

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The mass bias of 108

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Ag/107Ag ratio was corrected based on internal normalization with Pd

Pd/106Pd ratio) combined with standard-sample-standard bracketing approach without

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(using

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assuming identical mass bias for Pd and Ag.31, 32 Because the true value of the 108Pd/106Pd ratio was

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unknown, the certified value of 1.07638 for 107Ag/109Ag ratio of SRM 978a was first used to obtain 8


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Pd/108Pd ratio in two adjacent SRM 978a standard solutions, and then their

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mass bias-corrected

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average value was used for the correction of the

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calculation was based on the Russell exponential fractionation law:33

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Ag/107Ag ratio in the sample. The correction

f

RM = RT × i m

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(5)

mj

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where RM and RT represent the measured and mass bias-corrected ratios, respectively; mi and

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mj represent the absolute masses of the isotopes of interest; f is the mass bias correction factor.

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106

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method described by Lu et al.30 The sample digestion process and matrix effects in MC-ICP-MS

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measurement have also been evaluated as in our previous work.30 The method uncertainty was 0.04‰

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(2SD).

Cd and

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Cd interferences to

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Pd and

108

Pd was mathematically removed using an iterative

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3. Results and discussion

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3.1 Ag+ release experiment of silver nanoparticles

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In this study, we used PVP-coated AgNPs as a model type of engineered AgNPs. The

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PVP-coated AgNPs were characterized by TEM and UV-vis absorption spectroscopy. As shown in

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Fig. 1, TEM measurement showed that the AgNPs were typically spherical and well dispersed in the

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solution with an average diameter of 17.0 ± 4.2 nm. The UV-vis spectra showed a maximum

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absorption band (λmax) at ~400 nm, corresponding to the characteristic absorption of AgNPs.

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In previous studies on the environmental behaviors of AgNPs, environmental factors such as pH,

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salinity, DOM, and photoradiation have been repeatedly considered, but their roles in determining

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the fate of AgNPs are still controversial. Here we attempted to get new insights into the effect of

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DOM on the environmental behaviors of AgNPs. In order to study the persistence of AgNPs in

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aquatic environment, we first monitored the variation in the absorption spectra of AgNP suspension

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in the presence of different concentration of HA under sunlight. The concentration range of HA was

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set to 0-50 mg/L, which was accordant with the naturally occurring levels.34 Fig. 2a shows that in

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the absence of HA the characteristic absorption peak of AgNPs at 400 nm (i.e., λmax) gradually

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declined and almost disappeared after 336 h, clearly indicating the dissolution of AgNPs to release

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Ag+ ions in the solution. A broad peak at ~500 nm appeared after 168 h photoradiation probably due

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to the slight aggregation of AgNPs.35, 36 The aggregation of AgNPs also caused the color of the

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solution to turn from yellow into light purple (Fig. S1). While in the presence of HA beyond 1 mg/L

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(e.g., Fig. 2b and 2c), the λmax absorbance of AgNPs showed no evident decrease compared with the

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control without HA, suggesting that the presence of HA, even at a low concentration, could strongly

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suppress the dissolution of AgNPs and the release of Ag+. The results for more HA concentrations

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are given in Fig. S2. During the incubation process, the pH of the solution could keep relatively 10


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stable (see Fig. S3). We also repeated the experiments at a lower concentration (1.0 mg/L; see Fig.

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S4) and with a larger particle size of AgNPs (39.0 ± 3.3 nm; see Fig. S5 and S6). Results showed

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that the λmax absorbance of AgNPs at low concentration decreased more quickly within the first 24 h

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than that at high AgNP concentration (Fig. S4), and that large AgNPs had a higher persistence than

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small particles in the presence of HA (Fig. S6). We also found that the stabilizing ability of HA for

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AgNPs could be maintained even in the presence of 10 mM Na+ or Ca2+ ions (Fig. S7).

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We further characterized the AgNP suspension after photoradiation for periods of time up to 336

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h by TEM. The results are shown in Fig. 3. TEM measurements showed that AgNPs in the absence

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of HA intensively aggregated after 336 h incubation (Fig. 3c), which could account for the broad

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peak appeared at ~500 nm in UV-vis absorption spectra (Fig. 2a). In the presence of low

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concentration of HA (e.g., 1 mg/L), the aggregation was effectively inhibited with a large number of

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small particles remaining in the solution (Fig. 3g). With further increasing the HA concentration to

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higher levels (e.g., 20 mg/L), the aggregation of AgNPs were thoroughly suppressed, and no

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obvious change in particle size (from 18.3 ± 4.4 nm to 18.8 ± 3.4 nm) or shape was observed after

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336 h incubation (Fig. 3k and 3l), suggesting that high concentration of HA could stabilize the

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AgNPs for long term in aquatic environment.27

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Noteworthily, we found that in the presence of low concentration (e.g., 1 mg/L) of HA, a large

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number of small “satellite” particles with mean particle size of 3.3 ± 1.1 nm appeared and closely

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surrounded the original AgNPs after 24 h incubation (Fig. 3e). When further prolonging the

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incubation time to 336 h, the “satellite” particles surrounding the AgNPs disappeared; instead, many

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small particles were observed in the bulk solution. These small particles showed a slight growth in

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mean particle size compared with the “satellite” nanoparticles from 3.3 ± 1.1 nm to 6.8 ± 1.3 nm

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(Fig. 3f and 3h). This eventually resulted in a redistribution in particle size in the solution. As 11



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shown in Fig. 3h, a portion of particles (23.7%) kept the same diameter as the original AgNPs (17.4

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± 2.8 nm), while the other portion of particles (76.3%) were in a diameter of 6.8 ± 1.3 nm.

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Furthermore, high-resolution TEM and EDS measurements were used to characterize the AgNPs,

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and results confirmed that the newly formed particles were Ag nanoparticles (Fig. S8 and S9). The

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TEM results obtained at more HA concentrations (2, 4, 8, and 16 mg/L) are given in Fig. S10.

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In order to obtain quantitative data, the release of Ag+ from AgNPs was also monitored by using

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ICP-MS. The PVP-coated AgNPs were used immediately after washing with ultrapure water for

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several times to remove excess Ag+ ions, so the initial Ag+ was regarded as zero. In the AgNP

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solution with no HA (Fig. 4a), the mass percentage of the released Ag+ increased rapidly during the

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first 48 h (reach 23%) and then leveled off.37-39 The slight decrease of Ag+ after 192 h might be

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ascribed to the re-adsorption of Ag+ to AgNPs and photoreduction of Ag+ ions,30 which generated

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some precipitates with metallic luster at the inner wall of container and the surface of the solution.

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In the presence of 1 mg/L HA, the release profile of Ag+ was similar to that without HA, but the

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release amount was lower (the maximum release percentage was 18%). While the concentration of

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HA increased to higher levels, the release amount of Ag+ was greatly reduced to very low levels

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(e.g., 6% for 10 mg/L HA and 5% for 20 mg/L HA). Furthermore, we also monitored the Ag+

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release at a lower AgNP concentration (1.0 mg/L). The release curve (Fig. 4b) showed that the mass

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fraction of the released Ag+ at lower AgNP concentration was higher (the maximum release

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percentage reached 43%) than that at higher AgNP concentration. Because the Ag+ release with high

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concentration of AgNPs can be inhibited by multiple factors, such as reaction-induced proton

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depletion, oxygen depletion, and surface reaction inhibition, the first-order rate constant of Ag+

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release decreased with the initial AgNP concentration increasing.14 After 72 h incubation, the Ag+

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release began to decrease was due probably to the reduction of Ag+ by HA. 12


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We also carried out control experiments in the dark (see Fig. S11) and using Ag+ instead of

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AgNPs (see Fig. S12). With 1 mg/L HA in dark condition, Ag+ was continuously released from

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AgNPs with no small particles being formed in the solution (Fig. S11). When using Ag+ instead of

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AgNPs, the Ag+ could be reduced to form particles with the characteristic absorption peak of

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AgNPs (Fig. S12), which was consistent with that observed in previous reports.21, 22, 30

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3.2. Mechanism for the HA-dependent persistence of AgNPs in water

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Based on the results mentioned above, we propose the possible mechanism for the

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HA-dependent persistence of AgNPs in water as depicted in Fig. 5. Depending on the concentration

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of HA, AgNPs show distinctly different behaviors and fates. Generally, three types of mechanisms

261

may be responsible for the observed phenomena:

262

(1) With no HA or extremely low concentration of HA, a portion of AgNPs in the solution can

263

be dissolved into Ag+ ions by photo-oxidation as indicated by the Ag+ release curve in Fig. 4a. The

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other portion of AgNPs can aggregate into large particles (as illustrated by Fig. 3c) or transform into

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Ag0 precipitates via the photoreduction of Ag+ ions and almost no AgNPs ultimately remain in the

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solution (as indicated by the absence of characteristic absorption peak of AgNPs after 336 h in Fig.

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2a).

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(2) With low concentration of HA (e.g., 1 mg/L), the Ag+ release and aggregation of AgNPs can

269

be partly inhibited and AgNPs can persist in the aquatic environment for a relatively long period of

270

time. Interestingly, we observed the formation of numerous small particles and the redistribution of

271

particle size of AgNPs (see Fig. 3e-h). As mentioned above, HA can reduce Ag+ ions to form AgNPs

272

under natural conditions.20-22 Therefore, we suggest that the small particles observed in Fig. 3e and

273

3g should be secondary AgNPs formed by the reduction of Ag+ ions with HA induced by sunlight.

274

That is, the Ag+ released from primary AgNPs can be reduced back with HA under sunlight to 13



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re-generate AgNPs (i.e., secondary AgNPs). This mechanism can well explain the TEM results in

276

Fig. 3e and 3g. The photo-oxidative release of Ag+ can cause a concentration gradient of Ag+ around

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the AgNPs, and the Ag+ concentration near the surface of AgNPs should be higher than that in the

278

bulk solution, so the secondary particles are formed first around the primary AgNPs (i.e., the

279

“satellite” particles shown in Fig. 3e), and then dispersed to the bulk solution with a slight size

280

growth. The formation of new nanoparticles also explained the slight increase of λmax absorbance of

281

AgNPs with 1 mg/L HA in Fig. 2b.

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(3) With high concentration of HA (e.g., 20 mg/L), HA can tightly wrap the primary AgNPs due

283

to the high affinity of AgNPs for HA40, 41 and greatly suppress the Ag+ release. Thus, the primary

284

particles can keep stable in the solution for long term with no significant change in particle size or

285

morphology.

286

These mechanisms were also supported by the control experiments in dark condition (see Fig.

287

S11) and with Ag+ instead of AgNPs (see Fig. S12). In dark condition, Ag+ was continuously

288

released from AgNPs and no secondary particles were found in the solution, also suggesting that the

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secondary AgNPs were produced by the sunlight-induced Ag+ reduction with HA; and with Ag+

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only, the formation of nanoparticles proved the reducing ability of HA toward Ag+ under natural

291

conditions.

292

From the discussion above, we can learn that there are two critical HA concentrations that affect

293

the long-term persistence of AgNPs in the HA-containing water (see Fig. 5). The first critical

294

concentration (CC1) determines whether the AgNPs can persist in the water; and the second critical

295

concentration (CC2) determines whether the AgNPs can maintain their particle size. Specifically, for

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HA concentration (CHA) < CC1, AgNPs will ultimately transform into other forms of Ag via

297

photo-oxidation release of Ag+ ions; for CC1 < CHA < CC2, AgNPs can persist in the water due to 14


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the secondary formation of AgNPs mediated by HA and sunlight, but the particle size of AgNPs

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may rearrange; and for CHA > CC2, the AgNPs can persist in the water for long term with no

300

significant change of particle size. To the best of our knowledge, this may be the first demonstration

301

of the role of secondary particle formation in the persistence of AgNPs in an aquatic system. It

302

should be noted that the values of the two critical HA concentrations are closely dependent on the

303

intrinsic properties of AgNPs (e.g., particle size) and environmental factors (e.g., pH, salinity,

304

sunlight, etc.). In this study, the CC1 and CC2 were regarded to be 1 and 20 mg/L, respectively, and

305

their values may vary in other systems.

306

3.3 Evidence from Ag isotope fractionation during the Ag+ release process of AgNPs

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Although the characterization techniques used above such as absorption spectrometry and TEM

308

can well characterize the nanoparticles, they do not give direct evidence for the processes.

309

Therefore, to verify the proposed mechanisms in Fig. 5, we further investigated the Ag isotope

310

fractionation during the Ag+ release process of AgNPs in the presence of different concentrations of

311

HA. In our previous work,30 we have demonstrated that natural stable isotope fractionation can be

312

used as a new means to study environmental processes of nanoparticles. The different

313

environmental processes of AgNPs can cause different Ag isotope fractionation signatures, which

314

can be used as unique tracers to indicate the processes that the nanoparticles have undergone. For

315

instance, the HA-mediated reduction of Ag+ can cause a significant enrichment of

316

produced AgNPs with an isotopic enrichment factor (ε) up to 0.86‰.30 This value is at the same

317

level as the largest variations in the

318

the natural isotopic information is independent of the concentration of element or particle size, and

319

thus it can provide additional evidence for the studied processes. For this consideration, here we

320

also used Ag isotope fractionation as a tracer to study the processes in the persistence of AgNPs at

109

109

Ag in the

Ag/107Ag ratio found in terrestrial samples.31, 32, 42 Notably,

15



321

Page 16 of 34

different HA concentrations.

322

The Ag isotope fractionation during the dissolution of AgNPs with different concentrations of

323

HA is shown in Fig. 6. All the data followed isotope mass balance (see Table S1-S3). In the AgNP

324

solution without HA (Fig. 6a), the relative fraction of Ag+ released was in the range of

325

12.93-21.36%. At first 6 h, the dissolution of AgNPs caused a significant Ag isotope fractionation

326

(∆109/107Agrelease = 0.18‰; *P < 0.001) with

327

phenomenon was consistent with that observed previously,30 and could be explained by the

328

formation of an intermediate metal oxide phase (e.g., Ag2O) at the surface of AgNPs via

329

photo-oxidation.43, 44 Following the equilibrium isotope fractionation rules,45 the heavier isotope

330

was prone to being enriched in a phase with a higher oxidation state. Thus, the metal oxide phase

331

should be enriched in

332

time increased to 96 h, the degree of Ag isotope fractionation was mitigated compared with that at 6

333

h (∆109/107Agrelease = 0.07‰ for 96 h;

334

incubation time was suggested to be caused by the photoreduction of Ag+ ions, in which 109Ag was

335

prone to be enriched in the Ag0 precipitates30, thus depleting the 109Ag in the released Ag+. This was

336

also supported by the control experiment in the dark (Fig. S13). It should be noted that the degree of

337

the Ag isotope fractionation is affected by the relative fraction reacted.45 From 96 h to 168 h,

338

although the relative fraction of released Ag+ did not change much (22.5% to 21.4%), the residual

339

AgNPs kept dissolving and a large portion (56.8%) of total Ag was transformed into Ag0

340

precipitates (see Table S1), which caused the δ109/107Ag value of the residual AgNPs to continue to

341

shift negatively (∆109/107Agresidual = -0.13‰). Thus, to maintain the isotopic mass balance in the

342

system (see Table S1), the δ109/107Ag value of the released Ag+ increased again.

343

109

109

Ag being enriched in the released Ag+. This

Ag, causing the released Ag+ to be enriched in

109

Ag. As the incubation

**

P < 0.01). The mitigation of the fractionation at longer

In the case of 1 mg/L HA (Fig. 6b), a large Ag isotope fractionation was obtained at 6 h with 9.7% 16


Page 17 of 34


344

of AgNPs being dissolved and the released Ag+ being significantly enriched in

345

(∆109/107Agrelease = 0.30‰; *P < 0.001). When the incubation time increased to 96 h, the reactive

346

fraction of released Ag+ increased to 17.8%, and the δ109/107Ag value of released Ag+ greatly shifted

347

negatively (∆109/107Agrelease = 0.08‰) compared with that at 6 h (**P < 0.001), indicating a

348

remarkable loss of

349

that obtained without HA (Fig. 6a) and could not be ascribed to the photoreduction of Ag+ ions.30 As

350

mentioned above, the HA-mediated reduction of Ag+ to generate AgNPs can cause a significant Ag

351

isotope fractionation with the formed AgNPs being enriched in

352

enriched in 107Ag.30 This effect highly matched the variation in δ109/107Ag of the released Ag+ upon

353

the incubation time increasing, suggesting the occurrence of the HA-mediated reduction of Ag+ in

354

the system. Therefore, the Ag isotope fractionation provided a clear evidence for the HA-mediated

355

secondary formation of AgNPs from the released Ag+. After 168 h, the δ109/107Ag of released Ag+

356

(∆109/107Agrelease = 0.10‰) and the relative fraction of Ag+ released (from 17.8% to 17.4%) kept

357

unchanged compared with that at 96 h (***P = 0.38), suggesting that the experimental system had

358

reached an equilibrium within 96 h.

109

109

Ag

Ag in the released Ag+. Such a variation in δ109/107Ag was much larger than

109

Ag and the residual Ag+ being

359

In the presence of high concentration of HA (e.g., 20 mg/L; Fig. 6c), the Ag+ release was

360

strongly inhibited and the relative fraction of released Ag+ was kept to very low levels (4.1% for 96

361

h and 4.8% for 168 h). No significant Ag isotope fractionation was observed during the whole

362

incubation process (∆109/107Agrelease = 0.05‰ for 96 h (*P = 0.19) and 0.05‰ for 168 h (**P = 0.14)).

363

This result demonstrated that AgNPs were in a relatively steady state in the presence of high

364

concentration of HA. Specifically, the absence of Ag isotope fractionation suggested that the Ag+

365

release was directly suppressed by the high concentration of HA rather than being in situ reduced,

366

otherwise the HA-mediated reduction of Ag+ would cause significant Ag isotope fractionation. 17



367

Overall, the results of Ag isotope fractionation in Fig. 6a-c strongly evidenced the proposed

368

mechanism in Fig. 5 for the persistence of AgNPs in the HA-containing water.

369

3.4 Environmental implications

370

The environmental fate of AgNPs is highly relevant for their risk assessment for ecosystems and

371

human health. So far, engineered AgNPs discharged to environmental sewer water system are

372

usually thought to end up in wastewater treatment plants (WWTPs),46, 47 deposit in sludge, and

373

finally transformed to Ag2S as stable form and transport in ecosystems.48, 49 However, recent studies

374

also found that Ag2S can further transform into other Ag species or even regenerate AgNPs in some

375

environmental conditions, such as in the presence of environmentally relevant concentration of

376

Fe(III).50 This calls for reconsideration of the lifecycle of AgNPs in the environment. Furthermore,

377

a considerable mass of engineered AgNPs is being directly discharged into natural waters besides to

378

waste management systems.51 The present study demonstrates that AgNPs may persist in

379

HA-containing waters for longer periods of time than expected, even with low concentration of HA.

380

The detailed interpretation on the role of HA in the persistence of AgNPs will help understand the

381

environmental fate and assess the impacts of AgNPs on the ecosystems. For instance, the particle

382

size redistribution caused by the secondary particle formation should be considered in the risk

383

assessment of AgNPs for aquatic ecosystems because the toxicity of AgNPs is size-dependent.52

384

It is worth noting that the HA-mediated secondary particle formation of AgNPs relies on the

385

dissolved oxygen and sunlight. Despite the high affinity of Ag for reduced sulfur, this reaction is

386

highly possible to occur in surface waters, because surface waters are usually saturated with oxygen

387

absorbed from the atmosphere and from the photosynthesis of algae. The dissolved oxygen can

388

oxidize the reduced sulfur53 and facilitate the HA-mediated reduction of Ag+ (see Fig. 5). It should

389

also be noted that the transformations of AgNPs in real aquatic environments are extremely 18


Page 18 of 34

Page 19 of 34


390

complicated and affected by many factors. For instance, at naturally occurring concentration of Cl,

391

Ag+ can be transformed into AgClx(x-1)- soluble species,54 which may mitigate the reduction of Ag+.

392

This study only highlights the role of HA. The roles of other environmental factors and how the

393

critical concentrations of DOM are affected by other factors need to be clarified in future studies.

394

Another significance of this study is the successful use of stable isotope fractionation as a new tool

395

to unravel the complex environmental processes of AgNPs. On one hand, isotopic analysis can

396

provide extra information in addition to concentration of element and particle size; on the other

397

hand, the isotopic mass balance calculation ensures a more reliable investigation on the studied

398

system. This tool requires no artificial or radioactive tracers, and thereby it is particularly suitable

399

for studying real environmental systems. However, the application of this tool in real environment

400

is still limited by the insufficient sensitivity of instrument (i.e., MC-ICP-MS) and analytical

401

capability for complex samples. Therefore, the power of this tool in environmental nano research

402

needs to be further explored by continued efforts on technical improvement in the instrument and

403

expanded applications to more types of nanoparticles.

404

19



405

Acknowledgements

406

This work was financially supported by the Chinese Academy of Sciences (No. XDB14010400,

407

QYZDB-SSW-DQC018), the National Basic Research Program of China (2015CB931903,

408

2015CB932003), and the National Natural Science Foundation of China (No. 91543104, 21377141,

409

21422509). Q. Liu acknowledges the support from the Youth Innovation Promotion Association of

410

CAS.

411

Supporting information

412 413 414

The Supporting Information is available free of charge on the ACS Publications website.

Competing financial interests The authors declare no competing financial interests.

415

20


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Page 21 of 34


416

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(Ag2S-NPs) in the aquatic environment: Photoinduced transformation of Ag2S-NPs in the presence

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51. Blaser, S. A.; Scheringer, M.; MacLeod, M.; Hungerbuhler, K. Estimation of cumulative

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Titma, T.; Heinlaan, M.; Visnapuu, M.; Koller, D.; Kisand, V.; Kahru, A. Size-dependent toxicity of 26


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silver nanoparticles to bacteria, yeast, algae, crustaceans and mammalian cells in vitro. PLoS One

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559

Effect of chloride on the dissolution rate of silver nanoparticles and toxicity to E. coli. Environ. Sci.

560

Technol. 2013, 47 (11), 5738-5745.

561 562

27



563

Figures

564 565

Fig. 1. Characterization of PVP-coated AgNPs. (a) A typical TEM image. (b) Particle size

566

distribution. The average diameter (D) of AgNPs is 17.0 ± 4.2 nm. The particle size distribution

567

statistics were counted based on more than 200 particles.

568

28


Page 28 of 34

Page 29 of 34


569 570

Fig. 2. UV-vis absorption spectra of AgNP suspension after incubation with HA for 336 h

571

under sunlight. (a) 0 mg/L HA, (b) 1 mg/L HA, and (c) 20 mg/L HA. AgNP concentration: 17.0

572

mg/L. The UV-vis absorbance of the solution was recorded at an interval of 24 h.

573

29



Page 30 of 34

574 575

Fig. 3. Typical TEM images and corresponding particle size distribution of AgNPs incubated with 0, 1, and 20 mg/L of HA under sunlight for

576

24 h and 336 h. (a-d) AgNP solution incubated with no HA for 24 h (a, b) and 336 h (c, d); (e-h) AgNP solution exposed to 1 mg/L HA for 24 h (e, g)

577

and 336 h (f, h); AgNP solution exposed to 20 mg/L HA for 24 h (i, k) and 336 h (j, l). AgNP concentration: 17.0 mg/L. D: average diameter of

578

particles. The particle size distribution statistics were counted based on more than 200 particles. The inset in e shows the high-resolution image for the

579

secondarily formed “satellite” particles surrounding the primary AgNPs.

30


Page 31 of 34


580

581

Fig. 4. The Ag+ release kinetics from AgNPs during 336 h incubation under sunlight. (a) Effect

582

of HA concentration on the Ag+ release kinetics of AgNPs. AgNP concentration: 17.0 mg/L. (b)

583

Effect of AgNP concentration on the Ag+ release kinetics of AgNPs. HA concentration: 1 mg/L. The

584

error bars represent 1 SD from three parallel experiments (n = 3).

585

31



586 587

Fig. 5. Scheme showing the possible mechanisms for the HA-dependent persistence of AgNPs

588

in water.

589 590

32


Page 32 of 34

Page 33 of 34


591 592

Fig. 6. Ag isotope fractionation during the Ag+ release process from AgNPs under sunlight in the presence of HA. (a) HA = 0 mg/L, *P < 0.001,

593

**

594

δ109/107Ag value of the initial AgNPs. The percentages in parentheses represent the mass percentages of released Ag+ in the total Ag. The data are also

595

given in Table S1-S3 in SI to make a mass balance calculation. The error bars represent 1SD from three parallel experiments (n = 3).

P < 0.01,

***

P < 0.001. (b) HA = 1 mg/L, *P < 0.001, **P < 0.001,

***

P = 0.38. (c) HA = 20 mg/L, *P = 0.19, **P = 0.14. The bars start from the

33



TOC 219x128mm (150 x 150 DPI)


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Role of Secondary Particle Formation in the Persistence of Silver

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