Controlled Method for Silica Coating of Silver Colloids. Influence of

Theory Comput. .... A diffusion coefficient of 10-12 cm2 s-1 is estimated for cyanide ion diffusing through ... ACS Applied Materials & Interfaces 201...
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Langmuir 1998, 14, 3740-3748

Controlled Method for Silica Coating of Silver Colloids. Influence of Coating on the Rate of Chemical Reactions Thearith Ung, Luis M. Liz-Marza´n,*,† and Paul Mulvaney* Advanced Mineral Products Special Research Centre, School of Chemistry, University of Melbourne, Parkville 3052, Australia, and Department of Physical Chemistry and Organic Chemistry, Vigo University, Apdo. 874, 36200 Vigo, Spain Received January 5, 1998. In Final Form: April 1, 1998 The preparation and chemical reactivity of Ag@SiO2 particles have been investigated using UV-vis spectroscopy, laser Doppler electrophoresis, and electron microscopy. The factors governing the deposition of silica onto silane-primed silver particles have been investigated and the deposition conditions (pH and reagent concentrations) optimized. It is demonstrated that the thin silica shells deposited from aqueous/ ethanolic sodium silicate solutions are porous. Silver cores dissolve in alkaline cyanide solutions, with the rate depending on the shell thickness. Ag2S@SiO2 can be prepared following exposure to sulfide ion, while silica-coated alloys of silver and gold can be prepared by reaction of Ag@SiO2 with AuCl4-. The kinetics of these core reactions with solution reagents is governed by their rate of diffusion through the shell. A diffusion coefficient of 10-12 cm2 s-1 is estimated for cyanide ion diffusing through non-heattreated silica shells.

Introduction The wet chemical synthesis of small semiconductor and metal particles offers the most economic route for the preparation of nanostructured materials. A variety of applications have already been proposed for such semiconductor particles including solar cells,1 electrochromic devices,2 electroluminescent films,3,4 nonlinear optical switches,5 and high-density information storage systems.6 The demonstration of single electron tunneling7 and Coulomb staircases with ultrafine metal particles by STM8 likewise suggests that extremely small electronic capacitors and electrical switches could be created from cheap, chemically prepared metal particles. A fundamental difficulty to be overcome is the transfer of these materials out of solution while retaining their size-dependent properties. In solution the particles are mobile and will coalesce due to van der Waals forces unless they are protected. Consequently, the synthesis of small particles involves rapid nucleation, homogeneous growth, and a final encapsulation stage with polymers, ions, complexing ligands or surfactants to prevent the growth of larger, bulk crystals. Covalently bonded capping ligands are usually employed with both semiconductor9-11 and metal particles.12-14 These ligands are chemisorbed to the particle surface. They terminate crystal growth and †

Vigo University.

(1) O’Regan, B.; Gra¨tzel, M. Nature 1991, 353, 737. (2) Bedja, I.; Hotchandi, S.; Kamat, P. V. J. Phys. Chem. 1993, 97, 11064. (3) Dabbousi, B. O.; Bawendi, M. G.; Onitsuka, O.; Rubner, M. F. Appl. Phys. Lett. 1995, 66, 1316. (4) Colvin, V.; Schlamp, M.; Alivisatos, A. P. Nature 1994, 370, 354. (5) Neeves, A. E.; Birnboim, M. H. J. Opt. Soc. Am. B 1989, 6, 787. (6) Micheletto, R.; Fukuda, H.; Ohtsu, M. Langmuir 1995, 11, 3333. (7) Scho¨n, U.; Scho¨n, G.; Schmid, G. Angew. Chem., Int. Ed. Engl. 1993, 32, 250. (8) Scho¨nenberger, C.; van Houten, H.; Donkersloot, H. C. Europhys. Lett. 1992, 20, 249. (9) Hayes, D.; Micic, O.; Nenadovic, M. T.; Swayambunathan, V.; Meisel, D. J. Phys. Chem. 1989, 93, 4603. (10) Nosaka, Y.; Yamaguchi, K.; Miyama, H.; Hayashi, H. Chem. Lett. 1988, 605. (11) Dance, I. G.; Choy, A.; Scudder, M. L. J. Am. Chem. Soc. 1984, 106, 6285. (12) Giersig, M.; Mulvaney, P. Langmuir 1993, 9, 3408.

simultaneously confer steric or electrosteric stabilization against coagulation. Such particles can often be dried and redispersed in solvents without coalescence.15 There have already been a number of reports demonstrating that these capped materials in the form of 2D lattices or 3D networks possess unusual electronic properties.16,17 An unfortunate complication is that the organic capping agents are prone to chemical oxidation, especially under photolysis. Mercaptans in particular, which have a strong affinity for metal chalcogenides and soft sp metals such as gold or silver, are readily oxidized. Devices based on these capping functionalities are likely to be susceptible to chemical degradation. Instead we have recently turned to the creation of coreshell particles as a means to create stable nanoparticles which can be incorporated into thin films or polymers. Small, organic capping agents are often still indispensable as a means to control the initial size distribution of the chemically prepared materials. However, these ligands can also be utilized as a primer for the growth of secondary shells around the core particle. Based on this idea, a method for the synthesis of gold-silica core-shell colloidal particles has recently been reported.18-20 The method makes use of the silane coupling agent 3-aminopropyltrimethoxysilane (APS) as a surface primer to render the surface of colloidal gold vitreophilic, that is, receptive (13) Brust, M.; Bethell, D.; Schiffrin, D. J.; Kiely, C. J. Adv. Mater. 1995, 7, 795. (14) Motte, L.; Billoudet, F.; Pileni, M. P. J. Phys. Chem. 1995, 99, 16425. (15) Resch, U.; Eychmu¨ller, A.; Haase, M.; Weller, H. Langmuir 1992, 8, 2215. (16) Badia, A.; Gao, W.; Singh, L.; Demers, L.; Cuccia, L.; Reven, L. Langmuir 1996, 12, 1262. (17) Terrill, R. H.; Postlethwaite, T. A.; Chen, C.-H.; Poon, C.-D.; Terzis, A.; Chen, A.; Hutchison, J. E.; Clark, M. R.; Wignall, G.; Londono, J. D.; Superfine, R.; Falvo, M.; Johnson, C. S., Jr.; Samulski, E. T.; Murray, R. W. J. Am. Chem. Soc. 1995, 117, 12537. (18) Liz-Marza´n, L. M.; Philipse, A. P. J. Colloid Interface Sci. 1995, 176, 459. (19) Liz-Marza´n, L. M.; Giersig, M.; Mulvaney, P. J. Chem. Soc., Chem. Commun. 1996, 731. (20) Liz-Marza´n, L. M.; Giersig, M.; Mulvaney, P. Langmuir 1996, 12, 4329.

S0743-7463(98)00047-X CCC: $15.00 © 1998 American Chemical Society Published on Web 06/10/1998

Silica Coating of Silver Colloids

toward silica monomers or oligomers. This permits the deposition of thin, dense silica coatings in aqueous solution. The use of silica coatings itself is certainly not new. Iler pioneered this field of research and demonstrated that thin silica shells reduced the bulk conductivity of pressed metal powder disks and prevented the photocatalytic degradation of polymeric stabilizers by pigment particles.21 However, in all the earlier work, the particles employed were much larger, and in general, considerable particle clumping and coalescence took place during silica deposition. An electron micrograph of such silica-coated titania aggregates is shown in Furlong’s review article on silica.22 It is therefore important to note that the procedures outlined here are designed to minimize any coagulation of the core particles, which distinguishes them from preparations introduced in the earlier research work by Iler,21 Furlong,22 and James and co-workers.23 Mention is also made of the extensive contributions by Matijevic and co-workers24 and by Philipse and colleagues,25 who have reported a number of procedures for coating dispersions of metal oxides or even latex particles, but these methods do not lend themselves to vitreophobic materials such as gold or silver sols. There are several advantages to using silica shells instead of organic stabilizers. Silica is chemically inert and does not affect redox reactions at the core surface, except through physical blocking of the surface. Second, the silica shell is optically transparent, so that chemical reactions can be monitored spectroscopically. The shell can also be used to modulate the position and intensity of colloidal metal surface plasmon absorption bands, that is, the color of the metal sol.26 Finally, and most obviously, the shell prevents coagulation during chemical reactions, and concentrated dispersions of nanosized semiconducting, magnetic, or metallic materials can be created. Such core-shell particles may be further functionalized, enabling their incorporation into nonpolar solvents, glasses, or polymeric matrixes. An added advantage is that silica is well-known to form ordered colloid crystals, so silica coating may be a useful precursor to the creation of 2D and 3D arrays of any type of nanoparticle system. Because we feel the technique has some general utility, in this article we examine the factors that influence silica deposition and how they may be optimized. The single, most important point to note is that control of the electrical double layer is paramount if coalescence is to be avoided. At the typical concentrations employed here for particle preparation, diffusion-limited coalescence to form particle dimers will occur in about 10 ms. Maintaining high surface charge or potential is a prerequisite to proper coating. For clarity, the primary focus in this paper is on the coating of colloidal silver. With some modifications, the coating process can be extended to include Pt and CdS27 particles, both of which have their own surface chemical idiosyncrasies. Interestingly, the silica shell does not always render the core impervious to chemical depredation. Instead, the silica layer acts as a selective membrane (21) Iler, R. K. U.S. Patent No. 2,885, 366, 1959. (22) Furlong, D. N. In The Chemistry of Colloidal Silica, Bergna, H. E., Ed.; Advances in Chemistry Series 234; American Chemical Society: Washington, DC, 1994; p 535. (23) James, R. O.; Berucci, S. J.; Oltean, G. L. U.S. Patent No. 5,252,441, 1993. (24) Kawahashi, N.; Matijevic, E. J. Colloid Interface Sci. 1990, 138, 534. (25) Philipse, A. P.; van Bruggen, M. P. B.; Pathmamanoharan, C. Langmuir 1994, 10, 92. (26) Mulvaney, P. Langmuir 1996, 12, 788. (27) Correa-Duarte, M. A.; Giersig, M.; Liz-Marza´n, L. M. Chem. Phys. Lett. 1998, 286, 497.

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that controls the rate of chemical reactions involving the metal cores. In some cases, almost complete inhibition of reaction with solution reagents can be achieved; in other cases, the rate is simply retarded, and the rate can then be easily controlled by varying the silica shell thickness. The most important conclusion is that nanometer thick silica coatings provide an efficacious and useful method for stabilizing small particles against coalescence. Experimental Section I. Materials. (3-Aminopropyl)trimethoxysilane or APS from Aldrich, sodium silicate solution (Na2O(SiO2)3-5, 27 wt % SiO2) (Aldrich), AgClO4‚H2O (Sigma), NaBH4 (BDH), and trisodium citrate dihydrate (Normapur) were used as received. Analytical grade ethanol and Milli-Q grade water were used in all the preparations. All other chemicals were analytical grade and used as supplied by the vendor. Transmission electron microscopy (TEM) was carried out with a Philips CM10 microscope; particle size and silica shell thickness were measured from several negatives of each sample. UV-vis spectra were measured with a Hitachi U-2000 or HP Diode Array 8453 spectrophotometer. II. Colloid Preparation. The preparation of reasonably monodisperse colloids is more difficult for silver than gold. Most of the standard recipes for “monodisperse” sols make use of polymers as stabilizers. However, when poly(acrylic acid) (MW 2000, 0.1 mM) was used as a stabilizer, there was no silica deposition, probably due to ineffective adsorption of APS on the surface. Subsequently, colloidal silver was prepared by rapidly adding 1 mL of 0.01 M AgClO4 to 99 mL of a vigorously stirred, ice-cold solution containing 1 mM NaBH4 and 0.30 mM sodium citrate. The solution turned bright yellow immediately. Trisodium citrate is not a very efficient stabilizer of silver sols, and it is readily displaced by the silane coupling agents used to activate the surface toward silicate ion. The particles so prepared had a mean diameter of 10 nm. The preparation of Ag@SiO2 colloids involves three basic steps. First, the surface is activated with APS to generate siloxy groups receptive to silicate ion deposition. Then addition of sodium silicate ion at supersaturated concentrations leads to an initial nanometer thick mantle, and finally controlled precipitation of residual silicate by addition of ethanol creates a homogeneous shell.

Results The paper is broken up into the following sections. In section A, the effects of various solution parameters on the silica deposition process are determined. In doing so, the design of strategies for depositing silica onto other colloidal materials of interest should be made easier. The parameters investigated were in this order: (i) the pH of the colloid before APS addition, (ii) the pH of the colloid prior to the addition of the silicate solution, following the APS addition, (iii) [SiO32-], (iv) dialysis of uncoated silver sols, and (v) EtOH/water ratio on secondary deposition. The aim of the final step was to create larger particles without resort to the Sto¨ber growth28 used earlier, thereby simplifying the preparation. In section B, we report the results on chemical reactions of metal cores with solution reagents, and some of the types of core reactions which are possible. A brief account of some of this work has already appeared.29 We demonstrate that one way to create semiconductor cores is via reaction of metal core particles with suitable complexing agents. A.I. Effect of pH on APS Adsorption. The adsorption of APS onto surfaces is known to be affected by the surface charge on the substrate, the solution pH, and the APS concentration. To optimize the adsorption process, (28) Sto¨ber, W.; Fink, A.; Bohn, E. J. Colloid Interface Sci. 1968, 26, 62. (29) Giersig, M.; Ung, T.; Liz-Marza´n, L. M.; Mulvaney, P. Adv. Mater. 1997, 9, 570.

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Figure 1. Absorption spectra of the coated colloid particles prepared by activation with APS at different pH values. [Ag] ) 0.1 mM; [APS] ) 2.5 µM. After addition of APS at various pH values, the pH is raised to 9 and 0.02% silicate ion added. Ethanol is then added slowly until a 4:1 v/v ratio is reached.

the pH of silver sols was adjusted to various pH values between 3 and 12. APS solution was then added to each of the solutions, amounting to just less than monolayer coverage (2.5 µM for a 0.1 mM Ag colloid). This was calculated assuming 40 Å2 per APS molecule.30 Then addition of silicate solution took place 15 min later. A day later, ethanol was added in a 4:1 volume ratio to precipitate the remaining silicate ions from solution. The effectiveness of the APS adsorption at each pH was then determined by the TEM images of the resulting coated particles. Between pH 3 and 5, the colloid was more thickly and homogeneously coated. At pH 3, there was a relatively large amount of free silica clustered together to form a matrix, in which silver particles were embedded. However, at pH 5, not only were the particles more homogeneously coated with a thickness of 4-5 nm in diameter but they were also well dispersed on the TEM grid. At pH > 5, the majority of the particles were not coated. The absorption spectra of the silver sols also changed during the deposition process, as shown in Figure 1, where the position of the surface plasmon band is shown after silica deposition onto particles activated at different pH values. The position prior to silica deposition but after APS activation was 400 nm; it red-shifted for sols activated at pH 3, 4, and 5 but less so at higher pH values with virtually no change in position observed at pH 9. The red-shift is caused by the increased refractive index around the colloid particles after silica deposition. The larger the shift, the thicker the dielectric coating around the particles.20 Activation with APS at pH 5 appeared optimal. A.II. Effect of pH on the Adsorption of SiO32-. Having established that the optimum pH for APS adsorption is at pH 5, the influence of pH on the deposition of silicate ions onto the primed silver sol was examined. In this case, the pH of each solution was lowered to 5 and APS was then added. The pH was then readjusted to various values, followed by addition of silicate solution. A day later, a constant amount of ethanol was added to all the solutions. Figure 2 shows the thickness and maximum position after the deposition of silica at various pHs. The absorption band red-shifted with increasing pH of deposition. Again we attribute these spectral shifts to silica deposition onto the silver surface. This is also supported by TEM images. The increase in pH influences the quality of silica deposition in several ways. With (30) Plueddermann, E. P. Silane Coupling Agents, 2nd ed.; Plenum Press: New York, 1991.

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Figure 2. Effect of solution pH on the silicate deposition. After activation with APS at the optimal pH of 5, the colloid pH was adjusted to the values shown and silicate ion added. After 24 h, 4:1 v/v ethanol was added to precipitate out the remaining silicate ion. Spectra were recorded after the ethanol addition.

increasing pH, the silver sol becomes more homogeneously and thickly coated, while the amount of free silica decreases. At pH 12, the shell thickness is greatest, with 90% of the particles being coated silver particles, while the remaining 10% are free silica particles. In addition, the coated and free silica particles are reasonably monodisperse. However, one very unfortunate drawback at this pH is that almost every particle contains multiple silver cores. This is attributed to the high ionic strength of the solution. If deposition is carried out at lower pH values (pH 5-10), there is significantly more free silica generated. Addition of ethanol to enforce precipitation of soluble silica causes the particle size to increase, but the final size is determined by the pH at which the initial coating is carried out. As shown in the inset of Figure 2, in which the shell thickness is plotted against pH, the thickness remains almost constant from pH 3 to 9; it then increases drastically toward pH 12. The spectral shifts observed in the absorption spectrum again support the TEM results, with the red-shift being largest for the sols prepared at pH 12. At pH g 9, silica growth is controlled by Ostwald ripening,21 in which larger particles grow while the smaller ones dissolve. Thus, the system becomes more monodisperse with increasing pH. While the thickness of the silica shell is maximal at pH 11.5-12 and the system is quite monodisperse, the formation of multiple cores of colloidal silver is a decided disadvantage. However the silver is strongly stabilized against coalescence, and for some purposes the procedure outlined above will be sufficient. On the other hand, at pH 9, the shell thickness is about 1 /5 that formed at pH 12, but it is homogeneous and the formation of multiple cores does not take place. A.III. Effect of Sodium Silicate Concentration. Silica deposition as a function of silicate concentration was examined at pH 9 and 12. In these experiments the optimum conditions found in sections A.I and A.II were applied. Figure 3 shows the spectral changes observed after deposition of silica at various concentrations at pH 12. Below 0.01% silicate, the absorbance at all wavelengths is reduced and two absorption peaks are formed, with one centered around 400 nm and the other between 450 and 550 nm. In addition, the lower the silicate concentration, the more the peak on the right-hand side red-shifts. The formation of two peaks is due to the production of multiple core particles of Ag@SiO2. The multiple cores are coupled optically, their absorbance resembles that of colloidal silver rods, and the surface plasmon band in the coupled system is red-shifted as the

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Figure 3. Effect of silicate ion concentration on the spectra of colloidal silver. The peaks at longer wavelengths are due to multiplet formation. The pH prior to silicate deposition was 12.

axial ratio increases.31 Shown in Figure 4 is a series of micrographs of the particles coated with various amounts of silica at pH 12. At 0.02% SiO32-, the particles are all homogeneously, thickly coated. The proportion of silica shells containing multiple silver cores at this silica concentration is 30-40%. At 0.01% SiO32-, all the shells contain more than four silver cores, in the shape of long rods; the structure resembles the “Purple of Cassius” images reported by Thiessen, found when colloidal gold is precipitated by colloidal tin dioxide.32 At 0.005% SiO32-, the shell is too thin, and the silver cores are embedded as clusters in a silica matrix. A similar trend was observed at pH 9, with the difference that at low SiO32- concentrations, the shell thickness decrease results in patchy shells rather than a silica matrix. In summary, the critical concentration of SiO32- found is 0.02% for both pH 9 and 12 for 0.1 mM silver sols of a mean size of 10 nm. Below this limit, there is insufficient precipitation to obtain homogeneous shells. This concentration is sufficiently high that the rate of precipitation and the magnitude of the large double layer charge maintain the integrity of the individual particles. It is important to note that citrate ions adsorbed on the particles prior to coating provide some electrosteric stabilization of colloidal silver,33 and the actual barrier to coalescence is some 2-4 nm thick, despite the small ion size. The colloidal silica shell must reach this thickness rapidly in order to reduce the van der Waals interactions between the metal cores; monolayer coverage is insufficient to prevent particle coalescence, as clearly shown by the multiple-core particles in Figure 4. The individual particles are each coated with 1-3 nm of silica, and they have not coagulated before coating but following the initial deposition. If the shell formation is too slow and only 1-3 nm thick, the dispersion forces are still strong enough to cause coalescence of the nascently coated particles, leading to multiple core formation. Conversely, if the [SiO32-] is raised too far, it is impossible to prevent silica nucleation. The formation of these separate silica particles will not always be important but for formation of ordered structures, their removal is essential. A.IV. Role of Dialysis. The condensation of silicate ions to form polymeric and colloidal silica is known to be affected by the nature of the cations present in solution.21 To determine whether the solution ions affect the coating (31) Quinten, M.; Kreibig, U. Surf. Sci. 1986, 172, 557. (32) Thiessen, P. A. Kolloid Z. 1942, 101, 241. (33) Biggs, S.; Mulvaney, P. In Surfactant Adsorption and Surface Solubilization; Sharma, R., Ed.; ACS Symposium Series 615; American Chemical Society: Washington, DC, 1995; Chapter 17.

Figure 4. Electron micrographs of silica-coated silver particles showing the effect of increasing silicate ion concentration at pH 12: (top) 0.02%; (middle) 0.01%; (bottom) 0.005%. Scale bars: 100 nm.

process, the silver sols were dialyzed prior to deposition. More than 80% of ions were removed from a silver sol by dialyzing it for 2 days against Milli-Q water. APS was added at pH 5, followed by addition of active silica at pH 5, 9, 10.5, 11.5, and 12. The remaining procedures were as outlined above. While the same improvement in shell quality is observed as the pH of deposition increases, at pH 11.5, there is a smaller amount of residual silica and all the colloid particles are homogeneously coated, with virtually no multiple cores. Thus the removal of spectator ions from the solution improves the coating quality

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Figure 5. Effect of dialysis of the silver colloid prior to activation with APS and silica deposition. The particles are more homogeneously coated, and there is substantially less free colloidal silica. Preparation conditions as per Figure 1. Scale bar: 100 nm.

Figure 6. Effect of ethanol/H2O volume ratio on the spectra of silica-coated silver colloids. There is an increase in band intensity and a red-shift of the peak position with increasing ethanol content. The inset shows the thickness of the silica shell as a function of the added ethanol content.

dramatically. Ag@SiO2 particles produced from dialyzed silver sols are shown in Figure 5; these are the most homogeneous Ag@SiO2 particles we have been able to synthesize to date. A.V. Effect of Ethanol Concentration. The effect of ethanol concentration on the deposition of silica on a dialyzed silver sol was examined. The pH of the dialyzed sol was adjusted to 5, APS was added, the pH was adjusted again to 11.5, and finally silicate was added. After 24 h, ethanol was added in various ratios. As shown in Figure 6, the absorption bands of the silica-coated silver sol redshift and increase in absorbance with ethanol concentration, and maximum shell thickness is attained at an ethanol/water volume ratio of 4:1. At this ethanol concentration, virtually all of the silicate is forced to precipitate out of the solution onto the thinly precoated silver particles. Moreover, with increasing ethanol concentration, less and less matrix silica is found on the TEM grid; that is, the silica precipitates onto the existing colloid particles and does not lead to nucleation of fresh silica particles. In Figure 7, a micrograph of such colloids with silica shells some 25 nm thick is shown. The advantage of this method is that silica growth up to 25-30 nm can be controllably achieved without resorting to the Sto¨ber method.28 However, despite dialysis and pH optimization,

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Figure 7. Micrograph of silica-coated silver with a shell thickness of 25 nm. Scale bar: 200 nm.

Figure 8. Absorption spectra of Ag@SiO2 particles during oxidation by cyanide ion. The silica shell thickness was 15 nm. [Ag] ) 0.1 mM; [CN-] ) 1.7 mM; pH ) 12. Shown in the inset are kinetic traces as a function of shell thickness: (9) 0; (2) 5 nm; (b) 15 nm; ([) 23 nm.

there are invariably a few percent of silica particles which can only ever be removed by controlled centrifugation. B.I. Chemical Reactions of Core-Shell Particles. High-resolution TEM gives no indication of porosity in the nanometer thick silica shells, though it is well documented that the silica colloids formed from silicate condensation or by TES hydrolysis have densities substantially below the bulk values. BET gas adsorption isotherms consistently show the presence of micro- and mesoporous cavities in such colloid particles.34,35 In Figure 8, we show the absorption spectra obtained from a 0.1 mM silver sol coated with silica at various times after addition of 1.7 mM KCN at pH 12. The metallic silver is attacked by cyanide in the presence of air, oxidizing it to Ag(CN)2- according to

4Ag + 8CN- + O2 + 2H2O f 4Ag(CN)2- + 4OH(1) This process ultimately leads to the formation of a hollow silica shell, as has been shown for the same reaction on silica-coated Au nanoparticles.29 Note that the dissolution of the core requires the transport through the silica shell of both molecular oxygen and cyanide ions, and for (34) van Blaaderen, A.; Vrij, A. J. Colloid Interface Sci. 1993, 156, 1. (35) van Blaaderen, A.; Vrij, A. J. Non-Cryst. Solids 1992, 149, 161.

Silica Coating of Silver Colloids

Figure 9. Spectra of Ag@SiO2 ([Ag] ) 0.1 mM) as a function of time after addition of 50 µM AuCl4-.

complete dissolution outward diffusion of silver cyanide complex anions. The decrease in the absorption of light is shown as the solution changes from yellow to colorless. The surface plasmon band blue-shifts with decreasing particle diameter, which is consistent with the gradual dissolution of the particles.36 In the inset of Figure 8 we show how the dissolution rate decreases with increasing silica shell thickness. For the thicker shell used (23 nm), the sample was boiled before the kinetic run was performed, which led to a substantially slower dissolution rate. B.II. Colloidal Alloy Formation inside silica shells. We have also examined the conversion of the silver core into the more noble metal gold via oxidation with AuCl4-. This process was expected to be quite facile given the rapid kinetics for cyanide oxidation. However the deposition of gold within the core via

AuCl4- + 3Ag(core) f Au(core) + 3AgCl(core) + Cl(2) was found to be very sluggish. It seems likely that the AgCl formed precipitates out within the core volume and passivates the silver surface; however, since the passivation layer is expected to be very thin (2-5 Å), the expansion and explosion of the shell do not take place as found for the reaction with I2.29 It can also happen that the large size of the tetrachloroaurate ion reduces its diffusion through the pore structure of the silica shell. The spectral changes observed from one batch after addition of 50 µM AuCl4- to 0.1 mM Ag@SiO2 are shown in Figure 9. With uncoated particles, complete oxidation is observed after some 48 h with the colloid changing from yellow to red and the surface plasmon band of the colloidal silver shifting from 395 to 530 nm. When the silver is silica coated, the kinetics are much slower. After 4 days, a weak plasmon resonance emerges at 520 nm, but the silver band is not completely absent. Addition of higher concentrations of AuCl4- accelerated the rate of conversion marginally. B.III. Effect of Sulfide Ion and Ammonia. Sulfide ion is particularly aggressive toward silver metal, and oxidation of the naked sols is instantaneous. In Figure 10, we show the spectral changes observed when Ag@SiO2 is exposed to Na2S (0.12 mM, pH 10.5). The solutions are air-saturated, and there is a very rapid change in absorption. The final spectrum resembles that of colloidal (36) Doremus, R. H. J. Chem. Phys. 1965, 42, 414.

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Figure 10. Effect of sulfide ion on the spectrum of Ag@SiO2 colloids. pH ) 10.5; [Na2S] ) 0.12 mM; [Ag] ) 0.1 mM. Shell thickness was 15 nm.

Figure 11. Effect of ammonia on the rate of oxidation of colloidal Ag@SiO2 colloids as a function of the silica shell thickness. [NH3] ) 0.1M; [Ag] ) 0.1 mM.

Ag2S, which has been reported by Henglein et al.37 Silver sulfide is a difficult colloid to prepare in solution. It has a small band gap and readily coalesces. Not only does the permeable silica shell allow the semiconductor to be synthesized directly as a core but it also acts to control the particle size simultaneously. It may be possible to synthesize a number of semiconducting materials as quantized particles in aqueous solution by admission of chalcogenide gases to silica-coated metal or metal oxide precursors via this route. One reason for avoiding the Sto¨ber method28 to coat silver particles is that ammonia appears to catalyze the oxidation of silver by air, as shown in Figure 11. The reaction, which is noticeably slower than that for cyanide oxidation even though a higher concentration (0.1 M) was used, can be represented as

4Ag + 8NH3 + O2 + 2H2O f 4Ag(NH3)2+ + 4OH(3) Discussion The adsorption of APS is central to the coating process, and it is necessary to consider why a pH of 5 appears optimal for the priming of the silver surface. The orientation of APS at both silica and metallic iron surfaces has been studied by FTIR. However, the APS concentrations employed were very high, and multilayers containing (37) Spanhel, L.; Weller, H.; Fojtik, A.; Henglein, A. Ber. BunsenGes. Phys. Chem. 1987, 91, 88.

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and the subsequent displacement of citrate groups. However, when the amine is in the unprotonated form, APS forms a cyclic hydrogen-bonded zwitterion between the amino and silanol groups, effectively locking up the amine moiety and preventing it from surface complexation.30 The optimal pH of 5 reflects these competing requirements, with the charge on the APS sufficient to maintain colloid stability but not so negative that adsorption is retarded electrostatically and also still low enough to prevent zwitterion formation, and we suggest that the active complexing form is +H3NC3H6SiOH(O-)2. The optimal silicate ion concentration of 0.02% (∼3 mM) and the optimal pH for coating (pH ∼ 11.5) have been determined by electron microscopy. These values correspond quite well with those found for the coating of micron-sized particles of titania.22 The exposed surface areas used in this work were considerably lower (1 m2/L compared with 40 m2/L of titania particles), but in both cases, silica deposition is optimal at just above its solubility limit in water. The optimal pH is again a compromise between a number of complex, competing, pH-dependent processes. Iler showed that smooth homogeneous deposition of silica is best achieved from solutions of monomeric Si(OH)4, which has a solubility of 2-3 mM at 25 °C. Figure 12. Mobility measurements of citrate-stabilized gold sols (40 nm diameter) as a function of pH before and after the addition of one monolayer of APS. It is clear that APS helps stabilize the particles. Negative charge also shows that adhesion to the surface is not via the siloxy group, since the free amine group would be protonated and the ζ potential would be more positive. This confirms that amine groups are adhesive groups and attach to gold via dative bonding into vacant d orbitals of gold and are unprotonated when adsorbing. The ζ-pH dependence shows the citrate to be displaced, and the primed colloid has an IEP near 2, similar to that for silica sols. The top part of figure shows the various APS species present in water as a function of pH. At pH 5 the doubly ionized silane predominates.

a variety of species were found. However, pH was consistently found to be a major determinant of surface orientation.38,39 On addition to water the three alkoxy groups of the APS molecules hydrolyze to produce acidic OH groups. These have a pKa of about 2-3.30 The amino group of APS has a pKa of 9, so that APS exists in a number of forms in solution.30,38,39 +

H3NC3H6Si(OH)3 T +

H3NC3H6Si(O-)3 T H2NC3H6Si(O-)3 (4)

Around pH 2-4 APS is positively charged, between 4 and 9 it will be zwitterionic or mononegative, and above pH 9 it will be doubly or triply negatively charged. The silver colloid particles themselves are negatively charged due to adsorbed citrate ions. The pH dependence of surface and ligand charge is summarized in Figure 12. Initially, one expects that adsorption of the anionic forms will be retarded by the negative surface charge; however, adsorption of the cationic form at pH < 3 could easily flocculate the colloidal silver. The factors favoring high pH are that this maintains sol stability through increased ionization of the surface citrate carboxyl groups and that nucleophilic attack of the silver surface atoms by the amine moiety is facilitated in the unprotonated form, that is, at pH > 9. Further, a net negative charge on the APS is desirable with at least two of the siloxy groups ionized to ensure the sol is not destabilized by the adsorption of APS (38) Chiang, C.-H.; Ishida, H.; Koenig, J. L. J. Colloid Interface Sci. 1980, 74, 396. (39) Boerio, F. J.; Armogan, L.; Cheng, S. Y. J. Colloid Interface Sci. 1980, 73, 416.

SiO2 + 2H2O T Si(OH)4

(5)

The upper pH limit for deposition is controlled by the redissolution of silica:

SiO2 + 2OH- T SiO32- + H2O

pKa ) 10.5 (6)

Just below pH 10.5, reactions 5 and 6 are extremely facile, small silica nuclei continuously form and redissolve, and coating occurs primarily by monomer adsorption, which also leads to denser films. Lowering the pH will obviously cause more of the silicate to convert to silica. However, other processes sharply limit how much the pH can be lowered. The addition of monomeric silica to a nascent silica surface requires anionic surface sites:21

dSiO- + Si(OH)4 f dSiOSi(OH)3 + OH-

(7)

Here d denotes a surface site. The number of dSiOsites is very low and increases with pH, with the pKa being 9.5. So the rate of the deposition process is actually favored by high pH, with values between 8 and 10 normally being considered optimal. Furthermore, if two thinly coated particles approach each other at pH 4-9, there will be a mixture of both neutral and anionic sites on each surface. Unless there is sufficient charge to confer electrostatic stability, there will be rapid condensation and cementing of the two particles:

dSiOH + dSiO- f dSiOSi + OH-

(8)

This condensation of neutral and anionic silanes also aids formation of polymeric silica at pH values between 4 and 9, so at these pH values we find more free silica and more gelling of the coated silver particles into networks. Above pH 9, surface silanols are more ionized and the coated particles resist coalescence more easily. At these high pH values, there is less polymeric silica in dynamic equilibrium with the monomer, and as monomer precipitates out onto the surface, the polymer dissociates back to monomer. Surprisingly then, we find that the best coating conditions are at pH 11.5, where we would expect the silica to be redissolving! The silicate ion concentration is critical, with 1 mM SiO32- leading to colloid aggregation

Silica Coating of Silver Colloids

Langmuir, Vol. 14, No. 14, 1998 3747

due to slow shell formation and multiple core formation. At much above 3 mM, there is too much free silica produced. Thus the initial silicate ion concentration must be sufficient to generate 3-5 nm shells that afford stability against coalescence; yet the amount of silicate ion removed as a consequence of this adsorption step must be sufficient to reduce the concentration in solution to the point that the rate of free silica formation is immediately reduced. Increasing the silica concentration to just twice the solubility limit induces significant free silica formation. It is possible to increase the exposed metal surface area to offset this partially, but there is a clear limit imposed if free silica is to be avoided. If the metal particle and sodium silicate concentrations are too high, then the ionic strength increase will favor free silica nucleation. If the silica concentration is fixed, then raising the exposed surface area of metal particles will lead to thinner shells. Eventually the shell thickness will be insufficient to prevent flocculation by dispersion forces, because of the high Hamaker constant of silver. Finally, the improved coated particle quality after dialysis can be understood in terms of the competition between polymerization and particle growth. Both acid and salt reduce the charge on silica shells or on the polymers present, enabling silanol condensation to proceed. Sodium ion can bind strongly to surface-ionized silicate groups, effectively neutralizing them and allowing bridging. Dialysis reduces this effect of free sodium and reduces the amount of networking and gel formation evident in the electron micrographs. While the preparation of coatings at high pH values is supposed to give a dense, impervious shell layer, our experiments on core dissolution and transformation demonstrate conclusively that this is not the case. We comment here briefly on some aspects of the reaction kinetics not addressed above. Although it is not possible here to report categorically on the importance of reagent charge on the kinetics of core reaction, it is obvious that, in the micropores of the silica shell, electrostatics could play a key role in governing the actual rates of reaction. In the case of cyanide or sulfide oxidation, there must be inward diffusion of both complexing agent and molecular oxygen followed by product removal in anionic form. At pH 11 and an ionic strength of 1 mM, the inner surfaces of the micropores have a surface potential between -80 and -90 mV with respect to bulk solution.40 The Debye length will be about 100 Å, so that there is considerable double-layer overlap within micropores. The consequences of this electric potential are several. First, the local pH is no longer the bulk value but will be given by

pHpore ) pHbulk +

Fψpore 2.303RT

(9)

For a pore potential of -90 mV, the internal pH may well be as much as two pH units below the bulk solution value. For both CN- and HS-, the active species within the pores will be HCN or H2S and there can be active transport through the pores in the neutral form. Thus it is not clear whether the surface charge will aid or hinder the chemical reactivity of core materials with solution reagents. In the case of cations, there are also two possible scenarios. The local concentration within the pores may be enhanced due to the negative surface charge, or there could be electrostatic adsorption of the cations onto the silica walls, which may hinder their transport. It may well be possible to enhance the selectivity of the silica shell with respect (40) Larson, I.; Drummond, C.; Grieser, F. J. Am. Chem. Soc. 1993, 115, 11885.

to solution reagents, so that it serves as a membrane layer around the core, filtering certain reactions and accelerating others. The fact that gold sols insoluble in aqua regia can be prepared with quite thin shells suggests that strong selectivity operates in certain cases. The Kinetics of Core Oxidation. To quantify the dissolution of the metal cores by cyanide (and oxygen), we need to make a few simplifying assumptions. These are (i) that the reaction is diffusion-limited, (ii) that the shells around the core are homogeneous, and (iii) that the extinction coefficient of colloidal silver is independent of particle size, so that the absorbance is a direct measure of silver concentration. This last condition is only weakly met, since, for particle radius a < 5 nm, the extinction coefficient scales roughly as (a-1).36 The rate-limiting step with these restrictions in place is then diffusion of the active reagent through the silica shell. The diffusion of molecules through thin spherical shells has been analyzed by Barrer.41,42 For the Ag@SiO2 particles used here, we take the inner core radius to be a ) 5 nm and the outer radius b to be between 7 and 28 nm, corresponding to silica shell thicknesses of 2-23 nm. The addition of cyanide ion or a reagent with concentration C0, by rapid mixing, leads to the following initial boundary conditions:

c ) C0, t g 0, r > b

(10)

c ) 0, t g 0, r < a

(11)

c ) 0, t ) 0, a < r < b

(12)

and

Under these circumstances, the amount of material accumulating within the core in time t is given by

{

J ) 4πab(b - a)C0 R -

1 6

- 2π2

n)∞(-1)n

∑

n)1

n2

exp(-Rn2π2)

}

(13)

where

R)

Dt (b - a)2

(14)

A plot of the normalized accumulated mass as a function of R is shown in Figure 13. The solution is readily seen to consist of a steady-state flow which is linear in time and a transient term. At longer times when R > 0.4, a linear concentration gradient within the shell is established, and the flux through the shell is constant. During this buildup of the shell gradient, there is less reactant entering the core, and this results in the appearance of an induction period in plots of core dissolution versus time. As shown by Barrer, the intercept on the time axis, found by extrapolation of the steady-state portion of the dissolution curve, occurs at

τind )

(b - a)2 6D

(15)

For our colloids, the induction period 0.5-5 s as observed in the dissolution by cyanide ion, is only possible with the (41) Crank, J. The Mathematics of Diffusion, 2nd ed.; Clarendon Press: Oxford, 1975. (42) Barrer, R. M. Philos. Mag. 1944, 35, 802.

3748 Langmuir, Vol. 14, No. 14, 1998

Ung et al.

Figure 14. Simulated kinetic data for the oxidation of silver cores by oxygen and cyanide ion at pH 10 for various core and shell thicknesses. A rate of 1 means the core is completely dissolved in 1 s. Figure 13. Plot of the normalized rate of permeation of material with diffusion coefficient D into the core of a coated particle with inner radius a and outer radius b, as a function of time. Initial conditions are given by eqs 10-12. The intercept of the linear part of the curve shows the apparent induction time to be (b - a)2/6D and the slope of the linear, steady-state diffusion flux to be given by D/(b - a)2.

value of D ∼ 3 × 10-10 to 10-13 cm2 s-1. The real value may be somewhat higher than this, since there was a finite mixing time in these experiments, which also leads to an apparent induction period, but the rate of oxidation of particles with shells is clearly retarded compared with the rate of reaction of uncoated silver particles. Equation 13 also predicts that the actual rate of reaction once steady state is achieved will be slower for thicker shells at constant C0. This is because the concentration gradient, C0/(b - a), decreases with increasing shell thickness b. Thus the oxidation of silica-coated particles should show an induction period scaling as (b - a)2 and a rate of reaction such that J is proportional to ab/(b a). Both the slope of the reaction rate and the induction time prior to the onset of dissolution can be used to obtain values of D. This model is qualitatively in accord with the kinetic data, which show that the induction period does increase with increasing shell thickness and that the actual rate of dissolution is slower for thicker shells. We show in Figure 14 a plot of the simulated rate of dissolution per second as a function of shell thickness for different core sizes. For the calculations the following parameters were used: [CN-] ) 1.7 × 10-6 mol cm-3; [Ag] ) 1.0 × 10-7 mol cm-3; D ) 3 × 10-10 cm2 s-1; N ) 2 × 1012 particles cm-3. The curves allow us to state that a value of D which is about 104 times smaller than the solution phase values gives rates of decay similar to those observed experimentally (see Figure 8) and predicts a dissolution time that increases from about 8 s at a 5 nm shell to 12 s for a 15 nm shell and just 14 s for a 23 nm shell. It was observed that heating of the Ag@SiO2 colloid retards the oxidation process drastically, which means that the value of D is clearly very sensitive to the actual preparation conditions employed and that the value of 10-10cm2 s-1 applies only to fresh silica shells, that have

not been thermally treated to seal pores. A final point to note is that there is a distribution of core particle sizes. If we assume that all particles have an identical shell thickness, then the steady-state flux to the particles increases roughly quadratically with a, as is clear from eq 13. However the total amount required to dissolve the core still increases as a3; thus, as in the case of uncoated particles, these equations predict that smaller particles will dissolve more quickly than larger ones, with the dissolution time scaling linearly with the core radius. Conclusions Optimum conditions for coating a silver sol (10 nm diameter) have been achieved. The 0.10 mM sol can be best coated after dialysis for 2 days, by addition of a monolayer of APS at a pH of 5, with 0.02% SiO32- at pH 11.5, and by titration with ethanol to a volume ratio of 4:1. The shells are porous, and the cores can undergo a variety of chemical reactions to form semiconductor cores, hollow cores, or mixed metal cores. The controlled preparation by this route leads to fairly monodisperse, coated particles and facilitates their ordering into 2D or 3D arrays. The major difficulty facing scale-up of the procedure is simply that the solubility of silica at 0.02% limits the concentration of material that can be coated without significant free silica formation. Yet, the alternative route of direct TES coating in ethanol of primed particles has been unsuccessful to date. The lower dielectric constant of other solvents makes electrostatic stabilization of colloids difficult. Citrate has very low ionization in alcohols, and if an ionized capping agent is used, for example, a mercaptoalkylsulfonate, then its displacement by APS may be impossible. Conversely, the general preparation of concentrated coated metal and semiconductor particles via silica shell formation in aqueous solution does appear eminently feasible. Acknowledgment. The authors are grateful to R. Chapman and A. Russell for performing the mobility measurements. T.U. is grateful for receipt of an Australian Postgraduate Award. This work was supported by a grant from the Australian Research Council and by the Spanish Xunta de Galicia (Project No. XUGA30105A97). LA980047M