Kinetics of Halide-Induced Decomposition and Aggregation of Silver

Article pubs.acs.org/JPCC

Kinetics of Halide-Induced Decomposition and Aggregation of Silver Nanoparticles M. Gabriela Espinoza, Mallory L. Hinks, Alexandra M. Mendoza, David P. Pullman, and Karen I. Peterson* Department of Chemistry and Biochemistry, San Diego State University, 5500 Campanile Drive, San Diego, California 92182-1030, United States ABSTRACT: We present a kinetic study of the effects of NaF, NaCl, NaBr, and NaI on aqueous solutions of 5 nm silver nanoparticles. There are distinct differences between these halides, which we attribute to two competing, halide-induced processes: oxidative decomposition of the nanoparticle surfaces and aggregation of the nanoparticles. NaF essentially does not react with the nanoparticle surface, but at concentrations above about 75 mM induces aggregation, albeit erratically. NaCl reacts oxidatively at concentrations below 27 mM, but the reaction is very slow because of surface passivation. The distinct, and lower, onset of aggregation at 27 mM is explained by chloride ion forming a uniform layer, which lowers the charge on the nanoparticle surface to about 2/3 of its original value. Addition of NaI or NaBr is very different; the rate of oxidative decomposition is orders of magnitude faster than that of NaCl such that the onset of aggregation is less apparent. surface directly.9 Interestingly, the effect of chloride ion on the silver nanoparticle surface is also important in the synthesis of particular types of silver nanoparticles. Xia et al.10 have shown that chloride ions promote the growth of silver nanoparticles with specific sizes and shapes through etching of “seed particles”. Thus, we see that chloride ions exhibit two different rolesas a promoter of nanoparticle aggregation and as a chemical reactant in nanoparticle surface modification. The above two examples showcase chloride ion chemistry as an integral part of processes involving silver nanoparticles, but they also point to the disruptive potential of halide ions. If silver nanoparticles are to be considered for applications in biological systems, a thorough understanding of their interactions with halide ions is needed. In this paper, we will explore these interactions by studying the kinetics of the halide-induced decay of the surface plasmon resonance peak that is characteristic of silver nanoparticles. Specifically, we will determine the rate law and rate coefficients for the decay. This information will be used to better understand the mechanisms underlying silver nanoparticle decomposition and aggregation and to gain insight into the molecular details of the silver surface.

I. INTRODUCTION In the last few decades, an extensive and growing body of research has focused on the physical and chemical properties and potential applications of metal nanoparticles. Silver has held a prime place in this work because of the strong surface plasmon resonance absorption located in a convenient part of the spectrum (∼400 nm), its relatively unreactive nature, and the ease with which nanoparticles can be synthesized. These properties have encouraged the application of silver nanoparticles in areas such as solar cell enhancement,1,2 biosensing,3 and surface-enhanced Raman spectroscopy (SERS).4,5 Furthermore, silver metal has long been known to be a potent antimicrobial agent,6 and this has led to substantial interest in using silver nanoparticles in medical applications.7 The practical application of silver nanoparticles requires an understanding of how they interact with their environment. In some cases, this interaction is directly related to their activity and function. For example, SERS signals obtained with silver nanoparticles are found to be greatly enhanced when the particles are allowed to aggregate by the addition of soluble salts. The surfaces of the nanoparticles are generally surrounded by capping agents (e.g., citrate ions) producing a charged layer that serves as an electrostatic barrier to aggregation. Increasing the ionic strength of the solution, by adding ionic species, decreases the effective electric field between the charged nanoparticles and, consequently, lowers this barrier. Although simple aggregation is largely responsible for SERS enhancement, additional enhancement factors seem to come into play for chlorides.8 Considerable research has focused on the effect of salts, and it is apparent that chloride ions not only change the ionic strength of the solution but also affect the nanoparticle © 2012 American Chemical Society

II. EXPERIMENTAL METHOD A. Materials. Silver nitrate (99.9%), sodium fluoride (99.5%), sodium bromide (99.5%), sodium chloride (99.8%), and trisodium citrate dihydrate (>99.5%) were purchased from Received: February 6, 2012 Revised: March 14, 2012 Published: March 27, 2012 8305

dx.doi.org/10.1021/jp3011926 | J. Phys. Chem. C 2012, 116, 8305−8313

The Journal of Physical Chemistry C

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a Hewlett−Packard Diode Array Spectrophotometer (HP8452A). Nanoparticle images were collected using a FEI Technai 12 transmission electron microscope. D. Kinetics Measurements. Reactions were initiated by adding sodium halide solutions to the nanoparticle solutions while mixing well. The decay of the surface plasmon resonance peak was monitored as a function of time by removing aliquots from the reaction mixture and measuring the UV−visible spectrum from 200 to 800 nm. A fresh aliquot in a cleaned cuvette was used for each measurement for two reasons. First, the light beam from the spectrometer caused a film of silver to slowly develop on the cuvette, interfering with the apparent solution absorption. Second, the spectrometer light source appeared to perturb the reaction. This was tested by measuring the decay of the plasmon resonance peak with multiple spectra taken of a single aliquot and comparing this with the measured decay using separate aliquots. The decay rate was found to be significantly faster in the former case. This is possibly due to electrons being more readily removed from the silver surface when light is absorbed. The need for using fresh aliquots increased the time between spectral measurements, thus limiting the experiment to reactions with half-lives of more than about 2−3 min.

Fisher Scientific. Sodium borohydride (>98%) was purchased from MPBiomedicals, and sodium iodide (>99%) from Acro̅s Organics. All chemicals were used without further purification. Deionized water of 5 MΩ or higher was used throughout. B. Synthesis of Silver Nanoparticles. Solutions of nanoparticles were synthesized by dropwise addition of 0.45 mM NaBH4 to aqueous mixtures of silver nitrate and trisodium citrate at room temperature (typically, 21 °C). Upon addition of the NaBH4, the solutions immediately changed from colorless to bright yellow, corresponding to the appearance of the well-known surface plasmon resonance around 393 nm. After the particles were allowed to age for 3−7 days, sodium halide solutions were added to start the decay reaction. The amounts of reactants were chosen to produce the following concentrations in the reaction solution (before reaction): [Agnet] = 0.113 mM, [Na3C6H5O7] = 0.113 mM, and [NaBH4] = 0.028−0.056 mM. Note that [Agnet] refers to the molar concentration of silver atoms, regardless of their charge or form. Thus, if all the silver is contained, in its reduced form, in the nanoparticles, [Agnet] = n[Agn]o, where [Agn]o is the initial concentration of silver nanoparticles containing, on average, n atoms. A transmission electron microscope image of nanoparticles deposited from a typical solution is shown in Figure 1.

III. RESULTS AND DISCUSSION We report the effects of four halides, NaF, NaCl, NaBr, and NaI, on the silver nanoparticle surface plasmon resonance near 400 nm. Halide concentrations between 0.03 and 110 mM were studied and are interpreted in terms of two reaction pathways: oxidative decomposition (etching), which dominates at lower concentrations, and aggregation, which becomes important at higher concentrations. The mechanism of each is likely complex, so we will consider simplified mechanisms that capture the essential features. Salient features of the SPR peak shape and of its decay upon addition of the halides will be presented first, followed by a kinetics analysis and modeling of the chloride-induced etching and aggregation. The concentration dependence of NaCl shows the most distinct separation between the pathways. The NaF and mixed NaCl/NaF results will follow, and these will allow us to come to conclusions about the nanoparticle surface charges. Finally, the NaI and NaBr results will be presented. A. The Surface Plasmon Resonance Peak Shape. The yellow colloidal solution prepared as described in the previous section has an absorption peak at 393 ± 2 nm with a full width at half-maximum of 58 ± 3 nm. Within a few seconds following the addition of NaCl, NaBr, or NaI, the surface plasmon resonance peak changes shape. Spectra taken 1 min after addition are shown in Figure 2 for 1.25 mM NaX. The differences among the halides are readily apparent. Addition of NaCl increases the absorbance of the surface plasmon resonance peak by as much as 35% while decreasing the width by 15−30%. The peak position shows a slight blue-shift, decreasing by about 1−2 nm. The magnitude of these changes is dependent on the NaCl concentration, with the maximum change occurring at about 1 mM NaCl. Most of the change takes place within the first few minutes after addition of the NaCl, but minor changes continue to occur over a few hours. NaI causes the peak absorbance to decrease while increasing the peak width substantially. The wavelength shift is more dramatic and to the red. NaBr addition exhibits an intermediate effect, with the peak broadening slightly and shifting slightly to the red. NaF does not affect the peak at all indicating that, as

Figure 1. TEM image of silver nanoparticles deposited from a typical solution.

This image was taken by air-drying a 20 uL droplet of the nanoparticles on a 100 mesh, Formvar-coated copper TEM grid. The average diameter of 130 nanoparticles measured in two images is 5.3 ± 1.8 nm. Thus, assuming that the number density of the atoms in the nanoparticles is the same as in the bulk metal, an average nanoparticle contains 4.7 × 103 silver atoms, so the nanoparticle concentration, [Agn]o, is about 2.5 × 10−8 M. The strong optical absorption by the nanoparticles is due to excitation of the free electrons in the particles, and the area of the absorbance peak in the frequency spectrum is proportional to their concentration, [Agn]. Since each silver atom donates one electron to the metal, the area is also proportional to the concentration of silver atoms in the solution, [Ag1]. The peak area for unreacted nanoparticles, Ap,o, is fairly reproducible from sample to sample. The average of 21 random samples is (1.40 ± 0.07) × 1014 AU s−1, where AU refers to absorbance units. This value and the above initial concentration are used to determine the silver atom concentration during the reactions. C. Instrumentation. All UV−visible absorption spectra were recorded using either a Jasco V-670 Spectrophotometer or 8306



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Figure 2. Effect of 1.25 mM F−, Cl−, Br− and I− on the surface plasmon resonance peak. Spectra were taken 1 min after addition of the halide. Note that the F− spectrum overlaps the “no halide” peak.

expected, fluoride does not replace the surface species (oxide or citrate).11 Previous studies have shown that silver nanoparticles formed in an aerobic solution develop an oxide layer (Ag2O) which, relative to anaerobically produced particles, broadens the surface plasmon resonance peak from about 25 to 50 nm and red-shifts it from 382 nm to about 394 nm.12−15 Experiments by Henglein12 have shown that these changes can be attributed to silver atoms oxidized to ions on the nanoparticle surface. Yin et al.13 later refined this idea by including the effect of the dielectric properties of surface oxide film on the plasmon peak. Their calculations explain both the red-shift of the peak and the increase in width quantitatively. Continuing along these lines, Futamata and Maruyama14 found that a layer of AgCl on otherwise bare nanoparticles also shifts the peak to the red, although less so than Ag2O, but does not change the peak width significantly. They attribute this to the transparency of AgCl in the visible region of the spectrum compared to Ag2O. This explanation agrees with our observations. Since we synthesize the nanoparticles in air, the surface is at least partially covered with an oxide layer, and the resonance peak width and position we observe agree well with previous work. When NaCl is added, the oxide layer is replaced by a chloride layer, and the peak width and position change accordingly. Note that other surface species are also likely to be present. In particular, citrate ion is the known capping agent in our synthetic method, and borates, which are produced when sodium borohydride is oxidized, can affect the surface. In addition, sulfidation of silver nanoparticles is known to occur when exposed to laboratory air,16 although very slowly because of the very low ambient concentrations of sulfur-containing compounds. These all can contribute to the initial width of the SPR peak. Even so, the peak narrowing due to the addition of NaCl clearly indicates that the surface species are being replaced by chloride ions. Silver bromide and, particularly, silver iodide differ from silver chloride in having a significant optical absorption in the region of the surface plasmon resonance peak.17 Thus, it is not surprising that these halides broaden the SPR peak rather than narrow it. This happens immediately upon adding the halide, much more so for I− than for Br−. In time, AgI and AgBr peaks begin to appear in the spectrum until they completely replace the plasmon resonance peak (see parts a and b of Figure 3). Mulvaney15 has reported experiments in which small amounts of iodide were added to nanoparticle solutions, and he too

Figure 3. Time dependence of the surface plasmon resonance spectrum with (a) NaI added to obtain [NaI] = 10 mM and (b) NaBr added to obtain [NaBr] = 12.5 mM.

observed broadening. In his work, silver iodide peaks were not observed because the amount of added I− was too small to produce a sufficiently thick surface layer. His calculations, as well as those by Yin et al.,13 indicate that the broadening of the SPR peak is not due to an underlying AgI peak growing in, but rather to an interaction between the surface plasmon resonance energy levels and surface−adsorbate states. In agreement with this, Kinnan and Chumanov18 have reported strong evidence for the importance of plasmon-induced electronic coupling in SERS studies. B. SPR Peak Decay Kinetics. After the initial change in peak shape, the surface plasmon resonance peak decays at a rate which has an interesting dependence on NaX. This is shown in Figure 4 where the common logarithms of the half-lives for the reactions are plotted as a function of NaX concentration. The half-life is determined directly from the time dependence of the spectra except for the very slow reactions. Reactions with halflives greater than about 3 × 105 s (∼3 days) were only followed to about 10−20% completion, so the half-life was determined by fitting the data with a 2/3 order rate equation. The reason for choosing this rate law will be explained in more detail below. The reactivity of the nanoparticles with NaCl is the most distinctive. Following an order of magnitude increase in rate between 0 and 3 mM, the reaction becomes very slow up to 27 mM, at which point it increases abruptly. The abrupt increase in the rate at 27 mM has a well-established explanationit indicates the onset of aggregation induced by addition of an electrolyte.19 The decay rate of the nanoparticles when NaF is added is slow through a wider concentration range, from 0 to 8307



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of surface oxidation, and the order m will be 2/3 (this will be derived later). If aggregation dominates, and the slow step is the combination of two nanoparticles, the apparent loss of silver nanoparticle atoms is due to the loss of silver monomers, and m will be 2. C. Etching of Nanoparticles by NaCl. Oxidative decomposition is the dominant decay mechanism for NaCl concentrations up to about 27 mM. The evidence for this is the order of the reaction, the consistently narrow width of the surface plasmon resonance peak, and the observation of a reddish precipitate (a gray precipitate would be expected if aggregation dominates). In most of this concentration region, the reaction is too slow to determine the order, but for concentrations near the maximum at 1.0 mM (between 0.75 and 2.0 mM NaCl), plots of ln rate vs ln Ap could be made. An example of one of these plots is shown in Figure 5. The average

Figure 4. Common logarithm of the half-lives, in seconds, plotted as a function of halide concentration. The half-lives were determined directly from the time dependence of the spectra except for [NaCl] = 3−35 mM and [NaF] < 60 mM. The values of t1/2 were not reproducible for NaF > 75 mM, so these were not included.

75 mM, with no local maximum apparent. Above 75 mM, aggregation occurs, but the half-life for a given concentration is erratic, varying over many orders of magnitude. The differences in the onset concentration and the distinctiveness of this onset for these two ions implies that the effect of the electrolyte cannot entirely be explained by the change in the solution ionic strength; additional factors must be important. The reactivities of NaI and NaBr are very different from those of NaCl and NaF. The decay reaction is relatively fast even at low concentrations. Still, a detailed analysis of the kinetics shows (see below) that the high rate is due predominantly to an oxidative decomposition mechanism, with aggregation only contributing at higher concentrations. The rates of the reactions are followed by measuring the time dependence of the absorption spectrum. The area under the surface plasmon resonance (SPR) peak, Ap, in a frequency spectrum is proportional to the number of silver atoms in the nanoparticles, Ap ∝ [Ag1],15 but reaction products can interfere by absorbing in the general region of the nanoparticle monomers. For example, aggregation causes the SPR peak to broaden because the aggregated particles absorb at higher wavelengths.20,21 Although these underlying absorptions cannot be eliminated entirely, the interference is minimized by linearly adjusting the baseline, setting the absorbance to zero for wavelengths outside the peak area. If neutral silver atoms in the nanoparticles are the only reactants whose concentration changes significantly in time, the rate of decay can be given by the mth order equation decay rate = −

d[Ag1] dt

∝−

dA p dt

= kA pm

Figure 5. ln rate vs ln Ap for [NaCl] = 1.00 mM. The rate is calculated as the slope of the line between adjacent data points, and Ap is taken as value of the peak area midway between the points. A linear fit of the data gives a slope equal to the order of the reaction.

of the slopes of these plots give a reaction order of 0.67 with a standard deviation of 0.06. This is the value expected for the decomposition of a spherical particle, as shown in the next paragraph. The loss rate of silver atoms is dependent on [S], the available surface area per unit volume of solution: d[Ag1] dt

= −kS,eff [S]

(2)

The effective rate constant, kS,eff, is dependent on the concentrations of NaX and on other reactive species in the solution, but, as long as these remain essentially constant in time, they will not affect the apparent reaction order. For nanoparticles of radius r and concentration [Agn], [S] = 4πr2[Agn]NA, where NA is Avogadro’s number. In a simple decomposition process, the concentration of nanoparticles remains at its initial value, [Agn] = [Agn]o, but r does change, and this changes the concentration of silver atoms, [Ag1]. The radius can be related to [Ag1] by computing the volume of one nanoparticle, whose density, ρ, is assumed to be that of bulk silver: 4πr3/3 = ([Ag1]M)/([Agn]oρNA), where M is the atomic weight of silver. Combining these equations, eq 2 becomes

(1)

This equation only includes the contribution from a single reaction pathway and assumes that the nanoparticles are monodisperse. If two or more processes compete, the overall decay rate will contain more terms, each with a potentially different order. Assuming, for the moment, that one reaction pathway dominates, the order of the reaction can be determined from the slope of ln rate vs ln Ap, where the rate is calculated from the numerical derivative of the peak area as a function of time. We will be considering two decay pathways: oxidative decomposition and aggregation. If oxidative decomposition dominates, the nanoparticles lose silver atoms because

d[Ag1] dt

= −kS,eff (4π[Ag n]o NA ′ [Ag1]2/3 = −kS,eff

8308

⎞2/3

⎛

3[Ag1]M ⎟ )1/3 ⎜ ⎝

ρ

⎠ (3)



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Thus, the observed 2/3 order can be explained by surface etching of the nanoparticles. Between 5 and 27 mM NaCl, the decay is very slow, with reaction half-lives of 1−2 weeks. This is indicative of surface passivation in which products form a surface layer that inhibits reactants from combining further. In this concentration region, the decomposition was not followed long enough to obtain t1/2 directly, so the values were determined assuming the above 2/3 order rate law. Integrating eq 2 and setting [Ag1]1/3 = ([Ag1]o/ 2)1/3 gives ′ t1/2 = 0.62[Ag1]1/3 o / kS,eff

(4)

This is expected to be a lower limit for t1/2 because surface passivation inhibits the reaction, effectively causing the rate constant to decrease in time. A plausible mechanism for the decomposition, or etching, of the nanoparticles is oxidative corrosion, which involves reduction of oxygen at cathodic sites on the nanoparticles and oxidation of silver to form silver chloride at anodic sites. The same mechanism applies in the case of NaI and NaBr. Here, we specify the analysis to NaCl. cathode:

order with respect to the SPR peak absorbance; dAp/dt = −k2,obsAp2. Consistent with the second order rate law, 1/Ap depends linearly on time, after an induction period of about 200 s. An example is given in Figure 7 for 45 mM NaCl; the

kc

(1/2)O2 + H2O + 2e− → 2OH−

E° = 0.401 V

anode:

Figure 6. Time dependence of the surface plasmon resonance spectrum with NaCl added to obtain [NaCl] = 27 mM.

(5)

ka

2Ag + 2Cl− → 2AgCl(s) + 2e−

E° = −0.2223 V

(6)

With an insoluble salt produced, the overall reaction is favorable: E° = 0.179 V. If we assume that the chemical reaction at the anodic sites is rate limiting and is not affected by deposition of AgCl, the overall rate for the reactions above is given by the equation22 rate =

4kcka[S][O2 ][Cl−] 4kc[O2 ] + ka[Cl−]

Figure 7. Plot of 1/[Agn] vs t for [NaCl] = 45 mM. k2,obs = 2.79 × 10−3 nM−1 s−1. The conversion from peak area to concentration was 1.4 × 1014 AU s−1 = 25 nM.

(7)

If 4kc[O2] ≫ ka[Cl−], the rate is dependent on [Cl−]: rate = ka[S][Cl−]. This explains the increase in rate observed up to 1.0 mM NaCl. At high concentrations of Cl−, such that 4kc[O2] ≪ ka[Cl−], the rate becomes independent of [Cl−] and dependent on [O2]: rate = 4kc[S][O2]. Since the solution is open to air, [O2] is relatively constant throughout the reaction, so the rate should level off. Instead, we see that it decreases above 1.0 mM NaCl. A plausible explanation is that a layer of AgCl is interfering with the diffusion of reactants to the silver. The thickness of this layer is time dependent, so eq 7 no longer applies. D. NaCl-Induced Aggregation of Nanoparticles. The decay rate of the nanoparticles remains consistently slow until about 27 mM NaCl, at which point the rate increases abruptly. We also observe differences in the time dependence of the shape of the SPR peakthe peak initially decreases in width but then broadens, with a new peak temporarily appearing around 600 nm and eventually flattening out (see Figure 6). As the NaCl concentration increases, this evolution occurs much more quickly, and the intermediate peak at 600 nm is not observed. The precipitate that develops above 27 mM NaCl is a dark gray color as would be expected if it consists predominantly of silver metal. The increase in the decay rate above 27 mM NaCl coincides with a change in the reaction order; the rate becomes second-

slope is equal to k2,obs. The observed second-order kinetics suggests that the rate-limiting step in the aggregation process is the combination of two particles, analogous to a bimolecular reaction. Ag n + Ag n → Ag 2n

(8)

Note that for this process, the decrease in Ap is due to loss of nanoparticle monomers, which is directly proportional to the loss of those silver atoms contributing to the SPR peak at 400 nm. Thus, Ap ∝ [Agn], but in order for this proportionality to hold, aggregated particles should not contribute significantly to Ap. This is supported by evidence that dimer particles, at least, have a much higher absorption wavelength.20,21 We do observe a rise in the baseline as aggregation takes place, but this contribution is reduced by subtracting the baseline. The second-order rate constants are plotted in Figure 8. The rate increases by more than 4 orders of magnitude, which is expected for charged particles in a solution of increasing ionic strength. We will evaluate this dependence in an analysis using the Arrhenius equation, k2 = Afreqe−Ea/kT, with the frequency factor, Afreq, estimated from diffusion theory, and the activation energy calculated by finding the maximum in the interaction potential provided by the Derjaguin, Landau, Verwey, and Overbeek (DLVO) theory of colloid stability.23,24 8309



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= 34 for AH = 3 × 10−20 J, with the former giving a better fit to the data. The calculated and observed rate constants are shown in Figure 8. There was some difficulty in fitting both the higher and lower concentrations of NaCl together, so the fit was weighted toward the lower concentrations. Rather than attributing this difficulty to problems with the theory, we feel that it may point to an incomplete modification of the surface by the chloride ions at higher NaCl concentrations resulting in the surface charge being lowered by a smaller amount. Although addition of NaCl immediately changes the SPR peak shape, small changes continue to occur over a few hours. Therefore, the faster aggregation rates may not allow the particle surface enough time to reach a steady state charge. This would lead to anomalously low values of the rate constant. E. NaF Effects. In contrast to NaCl, NaF does not enhance the oxidative decomposition of the silver nanoparticles; the half-life is consistently long in the concentration region between 0 and 75 mM NaF. This can be explained by the fact that the solubility of AgF is much higher than that of the other halides; F− is highly solvated in aqueous solutions. Thus, the anode reaction of the electrochemical process on the surface is not enhanced by complexation of the silver ion, and oxidative decomposition is not favorable: E° = −0.3986 V. The lack of enhancement of nanoparticle oxidation, along with the observation that the SPR peak shape does not change when NaF is added, strongly suggests that the nanoparticle surface is not significantly affected by the presence of fluorine ion in solution. The measured half-lives of the nanoparticles were not reproducible for NaF concentrations between about 75 and 110 mMthey varied inconsistently, from values less than 30 s (the lower limit of our method) to much longer times, on the order of days. This striking behavior was verified many times. In addition, the decay rates sometimes seemed to change through the course of a particular run. Since the aggregation rate is very sensitive to the charge (a 10% change in Z will change k2 by an order of magnitude), this suggests that this parameter is not uniform and consistent among the batches of synthesized nanoparticles. Since this erratic behavior was not observed in the NaCl experiments, it appears that chloride ion improves the uniformity of the surface. The onset of aggregation induced by NaF is not well-defined but appears to occur near 75 mM. Above this concentration, the rate is consistently faster, with t1/2 < 105 s. Applying the DLVO theory and assuming that the aggregation onset occurs when the barrier, Umax, is about 15 kT, the charge on the nanoparticles is determined to be 65 for AH = 3 × 10−19 J and 47 for AH = 3 × 10−20 J. This is about 50% higher than the charges determined for the NaCl experiments. Thus, it is evident that the chloride ion lowers the surface charge of the nanoparticles in solution. To investigate this further, we carried out a set of experiments where NaCl, at a concentration well below the onset of aggregation, was first added to the nanoparticle solution, and then NaF was added 1−2 h later. If the Cl− does indeed change the surface charge, then an initial dose of Cl− followed by addition of F− should cause the onset of aggregation to be comparable to that of simple Cl− addition. The resulting lifetimes are plotted in Figure 9 along with the NaCl and NaF data; there is good overlap of the two sets of data. Therefore, we can conclude that the difference between the effects of F− and Cl− on the aggregation of the nanoparticles is due to modification of the surface by Cl−, which lowers the surface charge.

Figure 8. Observed (points) and calculated (lines) values of ln k2,obs vs [NaCl]. For the calculated values, a = 2.7 nm and Afreq = 7.42 × 109 M−1 s−1.

The Arrhenius frequency factor is equal to the encounter rate for two nanoparticlesthe rate at which they come close enough to react. This is governed by their diffusion rate through the solution, and this, in turn, depends on their hydrodynamic radius. If the encounter distance is simply twice the hydrodynamic radius, then the frequency factor is Afreq ≈ 8RT/3η, where η is the viscosity of the solvent. For water, at 298 K, Afreq = 7.42 × 109 M−1 s−1. Since the hydrodynamic radius is almost certainly larger than the particle radius, this equation gives an upper limit for Afreq. The DLVO potential combines an attractive van der Waals component and a repulsive Coulomb component: ⎤ ⎛ A ⎡ 2a2 2a2 4a2 ⎞⎥ ⎜⎜1 − ⎟⎟ U (r ) = − H ⎢ 2 + + ln 6 ⎢⎣ r − 4a2 ⎝ r2 r 2 ⎠⎥⎦ +

2 e 2Z2 ⎡ e κa ⎤ e−κr ⎢ ⎥ 4πε ⎣ 1 + κa ⎦ r

(9)

r is the distance between the centers of the spheres, a is the particle radius (2.7 nm), AH is the Hamaker constant, Z is the particle charge number, and e is the elementary charge. κ is the inverse of the Debye−Hückel screening length and depends on the ionic strength, I, of the solution: κ2 =

4e 2I 2 εkT

(10)

Here, ε is the permittivity of water, k is the Boltzmann constant, and I is the ionic strength, which is defined predominantly by the NaCl concentration. The Hamaker constant is a property of the metal and is calculated to be about 3 × 10−19 J for silver,23,25 but atomic force microscopy studies on silver metal surfaces have shown that AH may be an order of magnitude lower when molecules are adsorbed on the surface. The nanoparticle charge, Z, is unknown; there is undoubtedly some distribution in Z, so the single value that we obtain is considered as only representative of the average charge. The rate constants, k2, were calculated for a set of NaCl concentrations between 25 and 50 mM by equating the Arrhenius activation energy, E a , to U max (r) for each concentration. Two different values of AH were tried, and, for each, |Z| was varied to obtain the best fit to the observed rate constants. The results are |Z| = 44 for AH = 3 × 10−19 J and |Z| 8310



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Figure 9. Experiment in which NaCl is first added to the nanoparticle solution and then NaF is added after 2 h (solid points). The final concentration of NaCl is 8−12.5 mM. Addition of NaF defines t = 0. [NaX] = [NaCl] + [NaF]. The data for NaF and NaCl are superimposed as open points.

F. NaI Effects. When NaI is added to the nanoparticle solution, the color changes from yellow to orange because of the broadening of the SPR peak. As the reaction progresses, the peak narrows and, after about one half-life, the solution becomes opalescent white. Due to the distinctive absorption peak that appears,17 we attribute the opalescence to a colloidal suspension of AgI that then settles out over the course of a few days (see Figure 3a). The order of the reaction is fairly consistent throughout the concentration range with an average value of m = 0.7 ± 0.2. As explained above in the case of NaCl, this is the value expected for the decomposition of a spherical particle. In contrast to NaCl, we see in Figure 4 that the decay rate levels off at higher concentrations of NaI as predicted by eq 7 for 4kc[O2] ≪ ka[I−]. Therefore, in this case, passivation of the surface is not evident. The oxidation process is more favorable for NaI (E° = 0.5532 V) so a higher rate of oxidation is expected, but the lack of passivation indicates that the effect of I− on the particle surface is very different. One possibility is that AgI does not deposit on the silver surface at all, but this would be unexpected since AgI is even less soluble than AgCl. A more likely possibility is that the AgI layer is more porous, presenting less of a barrier to diffusion of the reactants to the surface. The iodide ion has a larger radius than the chloride ion (220 vs 181 pm), the latter being closer to the radius of the silver atom (144 pm), so it is reasonable to expect that it does not fit on the surface as well, causing the AgI layer to be less uniform. Above 60 mM NaI, the SPR peak begins to show some broadening in time, suggesting that aggregation is contributing to the decay rate, but, since the decomposition reaction is so fast, the two processes are not easily separable. At 84 mM NaI, a very broad absorption quickly grows in to the red of the SPR peak (see Figure 10a), with no evidence of AgI peaks developing. At this concentration, aggregation is likely the dominant process. G. NaBr Effects. Addition of NaBr to the nanoparticle solution causes the SPR peak to first broaden by about 10% less than the broadening seen when NaI is added. As the reaction proceeds, the peak narrows back to its initial width, and, with [NaBr] < 40 mM, a peak near 320 nm grows in. This can be seen in Figure 3b, which shows the time dependence of the UV−visible spectrum of the nanoparticles with 12.5 mM

Figure 10. Time dependence of the surface plasmon resonance spectrum with (a) NaI added to obtain [NaI] = 84 mM and (b) NaBr added to obtain [NaBr] = 65 mM.

NaBr. The peak at 320 nm is attributed to AgBr, which has a weak absorption in this region.17 At about 35 mM NaBr, the nanoparticle monomer peak begins to show some broadening in time, suggesting that aggregation is becoming important. Above 35 mM, the broadening becomes increasingly obvious, and the peak at 320 nm becomes less pronounced. This can be seen by comparing the spectra for 65 mM NaBr (Figure 10b) with those for 12.5 mM NaBr (Figure 3b). Above 50 mM, the decay rate increases sharply indicating that aggregation becomes the dominant mechanism at the higher concentrations. Although the general appearance of the data shown in Figure 4 is similar for NaI and NaBr, the decomposition rate induced by NaBr is about 10 times slower. Also, we have found that the order of reaction is less distinct and is much higher; for concentrations between 5 and 40 mM NaBr, the average value of m is found to be 1.8 ± 0.4. The oxidation process is less favorable (E° = 0.3297 V) when Br− is the complexing agent, so this could account for the slower rate, but the higher reaction order is puzzling. This order suggests a bimolecular process, but AgBr is still observed as a product, so we believe that oxidative decomposition of the particles is still the dominant process. One possible explanation could involve the growth of a halide layer forming on the silver nanoparticles. An interfering layer produces a time-dependent barrier to decomposition (the barrier increases as the layer thickens), and this can lead to an effective order which is not the true order for a bare surface. There is some indication of passivation at concentrations between 20 and 50 mM NaBr; i.e., the rate decreases somewhat before the onset of aggregation occurs. This suggests that the 8311



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Figure 11. Illustration of the two competing mechanisms proposed to explain the decay in the surface plasmon peak when sodium halide is added to silver nanoparticle solutions. In path A, the peak decreases due to loss of silver atoms through oxidative decomposition. In path B, the peak decreases due to loss of monomer nanoparticles through aggregation.

unhindered through a AgI layer, this would explain the lack of passivation observed when NaI is added to the nanoparticle solutions. Although the diffusion rate of Ag+ through AgBr or AgCl is much slower, the conductivity of AgBr is higher than that of AgCl at a given temperature,28 so a greater degree of passivation occurring for the latter is reasonable. It is unclear, though, whether this fact alone can explain the extreme passivation observed. We suspect there is an additional factor related to the smaller size of the chloride ion that comes into play. Recently a scanning tunneling microscope study was done, showing that a silver monolayer deposited on gold was so significantly stabilized after treatment with chloride that the silver chloride layer persevered against oxidation even after annealing at 1000 K.29 This stability was attributed to a low barrier to chlorine translocation along the metal surface, allowing a uniform layer to form which prevented oxygen from reaching the surface. This same mechanism could explain the enhanced passivation we see for NaCl addition. The onset concentration for aggregation also varies with the type of halide. In particular, it is significantly lower for NaCl than for the other halides. This has been discussed in terms of a distinct effect on the nanoparticle surface, in which the surface charge is lowered by almost a factor of 2. How this is accomplished is not clear. One possibility is that a chloride layer reduces the number of adsorption sites for the highly charged citrate ion. Alternatively, the chloride ion may replace citrate entirely but then form AgCl2− rather than AgCl, thereby retaining a negatively charged surface but with a lower value.

AgBr produces a more significant barrier to the reaction than does AgI, but is not as effective as AgCl. Above 50 mM NaBr, the aggregation pathway begins to dominate, as is observed for NaI. The SPR peak broadens, and the AgBr absorption becomes less evident in the spectrum (see Figure 10b).

IV. CONCLUSIONS In this report, we have considered two pathways, oxidative decomposition and aggregation, which can contribute to the decay of the surface plasmon resonance peak when sodium halides are added to aqueous silver nanoparticle solutions. These pathways are shown schematically in Figure 11. Although the pathways compete with each other, in most concentration ranges one of them dominates. At lower halide concentrations, oxidative decomposition is the dominant decay mechanism. The rate is dependent on the identity of the halide (Cl− < Br− < I−) and increases as the standard electrochemical potential of the decomposition reactions becomes more negative. The decay rate is also affected by the degree of surface passivation. Cl− forms a layer that presents a particularly high barrier to penetration of reactant species. Indeed, between NaCl concentrations of 3 and 25 mM, almost no reaction occurs at all. Br− also seems to form a passivating layer, but it is much less effective than the layer formed by Cl−. We suggest that the observed rate constant for the reaction of Br− is time dependent, decreasing as the layer thickens, and leading to an anomalous reaction order. I− does not show any indication of forming a passivating layerthe rate law is consistently 2/3 order for most of the [NaI] region. Since an immediate change in peak shape is observed when NaCl, NaBr, or NaI is added to the nanoparticle solutions, an initial halide surface layer of unknown structure must form very quickly. The subsequent changes in the UV/visible spectra suggest that this layer then develops into a silver halide layer. The rate of oxidative decomposition depends on the rate of diffusion of reactants to the surfacemolecular oxygen and electrons to the cathodic sites, and silver ions and halide ions to the anodic sites. Without a passivating layer on the surface, O2 diffusion through the solution is probably rate limiting. When a surface layer is present, at least one of the reactants at each site (O2 or electrons and Ag+ or X−) must diffuse through this in order to be able to react. Silver ions and electrons diffuse more rapidly than halide ions and oxygen through silver halide,26 and the rate should be dependent on the movement of these reactants. In addition, silver ion diffusion is known to vary for the different silver halides and is substantially faster through AgI than AgBr or AgCl.27 If the diffusion of Ag+ is essentially

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Phone: 619-594-4507. Fax: 619-594-4634. 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 recording TEM images of the nanoparticles. The TEM work was supported by NSF grant DBI-030829.

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REFERENCES

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Kinetics of Halide-Induced Decomposition and Aggregation of Silver

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