Modeling the Primary Size Effects of Citrate-Coated Silver

ARTICLE pubs.acs.org/est

Modeling the Primary Size Effects of Citrate-Coated Silver Nanoparticles on Their Ion Release Kinetics Wen Zhang, Ying Yao, Nicole Sullivan, and Yongsheng Chen* School of Civil and Environmental Engineering, Georgia Institute of Technology, Atlanta, Georgia, 30332.

bS Supporting Information ABSTRACT: Ion release is an important environmental behavior of silver nanoparticles (AgNPs), and characterization of Agþ release is critical for understanding the environmental fate, transport, and biological impacts of AgNPs. The ion release kinetics of AgNPs with three primary diameters (20, 40, and 80 nm) were studied by dispersing them in quarter-strength Hoagland medium at two initial concentrations (300 and 600 μg/L). Agþ release rates were found to depend on primary particle size and concentration, when other environmental factors (e.g., dissolved oxygen and protons) were kept constant. A kinetic model was developed to describe the Agþ release based on the hard sphere theory using the Arrhenius equation. The model fitted the experimental data well with correlation coefficients of 0.97
0.99, and the model usefully interpreted the dependence of ion release kinetics on the primary particle size and concentration. Moreover, the effects of environmental factors (e.g., dissolved oxygen, pH, temperature, and salinity) potentially can be interpreted as well. This model provides fundamental insight into the ion release kinetics of AgNPs in aqueous environments, allowing us to better understand and predict the nanotoxicity of AgNPs.

’ INTRODUCTION The widespread use of silver nanoparticles (AgNPs) in food packaging, clothing, and other household products inevitably leads to AgNP release into the environment.1
3 Much evidence has shown that AgNPs are toxic to bacteria 1 and some aquatic organisms 2,3 and accumulate in phytoplankton and marine invertebrates.4
6 After entry into the aquatic environment, AgNPs likely will release silver ions (or Agþ) and coexist with their ionic species.7,8 AgNPs and Agþ exhibit different physiochemical properties and subsequent biological potency.9
11 For example, AgNPs cause cell death by pitting bacterial cell membranes and increasing permeability;12 they also exhibit size- and shapedependent toxicity.1,13 In contrast, Agþ interacts with the thiol groups of proteins,14 resulting in inactivation of vital enzymes; disrupts bacterial membrane integrity and permeability;15 and likely affects DNA replication.16 When compared on the basis of total mass added, the antibacterial activity of AgNPs was found to be lower than that of Agþ.17,18 However, AgNPs are reported to exhibit the Trojan-horse mechanism, potentially disrupting cellular functions inside the cells.19 Thus, the ratio of AgNPs to Agþ is a critical factor in understanding and assessing the toxicity of AgNPs.2,3,9,20 Previous studies on ion release kinetics of AgNPs primarily focused on environmental factors, such as dissolved oxygen (DO), pH, temperature, salinity (or ionic strength), and the presence of natural organic matter (NOM).9,21 Kittler et al. also studied surface coating effects on the kinetics of Agþ release.21 r 2011 American Chemical Society

Unfortunately, the effects of intrinsic particle properties (e.g., size) on ion release kinetics have not been fully investigated even though this may provide direct evidence correlating the particle properties with toxicity.11,22,23 This information is also essential for manufacturers who wish to produce environmentally benign nanomaterials. More importantly, the solution chemistries (e.g., deionized (DI) water with oxygenation) used in some studies are not environmentally relevant. Lastly, the comprehensive modeling description for ion release kinetics, which is pivotal for establishing correlations between ion release kinetics of AgNPs and their toxicokinetics, is still inadequate.24,25 Although Kittler et al. developed an empirical model using first-order reaction kinetics, the model parameters neither carry mechanistic meanings nor explain the size dependence of ion release. Thus, the need to advance ion release kinetics modeling of AgNPs is evident and is also essential for decision-makers when establishing regulations regarding nanosilver manufacturing and disposal.26 Citrate-coated AgNPs are used in this study because they are the most popular silver colloids that are lab-synthesized for general and specific applications;27,28 the citrate coating serves as both the reducing agent and a stabilizer by decorating the particles with negative charges.24,25 Citrate-coated AgNPs are the Received: December 15, 2010 Accepted: April 13, 2011 Revised: April 5, 2011 Published: April 22, 2011 4422

dx.doi.org/10.1021/es104205a | Environ. Sci. Technol. 2011, 45, 4422–4428

Environmental Science & Technology most widely used in commercial products and are recommended by the Organization for Economic Co-operation and Development (OECD) for standard use in toxicology.9,20 This study aims to elucidate the fundamental mechanisms of Agþ release kinetics through experimental and modeling approaches. The ion release experiments were conducted in glass media bottles filled with quarter-strength Hoagland medium, a common hydroponic nutrient medium for growing plants.29 The electrolytic nutrients of this medium simulate relevant ionic conditions in soils,30 sediments,31 and groundwater systems,32,33 environmental matrices that AgNPs are likely to enter. In this study, we focused on the effects of primary particle size and concentration on ion release kinetics, because previous research has indicated that factors such as DO, pH, salinity, and temperature likely affect ion release from AgNPs.9,21 To quantitatively describe the ion release kinetics, we derived for the first time a kinetic model, which provides insight into the mechanisms of ion release kinetics of AgNPs in aqueous environments.

’ MATERIALS AND METHODS AgNPs. Water suspensions of citrate-coated AgNPs (20, 40, and 80 nm) without dissolved citrate were purchased from Ted Pella, Inc. (Redding, CA, USA). According to the vendor, the AgNPs will remain monodispersed in the stock suspension without aggregation for a minimum of one year when stored in a sealed bottle at 4 °C. The total AgNP concentrations of 20, 40, and 80 nm AgNPs in the stock suspensions were 6.47 ( 0.02, 6.56 ( 0.10, and 6.21 ( 0.29 mg/L respectively, and the stock suspensions contained no Agþ as measured by inductively coupled plasma-mass spectrometry (ICP-MS, Elan DRC II, PerkinElmer, USA), which is described below. Characterization of AgNPs. The morphology and sizes of AgNPs were determined by transmission electron microscopy (TEM) on a Philips EM420 at 120 kV. TEM samples were prepared by placing 5 μL of the aqueous AgNP water suspension on copper grids with a continuous carbon film coating, followed by solvent evaporation at room temperature. The mean particle diameters and size distribution were computed with at least 30 particles from each TEM micrograph. The crystallography of AgNPs was analyzed with X-ray diffraction (XRD) using an X-ray powder diffractometer (Philips PW 1800, PANalytical Almelo, The Netherlands) with a Cu KR radiation (λ = 1.54 Å). XRD analysis was recorded at a scanning rate of 0.02°/s in the 2θ range of 35
45°. The particle size distribution (PSD) and hydrodynamic radius were determined by dynamic light scattering (DLS) on a Zetasizer Nano ZS instrument (Malvern Instruments, UK) using 1.5 mL of 600 μg/L AgNP suspension diluted with DI water in a standard macro-cuvette (pass length: 10 mm). The temperature was maintained at 25 °C, and the scattering angle was 173°. A refractive index (RI) of 1.07 and an absorption value of 0.01 were used for the AgNPs.34 Prior to ICP-MS analysis, all glassware and digestion vessels were carefully cleaned and soaked in 10% (v/v) HNO3 for 24 h, and finally rinsed five times with 18 MΩ DI water. AgNP suspensions were digested as follows: 1 mL of sample suspension was mixed with 2 mL of trace-metal grade HNO3 (67
70%, w/w, Fisher Scientific) and heated on a hot plate at 120 °C for 20 min. Then, the digested solutions were diluted and analyzed by ICP-MS.

ARTICLE

Agþ Release Experiment. The ion release experiments were

conducted in 125 mL glass media bottles containing 100 mL of modified quarter-strength Hoagland medium. A detailed experimental schematic is shown in Figure S1 of the Supporting Information. The medium has the following salts with the concentrations (mg/L) shown in parentheses: KH2PO4 (52), Ca(NO3)2 (40.25), KNO3 (34.25), CaSO4 (72.25), MgSO4 (117.25), and K2SO4 (40.25). Some common anions, such as chloride, were not present in the medium to minimize the potential precipitation with Agþ. Before AgNPs were dispersed into the medium, it was equilibrated with the air; the equilibrated DO and pH of the medium were approximately 7.8 ( 0.2 mg/L and 5.6 ( 0.1, respectively. Two concentrations (300 and 600 μg/L) of AgNPs were added to the medium, and then the glass media bottles were placed in the dark with no stirring (and thus no hydrodynamic effects on ion release kinetics). Every 12 or 24 h, 4 mL of the medium was sampled from the bottle, and the released Agþ was separated from the AgNPs using Amicon Ultra-4 centrifugal filter units with pore diameters of 1
2 nm (Amicon Ultracel 3K, Millipore, USA). After centrifugation for 40 min at 5000  g (5430R, Eppendorf, German), 3 mL of the filtrate was collected and mixed with 2 mL of 67% nitric acid for ICP-MS analysis (Figure S2 of the Supporting Information shows the potential adsorption of Agþ onto the filter). Monitoring of Agþ concentrations in the medium lasted approximately 350 h with pseudoequilibrium observed. The Agþ release experiments were repeated three times. Aggregation of AgNPs in Quarter-Strength Hoagland Medium. The aggregation experiments were performed using time-resolved-dynamic light scattering (TR-DLS) on the same Zetasizer instrument. The same concentrations of AgNPs as the ion release experiments were dispersed into quarter-strength Hoagland medium in plastic cuvettes (BrandTech, VWR, PA). The cuvettes were then sealed by parafilm with a few punctures to allow air to enter while avoiding water evaporation. The DLS measurement began immediately. The DLS measurement for AgNPs without exposure to air used the same settings as above except that the cuvettes were sealed airtight by parafilm. DLS measurements were performed for 300 h, and for each sample the DLS measurement was repeated three times to confirm the aggregation observations.

’ RESULTS AND DISCUSSION Characterizations of AgNPs. Parts a
c of Figure 1 show TEM micrographs of the morphology of the AgNPs used in this study. The AgNPs are close to spherical, and the three different primary diameters were 19.9 ( 0.5, 40 ( 0.6, and 80 ( 0.7 nm, which were consistent with the manufacturer-reported values. Section S3 provides the particle size distribution diagrams computed from TEM images and additional comparison with the particle sizes calculated from Scherer’s equation. The X-ray diffraction patterns in parts d
f of Figure 1 indicate that the AgNPs have a typical face-centered cubic (FCC) structure, and a lattice parameter of 0.407 nm, which is the same as the reference data (0.4085 nm).24 Moreover, the multiringed electron diffraction patterns in Figure S4 of the Supporting Information further confirm that the AgNPs are FCC polycrystalline. Parts g
i of Figure 1 show PSD diagrams of AgNPs dispersed in DI water, which indicate that the three sizes of AgNPs had distinct hydrodynamic diameters. Polydispersivity index (PDI) is a dimensionless measure of the broadness of the size distribution 4423

dx.doi.org/10.1021/es104205a |Environ. Sci. Technol. 2011, 45, 4422–4428

Environmental Science & Technology

ARTICLE

Figure 1. (a
c) TEM images of 20, 40, and 80 nm AgNPs on a carbon-coated grid. The white bar on the bottom left equals 100 nm. (d
f) X-ray diffraction patterns. (g
i) PSD diagrams in terms of volume percentage of AgNPs dispersed in DI water. The number and intensity PSDs are provided in Figure S5 of the Supporting Information.

reported from the Zetasizer Nano ZS instrument. The PDI values are 0.038, 0.184, and 0.166 for 20, 40, and 80 nm AgNPs, respectively. As all of these are less than 0.25, the AgNPs are considered well dispersed in the suspension without aggregation. Model Development with the Arrhenius Equation. We derived the Agþ release rate (γAgþ) on the basis of the hard sphere collision theory that has been used previously for modeling the nanoparticle dissolution kinetics.35,36 The proposed oxidation reaction stoichiometry of AgNPs is:9,23 1 þ AgðsÞ þ O2ðaqÞ þ 2Hþ ðaqÞ T 2AgðaqÞ þ H2 OðlÞ 2 Because of their small size, AgNPs may act as soluble reactants and the oxidation reaction can be described by first-order reaction kinetics. Accordingly, γAgþ can be expressed by the Arrhenius equation: γAgþ ¼ kCAgNPs ½O2 0:5 ½Hþ 2  
Ea ¼ f exp ½AgNPs½O2 0:5 ½Hþ 2 kB T f ¼

πσ 2AB

  8πkB T 1=2 NA μAB

ð1Þ

ð1-1Þ

σAB ¼ σA þ σ B

ð1-2Þ

mA mB mA þ mB

ð1-3Þ

μAB ¼

where γAgþ is the Agþ release rate (mol/(L 3 h)); k is the reaction rate constant (mol/h); [AgNPs], [O2], and [Hþ] are the molar concentrations (mol/L) of AgNPs, DO, and proton, respectively; f is the frequency factor for the reaction; Ea is the activation energy (J); T is temperature (298 K); kB is the Boltzmann constant (1.38  10
23 J/K); NA is Avogadro’s number (6.02  1023/mol); σAB is the collision radius (nm); σA and σB are the molecular radii of the reactants A and B (nm); μAB is the reduced mass (g/mol); and mA and mB are the molecular weights of reactants A and B (g/mol). For the silver oxidation reaction, reactant A is AgNPs and reactant B is either oxygen or protons. The radius of oxygen or protons is much smaller than that of an AgNP (r). Likewise, the molecular weights of oxygen and protons are much lower than that of AgNP. Therefore, eqs 1
2 and 1
3 can be simplified to σAB = r and μAB = mB. Rearranging eq 1
1 yields f = πr2((8πkBT)/(mB))1/2NA, and thus γAgþ can be written as:  1=2  
Ea 2 8πkB T γAgþ ¼ πr NA exp ½AgNPs½O2 0:5 ½Hþ 2 mB kB T ð2Þ The molar concentration ([AgNPs]) of AgNPs can be replaced by the mass-based concentration ([Ag]): ½AgNPs ¼

½Ag 4 NA F πr 3 3

ð3Þ

where the initial [Ag] values were 300 and 600 μg/L, r is the AgNP radius (nm), and F is the density of AgNPs 4424

dx.doi.org/10.1021/es104205a |Environ. Sci. Technol. 2011, 45, 4422–4428

Environmental Science & Technology

ARTICLE

Figure 2. Agþ release kinetics in quarter-strength Hoagland medium; the model fits are shown by the black dashed lines.

(∼10.5 g/cm3 ). Combining eqs 2 and 3 yields:     3 8πkB T 1=2
1
Ea F exp γAgþ ¼ ½Agr
1 ½O2 0:5 ½Hþ 2 4 mB kB T ð4Þ eq 4 reveals that γAgþ could increase monotonically with increasing [Ag], [O2], or [Hþ] and is inversely proportional to r. Furthermore, increasing the temperature could significantly increase γAgþ. All of these predictions agree with previous findings.9,17,21 To describe the experimental results with the model, we derived the released Agþ concentration ([Agþ]released) as a function of time (t) with eq 4. Because γAgþ = (d[Agþ]released)/(dt) and [Agþ]released = [Ag]initial
[Ag], [Agþ]released can be solved by integrating eq 4: ½Agþ released ¼ ½Aginitial
expðb
atÞ

ð5Þ

In eq 5, a and b are the fitting parameters, and a = (3)/(4)(8kBT)/(mB)1/2F
1 exp (
Ea)/(kBT)r
1[O2]0.5[Hþ]2. When t is zero, [Agþ]released is zero, thus b = ln([Ag]initial). [O2] and [Hþ] are treated as constants over time, an assumption that is verified in the following sections. Thus, eq 5 can be rewritten as: ½Agþ released ¼ ½Aginitial ½1
expð
atÞ

ð6Þ

eq 6 is used to fit the experimental results with a and b or [Ag]initial as fitting parameters using the nonlinear least-squares curve fitting technique. More details about the model deviation and data fitting are in section S5. Ion Release Kinetics and Model Fitting. Figure 2 shows the concentrations of the released Agþ over time in the medium. For all sizes of AgNPs, the released Agþ concentrations increased almost linearly within the first 8
12 h, and some slight fluctuation in Agþ concentrations was observed for 40 and 80 nm NPs after the equilibrium was reached. The initial Agþ concentrations were nonzero and slightly different for each size, which indicates that oxidation may have occurred for each size at time = 0, probably due to unintended exposure to oxygen during the experimental setup. Increasing the initial AgNP concentration from 300 to 600 μg/L significantly increased the ion release rates for each size of AgNPs, which agrees with our model prediction. Small AgNPs (e.g., 20 nm) took longer (approximately 300 h) to reach reaction equilibrium, whereas large NPs (e.g., 80 nm) reached equilibrium after 100 h. The relative amounts of residual

AgNPs and released Agþ over time are calculated and summarized in Table S3 of the Supporting Information. In all cases (three sizes of AgNPs at two initial concentrations), AgNPs did not completely dissolve as a result of oxidation through the entire experimental period, which was also found by Kettler et al. and Liu et al.9,21 Thermodynamic analysis predicts that only in the long term should AgNPs be completely dissolved in aqueous environments containing sufficient DO.9 Thus, our observed equilibrated state of the Agþ release kinetics within 350 h does not correspond to the physicochemical equilibrium in which the released silver concentrations should be independent of the amount of AgNPs present. Obviously, Figure 2 reveals that the equilibrated concentrations of Agþ depend on the primary particle size and the initial AgNP concentration. This can be explained by the different thermodynamic properties of AgNPs as compared with their bulk phase.37,38 Small AgNPs have higher specific surface areas and higher enthalpies of formation compared to the same amount of bulk silver.23 In addition, the faster ion release rate for smaller AgNPs agrees with previous studies on minerals that are well interpreted with the Kelvin equation,39,40 which indicated that small particles had higher solubility than their larger counterparts. Sotiriou et al. also observed high ion release rates for small AgNPs, which may contribute to greater antibacterial activity.17 Speciation of Agþ was proven not to influence the ion release kinetics significantly, because thermodynamic calculations using MINEQL 4.5 estimated that more than 99% of the silver is present as soluble Agþ at equilibrium.9 The black dashed lines in Figure 2 show the model fit using eq 6. Table S1 of the Supporting Information summarizes the correlation coefficients (R2), which ranged from 0.97 to 0.99 for the three particle sizes and two initial concentrations, suggesting excellent agreement between the model fitting and experimental data. The fitting parameters, a and b, are shown in Table S1 of the Supporting Information. If the activation energy (Ea) is assumed to be the same for different sizes of AgNPs, the values of a should be in the order 20 nm > 40 nm > 80 nm, as a is inversely proportional to r. However, the values of a for 20 nm AgNPs are neither twice those for 40 nm AgNPs, nor four times those for 80 nm AgNPs, which suggests that Ea may vary with particle size. Considering the size dependence of surface energy, crystallinity, and crystallographic plane distribution 25,41 for AgNPs of different sizes, surface reactivity and Ea can be size-dependent too. Moreover, even with the same total surface area, small NPs have a greater fraction of atoms at edges and corners than large particles, which makes the surface more reactive.40 The fitted values of b for 20 and 40 nm AgNPs are similar to each other, whereas the values for 80 nm AgNPs are slightly lower but still in a comparable order of magnitude. The fitted values of [Ag]initial are much lower than the actual applied concentrations of AgNPs (300 and 600 μg/L), as shown in Table S1 of the Supporting Information. As mentioned above, this is because the model is fitted with the data on Agþ release, which did not correspond to a complete dissolution of AgNPs. We also thoroughly compared our experimental conditions and kinetic parameters with those in previous studies in section S6. The kinetic parameter (k) determined from our experimental data is in the same order of magnitude as those of Liu et al.,9 whereas those from Kittler et al. 21 are less than ours by 1 order of magnitude, probably due to the insufficient DO, which was not mentioned. 4425

dx.doi.org/10.1021/es104205a |Environ. Sci. Technol. 2011, 45, 4422–4428

Environmental Science & Technology

ARTICLE

Figure 3. Aggregation kinetics of 20 nm AgNPs in quarter-strength Hoagland medium with two initial concentrations (b 300 μg/L, O 600 μg/L) in the presence of DO (a) and in the absence of DO (b). (c) and (d) show the linear growth of the aggregation of AgNPs.

Environmental Factors for Agþ Release Kinetics. As men-

tioned previously, [O2] and [Hþ] are assumed to be constant over time during the experimental period, which allows us to derive [Agþ]released as a function of time. These assumptions are verified by comparison between the consumption and mass transfer rates of DO and protons in the oxidation reaction. When the medium was exposed to the air, the transfer of oxygen from air to the aqueous phase (O2(g)hO2(aq)) continuously sustained DO. Likewise, the transfer of carbon dioxide
(CO2(g) h CO2(aq) þ H2O h H2CO3(aq) h Hþ (aq) þ HCO3 ) and its speciation provide protons and a buffer capacity for pH. To calculate the consumption rates of DO and protons, the ion release rates are first calculated using eq 5; these are shown in Table S6 of the Supporting Information: γAgþ ¼

½Agþ t þ Δt
½Agþ t d½Agþ  ¼ dt Δt

ð7Þ

where [Agþ]tþΔt and [Agþ]t are the concentrations of released Agþ at time t and t þ Δt, which are obtained from Figure 2. Then, the consumption rates of DO and protons at different reaction times also can be determined based on the reaction stoichiometry. Table S6 of the Supporting Information shows that the Agþ release rates were highest at the beginning of the reaction and gradually decreased. Interestingly, high initial AgNP concentration (600 μg/L) did not significantly increase the ion release rates for any of the three sizes; instead, the ion release seemed to be slightly inhibited within the initial 4 h, after which the rates began to increase and eventually exceeded those for 300 μg/L. This probably occurred for two reasons. First, the increasing concentration of AgNPs may temporarily limit the diffusivity of oxygen and protons to reaction sites on silver particle surfaces according to the Stokes
Einstein equation 42,43 and thus slow the oxidation process. Second, AgNPs aggregated rapidly during the oxidation at the larger concentration of AgNPs, as shown later, because aggregation is reported to have pseudosecondorder kinetics.44 Thus, AgNPs at the initial concentration of 600 μg/L may aggregate faster and thus have less surface area

available for oxidation reactions in comparison to samples containing 300 μg/L. To compare the reaction consumption rates with the mass transfer rate of oxygen, we determined the overall mass transfer coefficient (KLa) for oxygen in quarter-strength Hoagland medium using the standard clean water test method.45 Figure S6 of the Supporting Information shows the DO measurement, and the detailed calculation of KLa is provided in section S7. On the basis of KLa, the oxygen mass transfer rate was 6.01 μmol/(L 3 h), which is much higher than the oxygen consumption rates in Table S6 of the Supporting Information, indicating that DO was not a rate-limiting factor in this study. Similarly, the pH was monitored during different experimental stages and Figure S7 of the Supporting Information shows that the medium’s pH was stable at approximately 5.6. Minor changes in pH occurred after nitrogen purging, exposure to the air, or the addition of AgNPs, probably because the mass transfer of carbon dioxide and speciation provided sufficient protons and buffer capacity for the oxidation of AgNPs. Therefore, the assumptions in our model derivations are justified. Aggregation Effects on Ion Release Kinetics. Because of the decrease in electrostatic repulsion between AgNPs in the presence of electrolytes,46 AgNPs may be unstable, and particle aggregation could occur concurrently with Agþ release. These two physiochemical processes may affect each other. Figure 3 shows the TR-DLS results for AgNPs (original diameter, 20 nm) at the two initial concentrations of 300 and 600 μg/L with or without exposure to the air. Over 300 h of monitoring, the hydrodynamic radii of the AgNPs increased to up to 3000 nm in the presence of DO, as shown in part a of Figure 3. In contrast, part b of Figure 3 shows that in the absence of DO, AgNPs aggregated steadily and reached the equilibrated hydrodynamic sizes of approximately 400
600 nm. During the initial 7
8 h the hydrodynamic radii increased almost linearly, and comparison of the slopes of the linear regression equations indicates that the aggregation rates of 600 μg/L (the red circles) are higher than those of 300 μg/L (black dots), which is consistent with previous 4426

dx.doi.org/10.1021/es104205a |Environ. Sci. Technol. 2011, 45, 4422–4428

Environmental Science & Technology studies.47,48 After the linear growth, the hydrodynamic radii in the presence of DO changed dramatically in random ways. The PDI values became greater than 0.25, indicating that the particle suspension became polydispersed. After approximately 50 to 150 h, the hydrodynamic radii in part a of Figure 3 began to decline, whereas the sizes of aggregated AgNPs did not decrease without DO (part b of Figure 3). The random changes in PSD and the decrease in particle sizes may be triggered by several physiochemical processes that occurred simultaneously during the aggregation,23 such as Agþ release, Agþ speciation, and precipitation, as illustrated in the insert in part a of Figure 3. The relatively stable sizes of AgNP aggregates without DO are attributed only to the aggregation, and clearly, the decrease in hydrodynamic sizes in the presence of DO should be attributed mainly to the ongoing oxidation. The effect of sedimentation of aggregated AgNPs was negligible by comparison with the random Brownian motion energy (section S9 for details). Clearly, the aggregation altered the PSD of AgNPs, which partially explains the discrepancy between the model fit and the experimental data. The kinetics model in eqs 4
6 assumes that all particles have a homogeneous and uniform particle size (e.g., 20 nm) throughout the oxidation process. In reality, the silver NP dispersion has a narrow PSD rather than a uniform size. To investigate the aggregation effects on the model fitting quality, we calculated the correlation coefficients for the model fits in different periods (e.g, from 8 to 96 h) and show the results in Table S2 of the Supporting Information. As the experimental time elapsed, the correlation coefficients decreased, and the model fitting quality became poor in some cases. Thus, aggregation influenced the ion release kinetics, largely because of the changes in the available surface areas of AgNPs. Considering the particle aggregation and the mass transfer of oxygen or carbon dioxide, the effect of solution salinity on the ion release kinetics should be discussed, because high ionic strength promotes particle aggregation,9 decreases oxygen solubility, and consequently affects the ion release kinetics according to our model. Our model potentially can be modified to incorporate the salinity effect by changing [O2] as a function of salinity. This again emphasizes the importance of employing environmentally revelent water chemistries for environmental fate studies of AgNPs because the behavior of AgNPs in the environment is certainly different from that in pure water. Looking at Figure 2 and Figure 3, it is striking that even though AgNPs aggregated, the primary particle size, rather than the aggregated sizes, played the governing role in ion release kinetics. This seems reasonable because despite considerable particle aggregation and increased hydrodynamic radius, aggregates preserve most of the surface area available for the heterogeneous reaction with oxygen.9 The ion release behaviors of NPs may be linked to their fate, transport, and even biological impacts. Better predictive models are clearly needed to understand their environmental behaviors and biological impacts. Although our current study is oriented toward a quantitative understanding of particle size and concentration effects on ion release kinetics, additional model simulations on the effects of temperature, DO, and pH were conducted and are compared with previous studies in section S10. The kinetic model is capable of interpreting the ion release kinetics as a function of reaction time, primary particle size, particle concentration, DO, pH, and temperature. These model equations potentially can predict the metal ion release kinetics of any type of metal-containing NPs as a result of chemical reactions

ARTICLE

(e.g., oxidation). This article used relatively simple water chemistries to facilitate the modeling and analysis of primary particle size and concentration effects on ion release kinetics. The present results, both modeling and experimental, suggest that the primary particle size of citrate-coated AgNPs, rather than the aggregated sizes, will dramatically govern the Agþ release kinetics. The ion release kinetics modeling may lay the groundwork for developing appropriate models to describe the kinetic behaviors of NPs in environmentally relevant solution chemistries.

’ ASSOCIATED CONTENT

bS

Details about the Agþ release experiments, characterizations using XRD and electron diffraction patterns, data fitting, kinetic parameters and model simulation of other factors (temperature, DO, and pH), and monitoring of DO and pH in quarter-strength Hoagland medium. This material is available free of charge via the Internet at http:// pubs.acs.org. Supporting Information.

’ AUTHOR INFORMATION Corresponding Author

*E-mail: [email protected]; phone: (þ1) 404-894-3089; fax: (þ1) 404-894-2278.

’ ACKNOWLEDGMENT This study was partially supported by the U.S. Environmental Protection Agency Science to Achieve Results Program Grant RD-83385601 and Engineering Research Center (ERC)/Semiconductor Research Corporation (SRC)/ESH grant (425.025). Wei Lin and Zhuo Li from the School of Materials Science and Engineering at Georgia Institute of Technology are thanked for the XRD analysis. ’ REFERENCES (1) Morones, J. R.; Elechiguerra, J. L.; Camacho, A.; Holt, K.; Kouri, J. B.; Ramirez, J. T.; Yacaman, M. J. The bactericidal effect of silver nanoparticles. Nanotechnology 2005, 16 (10), 2346–2353. (2) Navarro, E.; Piccapietra, F.; Wagner, B.; Marconi, F.; Kaegi, R.; Odzak, N.; Sigg, L.; Behra, R. Toxicity of silver nanoparticles to Chlamydomonas reinhardtii. Environ. Sci. Technol. 2008, 42 (23), 8959–8964. (3) Asharani, P. V.; Wu, Y. L.; Gong, Z. Y.; Valiyaveettil, S. Toxicity of silver nanoparticles in zebrafish models. Nanotechnology 2008, 19 (25), 8. (4) Fisher, N. S.; Wang, W.-X. Trophic transfer of silver to marine herbivores: A review of recent studies. Environ. Toxicol. Chem. 1998, 17 (4), 562–571. (5) Dhawan, A.; Sharma, V. Toxicity assessment of nanomaterials: Methods and challenges. Anal. Bioanal. Chem. 2010, 398 (2), 589–605. (6) Shaw, J. R.; Wood, C. M.; Birge, W. J.; Hogstrand, C. Toxicity of silver to the marine teleost (Oligocottus maculosus): Effects of salinity and ammonia. Environ. Toxicol. Chem. 1998, 17 (4), 594–600. (7) Farre, M.; Gajda-Schrantz, K.; Kantiani, L.; Barcelo, D. Ecotoxicity and analysis of nanomaterials in the aquatic environment. Anal. Bioanal. Chem. 2009, 393 (1), 81–95. (8) Klaine, S. J.; Alvarez, P. J. J.; Batley, G. E.; Fernandes, T. F.; Handy, R. D.; Lyon, D. Y.; Mahendra, S.; McLaughlin, M. J.; Lead, J. R. Nanomaterials in the environment: behavior, fate, bioavailability, and effects. Environ. Toxicol. Chem. 2008, 27 (9), 1825–1851. 4427

dx.doi.org/10.1021/es104205a |Environ. Sci. Technol. 2011, 45, 4422–4428

Environmental Science & Technology (9) Liu, J. Y.; Hurt, R. H. Ion release kinetics and particle persistence in aqueous nano-silver colloids. Environ. Sci. Technol. 2010, 44 (6), 2169–2175. (10) Panyala, N. R.; Pena-Mendez, E. M.; Havel, J. Silver or silver nanoparticles: a hazardous threat to the environment and human health? J. Appl. Biomed. 2008, 6 (3), 117–129. (11) 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. 2009, 43 (19), 7285–7290. (12) Shahverdi, A. R.; Fakhimi, A.; Shahverdi, H. R.; Minaian, S. Synthesis and effect of silver nanoparticles on the antibacterial activity of different antibiotics against Staphylococcus aureus and Escherichia coli. Nanomedicine: NBM 2007, 3 (2), 168–171. (13) Pal, S.; Tak, Y. K.; Song, J. M. Does the antibacterial activity of silver nanoparticles depend on the shape of the nanoparticle? A study of the gram-negative bacterium Escherichia coli. Appl. Environ. Microbiol. 2007, 73 (6), 1712–1720. (14) Matsumura, Y.; Yoshikata, K.; Kunisaki, S.; Tsuchido, T. Mode of bactericidal action of silver zeolite and its comparison with that of silver nitrate. Appl. Environ. Microbiol. 2003, 69 (7), 4278–4281. (15) Feng, Q. L.; Wu, J.; Chen, G. Q.; Cui, F. Z.; Kim, T. N.; Kim, J. O. A mechanistic study of the antibacterial effect of silver ions on Escherichia coli and Staphylococcus aureus. J. Biomed. Mater. Res. 2000, 52 (4), 662–668. (16) Neal, A. L. What can be inferred from bacterium-nanoparticle interactions about the potential consequences of environmental exposure to nanoparticles? Ecotoxicology 2008, 17 (5), 362–371. (17) Sotiriou, G. A.; Pratsinis, S. E. Antibacterial activity of nanosilver ions and particles. Environ. Sci. Technol. 2010, 44 (14), 5649–5654. (18) Jin, X.; Li, M.; Wang, J.; Marambio-Jones, C.; Peng, F.; Huang, X.; Damoiseaux, R.; Hoek, E. M. V. High-throughput screening of silver nanoparticle stability and bacterial inactivation in aquatic media: Influence of specific ions. Environ. Sci. Technol. 2010. (19) Park, E. J.; Yi, J.; Kim, Y.; Choi, K.; Park, K. Silver nanoparticles induce cytotoxicity by a Trojan-horse type mechanism. Toxicology in Vitro 2010, 24 (3), 872–878. (20) Damm, C.; Munstedt, H. Kinetic aspects of the silver ion release from antimicrobial polyamide/silver nanocomposites. Appl. Phys. A: Mater. Sci. Process. 2008, 91 (3), 479–486. (21) Kittler, S.; Greulich, C.; Diendorf, J.; Koller, M.; Epple, M. Toxicity of silver nanoparticles increases during storage because of slow dissolution under release of silver ions. Chem. Mater. 2010, 22 (16), 4548–4554. (22) Choi, O.; Hu, Z. Q. Size dependent and reactive oxygen species related nanosilver toxicity to nitrifying bacteria. Environ. Sci. Technol. 2008, 42 (12), 4583–4588. (23) Marambio-Jones, C.; Hoek, E. M. V. A review of the antibacterial effects of silver nanomaterials and potential implications for human health and the environment. J. Nanopart. Res. 2010, 12 (5), 1531–1551. (24) Wiley, B.; Sun, Y. G.; Xia, Y. Synthesis of silver nanostructures with controlled shapes and properties. Acc. Chem. Res. 2007, 40 (10), 1067–1076. (25) Wiley, B.; Sun, Y. G.; Mayers, B.; Xia, Y. N. Shape-controlled synthesis of metal nanostructures: The case of silver. Chem.—Eur. J. 2005, 11 (2), 454–463. (26) Hassellov, M.; Readman, J. W.; Ranville, J. F.; Tiede, K. Nanoparticle analysis and characterization methodologies in environmental risk assessment of engineered nanoparticles. Ecotoxicology 2008, 17 (5), 344–61. (27) Tolaymat, T. M.; El Badawy, A. M.; Genaidy, A.; Scheckel, K. G.; Luxton, T. P.; Suidan, M. An evidence-based environmental perspective of manufactured silver nanoparticle in syntheses and applications: A systematic review and critical appraisal of peer-reviewed scientific papers. Sci. Total Environ. 2010, 408 (5), 999–1006. (28) Pillai, Z. S.; Kamat, P. V. What factors control the size and shape of silver nanoparticles in the citrate ion reduction method? J. Phys. Chem. B 2003, 108 (3), 945–951.

ARTICLE

(29) Griffin, G. J.; Hale, M. G.; Shay, F. J. Nature and quantity of sloughed organic matter produced by roots of axenic peanut plants. Soil Biol. Biochem. 1976, 8 (1), 29–32. (30) Gibeaut, D. M.; Hulett, J.; Cramer, G. R.; Seemann, J. R. Maximal biomass of Arabidopsis thaliana using a simple, low-maintenance hydroponic method and favorable environmental conditions. Plant Physiol. 1997, 115 (2), 317–319. (31) Reddy, K. R.; Tucker, J. C. Growth and nutrient uptake of Pennywort (Hydrcotyle umbellata L.), as Influenced by the nitrogen concentration of the water. J. Aquat. Plant Manage 1985, 23, 35–40. (32) Corseuil, H. X.; Moreno, F. N. Phytoremediation potential of willow trees for aquifers contaminated with ethanol-blended gasoline. Water Res. 2001, 35 (12), 3013–3017. (33) Annandale, J. G.; Jovanovic, N. Z.; Tanner, P. D.; Benade, N.; Du Plessis, H. M. The sustainability of irrigation with gypsiferous mine water and implications for the mining industry in South Africa. Mine Water Environ. 2002, 21 (2), 81–90. (34) Velikov, K. P.; Zegers, G. E.; van Blaaderen, A. Synthesis and characterization of large colloidal silver particles. Langmuir 2003, 19 (4), 1384–1389. (35) Ho, C.-M.; Yau, S. K.-W.; Lok, C.-N.; So, M.-H.; Che, C.-M. Oxidative dissolution of silver nanoparticles by biologically relevant oxidants: A kinetic and mechanistic study. Chem.
Asian J. 2010, 5 (2), 285–293. (36) Meulenkamp, E. A. Size Dependence of the Dissolution of ZnO Nanoparticles. J. Phys. Chem. B 1998, 102 (40), 7764–7769. (37) Luo, W. H.; Hu, W. Y.; Xiao, S. F. Size effect on the thermodynamic properties of silver nanoparticles. J. Phys. Chem. C 2008, 112 (7), 2359–2369. (38) Gilbert, B.; Huang, F.; Zhang, H.; Waychunas, G. A.; Banfield, J. F. Nanoparticles: Strained and stiff. Science 2004, 305 (5684), 651–654. (39) Stumm, W.; Morgan, J. J. Aquatic Chemistry, 3rd ed.; John Wiley & Sons, Inc.: New York, 1996. (40) Liu, J.; Aruguete, D. M.; Murayama, M.; Hochella, M. F. Influence of size and aggregation on the reactivity of an environmentally and industrially relevant manomaterial (PbS). Environ. Sci. Technol. 2009, 43 (21), 8178–8183. (41) Kittler, S.; Greulich, C.; Koller, M.; Epple, M. Synthesis of PVPcoated silver nanoparticles and their biological activity towards human mesenchymal stem cells. Materialwiss. Werkst. 2009, 40 (4), 258–264. (42) Mackay, M. E.; Dao, T. T.; Tuteja, A.; Ho, D. L.; Van Horn, B.; Kim, H.-C.; Hawker, C. J. Nanoscale effects leading to non-Einstein-like decrease in viscosity. Nat. Mater. 2003, 2 (11), 762–766. (43) Tarasenko, A.; Jastrabik, L. Influence of cytoskeleton on nanoparticle migration in biological cells. Appl. Phys. Lett. 2009, 95 (17), 173705. (44) Chen, K. L.; Mylon, S. E.; Elimelech, M. Aggregation kinetics of alginate-coated hematite nanoparticles in monovalent and divalent electrolytes. Environ. Sci. Technol. 2006, 40 (5), 1516–1523. (45) Tchobanoglous, G.; Burton, F. Wastewater Engineering Treatment, Disposal, and Reuse; Third ed.; Metcalf & Eddy, 1991. (46) Buettner, K. M.; Rinciog, C. I.; Mylon, E. S. Aggregation kinetics of cerium oxide nanoparticles in monovalent and divalent electrolytes. Colloids Surf., A 2010, 366 (1
3), 74–79. (47) Kallay, N.; Zalac, S. Stability of nanodispersions: a model for kinetics of aggregation of nanoparticles. J. Colloid Interface Sci. 2002, 253 (1), 70–76. (48) Trinh, L. T. T.; Kjoniksen, A. L.; Zhu, K. Z.; Knudsen, K. D.; Volden, S.; Glomm, W. R.; Nystrom, B. Slow salt-induced aggregation of citrate-covered silver particles in aqueous solutions of cellulose derivatives. Colloid Polym. Sci. 2009, 287 (12), 1391–1404.

4428

dx.doi.org/10.1021/es104205a |Environ. Sci. Technol. 2011, 45, 4422–4428