Article pubs.acs.org/est

Enhanced Formation of Silver Nanoparticles in Ag+‑NOM-Iron(II, III) Systems and Antibacterial Activity Studies Nathaniel F. Adegboyega,† Virender K. Sharma,*,‡ Karolina M. Siskova,§ Renata Vecerova,∥ Milan Kolar,∥ Radek Zbořil,§ and Jorge L. Gardea-Torresdey⊥ †

Chemistry Department, Florida Institute of Technology, 150 West University Boulevard, Melbourne, Florida 32901, United States Department of Environmental and Occupational Health, School of Rural Public Health, Texas A&M University, 1266 TAMU, College Station, Texas 77843-1266, United States § Regional Centre of Advanced Technologies and Materials, Department of Physical Chemistry, Faculty of Science, Palacky University, Slechtitelu 11, 78371 Olomouc, Czech Republic ∥ Department of Microbiology, Faculty of Medicine and Dentistry, Palacky University in Olomouc, Hnevotinska 3, 77146 Olomouc, Czech Republic ⊥ Department of Chemistry, The University of Texas at El Paso, 500 West University Avenue, El Paso, Texas 79968, United States ‡

S Supporting Information *

ABSTRACT: This work reports the role of iron redox pair (Fe3+/Fe2+) in the formation of naturally occurring silver nanoparticles (AgNPs) in the aquatic environment. The results showed that Fe3+ or Fe2+ ions in the mixtures of Ag+ and natural organic matter enhanced the formation of AgNPs. The formation of AgNPs depended on pH and types of organic matter. Increase in pH enhanced the formation of AgNPs, and humic acids as ligands showed higher formation of AgNPs compared to fulvic acids. The observed results were described by considering the potentials of redox pairs of silver and iron species and the possible species involved in reducing silver ions to AgNPs. Dynamic light scattering and transmission electron microscopy measurements of AgNPs revealed mostly bimodal size distribution with decrease in size of AgNPs due to iron species in the reaction mixture. Minimum inhibitory concentration of AgNPs needed to inhibit the growth of various bacterial species suggested the role of surfaces of tested Gram-positive and Gram-negative bacteria. Stability study of AgNPs, formed in Ag+-humic acid/fulvic acids-Fe3+ mixtures over a period of several months showed high stability of the particles with significant increase in surface plasmon resonance peak. The environmental implications of the results in terms of fate, transport, and ecotoxicity of organic-coated AgNPs are briefly presented.

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INTRODUCTION Silver nanoparticles (AgNPs) possess unique optical, catalytic, and sensing properties that lead to their application in various processes, such as direct propylene epoxidation, carbon dioxide electrolysis, electrochemical, and chemical sensors (probes).1,2 AgNPs have also shown antibacterial effects and, therefore, are applied in a large number of consumer products, such as textiles, sprays, laundry additives, medical devices, home appliances, food supplements, and paints.3−6 Current and future estimates on applications of AgNPs have raised concerns because of their possible association with health effects, which include cardiovascular diseases, argyria, and respiratory inflammation.7,8 AgNPs released into the environment may accumulate in marine invertebrates and phytoplankton and may also exert toxic effects to aquatic organisms.9−12 In the past few years, several workers have investigated fate, transport, dissolution, and bioavailability of engineered AgNPs to understand their potential risks to the ecosystem.9,10,13−16 Aggregation and mobility of engineered AgNPs depends on water chemistry (e.g., pH, ionic components, ionic strength, and organic matter) and nature of coating material (e.g., sulfide, © 2014 American Chemical Society

chloride, carbonate, borate, citric acid, polysaccharides, surfactants, proteins, and polymers).17−26 Toxic effects of engineered AgNPs on aquatic organisms under different water conditions are also forthcoming.27−31 Comparatively, very little information on the potentially formed AgNPs in aquatic environment is available in the literature. Recently, we have shown the possibility of direct formation of AgNPs under environmental conditions.32,33 Other workers have also shown the likely formation of AgNPs by reduction of Ag+ by humics and microbiological activities.2,9,34 AgNPs formed from the reduction of Ag+ ion by humic and fulvic acids were stable for several months and could possibly transport to a longer distance from their points of origin. Effect of metals on the formation and stability of these AgNPs is currently unknown and is the focus of the present study. Iron was chosen as a metal because it is one of the most interesting trace metals in Received: Revised: Accepted: Published: 3228

December 18, 2013 February 11, 2014 February 13, 2014 February 13, 2014 dx.doi.org/10.1021/es405641r | Environ. Sci. Technol. 2014, 48, 3228−3235

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terrestrial and aquatic environment.35 Iron exists as ferrous (Fe2+) and ferric (Fe3+) ions in natural waters. Redox conditions, pH, and concentration and nature of organic ligands control the speciation of iron.36 In the present study, we have demonstrated that the formation of AgNPs can be enhanced in the presence of either Fe2+ or Fe3+ ions without significantly altering the stability of AgNPs. Several studies have been performed to test the antibacterial properties of engineered AgNPs in the presence of natural organic matter (NOM) (e.g., humic acids),9,37 but similar studies of AgNPs, formed from the direct reduction of Ag+ by NOM are missing in the literature. Therefore, the antibacterial properties of these AgNPs were also tested against several bacterial species. The objectives of the present work are (i) to demonstrate the enhancement of the formation of AgNPs by adding Fe2+ and Fe3+ to Ag+-NOM system, (ii) to understand the effect of NOM and pH on the formation of AgNPs in Fe2+-Ag+-NOM system, (iii) to determine the stability of formed AgNPs with and without Fe3+, and (iv) to investigate antibacterial activity of formed AgNPs in Fe2+-Ag+NOM system against various gram negative (g−) and gram positive (g+) bacterial species.

Antimicrobial Study. Antimicrobial activities of the formed AgNPs in Ag+-SRHA and Fe2+-Ag+-SRHA mixtures were investigated at pH 7.0 using a standard microdilution method, which has been described in detail elsewhere.38 The detailed procedure of the antimicrobial tests is given in SI, Text S2.

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RESULTS AND DISCUSSION Effect of Fe(II)/Fe(III). In the initial experiment, Ag+ was mixed with SRHA at pH 6.0 and heated at 90 °C for 4 h. The characteristic intense yellow color due to surface plasmon resonance (SPR) of AgNPs appeared (Figure 1A). Control

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EXPERIMENTAL SECTION Formation of Silver Nanoparticles. Chemicals used and preparation of solutions of the study are described in Supporting Information, SI, Text S1. Silver nanoparticles were formed by reducing silver ions with HA and FA.33 Briefly, 3 mL of silver nitrate was mixed with 3 mL of HA or FA in a capped test tube under normal laboratory conditions, and the solution mixture was placed in a water bath kept at 90 ± 1 °C. The solution mixture was heated for 4 h. After this heating period, the samples tubes were equilibrated at room temperature (24 ± 1 °C), followed by UV−vis measurements to observe the SPR of AgNPs. Similar experiments, without heating the mixed solutions, were also performed at 24 °C and solutions were always kept in dark. Exposure of solution to light of the laboratory was during the spectroscopic measurements. An Agilent Technologies 8453 GA11034 spectrophotometer and Specord S600 (Analytic Jena) spectrophotometer were applied to collect UV−vis spectra using 1 cm optical path length cuvettes at room temperature. In the case of reaction mixtures of Ag+-FA or Ag+-HA also containing Fe2+ and Fe3+, trace amounts of Fe2+ and Fe3+ ions from the stock solutions were initially added into 3 mL of HA or FA, which were already adjusted to desired pH 4.0 and 6.0. The mixing was allowed to occur for 3 min. This mixed solution was then added into 3 mL silver nitrate. The size distributions of AgNPs were carried out by collecting dynamic light scattering (DLS) measurement using a Zetasizer Nano Series (Malvern Instruments) instrument. This instrument was also used to measure zeta potentials in aqueous dispersions. Morphological study was done on a JEOL JEM-2010 transmission electron microscope (TEM) equipped with a LaB6 cathode. TEM images were performed on a JEOL 2010 microscope at accelerating voltage of 160 kV. Sample dispersion was drop-coated on copper grid with holey-carbon film. The samples were dried at room temperature and were washed several times with deionized water before being subjected to analyses. Chemical mapping (EDS) of selected samples was performed with HRTEM FEI TitanG2 microscope.

Figure 1. UV−vis absorption spectra and DLS determined size distributions based on intensity fluctuation of AgNPs in Ag+-SRHA solutions with and without Fe2+ and Fe3+ at pH 6.0. (A) UV−vis spectra and (B) DLS measurements. ([Ag+] = 1 × 10−3 mol L−1, 40 mg L−1 SRHA).

spectra of SRHA and iron-SRHA are given in SI Figure SI-1. Next, Fe2+ in the solution was attempted by mixing Ag+ with Fe2+-SRHA solution at pH 6.0 and heated at 90 °C for 4 h. The SPR of AgNPs also appeared, but had enhanced absorbance (Figure 1A). Increase in concentration of Fe2+ had the slight effect of further enhancement of the absorbance. Significantly, when solution of Ag+ and Fe2+ without SRHA was mixed at pH 6.0, there was no SPR of AgNPs (Figure 1A). Similar results were observed by mixing Ag+ with Fe2+-SRHA solution at 24 °C and pH 6.0 (SI Figure SI-2). No growth of AgNPs was seen without Fe2+, but SPR of AgNPs appeared in solution containing Fe2+ in 7 days. Further growth in the formation of AgNPs occurred in next 7 days, but again no SPR of AgNPs appeared without Fe2+ (SI Figure SI-2). Increase in concentration of Fe2+ showed little increase in absorbance; similar to the formation of AgNPs at 90 °C. The results can be understood by considering the potential of the redox pairs in the system (Table 1).39−43 The direct 3229

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Table 1. Redox Potentials of Possible Reactions in the Ag+NOM-Fe Systema (T1) (T2) (T3) (T4) (T5) (T6) (T7) a

reaction

E0 vs NHE

reference

Ag + e ⇌ Ag Ag+ + Ago∞ + e− ⇌ Ago∞ Fe3+ + e− ⇌ Fe2+ Q + 2H+ + 2e− ⇌ HQ FA(ox) + e− ⇌ FA(Red) HA(ox) + e− ⇌ HA(Red) FeIII(HS) + e− ⇌ FeII(HS)

−1.8 V 0.8 V 0.77 V −0.699 V ∼0.5 V ∼0.7 V −0.20 to 0.30 V

39 40 41 40 42 43 41

+

−

o

matter could reduce silver ion to metallic silver through reactions 5 and 6.32 For example, quinone kinds of moieties in NOM44 have significant redox potential to result in the formation of silver cluster because of positive redox potential of the reaction, E0 ≈ 1.5 V; obtained from reactions T2 and T4 in Table 1. Positive redox potentials were also obtained when redox pairs of FA and HA were used (i.e., reactions T2, T5, and T6). Reaction 7 resulted in the formation of cluster.45 The presence of Fe2+ in the reaction mixture of Ag+ and SRHA can result in several reactions, which could lead to the formation and dissociation of complexes, as well as the formation of reactive oxygen species (ROS) (Table 2).46−48 The potential of Fe(II)-HA/Fe(III)-HA (i.e., reaction T7 of Table 1) indicates the feasibility of the formation of AgNPs in the presence of Fe2+ (reaction 8). Basically, complex formation of Fe2+ with NOM (reaction F1, Table 2) provides additional driving force to form AgNPs and thus enhanced the formation of AgNPs in the Ag+-Fe2+-SRHA mixed solution.

FA-Fulvic Acid; HA-Humic Acid; HS-Humic Substances.

reduction of Ag+ ion to isolated Ag0 by either Fe2+ or SRHA is not thermodynamically feasible because of the high negative potentials of the reactions (eqs 1 and 2). Ag + + Fe2 + ⇌ Ag o + Fe3 + E 0 = − 2.57 V

(1)

Ag + + HA (Red) ⇌ Ag o + HA (Ox) E 0 = − 2.50 V

(2)

Ag + + Ag o∞ + Fe II(HS) E 0 ≈ 0.5−1.0 V ⇌ Ag o∞ + Fe III(HS)

However, if Ag+ ions exist on stable silver clusters or onto a solid silver electrode, then formation is thermodynamically possible due to positive electrode potential (e.g., reaction T2 in Table 1). The organic matter plays a significant role in generating a silver cluster through a series of reactions (reactions 3−6). HS in eqs 5 and 6 represents the NOM. 2Ag + + 2OH− → Ag 2O + H 2O fast

(3)

Ag 2O + (Ag +)n → Ag 2O − (Ag +)n fast

(4)

Ag 2O‐(Ag +)n + HS ⇌ Ag 2O − (Ag +)n − HS K ad

(5)

The possible ROS in the reaction mixtures are O2•−, H2O2, and • OH (reactions F2−F5, Table 2), which may react either with Fe(II)-SRHA complex (reactions F4−F6) or with SRHA (reactions F7 and F8). The formed O2•− may also react with Ag+ to ultimately result in AgNPs (reactions G1−G4). However, there is also a possibility that O2•− may not be involved in any of the reactions, but self-decomposes to H2O2 (reaction F3), which has been shown to dissolve AgNPs.49 The enhanced formation of AgNPs due to Fe2+ suggests the dominance of the reactions that formed O2•−, and also the direct formation of AgNPs. It should be pointed out that the rate constants given in Table 2 are at 25 °C and temperature effect on these reactions would have different values at 90 °C. The influence of Fe3+ on the formation of AgNPs in the mixture of Ag+-SRHA at pH 6.0 and 90 °C was also studied. The results in Figure 1A demonstrate the enhanced formation of AgNPs; similar to the effect of Fe2+. This was surprising because the Fe3+-SRHA complex, formed in reaction F9 (Table 2), has no reducing capability to form metallic silver from Ag+.

Ag 2O‐(Ag +)n − HS → Ag 2O − (Ag +)n − 1 + Ag o k red + HS(ox) (6) o

o

Ag + Ag → Ag 2 fast

(8)

(7)

Initially formed colloidal Ag2O in reaction 3 could adsorb Ag+ (reaction 4).32 The functional groups present in the organic

Table 2. Possible Reactions and Their Rate Constants in the Ag+-NOM-Fe System at 25 °Ca reaction F1 F2 F3 F4 F5 F6 F7 F8 F9 F10 F11 F12 G1 G2 G3 G4 a

reference

Fe + L → Fe L FeIIL + O2 → FeIIIL + O2•− O2•− + O2•− + 2 H+ → H2O2 + O2 FeIIL + O2•− + 2H+ → FeIIIL + H2O2 FeIIL + H2O2 → FeIIIL + •OH + OH− FeIIL + •OH → FeIIIL + OH− L + O2•− + 2 H+ → L(Ox) +H2O2 L + •OH → L(Ox) + OH− FeIII + L → FeIIIL FeIIIL1 → FeIII + L1 FeIIIL2 → FeIII + L2 FeIIIL + O2•− → FeIIL + O2 Ag+ + O2•− → Ag0 + O2 Ag0 + O2•− → Ag0− + O2 Ag0− + O2 → Ag0 + O2•− Ag+ + Ag0− → Ag2 II

II

1

−1

−1

k = 1.5× 10 (g/L) s k = 1.0 × 102 M−1s−1 k = (0.2−5.0) × 104 M−1s−1 k = 1.0 × 107 M−1s−1 k = 3.1 × 104 M−1s−1 k = 5.0 × 108 M−1s−1 k = 1.2 × 101 (g/L)−1 s−1 k = 3.1× 107 (g/L)−1 s−1 k = 5.0 × 102 (g/L)−1 s−1 kd1 = 3.55 × 103 s−1 kd2 = 8.08 × 105 s−1 k = 1.2 × 108 M−1s−1 k = 6.4 × 101 M−1s−1 k = 1.0 × 1010 M−1s−1 k = 2.0 × 109 M−1s−1 k = 6.0 × 105 M−1s−1

48 48 47 48 48 48 48 48 48 46 46 48 47 47 47 47

L1 − NOM1; L2 − NOM2; L = SRFA 3230

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Figure 2. UV−vis absorption spectra of AgNPs in reduction of Ag+ by different organic matter with and without Fe3+ at pH 6.0. (A) Without Fe3+ and (B) with Fe3+ ([Ag+] = 1 × 10−3 mol L−1, [SRHA] = [SRFA] = [NLFA] = 40 mg L−1; [Fe3+] = 13 μM).

different rate constants, depending on the nature of ligand (reactions F10 and F11 in Table 2). Also, the rate constants of the reduction of Fe3+-organic complexes by O2•− to form Fe2+ will be different for HA and FA as ligands (reaction F12 of Table 2).46 Reactions involved in the presence of Fe3+ are affected by the nature of ligands and thus would give the observed growth of AgNPs (Figure 2B). Effect of pH. In this set of experiments, growth of AgNPs in Fe2+-Ag+-SRHA solution was studied in the pH range of 3.0− 6.0. At pH 3.0, there was no significant growth of AgNPs without Fe2+; while significant formation of AgNPs was seen at pH 4.0 (Figure 3). Control spectra are in SI Figure SI-4. However, growth was less than that at pH 6.0. Thus, Fe2+ enhanced the growth of AgNPs at pH 4.0; similar to the results at pH 6.0, (Figure 3). The effect of pH can be described by considering processes that favor the growth of AgNPs with increasing pH. First, the higher amount of OH− ions in solution would increase the rate of formation of Ag2O-(Ag+)n (reaction 4) particles by increasing the formation of colloidal Ag2O (reaction 3) and hence increase the growth of AgNPs through reactions 5−7. Second, the deprotonated species of HA increase in concentration compared to protonated species with increase in pH. The deprotonated species facilitated more reduction of Ag+ ion (reaction 6) than that of protonated species because deprotonated functional groups have relatively higher electron density compared to protonated species of HA. In the presence of Fe3+, an increase in pH favors the formation of hydroxo species (e.g., FeOH2+ and Fe(OH)2+), which are more likely to be reduced by HA than the Fe3+ species.50 The higher reduction rate of Fe3+ to form Fe2+ would lead to the observed growth of AgNPs, as expected from reaction 8. The effect of pH on size distribution of AgNPs was also studied (SI Figure SI-5). The mean HDD of the AgNPs (prepared at pH 4.0) was 132 nm without Fe2+, which is much smaller than that observed at pH 6.0 (i.e., 201 nm). Significantly, AgNPs were monodispersed at pH 4.0; suggested by the PI of 0.183 whereas polydispersity in particle sizes was observed at pH 6.0. However, the presence of Fe2+ resulted in a decrease in HDD, which was 63 nm at pH 4.0 and also the particles were polydispersed; similar to results at pH 6.0 (see Figure 1B). The presence of Fe2+ in Ag+-SRHA mixtures caused an increase in the formation of of AgNPs at both pH (see Figure 3), with a corresponding decrease in HDD of the particles. Size and shapes of formed AgNPs were further examined by TEM images (Figures 3B−E). TEM images of AgNPs formed in Ag+-SRHA at pH 4.0 revealed mostly very small nanoparticles ( SRHA > SRFA, consistent with the previous study.32 The difference in moieties of humic and fulvic acids such as semiquinones, thiols, phenolic, and carboxylic groups44 may explain the trend seen in growth of AgNPs in mixture of Ag+ and HA/FA (Figure 2A). In the case of Fe3+ present in the mixture of Ag+-HA/FA, an enhancement in the growth of AgNPs for all HA and FAs was observed (Figure 2B). The order of growth, SRHA > NLFA > SRFA, was somewhat different from the trend without Fe3+. This suggests the role of ligands, complexed with Fe3+, which will have different rates, depending on the nature and functional groups involved in the complexation. For example, the dissociation of Fe3+-organic complexes to yield Fe3+ have 3231

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of zeta potential did not change significantly with Fe2+ in the reaction mixture (−23 and −24 mV at pH 6.0 and 4.0, respectively). The negative functionality of the organics in humic acids would induce the negative character of the shell surrounding the AgNPs. The shift of SPR band maximum from ∼400 to ∼440 nm and its fwhm (full-with-at-half-maximum) increased in the case of AgNPs formed in Ag+-SRHA-Fe2+ mixture at pH 4.0, in comparison to those formed in Ag+-SRHA mixture at pH 4.0. This could also be described on the basis of TEM images (Figure 3C), where particles of different shapes (including faceted shapes) with diameters between 5 and 100 nm dominate. Significantly, it correlated well with a slightly increased polydispersity (0.291), determined by DLS for AgNPs in the Ag+-SRHA-Fe2+ system at pH 4.0. The enlarged particle size distribution and the presence of faceted shapes were most probably responsible for the shape of SPR band possessing UV-shoulder (∼350 nm) and red-tail (Figure 3A). Faceted shapes reveal two contributions to SPR band: transversal and longitudinal,51 hence the occurrence of the UV-shoulder and red-tail was not surprising. The latter feature could also be possibly induced by the presence of larger particles. The main characteristics of formed AgNPs in mixtures at pH 6.0 did not vary (Figures 3D,E). A special kind of corona surrounding the AgNPs in the presence of 26 μM Fe2+ at pH 6.0 was clearly visible in TEM images (Figure 3E). Energy-dispersive X-ray spectrometry (EDS) chemical mapping of the selected sample (26 μM Fe2+, pH 4.0) was also performed with the aim to identify the chemical nature of both ultrasmall and larger objects by focusing on silver and iron (SI Figure SI-6). As shown in SI Figure SI-6, both objects had silver origins (HAADF image and green colored silver map, SI Figure SI-6). Iron species were also clearly identified within the analyzed window in the form of irregular objects, coating both silver NPs and the grid surface (violet colored iron map, SI Figure SI-6). These iron species correspond to iron oxides, formed from soluble iron ions by precipitation during the sample drying. Antibacterial Studies. One of the important target sites by AgNPs is the bacterial cell wall that is significantly different in Gram-positive bacteria as compared with Gram-negative bacteria. Therefore, antibacterial tests were carried out on several strains of both types of bacteria (Table 3). Results showed that Gram-positive bacteria were more resistant (i.e., higher MIC values obtained) to AgNPs formed by SRHA and/ or SRHA-Fe2+ than those of Gram-negative bacteria. This may be as a result of the bacterial cell wall structure containing large amounts of peptidoglycan in the case of Gram-positive bacteria.

Figure 3. (A) UV−vis spectra of AgNPs at different pH in Ag+-SRHAFe2+ mixtures, TEM images: (B) no Fe2+, pH 4.0; (C) 26 μM Fe2+, pH 4.0; (D) no Fe2+, pH 6.0; and (E) 26 μM Fe2+, pH 6.0. ([Ag+] = 1 × 10−3 mol L−1, 40 mg L−1 SRFA).

SPR band located around ∼400 nm (Figure 3A). Particles with diameters exceeding 80 nm, having distinguishing shell of less contrast material could be seen (Figure 3B). These large-sized particles were likely stemming from SRHA and/or SRHAresidues in the mixture. The organic shell surrounding nanoparticles and their possible aggregates explain the discrepancy between average particle size determined by DLS and sizes observed in the TEM image. The negative zeta potential values of formed AgNPs indicate the formation of organic-coating around formed AgNPs. Values of zeta potentials of AgNPs were −23 and −30 mV without Fe2+ in the Ag+-SRHA mixture at pH 6.0 and 4.0, respectively. Values

Table 3. Minimum Inhibitory Concentrations (μg mL−1) of Silver Nanoparticles Formed in Fe2+-Ag+-SRHA Mixtures at Different pH AgNPs [Fe2+] 0 μM species

Ag+

Enterococcus faecalis CCM 4224 (g+) Staphylococcus aureus CCM 3953 (g+) Staphylococcus aureus (MRSA) (MRSA (g+) Staphylococcus epidermidis 1 (g+) Pseudomonas aeruginosa CCM 3955 (g−) Pseudomonas aeruginosa (g−) Klebsiella pneumoniae (ESBL) (g−)

21.2 10.6 21.2 10.6 10.6 5.3 10.6

26 μM

AgNPs [Fe2+] 98 μM

0 μM

pH 4.0 21.2 10.6 21.2 10.6 5.3 5.3 10.6 3232

21.2 10.6 21.2 10.6 5.3 5.3 10.6

26 μM

98 μM

pH 6.0 21.2 10.6 42.5 10.6 10.6 5.3 10.6

85 21.2 85 21.2 5.3 5.3 10.6

21.2 21.2 21.2 21.2 10.6 5.3 10.6

42.5 21.2 42.5 21.2 10.6 10.6 10.6

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SRFA mixtures.32,33 Basically, Fe3+ ions not only enhanced the formation of AgNPs, but also kept the stability of AgNPs in the Ag+-NOM-Fe3+ mixtures. The presence of high molecule weight components on the surface of AgNPs prevented aggregation of the particles. It seems that the surface coating on AgNPs did not alter in the presence of Fe3+ ions because zeta potential values were −18 mV and −23 mV with and without Fe3+, respectively, in the mixtures of Ag+-NOM. Furthermore, surfaces of AgNPs were sufficiently covered by coating of organic matter that did not allow dissolution of AgNPs in order to decrease their stability. Since transport of AgNPs is sensitive to surface properties, the potentially formed AgNPs from the interaction of Ag+ and NOM in the presence of Fe3+ ions would transport to a long distance from their points of origin. However, other solution parameters such as pH, ionic strength, and concentration of NOM would also determine the stability of AgNPs. Overall, many physicochemical factors such as size and concentration of particles, flow velocity, nature of NOM, functionalized surfaces, and chemistry of aqueous and solid phases would determine aggregation, dissolution, and transformation of AgNPs34 and hence their transport in the aquatic environment. The formed AgNPs in the presence and absence of Fe2+ showed toxic effects to different bacterial species. However, the toxicity differed depending on the type of species (i.e., Gram positive versus Gram negative) and hence AgNPs may result in significant different toxicity in the aquatic systems. The nature of organic matter coating on AgNPs may also influence the toxicity. Future studies may include other humic substances such as peat humic acids, sedimentary and soil humic acids, and other NOMs, alone or in combination, to learn the ecological importance of formed AgNPs in natural aquatic systems containing higher concentration of silver and iron ions as well as various kinds of organic matters (e.g., silver mine tailings, acid mine drainage, municipal wastewater effluents, and photographic wastewaters).53−56 The irradiation of light can greatly influence the formation and stability of AgNPs,32,57,58 therefore the effect of sunlight on the toxicity of AgNPs-coated with NOM may also be examined. The surface modification of the organic ligands-coated AgNPs due to irradiation of solar light would determine the toxicity to receptor organisms in the environment. Light also affects the redox transformation between Fe3+ and Fe2+ as free and organic complexes species, the mechanistic aspects on photochemical formation and subsequent fate and toxicity may therefore be examined in the future work. Finally, detailed proteomic and genotoxicity studies are needed to enhance understanding of toxic effects of naturally occurring organic matter-coated AgNPs on organisms.

In other words, the wall of Gram-negative bacteria is more complex (containing lipopolysaccharides, outer membrane with porines, inner membrane) and this property is likely to determine the higher susceptibility of Gram-negative bacteria (i.e., lower MIC values obtained) toward AgNPs in media of SRHA and/or SRHA-Fe2+ ions as was seen (Table 3). For instance, SRHA may act as an “active” coating of AgNPs improving, thus, the interaction with lipopolysaccharides of Gram-negative bacteria. This is significant when taking into account that Gram-negative bacteria are generally more resistant toward antibiotics.52 Comparing the MIC values of ionic silver (Ag+) and AgNPs formed by SRHA and/or SRHA-Fe2+ (Table 3), it could also be concluded that ionic silver is more toxic (i.e., lower MIC values) than the AgNPs toward Gram-positive bacteria. On the contrary, Gram-negative bacteria are more prone to the impact of both ionic silver as well as the AgNPs. Lower MIC values were obtained in comparison to those determined for Grampositive bacteria, however similar MIC values for Ag+ and AgNPs at pH 6.0 (Table 3). Furthermore, differences in MIC values in the case of Gram-positive bacteria were also apparent with change in pH. Values of MICs were clearly higher for AgNPs, formed at pH 6.0 than that at pH 4.0. This probably stems from the fact that the AgNPs at pH 6.0 were more optimal for the tested bacteria; approaching the physiological and natural water pH. Thus, the tested bacteria were more viable. Conversely, in more AgNPs-susceptible Gram-negative bacteria, the pH effect on MIC values was less pronounced (Table 3), which could be again related to the complex character of Gram-negative bacteria cell wall. The slight differences in MIC values, determined for the systems consisting of AgNPs generated by SRHA-Fe2+ (Table 3), can be caused by somewhat similar physicochemical characteristics of AgNPs such as average particle size (see Figure 1B), shapes of nanoparticles, and zeta potentials described above. Environmental Implications. The results showed that both forms of iron, Fe2+ and Fe3+, which exist in the environment, enhanced the formation of AgNPs in the interaction of Ag+ and NOM. The stability of the formed AgNPs was tested by monitoring their UV−vis spectra at room temperature in mixtures of Ag+-SRHA and Ag+-NLFA having Fe3+ ion (Figure 4). The SPR of AgNPs increased significantly during the period of seven months; suggesting high stability of the AgNPs in both mixtures. These results are similar to stability observed earlier in the mixtures of Ag+-SRHA and Ag+-

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ASSOCIATED CONTENT

S Supporting Information *

Chemicals; antimicrobial procedure; UV−vis spectra of control solutions at pH 6.0 (Figure SI-1); UV−vis absorption spectra of AgNPs in Ag+-SRHA solutions with and without Fe2+ (Figure SI-2); UV−vis spectra of control solutions at pH 6.0 (Figure SI3); UV−vis spectra of control solutions at pH 4.0 (Figure SI4); DLS determined size distributions based on intensity fluctuation of AgNPs in Ag+-SRHA solutions at different pH (Figure SI-5); and high-angle annular dark-field image of the representative sample (Figure SI-6). This material is available free of charge via the Internet at http://pubs.acs.org.

Figure 4. Aging of AgNPs at pH 6.0 and room temperature (24 °C). ([Ag+] = 1 × 10−3 mol L−1, [SRHA] = [NLFA] = 40 mg L−1 SRFA. [Fe3+] = 13 μM). 3233

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Chlamydomonas reinhardtii. Environ. Sci. Technol. 2008, 42, 8959− 8964. (12) Browning, L. M.; Lee, K. J.; Nallathamby, D.; Xu, X. H. N. Silver nanoparticles incite size- and dose-dependent developmentphenotypes and nanotoxicity in zebrafish embryos. Chem. Res. Toxicol. 2013, 26, 1503−1513. (13) Thio, B. J. R.; Montes, M. O.; Mahmoud, M. A.; Lee, D.; Zhou, D.; Keller, A. A. Mobility of capped silver nanoparticles under environmentally relevant conditions. Environ. Sci. Technol. 2012, 46, 6985−6991. (14) Holden, P. A.; Nisbet, R. M.; Lenihan, H. S.; Miller, R. J.; Cherr, G. N.; Schimel, J. P.; Gardea-Torresdey, J. Ecological nanotoxicology: Integrating nanomaterial hazard considerations across the subcellular, population, community, and ecosystems levels. Acc. Chem. Res. 2013, 46, 813−822. (15) Dobias, J.; Bernier-Latmani, R. Silver release from silver nanoparticles in natural waters. Environ. Sci. Technol. 2013, In Press. (16) Panacek, A.; Prucek, R.; Safarova, D.; Dittrich, M.; Richtrova, J.; Benickova, K.; Zboril, R.; Kvitek, L. Acute and chronic toxicity effects of silver nanoparticles (NPs) on Drosophila melanogaster. Environ. Sci. Technol. 2011, 45, 4974−4979. (17) Liu, J.; Legros, S.; Von Der Kammer, F.; Hofmann, T. Natural organic matter concentration and hydrochemistry influence aggregation kinetics of functionalized engineered nanoparticles. Environ. Sci. Technol. 2013, 47, 4113−4120. (18) Zhang, H.; Smith, J. A.; Oyanedel-Craver, V. The effect of natural water conditions on the anti-bacterial performance and stability of silver nanoparticles capped with different polymers. Water Res. 2012, 46, 691−699. (19) He, D.; Bligh, M. W.; Waite, T. D. Effects of aggregate structure on the dissolution kinetics of citrate-stabilized silver nanoparticles. Environ. Sci. Technol. 2013, 47, 9148−9156. (20) Furman, O.; Usenko, S.; Lau, B. L. T. Relative importance of the humic and fulvic fractions of natural organic matter in the aggregation and deposition of silver nanoparticles. Environ. Sci. Technol. 2013, 47, 1349−1356. (21) El Badawy, A. M.; Luxton, T. P.; Silva, R. G.; Scheckel, K. G.; Suidan, M. T.; Tolaymat, T. M. Impact of environmental conditions (pH, ionic strength, and electrolyte type) on the surface charge and aggregation of silver nanoparticles suspensions. Environ. Sci. Technol. 2010, 44, 1260−1266. (22) Li, X.; Lenhart, J. J. Aggregation and dissolution of silver nanoparticles in natural surface water. Environ. Sci. Technol. 2012, 46, 5378−5386. (23) Delay, M.; Frimmel, F. H. Nanoparticles in aquatic systems. Anal. Bioanal. Chem. 2012, 402, 583−592. (24) Gondikas, A. P.; Morris, A.; Reinsch, B. C.; Marinakos, S. M.; Lowry, G. V.; Hsu-Kim, H. Cysteine-induced modifications of zerovalent silver nanomaterials: Implications for particle surface chemistry, aggregation, dissolution, and silver speciation. Environ. Sci. Technol. 2012, 46, 7037−7045. (25) Huynh, K. A.; Chen, K. L. Aggregation kinetics of citrate and polyvinylpyrrolidone coated silver nanoparticles in monovalent and divalent electrolyte solutions. Environ. Sci. Technol. 2011, 45, 5564− 5571. (26) Piccapietra, F.; Sigg, L.; Behra, R. Colloidal stability of carbonate-coated silver nanoparticles in synthetic and natural freshwater. Environ. Sci. Technol. 2012, 46, 818−825. (27) Pokhrel, L. R.; Dubey, B.; Scheuerman, P. Impacts of select organic ligands on the colloidal stability, dissolution dynamics, and toxicity of silver nanoparticles. Environ. Sci. Technol. 2013, 47, 12877− 12885. (28) Chen, Z.; Porcher, C.; Campbell, P. G. C.; Fortin, C. Influence of humic acid on algal uptake and toxicity of ionic silver. Environ. Sci. Technol. 2013, 47, 8835−8842. (29) He, D.; Dorantes-Aranda, J. J.; Waite, T. D. Silver nanoparticlealgae interactions: Oxidative dissolution, reactive oxygen species generation and synergistic toxic effects. Environ. Sci. Technol. 2012, 46, 8731−8738.

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS Authors wish to acknowledge the support by project no. P108/ 11/P657 awarded by Grant Agency of the Czech Republic, the Operational Program Education for CompetitivenessEuropean Social Fund (CZ.1.07/2.3.00/20.0056), Operational Program Research and Development for Innovations − European Regional Development Fund (project CZ.1.05/ 2.1.00/03.0058 of the Ministry of Education, Youth and Sports of the Czech Republic) and grant project LF_2013_012. Authors wish to thank Klara Safarova for TEM measurements. Dr. Gardea-Torresdey acknowledges support from the National Science Foundation (NSF) and the Environmental Protection Agency (EPA) under Cooperative Agreement DBI-0830117. Any opinions, findings, and conclusions or recommendations expressed in this material are those of the authors and do not necessarily reflect the views of the NSF or the EPA. This work has not been subjected to EPA review and no official endorsement should be inferred. Authors thank anonymous reviewers for their comments, which improved the paper greatly.

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REFERENCES

(1) Gonzalez-Fuenzalida, R. A.; Moliner-Martinez, Y.; GonzalezBejar, M.; Molinas-Legua, C.; Verdu-Andres, J.; Perez-Prieto, J.; Campins-Falco, P. In situ colorimetric quantification of silver cations in the presence of silver nanoparticles. Anal. Chem. 2013, 85, 10013− 10016. (2) Sharma, V. K.; Siskova, K.; Zboril, R.; Gardea-Torresdey, J. Organic-coated silver nanoparticles in biological and environmental conditions: Fate, stability and toxicity. Adv. Colloid Interface Sci. 2014, 204, 15−34. (3) Chernousova, S.; Epple, M. Silver as antibacterial agent: Ion, nanoparticle, and metal. Angew. Chem., Int. Ed. 2013, 52, 1636−1653. (4) Dallas, P.; Sharma, V. K.; Zboril, R. Silver polymeric nanocomposites as advanced antimicrobial agents: Classification, synthetic paths, applications, and perspectives. Adv. Colloid Interface Sci. 2011, 166, 119−135. (5) Sharma, V. K.; Yngard, R. A.; Lin, Y. Silver nanoparticles: Green synthesis and their antimicrobial activities. Adv. Colloid Interface Sci. 2009, 145, 83−96. (6) Eckhardt, S.; Brunetto, P. S.; Gagnon, J.; Priebe, M.; Giese, B.; Fromm, K. M. Nanobio silver: Its interactions with peptides and bacteria, and its uses in medicine. Chem. Rev. 2013, 113, 4708−4754. (7) Auffan, M.; Rose, J.; Bottero, J.-.; Lowry, G. V.; Jolivet, J.-.; Wiesner, M. R. Towards a definition of inorganic nanoparticles from an environmental, health and safety perspective. Nat. Nanotechnol. 2009, 4, 634−641. (8) Quadros, M. E.; Pierson, R.; Tulve, N. S.; Willis, R.; Rogers, K.; Thomas, T. A.; Marr, L. C. Release of silver from nanotechnologybased consumer products for children. Environ. Sci. Technol. 2013, 47, 8894−8901. (9) Quigg, A.; Chin, W.-.; Chen, C.-.; Zhang, S.; Jiang, Y.; Miao, A.-.; Schwehr, K. A.; Xu, C.; Santschi, P. H. Direct and indirect toxic effects of engineered nanoparticles on algae: Role of natural organic matter. ACS Sust. Chem. Eng. 2013, 1, 686−702. (10) Levard, C.; JHotze, E. M.; Lowry, G. V.; Brown, J. G. E. Environmental transformations of silver nanoparticles: Impact on stability and toxicity. Environ. Sci. Technol. 2012, 46, 6900−6914. (11) Navarro, E.; Piccapietra, F.; Wagner, B.; Marconi, F.; Kaegi, R.; Odzak, N.; Sigg, L.; Behra, R. Toxicity of silver nanoparticles to 3234

dx.doi.org/10.1021/es405641r | Environ. Sci. Technol. 2014, 48, 3228−3235

Environmental Science & Technology

Article

(30) Levard, C.; Mitra, S.; Yang, T.; Jew, A. D.; Badireddy, A. R.; Lowry, G. V.; Brown, G. E. Effect of chloride on the dissolution rate of silver nanoparticles and toxicity to E. coli. Environ. Sci. Technol. 2013, 47, 5738−5745. (31) Xiu, Z. M.; Ma, J.; Alvarez, P. J. J. Differential effect of common ligands and molecular oxygen on antimicrobial activity of silver nanoparticles versus silver ions. Environ. Sci. Technol. 2011, 45, 9003− 9008. (32) Adegboyega, N. F.; Sharma, V. K.; Siskova, K.; Zbořil, R.; Sohn, M.; Banerjee, S. Interactions of Aqueous Ag+ with Fulvic Acids: Mechanisms of Silver Nanoparticle Formation and Investigation of Stability. Environ. Sci. Technol. 2013, 47, 757−764. (33) Akaighe, N.; MacCuspie, R. I.; Navarro, D. A.; Aga, D. S.; Banerjee, S.; Sohn, M.; Sharma, V. K. Humic acid-induced silver nanoparticle formation under environmentally relevant conditions. Environ. Sci. Technol. 2011, 45, 3895−3901. (34) Maurer, F.; Christl, I.; Hoffmann, M.; Kretzschmar, R. Reduction and Reoxidation of Humic Acid: Influence on Speciation of Cadmium and Silver. Environ. Sci. Technol. 2012, 46, 8808−8816. (35) Garg, S.; Rose, A. L.; Waite, T. D. Superoxide mediated reduction of organically complexed iron(III): Comparison of nondissociative and dissociative reduction pathways. Environ. Sci. Technol. 2007, 41, 3205−3212. (36) Fujii, M.; Rose, A. L.; Waite, T. D.; Omura, T. Oxygen and superoxide-mediated redox kinetics of iron complexed by humic substances in coastal seawater. Environ. Sci. Technol. 2010, 44, 9337− 9342. (37) Wang, H.; Wu, F.; Meng, W.; White, J. C.; Holden, P. A.; Xing, B. Engineered nanoparticles may induce genotoxicity. Environ. Sci. Technol. 2013, 47, 13212−13214. (38) Markova, Z.; Siskova, K. M.; Filip, J.; Cuda, J.; Kolar, M.; Safarova, K.; Medrik, I.; Zboril, R. Air stable magnetic bimetallic Fe-Ag nanoparticles for advanced antimicrobial treatment and phosphorous removal. Environ. Sci. Technol. 2013, 47, 5285−5293. (39) Henglein, A. Non-metallic silver clusters in aqueous solution: Stabilization and chemical reactions. Chem. Phys. Lett. 1989, 154, 473− 476. (40) Gentry, S. T.; Fredericks, S. J.; Krchnavek, R. Controlled particle growth of silver sols through the use of hydroquinone as a selective reducing agent. Langmuir 2009, 25, 2613−2621. (41) Rose, A. L.; Waite, T. D. Effect of dissolved natural organic matter on the kinetics of ferrous iron oxygenation in seawater. Environ. Sci. Technol. 2003, 37, 4877−4886. (42) Wilson, S. A.; Weber, J. H. A comparative study of numberaverage dissociation-corrected molecular weights of fulvic acids isolated from water and soil. Chem. Geol. 1977, 19, 285−293. (43) Struyk, Z.; Sposito, G. Redox properties of standard humic acids. Geoderma 2001, 102, 329−346. (44) Aeschbacher, M.; Graf, C.; Schwarzenbach, R. P.; Sander, M. Antioxidant properties of humic substances. Environ. Sci. Technol. 2012, 46, 4916−4925. (45) Stamplecoskie, K. G.; Scaiano, J. C. Silver as an example of the applications of photochemistry to the synthesis and uses of nanomaterials. Photochem. Photobiol. 2012, 88 (4), 762−768. (46) Jones, A. M.; Pham, A. N.; Collins, R. N.; Waite, T. D. Dissociation kinetics of Fe(III)- and Al(III)-natural organic matter complexes at pH 6.0 and 8.0 and 25 °C. Geochim. Cosmochim. Acta 2009, 73, 2875−2887. (47) Jones, A. M.; Garg, S.; He, D.; Pham, A. N.; Waite, T. D. Superoxide-mediated formation and charging of silver nanoparticles. Environ. Sci. Technol. 2011, 45, 1428−1434. (48) Rose, A. L.; Waite, T. D. Kinetic model for Fe(II) oxidation in seawater in the absence and presence of natural organic matter. Environ. Sci. Technol. 2002, 36, 433−444. (49) He, D.; Garg, S.; Waite, T. D. H2O2-mediated oxidation of zerovalent silver and resultant interactions among silver nanoparticles, silver ions, and reactive oxygen species. Langmuir 2012, 28, 10266− 10275.

(50) Skogerboe, R. K.; Wilson, S. A. Reduction of ionic species by fulvic acid. Anal. Chem. 1981, 53, 228−232. (51) Link, S.; El-Sayed, M. A. Spectral properties and relaxation dynamics of surface plasmon electronic oscillations in gold and silver nanodots and nanorods. J. Phys. Chem. B 1999, 103, 8410−8426. (52) 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: Nanotech. Biol. Med. 2007, 3, 168−171. (53) Liu, J.; Sonshine, D. A.; Shervani, S.; Hurt, R. H. Controlled release of biologically active silver from nanosilver surfaces. ACS Nano 2010, 4, 6903−6913. (54) Wen, L. S.; Santschi, P. H.; Gill, G. A.; Tang, D. Silver concentrations in Colorado, USA, watersheds using improved methodology. Environ. Toxicol. Chem. 2002, 21, 2040−2051. (55) Purcell, T. W.; Peters, J. J. Sources of silver in the environment. Environ. Toxicol. Chem. 1998, 17, 539−546. (56) Tappin, A. D.; Barriada, J. L.; Braungardt, C. B.; Evans, E. H.; Patey, M. D.; Achterberg, E. P. Dissolved silver in European estuarine and coastal waters. Water Res. 2010, 44, 4204−4216. (57) Hou, W. C.; Stuart, B.; Howes, R.; Zepp, R. G. Sunlight-driven reduction of silver ions by natural organic matter: Formation and transformation of silver nanoparticles. Environ. Sci. Technol. 2013, 47, 7713−7721. (58) Yin, Y.; Liu, J.; Jiang, G. Sunlight-induced reduction of ionic Ag and Au to metallic nanoparticles by dissolved organic matter. ACS Nano 2012, 6, 7910−7919.

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