Article pubs.acs.org/JPCC
Temperature Dependence and Mechanism of Chloride-Induced Aggregation of Silver Nanoparticles Karen I. Peterson,* Megan E. Lipnick, Luis A. Mejia, and David P. Pullman* Department of Chemistry and Biochemistry, San Diego State University, 5500 Campanile Drive, San Diego, California 92182-1030, United States S Supporting Information *
ABSTRACT: With the increasing occurrence of silver nanoparticles in commercial products and in the environment, it is important to understand the transformations that the nanoparticles undergo as a result of their interactions with other species. In this paper, we focus on interactions with the chloride ion, which is abundant in natural waters as well as biological systems. Chloride ion added to a solution of citrate-capped silver nanoparticles disturbs their stability by modifying the nanoparticle surface, enhancing dissolution of the particles, and increasing the ionic strength of the solution. Because of the surface modifications, aggregation occurs more rapidly than would be expected from the increase in ionic strength. This indicates that the nanoparticles are experiencing a decrease in surface charge. To elucidate the atomic-scale processes behind this behavior, we have studied the temperature dependence of the rate of decay of silver nanoparticles at a low NaCl concentration (10 mM), where dissolution but no aggregation occurs, and a higher concentration (40 mM), where aggregation is the dominant process. Particle dissolution was found to have a positive temperature dependence with an activation energy of 69 ± 6 kJ/mol. Conversely, the aggregation rate was inversely dependent on temperature but exhibited a lag time that increased with temperature. We develop an empirical model for chloride-induced aggregation, which is based on a time-dependent activation energy that arises from surface changes brought about by chloride reactions on the silver nanoparticle surface. The mathematical form of the model fits the data very well and provides insight into the molecular processes involved in aggregation in aqueous chloride solutions.
I. INTRODUCTION Many types of nanoparticles have the potential for novel applications in medical therapies.1 Nanoparticles combine the properties of solids and molecules because they are mobile in solution yet have a high surface area on which molecules can adsorb and react. Thus, they can transport materials as well as provide catalytic surfaces for in vivo reactions. In addition, they can be a direct source of an active agent. The most ubiquitous example is silver nanoparticles (AgNPs),2 whose activity as an antimicrobial agent is known to be due predominantly to the slow release of silver ion.3,4 In order to take full advantage of nanoparticles, information about their mobility, surface properties, and transformation in biological systems is needed. AgNPs are currently used as antimicrobial agents in wound dressings, water purification, clothing, and other commercial products.5,6 Therefore, it is important to understand the release rate of Ag+ in biochemical and other aqueous environments. There is an added issue of stability toward aggregation that can be important because the Ag+ release rate can depend on the aggregated state of the particles.7 Also, aggregation affects the mobility of the particles, particularly with respect to movement through membranes, and this can cause them to accumulate in areas where they are not useful.1 © XXXX American Chemical Society
Because AgNP dissolution and aggregation can influence each other, the interplay of these two processes must also be considered. Aggregation has been shown to decrease the rate of dissolution,8 and it is reasonable to expect that dissolution can change the nanoparticle surface in such a way as to affect the aggregation rate. Chloride ion is a major component of most natural aqueous environments, and it is known to promote both dissolution and aggregation of AgNPs. Therefore, we have focused on understanding the mechanisms of the interactions of chloride ion with AgNPs and the ensuing dissolution and aggregation that occurs. In order to stabilize nanoparticles in solution, adsorbed species called capping agents are generally included in the nanoparticle synthesis. In our work, we synthesize AgNPs in which citrate anions physically adsorb to the surface to provide a negatively charged Coulombic barrier to aggregation. Chloride ions subsequently added to the solution affect the aggregation rate of citrate-capped silver nanoparticles in two major ways: they increase the ionic strength, which reduces the Received: July 21, 2016 Revised: September 1, 2016
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DOI: 10.1021/acs.jpcc.6b07329 J. Phys. Chem. C XXXX, XXX, XXX−XXX
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Fisher Scientific. Sodium borohydride (>98%) was purchased from MPBiomedicals. The 200-mesh copper grids coated with Formvar, used for the transmission electron microscope (TEM) images, were purchased from Ted Pella, Inc. These were then coated with 0.1% w/v aqueous polylysine solution purchased from Sigma-Aldrich. All chemicals were used without further purification. Deionized water of 5 MΩ or higher was used throughout. Synthesis of Silver Nanoparticles. Solutions of nanoparticles were synthesized as before,10 by dropwise addition of 0.45 mM NaBH4 to aqueous mixtures of silver nitrate and trisodium citrate at room temperature (typically, 21 °C); the particles were allowed to age for at least 4 days before sodium chloride was added. The concentrations of reagents in the working solutions, if unreacted, would be [Ag+] = 0.113 mM, [Na3C6H5O7] = 0.113 mM, and [NaBH4] = 0.028 mM. Ag+ will be transformed during the reaction, but we can state that the number of silver atoms in some form or another starts at [Agnet] = 0.113 mM. On average, [Agnet] = n[Agn]0 where n is the average number of silver atoms per nanoparticle. The average diameter of the nanoparticles is 5.3 ± 1.8 nm, as determined previously by TEM measurements.10 From this value, and the concentration of silver in the solution, the nanoparticle concentration is estimated to be about 2.5 × 10−8 M. The surface plasmon resonance (SPR) peak of the silver nanoparticles in the solutions described above is located at 392 ± 2 nm with an absorbance of 1.3 ± 0.1. In the wavenumber spectrum, for a nanoparticle concentration of 2.5 × 10−8 M, the average area is (4.5 ± 0.3) × 103 AU cm−1, where AU refers to absorbance units. The area under the peak is a better representation of the nanoparticle concentration than is the peak absorbance,12 and we have found that the area is quite consistent from sample to sample for unreacted nanoparticles. Instrumentation. All UV−vis absorption spectra were recorded using either a Jasco V-670 spectrophotometer or a Hewlett-Packard diode array spectrophotometer (HP-8452A). Dynamic light scattering measurements were made with a Malvern Instruments, Inc., Zetasizer Model Nano-ZS. Nanoparticle images were collected using a FEI Technai 12 transmission electron microscope. Kinetics Measurements. The experimental method for monitoring the decay of the AgNP monomers is similar to that in previous work.10 The major difference is that the temperature was controlled by means of an orbital mixing chilling/heating dry bath. After addition of sodium halide solutions to the nanoparticle solutions, aliquots were taken from the reaction mixture at appropriate time intervals and the UV−vis spectrum was measured, usually from 200 to 800 nm. A new aliquot was used for each measurement to avoid significant light-induced perturbation of the AgNPs. Also, the cuvette was cleaned after each measurement to avoid the development of an interfering film of silver. These effects of light on the rate of NP decomposition in salt solutions also required us to keep the solution vials in the dark as much as possible. The silver nanoparticle SPR peak was used to monitor both aggregation and dissolution of the particles. When aggregation takes place, the peak absorbance near 392 nm decreases because of the loss of AgNP monomers, and dissolution causes the peak to decrease by loss of silver atoms in the particles. The area under the absorption spectrum with wavenumber as the abscissa is proportional to the AgNP concentration:
thickness of the electrical double layer around the particles, thereby lowering the barrier to aggregation, and they replace physisorbed citrate anions with a chemisorbed layer, which appears to reduce the surface charge, further lowering the repulsive barrier. In addition, the chloride ions promote dissolution of the nanoparticles through enhanced oxidation of the surface atoms to form AgCl. Thus, aggregation and dissolution occur simultaneously, and interpretation of silver nanoparticle aggregation behavior must take this into account. Some previous work has been done to explore the connection between chloride-induced dissolution and aggregation. He et al.7 looked at the effect of chloride-induced aggregation on the dissolution of citrate-capped nanoparticles, and they have found that the dissolution rate depends on both the extent of aggregation and the type (diffusion limited vs reaction limited). They have shown that the determining factor is the amount of surface area available. Li et al.9 observed dissolution and consequent formation of AgCl occurring during chlorideinduced aggregation of uncapped nanoparticles. Although the appearance of the aggregates is clearly affected by the AgCl formation, the overall aggregation rate did not appear to be greatly affected. Even so, in their experiments, the rates of the two processes are likely competitive perhaps due to the lack of protection from citrate on the surface. Dynamic light scattering (DLS) was used in these and many other aggregation studies, but this method has some limitation in the case of chloride solutions because it does not give information about the transformation of silver to silver chloride due to the lack of chemical specificity of the method. Monitoring the nanoparticles by UV−vis absorption spectroscopy is a better gauge of the amount of silver lost in the processes. In prior studies,10 we found that citrate-capped 5 nm silver nanoparticles respond to chloride ion in a manner which depends distinctly on the NaCl concentration. Below about 30 mM NaCl, particle dissolution is the primary loss mechanism for silver in the silver nanoparticles. Up to about 2 mM NaCl the dissolution rate increases with NaCl concentration, but then it levels off and decreases to a very slow rate that remains constant up to about 30 mM NaCl. We attributed this behavior to the formation of a AgCl layer that becomes passivating as it thickens. Therefore, there is a concentration region between about 5 and 25 mM NaCl where dissolution is very slow and no aggregation occurs. We find in the work presented here that this is true throughout the temperature range studies, 5−50 °C. We also found previously that aggregation takes over as the dominant particle decay process above 30 mM NaCl. This is consistent with the DLVO theoretical expectation, but preliminary experiments11 have indicated that the temperature dependence of the decay is negativethe rate of aggregation decreases with temperature. This argues against a bimolecular type of rate-determining step. Because of this anomaly, and the desire to find an overall mechanism for what we have observed, we have carried out a more careful temperature study of both the aggregation and dissolution aspects of the system. In this paper, we will show that NaCl-induced aggregation of silver nanoparticles is strongly influenced by modifications of the surface. We propose a mechanism that takes this into account and also explain the temperature dependence of the aggregation process.
II. EXPERIMENTAL METHOD Materials. Silver nitrate (99.9%), sodium chloride (99.8%), and trisodium citrate dihydrate (>99.5%) were purchased from B
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∫band A(ν)̃ dν ̃ = [AgNP]b ∫band ε(ν)̃ dν ̃
continues to narrow while the height decreases. The time evolution of the peak shape is demonstrated in Figure 1 with
(1)
where A(ν̃) is the absorbance at a particular wavenumber, b is the light path length, and ∫ bandε(ν̃) dν̃ is the integrated absorption coefficient. To minimize the contribution of aggregated particles to the monomer peak, the area under a baseline drawn from the absorbance at 20 000 cm−1 (500 nm; this is an arbitrary cutoff set to minimize the contributions from aggregates) to that around 31 450 cm−1 (318 nm; this is the minimum in the SPR band on the high-energy side, which shifts slightly as the reaction progresses) was subtracted from the peak area. Figure 1. Time dependence of the SPR peak with 10 mM NaCl: (a) the absorbance of the peak maximum and (b) the fwhm (full width at half-maximum).
III. RESULTS AND DISCUSSION In order to distinguish between NaCl-induced surface dissolution and aggregation, we rely on the changes in the peak shape and the kinetic order of the decay rate. Our confidence in this is based on our studies of the time dependence of the SPR peak at two different concentrations of NaCl: (1) 10 mM, for which dissolution occurs but is well below the onset of aggregation, and (2) 40 mM NaCl, for which aggregation is clearly the dominant process at temperatures below about 30 °C. In the 10 mM NaCl experiments, we consistently observe a decay rate of the SPR peak that is close to first order in conjunction with a narrowing of the SPR peak. In the 40 mM NaCl experiments, the decay rate is much faster and has a higher-order dependence on the monomer concentration (SPR peak area). The time dependence of the absorption spectrum is also markedly differentafter the initial peak narrowing which occurs when NaCl is first added, the full width at half-maximum (fwhm) does not change significantly in time, and a relatively flat absorbance at higher wavelengths develops. We also conducted TEM measurements of the effects of NaCl addition. After NaCl was added (40 mM total), TEM samples were prepared from the solution after 3, 8, and 17 min of reaction time. The preparation time was about 1 min for each sample. The TEM images are shown in Figure S1 of the Supporting Information. In these images, little aggregation is apparent after 3 and even 8 min, but after 17 min, aggregation is clearly evident. This agrees well with the UV−vis spectra of the same process; the SPR peak does not change significantly until after 10 min (see Figure S2). For comparison, a TEM image of the 10 mM sample, in which aggregation is not apparent in the UV−vis spectrum, is also shown in Figure S1. Kinetics of Particle Dissolution. Chloride ion promotes oxidation of silver through formation of insoluble AgCl that covers the surface.13,14 This process can change the surface charge of silver nanoparticles that have been stabilized by negatively charged citrate ions adsorbed on the surface. Since the onset of aggregation occurs at a much lower ionic strength for AgNPs in aqueous NaCl than in aqueous solutions of other ionic compounds, such a NaF (which is not expected to alter the nanoparticle surface significantly), the surface charge is apparently being lowered.10,15 Although a simple replacement of citrate ions with Cl− can explain a reduction in charge, there may be more complicated surface effects that come into play. Therefore, characterization of AgNP dissolution at concentrations below the onset of aggregation will be helpful in understanding aggregation at the higher concentrations. When NaCl is first added to the silver nanoparticle solution to reach a 10 mM concentration, the SPR peak immediately narrows and increases in height, and then, more slowly, it
graphs of the peak absorbance and fwhm vs time. This behavior of the peak shape is consistent throughout the temperature range that we studied, 5−50 °C. The initial abrupt change in the peak shape suggests a decrease in the damping of the surface plasmon resonance. This effect has been observed before by others, and it has been attributed to the formation of a film of AgCl, which replaces oxides on the nanoparticle surface.16 The reduced damping is due to the lower absorptivity of AgCl around 400 nm compared to Ag2O. Further support for immediate formation of AgCl is given by the appearance of a peak near 260 nm (see Figure 2), which can be attributed to
Figure 2. Time dependence of the SPR peak for 10 mM NaCl. Note the change in the broad peak near 260 nm.
AgCl.17−20 AgCl nanoparticles in microemulsions exhibit an absorption peak in this region, but there is some variation in the position. While Husein et al.17 observe the peak around 250 nm, which is close to the location expected for the direct band gap of AgCl, Wu et al.19 observe the peak near 290 nm, which is closer to the lower energy edge of the gap. The reason for this difference is not clear, but it may be due to differences in size.17 AgCl particles formed by laser ablation of Ag in aqueous NaCl solutions exhibit a broader peak in the region between 250 and 300 nm, similar to what we observe. Interestingly, the peak we observe shifts to the red and broadens in time, and this coincides with the continuing sharpening of the SPR peak over time (Figure 1). More examples of spectra showing how this region evolves are given in the Supporting Information (Figure S2); at lower temperature and higher NaCl concentration, the peak at 250 nm is generally sharper and does not shift as much. Although we attribute these observations to an immediate formation of AgCl on the AgNP surface followed by slower structural changes, we cannot rule out the formation of AgCl particles being a contributing factor. C
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would be faster and, perhaps, more significantly affect the aggregation process. Therefore, we have studied the temperature dependence of samples with 10 mM NaCl. An example of the decay rate at T = 30 °C is shown in Figure 4. The first-order
To further investigate the continuous, but slow, decrease in the peak width, dynamic light scattering (DLS) was used to follow the AgNP size in time. We expected to find a slow change in size as the nanoparticles dissolved or perhaps very little change if the silver transformed into silver chloride on the particle surface. Instead, we found that after only 10−15 min, large particles developed, overwhelming the DLS signal from the much smaller nanoparticles, and these new particles slowly grew in time (see Figure S3 in the Supporting Information as well as the first few time traces in Figure 3). Little change was
Figure 4. Time dependence of the area under the SPR peak for 10 mM NaCl at 37 °C.
rate constant is obtained from the slope of the linear portion of the curve. Rate constants obtained in the temperature range from 5 to 50 °C are shown in the Arrhenius plot given in Figure 5. The value of the rate constant at 25 °C, (6 ± 4) × 10−6 s−1, is Figure 3. Volume % DLS spectra of AgNP; [Ag] = 0.56 mM. Spectra were taken in the following time sequence: (a) no NaCl or NH3 added; (b) 7 min after addition of NaCl ([NaCl] = 1.0); (c) 11 min after addition of NaCl; (d) 30 s after addition of NH3 ([NH3] = 7.5 mM); (e) 3 h after addition of NH3; (f) 3.5 h after addition of NH3; (g) 50 h after addition of NH3. NaCl and NH3 aliquots increased the solution volume by less than 1%.
observed in UV−vis spectra of the same solution taken concurrently with the DLS measurements, showing that the concentration of the larger particles was relatively small. We suspected they were AgCl particles forming separately from the nanoparticles, and this was tested by adding a small amount of ammonia to the solution (final concentration 7.5 mM). Ammonia dissolves AgCl preferentially over silver metal, forming the silver ammonia complex Ag(NH3)2+(aq).21 We found that the large particles slowly disappeared, strongly suggesting that they are composed of AgCl rather than silver atoms (see Figure 3). Thus, it appears that as the silver nanoparticle surface oxidizes in 10 mM NaCl, AgCl may be produced on the surface, but it also dissolves and recrystallizes away from the silver particles, growing slowly in time. Therefore, the surface of the nanoparticles in a chloride environment may not involve a continuous buildup of a AgCl film. Evidence for precipitation of AgCl away from the silver surface has been reported previously in electrochemical oxidation of silver surfaces.13 The above observations suggest that the interaction of chloride with the AgNP surface is changing over a relatively long period of time after addition of the NaCl. Since this could be an important component of the aggregation process, we studied the dissolution kinetics in more detail. In previous work, we found that the dissolution rate follows 2/3-order kinetics, but to a good approximation, the data could be fit to a simple first-order rate law.10 The rate is very slow with a halflife on the order of days over the concentration range of NaCl up to about 25 mM. The temperature dependence was not measured, but we expected that at higher temperature the rate
Figure 5. Temperature dependence of the rate of AgNP dissolution by oxidative decomposition of the surface; [NaCl] = 10 mM.
comparable to that found in a number of other studies.22−24 The activation energy derived from the slope of the plot is 69 ± 6 kJ/mol. Dissolution in distilled water has also been found to increase with temperature,22 with an associated activation energy of 77 kJ/mol that is similar to our value. As we will see in the next section, the 30-fold increase in the dissolution rate in the temperature range studied results in a nanoparticle decay rate that becomes comparable to the aggregation rate at the higher temperatures, preventing a clear separation between the two processes. Kinetics of Aggregation. The onset of nanoparticle aggregation occurs at about 30 mM NaCl, and the rate increases rapidly with NaCl concentration. Above 40 mM, aggregation becomes too fast to measure with our present techniques, so we used 40 mM NaCl to study the aggregation process. Aggregation cannot be completely separated from the effects of dissolution, but we found in our previous work10 that at room temperature the dissolution rate does not increase with NaCl concentration from 10 mM until the onset of aggregation at about 25 mM. This was attributed to passivation of the surface, and although aggregation prevents direct observation, it is reasonable to expect that the dissolution rate will continue to remain steady at higher NaCl concentrations. Since the D
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required for aggregation; but it may also react with the surface in such a way as to initially inhibit aggregation and thus cause the observed lag time, with the length of time depending on the rate of the surface reaction. To test whether the lag time relates specifically to the presence of NaCl, we conducted experiments using Na2SO4 to induce aggregation. Like fluoride ion, sulfate ion does not replace citrate on the silver nanoparticle surface (and, thus, does not lower the surface charge), so the ionic strength needed to induce aggregation is much higher than it is for NaCl; the onset occurs at about 100 mM rather than 30 mM. When Na2SO4 was added to the nanoparticle solution, aggregation occurred, but no lag time in the rate was observed (see Supporting Information, Figure S4), thus supporting the hypothesis that the lag time is caused by chloride interactions with the nanoparticle surface. Further evidence was provided by experiments in which NaCl was first added to reach a concentration of about 10 mM, thereby coating the surface with chloride ion but not inducing aggregation and then, a few hours later, adding enough Na2SO4 to induce aggregation (∼18 mM). In this case, a lag time was observed, and the overall ionic strength required for the onset of aggregation was similar to the NaCl experiment. After the apparent lag in the decay, the SPR peak decreases at a rate which is approximately second or third order with respect to the AgNP concentration (it is clearly not first order, thus differentiating it from AgNP dissolution). An example is shown in Figure 7b where 1/[AgNP]2 vs time is plotted; the latter part of the curve appears to be linear showing third-order behavior. The decay rate due to aggregation decreases with temperature, in contrast to the positive temperature dependence of the dissolution rate. This is shown in Figure 8, where the reciprocal
aggregation rate at 40 mM NaCl is over 2 orders of magnitude faster than the dissolution rate, we can assume that aggregation is the dominant decay process, and we expect that the peak shape will reflect this. Indeed, we find that a broad, relatively flat, absorption indicative of large aggregates develops at higher wavelengths and the decay rate is second order or higher. An example of the changing spectrum is shown in Figure 6. Note,
Figure 6. Time dependence of the SPR peak after NaCl is added at 25 °C; [NaCl] = 40 mM.
also, that there is a lag time in the decay of the peak; the change in the first 600 s is less than in the next 400 s. This lag time, which is more apparent when the AgNP concentration is plotted as a function of time as in Figure 7, is consistently
Figure 8. Temperature dependence of the rate of decay of the SPR peak for [NaCl] = 10 mM, where dissolution dominates, and [NaCl] = 40 mM, where aggregation dominates. Figure 7. (a) Decay of the SPR peak at 25 °C; [NaCl] = 40 mM NaCl. The data are fitted by eq 6 with a = −15.22 kJ/mol, ΔE = 27.02 kJ/ mol, and τ = 498 s. (b) Time dependence of 1/[AgNP]2 for 40 mM NaCl at 25 °C. The [AgNP] is proportional to the area under the SPR: [AgNP] = (2.5 × 10−8 M)/(4500 cm−1)∫ bandA(ν̃) dν̃.
of the half-life for the decrease of the 400 nm SPR peak area, 1/ τ1/2, is used in place of the rate constant in order to avoid committing to a particular reaction order. Thus, a negative effective activation energy is apparent for the aggregation process. Although the rate law for aggregation could be second order with respect to monomer concentration, the negative temperature dependence of the AgNP monomer decay indicates that the rate-limiting step for the aggregation process is not a simple monomer−monomer association. A mechanism that is consistent with the temperature dependence is one in which the monomer particles are involved in a fast associative equilibrium with dimers, and this equilibrium would be expected to shift toward dimer production at colder temper-
observed and has been found to increase with temperature. We have also observed that at higher temperatures (above about 40 °C) the SPR peak appears to narrow in time, and the absorbance at higher wavelengths does not increase as much, indicating that the contribution of dissolution to the decay rate is becoming increasingly important. Thus, the chloride ion appears to have a complex effect on AgNP aggregation: it lowers the surface charge, thereby reducing the ionic strength E
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temperature-dependent viscosity of water. The goodness of the fit can be seen in the run, shown in Figure 7a, for data taken at 25 °C. The averages of the fitted parameters for each temperature are shown in Figure 9. The values of a and ΔE increase mildly
atures. Then, if the addition of a third particle is rate-limiting and dependent on the dimer concentrations, an effective negative activation energy could arise. These proposed initial steps toward aggregation are summarized here: step 1: 2AgNP ↔ (AgNP)2 K1 = e−(ΔH1− T ΔS1)/ RT
(2)
step 2: (AgNP)2 + AgNP → (AgNP)3 k 2 = A freq e−Ea / RT (3)
According to this mechanism, the dimer is in fast equilibrium with the monomers so that [(AgNP)2] = K1[AgNP]2. Then, with the second step rate-limiting, the rate of decay of monomers is d[AgNP] = −k 2K1[AgNP]3 dt
(4)
This rate law agrees with the linear part of the curve in Figure 7, but the observed lag time in the AgNP decay curve indicates that some additional time-dependent process must be included in the mechanism. The experiments at low [NaCl], described in the previous section, suggest that a relatively slow organization of the surface is occurring, and it is possible that this results in a changing barrier for the addition of a third particle, i.e., a timedependent rate coefficient. If so, the time dependence could be incorporated into the rate equation by considering an evolving dimer structure: (AgNP)2 → (AgNP)2*. (AgNP)2 would be a weakly bound dimer that is in equilibrium with the monomer and has a low probability for further aggregation. (AgNP)2* is perhaps a more strongly bound dimer that is not easily dissociated and has a lower barrier to addition of a third monomer. In the proposed mechanism, the second step (eq 3) is very slow until the concentration of (AgNP)2* becomes significant. We attempt to model the effect of this transformation on the rate by introducing a time-dependent activation energy that changes from a high value, E, at t = 0 when (AgNP)2 is dominant, to a lower value, E*, as (AgNP)2* is formed:
Figure 9. Temperature dependence of the parameters obtained by fitting the decay of the SPR peak to eq 3: a = ΔH1 − TΔS1 + E*; ΔE = E − E*; τ is the characteristic time for the transformation from (AgNP)2 to (AgNP)2*.
(AgNP)2 → → → (AgNP)2 * k 2 = A freq e
−E / RT
→ → → k 2* = A freq e
−E*/ RT
with temperature. From the definition of a, given in eq 6, the positive slope of the plot of a vs T (Figure 9) implies a negative value for ΔS1, which is reasonable for an association reaction. The intercept, which is equal to ΔH1 + E*, is negative, and this can be explained by a negative value of ΔH1 whose magnitude is greater than E*. The characteristic time, τ, increases exponentially with temperature, so it is plotted as the natural logarithm in Figure 9. Interestingly, this relationship correlates with the increase in the rate of dissolution of the particles in 10 mM NaCl. This mathematical analysis was based on a simple mechanism where the rate constant of the second step was taken to be time dependent due to a changing activation energy. One possible explanation is that there is a change in the form of the dimer, perhaps caused by a change in the nanparticle surface morphology. Of course, the monomers would also be undergoing surface changes, and this would affect the first step in the process−monomer association. Even so, the effect on the second step may be more apparent because the association of a third particle to a dimer is slower than the association of two monomers.26 DLVO theory predicts a higher barrier for side-on interaction between a monomer and dimer, and this effectively leads to an orientation parameter in the rate constant that is less than one.
(5)
A functional form which fits these criteria and also fits the data is Ea = E* + ΔEe−t/τ with ΔE = E − E* and τ interpreted as the time constant characteristic of the transformation. Thus, Ea = E at t = 0 and Ea → E* as t → ∞. Writing k2 in the Arrhenius form with the time-dependent activation energy, and also writing the equilibrium constant for step 1 in terms of the enthalpy and entropy change (eq 2), eq 4 can be written as −t/ τ d[AgNP] = −A freq e−(a +ΔEe )/ RT [AgNP]3 dt
a = ΔH1 − T ΔS1 + E*
(6)
Best-fit values of the parameters a, ΔE, and τ at temperatures between 280 and 325 K were obtained by least-squares fitting of [AgNP], the time-dependent AgNP monomer concentration, to the observed decay of [AgNP]. During each iteration of the fitting process, numerical integration of eq 6 was carried out with the fifth-order Runge−Kutta−Fehlberg algorithm25 in order to obtain [AgNP] as a function of time. To reduce the number of parameters, the frequency factor was set equal to the diffusion-controlled rate constant, 8RT/3μ, where μ is the F
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IV. CONCLUSIONS NaCl modifies the surfaces of silver nanoparticles, most likely through the formation of a AgCl layer. This is shown in our work by a narrowing of the SPR peak, slow dissolution of the particles, and a lowering of the ionic strength required for aggregation to occur. Other research supports this.9,13,14,27 Although the observable effects of chloride ion adsorption become evident very quickly (in less than 1 min), there appears to be a reorganization of the surface that occurs over a longer time period; the peak shape continues to evolve, albeit at a slower rate. With 40 mM NaCl, we assume that surface modifications and dissolution processes are occurring, but the monomer loss due to aggregation is so rapid that it dominates the decay of the SPR peak. Even so, aggregation is delayed by a few minutes indicating that surface changes or reaction may be affecting the onset of aggregation. Further evidence for this connection is found in the temperature dependencesboth the dissolution rate and the aggregation lag time increase with temperature, whereas the overall aggregation rate decreases with temperature. We have developed an empirical model that fits the 40 mM NaCl data in the temperature range studied. The equation is based on a mechanism in which the first step is a preequilibrium between monomer and dimer particles. This differs from the assumptions of kinetic analyses of dynamic light scattering experiments, which generally do not consider reversibility in the individual steps of the aggregation mechanism.28,29 The second step, which adds a third particle, is rate-limiting with an activation energy that is time dependent. Presumably, the time dependence is related to the changes which occur on the particle surface after chloride ion is adsorbed, but we can only conjecture as to the details. It appears that there is an initial fast adsorption of Cl− that only changes the surface charge a small amount. Then, the Cl− slowly oxidizes the surface while lowering the surface charge more significantly as it forms AgCl. The consequence of the lower surface charge is an increase in the aggregation rate through a lower associative activation energy. It is also possible that the AgCl surface layer promotes coalescence of the dimer, causing the first association step to become irreversible. TEM pictures of NaCl-induced aggregates, which show fused structures, supports this.13,27 An alternate interpretation, proposed recently by Lodeiro et al.30 for certain of their systems, is that production of AgCln1−n species near the surface can act as a protective layer toward aggregation. Although we cannot rule this out entirely, the concentration of NaCl that we are using probably does not support high concentrations of AgCln1−n. We are now investigating the surface modification caused by chloride ion in more depth through surface-enhanced Raman spectroscopy, which may elucidate changes in the Ag− Cl interactions on the nanoparticle surface.
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Figure S4: comparison of aggregation rate of AgNPs in NaCl/Na2SO4 solutions (PDF)
AUTHOR INFORMATION
Corresponding Authors
*(K.P.) E-mail
[email protected]; Tel 619-594-4507. *(D.P.) E-mail
[email protected]; Tel 619-594-5573. Present Addresses
M.A.L.: PolyPeptide Laboratories, San Diego, 9395 Cabot Drive, San Diego, CA 92126. L.A.M.: Health Advances USA, Inc., 696 Naples St., Chula Vista, CA 91911. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS We gratefully acknowledge Pure BioScience (El Cajon, CA) for financial support of this work. We thank Dr. Steven Barlow for assistance in recording TEM images of the nanoparticles. The TEM work was supported by NSF Grant DBI-030829.
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ABBREVIATIONS AgNP, silver nanoparticle; DLS, dynamic light scattering; SPR, surface plasmon resonance; DLVO, Derjaguin, Landau, Verwey, and Overbeek; TEM, transmission electron microscope.
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REFERENCES
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ASSOCIATED CONTENT
S Supporting Information *
The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpcc.6b07329. Figure S1: TEM images of AgNP with NaCl added; Figure S2: the time evolution of UV−vis spectra of AgNP solutions after NaCl was added; Figure S3: volume % DLS spectra of silver nanoparticle solutions; time evolution of the particle size when NaCl is added; G
DOI: 10.1021/acs.jpcc.6b07329 J. Phys. Chem. C XXXX, XXX, XXX−XXX
Article
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DOI: 10.1021/acs.jpcc.6b07329 J. Phys. Chem. C XXXX, XXX, XXX−XXX
