Spontaneous formation of silver particles in basic 2-propanol - The

J. Phys. Chem. 1993, 97, 11542-11550

11542

Spontaneous Formation of Silver Particles in Basic 2-Propanol Z.-Y.Huang,t G. Mills,’*+and B. Hajek* Department of Chemistry, Auburn University, Auburn University, Alabama 36849, and Agronomy and Soils Department, Auburn University, Auburn University, Alabama 36849 Received: June 22, 1993; In Final Form: September 3, 1993’

Silver ions are spontaneously reduced in basic, air-saturated solutions of 2-propanol to yield stable Ag colloids in the presence of powdered Nafion. The metal particles (7.4 nm mean diameter) react with oxygen to form Ag2O upon exposure to UV light. Changes in the absorption band of the Ag colloid are observed when the metal particles are in contact witii AgzO particles. The optical changes are explained in terms of surface effects of the metal particles that arise from metal-metal oxide interactions. Kinetic results indicate that the reduction of Ag+ ions is an autocatalytic process. Small metal particles deposited on the surface of larger oxide particles act as catalyst for reduction process. The kinetic observations are consistent with a simple autocatalytic model which correlates autoacceleration with increases in the number of Ag particles. It is proposed that reduction of Ag+ ions takes place on the Ag2O surface whereas oxidation of the solvent molecules occurs on the surface of metal particles.

Introduction Small particles of metals in the condensed phase have received increasing attention. This interest is related to the unique properties of small particles with diametersless than 20 nm, which result from size effects.’-3 The particle-size dependence of electronic susceptibility and magnetic resonance relaxation are examples of quantum effects of small metal particles.’ Enhanced catalytic and chemical activities are other typical size effect^.^,^ For example, extensive investigations on the size-dependent properties of Ag particles and clusters have been performed in a variety of matrices.l-8 The size effects have a profound influence in the chemical behavior of silver clusters and particles in solution. The concept of a size-dependentredox potential of the reduction of silver ions in solution9 has been verified e~perimentally.~ The reduction of Ag+ in water is known to proceed according to the following steps:

The redox potential of reaction 1 is dependent on the number of Ag atoms ( m ) in the particle. Negative redox potentials were estimated for values of m < 10, but it increases rapidly with increasing 111.3JO Also, a correlation between the kinetics of reduction and the size of the Ag particles has been found. It was shown that smaller particles will grow faster, by reduction and combination reactions, than larger partic1es.l0 Recent investigations have provided evidence that Ag clusters can be stabilized in aqueous solution with the aid of complexing agents.lIJ2 In Nafion membranes containing alcohol-water mixtures the growth processes are significantly retarded.” Ag clusters, or nonmetallic silver particles, contain only a small number of atoms and exhibit properties that are atypical of the bulk metal. The clusters have attracted considerable interest, since they appeared to be important intermediates in the photographic development p r o c e ~ s . ~ ~Size J * effects have also been considered to be important during the formation of Ag films by electroless plating processes.Is So far, most of the current information on the room temperature synthesis and chemical reactivity of Ag particles in solution has been obtained in water. Department of Chemistry, Auburn University. 8 Agronomy and Soils Department, Auburn University.

Abstract published in Advance ACS Absfructs,October IS, 1993.

0022-365419312097-11542$04.00/0

It seemed logical to expect some changes in the stability and chemical reactivity of the particles when organic solvents are utilized instead of aqueous solutions. An enhanced stability of the particles toward oxidation was expected in reducing solvents. These ideas were tested by attempting the preparation of Ag particles in several alcoholic solutions. The aim of this research was to develop alternative methods for the preparation of thin Ag films. It was found that Ag+ ions were spontaneously reduced by reaction with the solvents in basic solutions saturated with air. Small silver particles were formed via an autocatalytic process. This report is based on results obtained in 2-propanol solutions containing Nafion as particle stabilizer.

Experimental Section 2-Propanol(Fisher, ACS), NaOH (Fisher, ACS), and AgNO3 (Fisher, Aldrich) were used as received. The same results, within experimentalerror, were obtained with AgNO3 samples of either manufacturer. Stock solutions of Nafion were obtained from Aldrich ( d = 0.874 g mL-I). They contained 5% (wt) of powdered Nafion 117 perfluorinated membrane dissolved in a mixture of alcohols and 10% water. The polymeric powder is a strong acid with an equivalent weight of 1100 glequiv. In a typical preparation, 1 mL of Nafion stock solution was added to about 90 mL of a AgNO3 solution in 2-propanol. Next, 2.5 mL of a 2 X 10-2 M NaOH solution in 2-propanol was added and the volume was adjusted to 100 mL with alcohol. These additions were made within 1 min under vigorous stirring, and the final concentrations were 2.5 X 10-4 M AgNO3,5 X 10-4 M NaOH, and 4 X 10-4N Nafion. While most experiments were performed with fresh stock solutions of AgNO3 in alcohols, no variation in the results was observed with old stock solutions that were aged in the dark for several days. The experimental results were not influenced by the presence of ambient light. All glassware was cleaned with aqua regia; the water used in all preparations was obtained from a Milli-Q/RO system from Millipore. Optical absorption determinations were performed with a Hitachi U-2000 spectrophotometer. A PTI A-1010 S system with a 150-W Xe lamp was used for illumination experiments. The light was filtered through a 5-cm water filter and a 290-nm interference filter (Oriel), or a 0-51 filter from Kopp for irradiations with A 1 380 nm. X-ray diffraction (XRD) and transmission electron microscopy (TEM) were carried out with a Siemens 1-2 powder diffractometer and a JEOL 100-CX microscope, respectively. 0 1993 American Chemical Society

The Journal of Physical Chemistry, Vol. 97, No. 44, 1993 11543

Silver Particles in Basic 2-Propanol

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Two-Theta Figure 2. Diffraction patterns of Ag powders: (A) precipitated from a solution containing M AgNOs with 10-2 M NaOH; (B) obtained from a solution containing 2.5 X 10-4 M AgNO3, 10-3 M NaOH, and 4 X 10-4 N Nafion.

h (nm) Figure 1. Evolution of the absorption spectra with time in a 2-propanol solution containing 2.5 X 10-4 M AgNO3, 5 X 10-4M NaOH, and 4 X 10-4 N Nafion.

Results A. Colloid Characterization. A bright yellow solution was slowly formed when hydroxide ions were added to an air-saturated alcoholic solution of Ag+ ions. The evolution of the absorption spectra with time is presented in Figure 1. Only slight changes in absorption were detected during the first 5 min of reaction. Between 5 and 16 min a broad absorption centered at 410-425 nm developed, the exact position of this band varied slightly from one preparation to another. At longer reactions times the absorption band narrowed and shifted continuously to shorter wavelengths. The shifts in theabsorption band ceased after about 5 h, at which point a narrow and strong absorption band centered at 390-405 nm was obtained. Afterward, only a slow but continuous increase in the intensity of the 390 nm band was noticed. The evolution of the optical spectra was completed in about 10 h. It should be noted that while the exact A, of the early absorption band as well as of the final absorption band varied slightly from preparation to preparation, a fairly constant blue shift of about 20 nm was observed in all preparations. In addition, a weak and broad shoulder centered at about 267 nm developed with increasing reaction time. No significant changes in the absorption spectra were detected upon aging the solution in the dark for several days; only a small shift of the maximum to longer wavelengths was observed in solutions that were more than a week old. The final spectra (Figure 4) resembled that of aqueous colloidal Ag,16J7 and the absorption at A 1 230 nm followed Beer's law with an extinction coefficient of t = 7.6 X l o 3 M-1 cm-1 at A,,. Some information on the nature of products formed by similar reactions has been gathered in several investigations. Metallic particles and Ag mirrors have been produced at room temperature in basic water-alcohol mixtures.lEaOn the other hand, generation of white Ag2O particles occurred in ethanol containing OH- ions when mixing of the reagents was done at -45 OC, but the color of the particles changed to brown after heating the solutions to room temperature.18b In addition, it is conceivable that small metal aggregates, formed during the reduction processes, may react with oxygen to produce mixtures of oxide and metal particles. Therefore,a thorough identificationof the products was necessary. While X-ray diffraction can provide information on the overall composition of the samples, the results obtained from powders of the colloid were often complicated due to presence of Nafion and other products. An additional complication resulted from peak broadening that is characteristic of small crystallite size. In order to circumventthese problems, larger particles were analyzed initially and the results were compared with measurements from several colloidal solutions. Micrometer-size gray particles that precipitated after reacting 10-2 M AgN03 with M NaOH

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Figure 3. Histogram of a colloid prepared as described in Figure l.fis the fraction of particles with a diameter within a certain diameter range.

in Nafion-free 2-propanolwere cleaned several times with solvent prior to analysis. As shown in Figure 2A, the powder exhibited diffraction peaks that correspond to the 111 and 200 lattice plains of metallic Ag. Particles that precipitated from solutions containing 2.5 X 10-4 M AgNOp, le3M NaOH, and 4 X 1 V N Nafion gave similar results (Figure 2B). The red-brown powder used in the diffraction experiment was produced by evaporating the solvent to dryness, followed by repeated cleaning with solvent. The broadening of the diffraction peaks in Figure 2B is due to the smaller size of the metal particles in this sample and the diffraction of the perfluorinated polymer. Two broad peaks centered at about 20 = 17O and 39O were detected in XRD experiments with polymer powders obtained by evaporating the Nafion stock solution to dryness. These signals increased the overall background, broadening the peaks of the particles. It should be noted that the interferenceby the perfluorinatedpolymer existed in all colloidal samples prepared with Nafion, even in samples washed repeatedly with alcohol. Diffraction experiments were also carried out with powder samples from the yellow colloid described above. To avoid any possible growth of the colloidal particles, fresh colloids were freeze-dried at low temperature. Peaks corresponding to NaNO3 and Nafion overlappingwith Ag signals were identified on the patterns of the yellow powders. TEM measurements of the yellow colloid were performed on samples prepared by placing a drop of a fresh colloid on Cu grids covered with holey carbon films, followed by solvent evaporation under vacuum. The colloid consisted of nearly spherical particles; a histogram of particle diameters obtained by analyzing images from several regions of the grid is shown in Figure 3. While the most frequent particle sizewas 4.5 nm, a statistical averageparticle size of 7.4 nm was calculated. An average agglomeration number of n = 1.25 X 104 and a particle concentration of 2 X 10-8M were calculated using 7.4 nm as the average diameter and the density of Ag (d = 10.52 g cm-3). Electron diffraction patterns of the particles corresponded to cubic Ag, which is in agreement with the XRD observations. In addition, weak signals corresponding to the 111 and 220 lattice planes of AgzO were also detected.

Huang et al.

11544 The Journal of Physical Chemistry, Vol. 97, No. 44, 1993

(nm) Figure 4. Spectral changes during irradiation of colloidal Ag with 290nm light in the presence of air. Irradiation times (min) from top to bottom are 0, 7, 15, 31, 64, 100, 180,277, 448,662,691,and 810.

These diffraction peaks were not detected by XRD measurements, implying that the Ag20 crystallites were probably formed by air-oxidation of small Ag particles. Results similar to those reported in a previous study were obtained when thecolloidal synthesiswas attempted at 4 5 OC.leb The colorless 2-propanolsolutionsproduced under these conditions exhibited an absorption spectrum that resembled the spectrum measured after 2.5 min of reaction; see Figure 1. TEM analysis of the solutions prepared at low temperature revealed the presence of silver oxide particles with irregular shapes and with sizesgreater than 100nm, and also very few Ag particles. Heating the solutions to room temperature resulted in brown solutions that, according to XRD measurements,consisted of large metallic particles. Thus, agglomeration of the oxide particles seemed to take place when the cold solutions were heated to room temperature because the Ag crystallites formed during the subsequent reduction process were much larger than those generated in the normal synthesis of the colloid. B. Particle Reactions. Gray metal particles precipitated upon addition of 0.1 M H N 0 3to a colloid prepared as described in the Experimental Section. A similar experiment with HC1 resulted in the formation of a turbid suspension of AgCl and large Ag particles. The reaction of l e 3 M NaCl with 2.5 X 10-4M AgN03 in the presence of 4 X l e N Nafion resulted in a turbid solution of AgCl particles. In contrast, no reaction was detected when 10-3 M NaCl was added to the colloid. This results indicated that essentially all Ag+ ions were transformed to the elemental state during the reduction process. As in the case of aqueous colloids,17b the colloidal particles agglomerated in the presence of H202 to yield blue colloids. Illumination of evacuated colloids at 290 nm or irradiation of air-saturated colloids with photons of A 1 380 nm produced no change in the optical spectra of the particles. Different behavior was noticed when colloids were illuminated with light of 290 nm in the presence of air. As shown in Figure 4 the band centered at about 390 nm decreased in intensity and shifted to longer wavelengths with increasing irradiation time. The position of the maximum ceased to shift after about 64 min of irradiation, at which point A, was about 410 nm. Further illumination of the colloid resulted in a decrease of the optical density without any change in Am,,. The decay in optical density was exponential with time and a rate constant of 4.8 X 10-5 s-1 was determined from a first-orderplot of thedata. Micrometer-sizepolycrystalline aggregates of AgzO as well as small Ag particles were detected by TEM in a colloid illuminated for 8 10 nm. A statistical average diameter of 7.4 nm was obtained for the metal particles. As mentioned above, a red shift of A,, occurred during irradiation of colloidal solutions containing air. In contrast, the

h (nm) Figure 5. Effect of Nafion addition to a colloid during the early stages of formation: continuous lincs, colloid spectrum after 2.5,5, and 10 min of thermal reaction; broken line, spectrum after addition of Nafion to a solution reacted for 10 min.

m 2701110

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Figure 6. Optical changes between 250and 280 nm during the formation of a colloid. Experimental conditions as in Figure 1.

opposite shift of A, was observed during the colloid formation process; see Figure 1. During the first 16 min of reaction the maximum was centered at about 410 nm but shifted to about 390 nm at longer times. Since protons were generated during the reduction of the Ag+ ions (see below), it seemed possible that the shift in the absorption peak was related to the change in [H+] during the course of the reaction. This hypothesis was tested by allowing the thermal reaction to proceed for 10 min, followed by additionof 1 W N Nafionor HNOs. Theresultsoftheexperiment with Nafion are shown in Figure 5. Immediately after addition of Nafion the broad absorption band centered at about 410 nm changed into a narrower and stronger absorption band with maximum at about 390 nm and the reduction process was terminated. The intensity of the new band decreased slowly at longer times after the addition of Nafion, indicating that the Ag particles were oxidized by the added H+ ions. Similar observations were made when HN03 was used instead of Nafion. In contrast, addition of OH- ions to a similar solution led to a sharp increase in the rate of colloid formation and to a shift of A, to wavelengths above 410 nm due to agglomeration and precipitation of the Ag particles. C. Dynamics of Particle Formation. Figure 6 shows the variation of optical density between 250 and 280 nm with increasing reaction time in a 2-propanol solution containing 2.5 X 10-4 M AgN03, 5 X 10-4 M NaOH, and 4 X 10-4 N Nafion. At short reactions times an absorption maximum followed by a minimum were detectedindependent of the wavelength. However, the position of the maximum along the time coordinate was dependent on A. While the maximum at 250nm occurred without time delay, the absorption maxima at 260,270, and 280 nm were detected at about 3,5, and 6 min, respectively. In Contrast, the minimum in optical density (O.D.) was wavelength-independent and occurred at about 9 min. After the minimum, the optical density increased sharply until about 24 min. At that time the

The Journal of Physical Chemistry, Vol. 97,No. 44, 1993 11545


plotting ln(a/( 1 - a ) ) vs time. The fit was very good (r2 = 0.998, where r is the correlation coefficient) up to about 45% conversion; an observed rate constant of kobg = 1.7 X 10-3 s-1 was calculated from the slope of this plot. Similar kinetic results were obtained by using kinetic data obtained at different wavelengths, kob = 2 X l e 3 S-I (410 nm) and kob = 1.4 X 10-3 s-1 (390 nm); the kinetics of the first step was highly reproducible. At conversions higher than 44% the increase in optical density followed firstorder kinetics with an apparent rate constant of k,, = 3.9 X 10-4 S-1.

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Figure 7. Evolution of the absorption at 400 nm during colloid formation: curve a, 2.5 X lo" M AgNOo, 5 X lo" M NaOH, and 4 X 10-4 N Nafion; curve b, as in curve a plus 2.5 X lP5M Ag colloid.

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[AgNO,] (x lo4 M) Figure 8. (a, top) Autocatalytic plots of optical data obtained at several initial concentrations of Ag+. [NaOH] = 5 X lo" M; [Nafion] = 4 X lo" N; [As+] = (0)2.5 X lo" M, (0) 3.5 X lo" M, (A)4 X l p M, and (0) 5 X lo" M. (b, bottom) Effect of the initial [Ag+] on the observed autocatalytic rate constant in solutionscontaining [NaOH] = 5 X 10-4 M and [Nafion] = 4 X lo" N. optical density decreased slightly, but it increased smoothly at longer times. Different behavior was observed at longer wavelengths. The evolution of the optical density a t 400 nm is presented in Figure 7, curve a. Small changes were detected in the first 9 min but the absorbance i'ncreased drastically at longer times. The sigmoidal shape of the curve suggested that autocatalysis was involved.19 As shown in the Appendix, ln(a/( 1 - a ) ) is expected to change linearly with time in the case of simple autocatalysis ( a = O.D.,/O.D.,; O.D., and O.D., are the optical densities at times t and m, respectively). The results presented in Figure 8a confirm this assumption since a straight line was obtained by

The reduction of Ag+ ions was also investigated in solutions containing colloidal species at early stages of the reaction. For this purpose, 10 mL of the colloid was mixed with an AgNO3 solution, followed by addition of stabilizer as described in the Experimental Section. N o reduction of the Ag+ ions occurred under these conditions; in fact, the added colloid was slowly oxidized by protons released from Nafion. However, in the presence of stabilizer and base a fast reduction of the silver ions was noticed. The evolution of the optical density a t 400 nm as function of time is presented in curve b of Figure 7. The reduction process appeared to occur in a single autocatalyticstep, with kotg = 4.5 X 10-3 s-1. This result was obtained after correcting the data from the absorption of the colloid, that was added at the beginning of the reaction. While the reduction process was faster when colloid was initially present in the solution, unstable redbrown colloids were produced. The rate of colloid formation decreased with increasing concentration of H2O. For example, in a 50% water-50% 2-propanol mixture the reduction process was completed after only 24 h of reaction, but it was totally inhibited when pure water was used as a solvent. In contrast, no change in the kinetics of colloid formation was detected when the reaction was performed in 2-propanol containing 10-* M acetone. This observation suggests that acetone, which is the expected product when 2-propanol is oxidized, was not involved in the autocatalytic reaction. Experiments with 0 2 - or Ar-saturated solutions of alcohol produced the same results as those obtained with airequilibrated solutions. As mentioned before, two reactions steps were detected during the colloid formation. The initial process was autocatalytic and it became predominant with increasing concentration of AgNO3 or NaOH. At least 33% of thereduction process occurred via autocatalysis in most experiments. During the second reaction step the optical absorption increased exponentially with time. Furthermore, treatment of the results with the method of dimentionless parameters20yielded a reaction order close to 1. However, in several experiments the kinetics of this step was complex. In those cases, straight lines were obtained from both first order and autocatalytic plots of the data. These results suggest that the two processes occurred simultaneously in all cases, but that autocatalysis was initially the fastest reaction. The dependence of the reduction kinetics on the silver-ion concentration was studied between 7.5 X 10-5 I [AgNOj] 1 5 X lo4 M, with [NaOH] = 5 X 10-4 M and 4 X 10-4 N Nafion. Kinetic determinations at even higher AgNO3 concentrations were hampered by the precipitations of large Ag particles. At concentrations higher than 3 X 1 V M the color of the colloids turned dark red. No reaction was observed at [AgN03] = 7.5 X 10-5 M. However, the reduction occurred in a similar solution that was Nafion-free. In all other experiments plots of optical density vs time yielded sigmoidal curves similar to curve a in Figure 7. Typical examples of autocatalysis plots are presented in Figure 8a. The observed rate constant varied linearly with [AgN03]between 7.5 X 10-5 and 5 X 10-4 M. These results are illustrated in Figure 8b and the slope of the straight line is 11.3 M-f s-1, While the slopes of the straight lines shown in Figure 8a increased with increasing AgN03 concentration, the intercepts followed the opposite trend. A plot of e(-int) vs [AgNO3] yielded a straight line with a slope of 7.1 X lo4 M-I (Figure 9). The percent reduction via autocatalysis increased steadily with increasing [AgNOp] as shown in Figure loa. This implies that

Huang et al.


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Figure 9. Dependency of e(-int)on the initial [As+]. Same conditions as in Figure 8b.

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[Nafion] (x 1O4 N) Figure 11. Dependencyofthekon theNafionconcentrationinsolutions containing 2.5 X 10-4 M AgNO3 and 5 X 10-4 M NaOH. =-5 X 10-2+-1.5 X 10-210g[Nafion],wasobtained.Incontrast, all autocatalytic plots had a common intercept of int = -2.6, indicating that the value of the intercept is independent of the Nafion concentration. At [Nafion] C 4 X 1 W Nthe autocatalytic step accounted for at least 70% of the reduction, but unstable dark colloids were produced and no simple kinetic law was obeyed during the second step of the reduction. The reduction process was strongly dependent on the presence of base in the solutions, since it was not observed in the absence of OH-. However, it was not possible to study in a systematic way the [OH-] influence on the autocatalytic process. In solutions containing 2.5 X 10-4 M AgN03 and 4 X 10-4 N Nafion a fast precipitation of Ag particles occurred at [OH-] > 5 X 10-4 M, causing interferences in the optical determinations. Lowering the [OH-] to concentrations of less than 5 X 10-4 M resulted in lower reaction rates because most the hydroxide ions were neutralized by Nafion. Attempts were made to overcome these problems by lowering the concentrations of all reagents. For example, the rate of reduction was investigated using solutions containing 10-4 M AgNOs, 1.6 X 10-4 N Nafion, and several OH- concentrations. While the rate of reaction increased with increasing [OH-], no simple kinetic law was obeyed in most cases.

Discussion

[AgNO,] (X IO4 M) Figure 10. (a, top) Percent reductionof Ag+viaautocatalysisas function of the initial [AB+].Same conditionsas in Figure 8b. (b, bottom) Effect of the initial [Ag+]on the apparent first-order rate constant in solutions containing [NaOH] = 5 X 10-4 M and [Nafion] = 4 X 10-4 N. relative importance of the second step, the apparent first-order process, diminished as [Ag+] increased. The rate of the second step was also dependent on the concentration of AgN03. As depicted in Figure lob, kappincreased from 4.8 X 10-5 s-1 a t 1.5 X 10-4 M AgNO3 to 4.5 X 10-4 s-1 a t [AgN03] 1 3 X 10-4 M. Solutions containing 2.5 X 10-4M AgNO3,S X 10-4M NaOH, and varying concentrations of Nafion were used to study the influence of this chemical on the rate of reaction. Figure 11 showsthe dependency kobon [Nafion] . In the absence of polymer particles the reduction was fast with kob = 1.4 X 10-2 s-1. The slopes of the autocatalytic plots, and therefore kob, decreased continuously with increasing [Nafion] and no reduction occurred when the concentration of the polymeric compound was equal to [NaOH]. A plot of kob vs log[Nafion] yielded a straight-line (rZ = 0.99) relationship from which the empirical equation, kob

The spontaneous reduction of Ag+ ions in basic 2-propanol results in a fast formation of metallic particles. Spectroscopic and diffraction results (Figures 1 and 2) show that in the presence of Nafion the reduction is slower and that stable colloids of Ag are formed. While the standard reduction potential of Ag+ in 2-propanolis not known, it is probably similar to thevalueobtained in H2O since Eo values for this reaction in methanol, ethanol, and butanol are 0.764,0.749, and 0.801 V, respectively.2' The fact that no reduction of Ag+ ions takes place in base-free alcoholic solutions is an indication that, as in the case of HzO, the spontaneous formation of Ag, clusters (n C 10) through reaction 1 is unfavorable in pure 2-propanol. Significant changes in the reactivity of Ag+ ions occur in solutionscontaining base. In water, metal particles are formed when Ag2O is treated with concentrated KOH.22 Formation of metal in alcohols is spontaneous a t lower OH-concentrations because these solvents are stronger reductants than water. It is clear that hydroxide ions play a crucial role in the reaction of Ag+ ions with alcohols. This role appears to be 2-fold: providing the necessary thermodynamic conditions that allow the initiation of the reduction process, and preventing a fast corrosion of the very small particles a t early stages of the reaction. In the presence of Ag+ ions the base is expected to form Ag20. This assumption is based on the similarities between the optical spectrum of solutions prepared at low temperature and that measured shortly after mixing the reagents a t room temperature (Figure 1,spectrum after 2.5 min) and considering that Ag2O was identified by TEM



measurements as the predominant product of the low temperature synthesis. Formation of silver hydroxides is considered less likely since this type of compound is formed in solutions containing a large excess of base.= In a typical synthesis 2.5 X 10-4M AgN03, 5 X l P M NaOH, and 4 X 1 W N Nafion were used as initial concentrations. Deprotonation of the strongly acidic Nafion particles generates 4 X 10-4 M H+ ions, and the resulting hydroxide ion concentration after neutralization ([OH-IR) is only 1 W M. Hence, Ag+ is the excess reagent and the theoretical yield of AglO formula units is 5 X l t 5 M. 2Ag+

+ 20H- e A g 2 0 + H20

(3) The oxide is most probably in the colloidal form since no turbidity or precipitation is detected in the solutions. In water-alcohols mixtures of methanol and ethanol Ag2O is less soluble than in pure water and a solubility of 3.9 X l t 5 M has been reported for solutions containing 91% ethanol.23 A similar or even lower solubility is expected in 2-propanolsolutions. We thereforebelieve that Ag particle formation is initiated by adsorption of Ag+ ions from solution onto the surface of the large oxide particles. Adsorbed silver ions are then reduced by reaction with solvent molecules. This sequence of steps is similar to the reactions of physical development in the photographic process that induce fog formation.24 Evidence has been presented that reduction of Ag+ in the presence of AgBr particles occurs preferentially on the surface of the semiconductor.25 Also, the reduction of silver hydroxide complexes in basic water seems to proceed via surface reactions.22 It should be noted that the reduction reaction is substantiallyslower in mixturesof 2-propanol with H2O. In these mixtures the solubility of the oxide is higher than in neat 2-propanol. The rate of the reduction process is lower under these conditions because less Ag2O is formed. As shown in Figure 1, theabsorption bandof themetal particles exhibits a maximum at 410 nm during the early stages of the colloid formation. At longer reaction times the band shifts continuously to shorter wavelengths until the typical spectra of colloidal Ag are obtained. The existence of the absorption band with A,, = 410 nm can be correlated to surface effects. While quantum-size effects are not important for Ag particles in solid matrices when d,,, 2 2.2 nm, surface effects became noticeable at sizes of about 4 nm.8 For aqueous colloids a red shift of the colloid absorptionband is observed after adsorptionof Cd2+,PV+, and In3+ ions on the particle surface.16 This plasmon band arises from surface oscillations of the electron gas and it appears to be sensitive to modifications of the particle surface. Formation of small Ag particles on the surface of the large oxide particles will result in electronic interactions between the two types of materials because Ag2O has semiconducting properties.26 This means that thesurfaceeffectsofthe Ag particles are related to metal-semiconductor interactions. In the case of the aqueous colloids the red shift of the plasmon band has been rationalized in terms of a withdrawal of electron density from the Ag particles by the adsorbed metal ions.16 A similar interpretation can be used to explain the results shown in Figure 1. This interpretation requires a transfer of electron density from the metal to the oxide particles. While n-type as well as p-type Ag2O have been prepared, the contact between Ag and n-AgzO is ohmic whereas a Shottky barrier is formed when the metal is in contact with the p-type semiconductor.26 The rather high Shottky barrier that is formed in the latter case is expected to inhibit the transfer of electrondensity from themetal to thesemiconductor,Bbwhereas no limitation for this transfer is expected in the case of an ohmic contact. Thus, the shifts in the plasmon band of colloidal Ag in 2-propanol are explained under the assumption that transfer of electron density from the metal to the oxide takes place when small Ag particles are formed on the surface of n-Ag2O. As the number of Ag particles generated on the surface of Ag2O particles increaseswith time, the intensityof the absorption at 410 nm becomes stronger. This is seen in Figure 1 during the

initial 16 min of reaction, where ,A, remained unchanged but the intensity of the absorption band increased. Hence, at the beginning, the concentration of the Ag particles increases with time without a significant change in average particle diameter. However, the reduction of Ag+ ions by 2-propanol will generate protons according to the following overall stoichiometry: A g + + '/2(CH3),CHOH-Ag

+ H + + '/,(CH,),CO

(4)

Therefore, a slow dissolution of the oxide support is expected to occur with increasing proton concentration. As the Ag2O is dissolved, the metal-semiconductor interactions cease to exist and the absorption band is blue-shifted. The smooth shift of A, and slow increase in absorbance after 16 min (Figure 1) can be attributed to a rather slow kinetics of oxide particle dissolution. Colloidal Ag particles prepared in 2-propanol with Nafion are very stable toward agglomeration or oxidation by air despite of their small size. When the particles are illuminated with 290-nm photons they react with oxygen to form Ag2O according to the overall reaction: Ag,

+ do2+ hv

-

(Ag20),,

(5) whered = 0.25nand (Ag2O),!is an oxide particleof agglomeration number n' = 0.5n. In principle, reaction 5 can by initiated by photochemical processes of colloidal Ag, or by Ag2O particles that may have formed via thermal oxidation of small metal particles. Aqueous colloids of Ag are dissolved when illuminatd with UV light in the presence of N2O,Z7whereas Ag2O is a lightsensitivesemiconductor.26 In our experiments the photooxidation of colloidal Ag was first order with respect to [Ag,]. This kinetic result is consistent with photochemical processes initiated by metal particles instead of oxide particles. In the latter case the rate of metal particle oxidation is expected to increase with increasing time because of the larger amounts of oxide present as the reaction proceeds. During the photochemical oxidation of colloidal silver ,A, is shifted from 390 to 410 nm (see Figure 4). This shift in A, is in the opposite direction of the shift that was detected during the colloid formation (Figure 1). The changes in the colloid spectra depicted in Figure 4 are consistent with the assumption that the absorptionmaximumat 410nmisaconsequenceofmetalparticlc+ oxide particle interactions. These interactions arise when Ag2O is formed on the surface of the metal particles by photooxidation of surface Ag atoms, leading to red shifts in . ,A The shifts cease to occur when most of the metal particles are in contact with oxide. Further irradiation of the solution decreases the intensity of the band centered at 410 nm as the remaining metal is converted to Ag2O. The results of Figure 5 are also in agreement with the idea that surface effects have a strong influence on the position and shape of the Ag colloid plasmon band. Acids were added to Ag colloids formed at early stages of the reduction reaction in these experiments. An abrupt change in the optical spectrum and a shift in A, from 410 to 390 nm is observed after the solution is acidified with 10-4 N Nafion. The addition of acid induces partial dissolution of the oxide which, in turn, separates the Ag particles from the oxide surface. Under these conditions the metal-oxide interactions are lost and the absorption band is blue shifted to yield the characteristic spectra of colloidal Ag with a plasmon band at 390nm. The results of theseexperimentssupport the concept that shape and position of the Ag colloid plasmon band are influenced by the metal-oxide interactions and also provide evidence that the oxide particles participate in the autocatalyticreduction reaction. Dissolution of the Ag20 particles by the acid inhibits the reduction of Ag+ ions and the existing metal particles became unstable to corrosion reactions. Perhaps the most intriguing result of this investigation is the autocatalytic nature of the reduction process. In order to understand the sequence of steps that lead to metal particle formation, it is useful to inspect the evolution of absorbance as

Huang et al.

11548 The Journal of Physical Chemistry, Vol. 97,No. 44, 1993

function of time. According to the results depicted in Figure 6 , species absorbing between 250 and 280 nm are present in solution after mixing of the reagents is completed. These species survive for about 6 min; only minor changes are detected at longer wavelengths during that period of time. A fast growth of the band at 410 nm takes place when the cluster decay is completed (=9 min). As is evident from curve a in Figure 7, the initial 9 min corresponds to the induction period of the autocatalytic growth. Addition of Ag particles to the reaction mixture accelerates significantly the rate of Ag+ reduction and decreases the induction period to about 2 min (curve b, Figure 7). Obviously, the metal particles catalyze the reduction process, the initial 9 min being required to form a minimum concentration of catalyst needed for initiation of autocatalysis. This impliesthat the species absorbing in the range of 250-280 nm are precursors of the metal particles. A similar sequence of events involving UV-absorbing species, which decay to generate small Ag particles, have been observed during the radiation-chemical reduction of silver ions in aqueous Ag2S04 solutions."J*28The cluster Ag42+ has been identified as the species absorbing at 250-280 nm. This cluster decaysvia second-order sulfate-catalyzed agglomeration processes to form metal particles in a few milliseconds. In sulfate-free solutions Ag4*+decays in part to form other clusters, but it is stable for several hours in solutions containing stabilizers.11c In view of the similarities between our results and those from previous investigations we assume that the absorption in the 250280 nm range is due to clusters similar to Ag42+. Also, this is the smallest known cluster that is inert toward air-oxidation in solution at room temperature.llc We expect these species to be relatively stable under our experimental conditions because 2-propanol is a weaker oxidant than water. When A g ~ S 0 4is reduced inside Nafion membranes, the rate of agglomeration of Ag42+is lower than the rate in homogeneous solution.13 This effect is related to the slower diffusion of the clusters in the cavities of the membranes. Unlike the Nafion membranes, no cavities exist in the perfluorinated polymer particles, which are believed to consist of micelle-type structures with a core of polymer chains and the sulfonic groups facing the polar solvent.29 While electrostatic attractions between the clusters and the sulfonic groups of Nafion are possible, we believe that the rather long survival time of the species absorbing between 250 and 280 nm is a result of the relatively small amount of clusters and stabilization by the Ag2O surface. The stability of these species is enhanced becausecluster-cluster reactions are slow under these conditions. The sequence of events shown at early times in Figure 6 corresponds to the uncatalyzed formation of Ag particles. As mentioned previously, it is assumed that the uncatalyzed reduction takes place on the oxide surface. After enough metal particles are present, the reduction process becomes autocatalytic. The reduction of Ag+ ions in aqueous solution via autocatalysis is an important step in the photographic process that has been studied exten~ively.2~ Electrochemical and catalysis models explain the kinetic observationsby assuming that reduction of Ag+ ions occurs at the metal particle surface through reactions 1 and 2.24 According to these models the particle radius and surface area increase as the reaction proceeds. Ag colloids are efficient catalysts for the reaction and the increase of optical density (O.D.) with time follows the equation ( o . D . ~- (0.~.,)1/3 = kt

(6) where O.D.0 is the optical density a t t = 0. This equation is the solution to the rate expression for metal formation

d m l d t = k[Ag+] [R]S (7) where m is the mass of Ag produced, R is reducing agent, and S is the area of the Ag surface ( S 0: m2/3). However, all attempts to fit our data with eq 6 were unsuccessful. In addition, spontaneous metal ion reduction was not observed in base-free

solutions containing Ag+ ions and colloidal particles, even in the absence of Nafion. Thus, a mechanism based only on reactions 1 and 2 seems unlikely. Deviations from the usual autocatalytic process have been observed in a few occasions but no alternative mechanism has been suggested.30 On the other hand, the rate of metal particle formation by reduction of metal oxides is, in some cases, directly proportional to the metal mass.31 An example of this type of process is the high-temperature reduction of Ag20.32 We have assumed that the autocatalytic formation of Ag particles in 2-propanol follows a similar rate law d[Ag,]/dt = kTAg+] [2-propanol] [Ag,] (8) where k'is the rate constant for autocatalysis and n is the average agglomeration number. The rate of particle formation is independent of [OH-] because most of these ions are converted to Ag2O via reaction 3, and the concentration of the oxide is expected to remain nearly constant during the autocatalytic step. Since the alcohol is the solvent, [2-propanol] will not change significantly during the reaction and eq 8 reduces to d[Ag,lldt = k,[As+l[Ag,l (9) where k, is the measured rate constant for autocatalysis; k, = kTZpropanol]. The solution to eq 9 is 1 / ( [ ~ g + 1 0+ [

A ~ ~ II lJ n ( [ ~ g + I 0 / [ ~ g n l o+)

ln([~gnI,/[~g+Il= ) I kat (10) where [Ag+]o is the initial concentration of Ag+ ions, [Ag,]o is the concentration of metal particles needed to initiate the autocatalytic reduction, [Ag,lf is the concentration of particles after a time t, and [Ag+], is the [Ag+] remaining unreacted after a time t . The latter equation can be modified to allow the use of spectral data such as that depicted in Figure 7. It is shown in the Appendix that eq 10 can be transformed into eq 11 1n(a/(l - a ) ) = (k,I~g+Io)t-ln([~g+Io/n[~gnIo) (11) where a = O.D.f/O.D., and the optical densities were measured at 400 nm. According to eq 11 a plot of ln(a/( 1 - a ) ) vs time should yield a straight line, where the slope corresponds to the observed rate constant, kob. Examples of such plots are presented in Figure 8a for several initial concentrations of Ag+. From eq 11, kob is expected to vary linearly with the initial concentration of Ag+ since kobs = k,[Ag+]o. Figure 8b is a plot of kob vs the initial concentration of Ag+ and a rate constant for autocatalysis of k, = 11.3 M-1 s-l was determined from the slope of the straight line. At [Ag+] = 7.5 X M formation of metal particles occurred in Nafion-free solution but was inhibited by the presence of the perfluorinated polymer particles. A logical interpretation of these results is that at the low [Ag+] most silver ions are strongly bound to the polymer particles. Hence, no Ag+ ions are available in solution for formation of Ag2O and metal particles. At concentrations higher than 7.5 X M Ag+ enough silver ions exist in solution to react with base and form silver oxide. This interpretation implies that the mobility of Ag+ ions is severely depressed when bound to Nafion particles, which is consistent with the results from experiments using Nafion membranes." Equation 11 also relates the intercepts of the plots shown in Figure 8a with [Ag+]o since int = -ln([Ag+]o/n[Ag,]o). As expected, a straight line is obtained by plotting e-'"' vs [Ag+]o (Figure 9), with a slope of 7.1 X lo4 M-I. This slope is equal to (n[Ag,]o)-I and the concentration of reduced Ag+ ions that are needed to start autocatalysis ( c A ~ ) is equal to n[Ag,]o and corresponds to C A ~= 1.4 X le5M. It should be noted that C A ~ is independent of [Ag+]o,which explains the fact that the induction period was the same in these experiments. We have shown that the kinetics of the Ag+ ion reduction can be described by a simple autocatalysis model with a rate law

Silver Particles in Basic 2-Propanol given by eq 9. According to eq 9 autocatalysis occurs because the number, not the size, of Ag particles increases with time. However, the presence of As20 is also required for the reaction to take place. A simple mechanism that is consistent with these conditions can be postulated. According to this mechanism, small Ag particles are formed initially through an uncatalyzed reaction on the surfaceof large Ag20 particles. Solventmolecules adsorbed on the Ag particles are oxidized, and because of the physical contact between metal and oxide particles, the resulting electrons are transferred from the Ag particles to the oxide particles. The driving force for the oxidation reaction can be related, in part, to the fact that metal particles are deficient in electron density when they are in contact withAg20particles. As shown in Figure 1, the deficiency in electron density persists during the autocatalytic step, which implies a continuous transfer of electron density from the metal to the oxide. Electrons transferred to the semiconductor become delocalized in this material. Silver ions adsorbed on the surface of Ag2O are then reduced via reaction with the delocalized electrons. This mechanism is represented by the following reactions:

where Ag,-(AgzO) represents a metal particle in contact with an oxide particle, Agn2--(AgzO) is a charged metal particle in contact with Ag20, e-(AgzO) represents an electron delocalized in the semiconductor,Ag+,& and Aga&correspond to a surface adsorbed metal ion and metal atom, respectively. Diffusion and agglomeration of adsorbed metal atoms on the oxide surface results in formation of metal clusters. Unlike the case of the uncatalyzed reaction, where agglomeration of clusters seems to be the rate limiting step, growth of clusters is faster in the catalyzed process. Clusters can grow on the surface by capturing mobile metal atoms or by agglomerationvia coalescence or ripening.24.33 These processes will occur faster at the higher cluster concentrations that exist on the oxide surface when reactions 12-1 4 take place. In addition, growth of metal particles may also occur via the mechanism represented by reactions 1 and 2. The rate of solvent oxidation, which is probably the slowest step, is expected to increase as the number of metal particles in contact with Ag20 increases, and the reduction of silver ions becomes autoaccelerating. Results from pulse radiolysis experiments indicated the absorption band characteristic of metal particles is fully developed for Ag clusters that contain 14 Ag atoms.13.28b However, these clusters are only short-lived intermediates in a growth process that generates much larger particles. Growth steps involving intermediates with n < 14 are not observed because the development of the metal absorption band is very fast. Thus, the only process detected optically in our experiments is to the transformation of nonabsorbing Ag particles in particles with a fully developed metallic absorption band. According to the mechanism represented by reactions 12-14, protons are generatedat somedistanceaway from thesites where metal atoms and clusters are formed. Thus, proton-induced corrosion of these species'3 is not fast enough under these conditions to efficiently compete with cluster growth. Electrons delocalized in the semiconductorparticles can, in principle, be scavenged by oxygen to form peroxide radicals. This type of radical is able to induce fog in photographic emulsions by reducing Ag+,& on the surface of semiconductor grains.24 The rate of Ag+ ion reduction in 2-propanol is the same in the presence or absence of air. This implies that a similar reaction between silver ions and peroxide radical takes place in our system if 0 2 reacts with the delocalized electrons.


While autocatalysis is initially the predominant pathway, formation of metal particles also occurs via a second reaction pathway. The second reaction pathway follows an apparent firstorder rate law and predominates after a certain fraction of Ag+ ions is reduced. From the results of Figure 10a it is estimated that autocatalysis ceases to predominate in most cases when the concentration of Ag+ ions left unreduced is about (1-1.2) X 10-4 M. It should be noted that, according to reaction 3, this is roughly the amount of silver ions that is present as AgzO under our experimental conditions. A possible mechanism for the apparent first-order step involves desorption of Ag+,ds from the semiconductor surface and reduction of the metal ions on the surface of existing metal particles

This process is similar to the mechanism of solution physical development in photography, where Ag+ ions are liberated into solution via dissolution of the semicond~ctor.~~ The apparent first-order process becomes the predominant pathway when most of the free metal ions are reduced. At that point Ag20 is the principal source of Ag+,& ions, which result from the partial dissolution of the oxide when protons attack the semiconductor lattice. We believe that both processes occur simultaneously but that autocatalysis is initially the major pathway since kob 1 4kapp for all initial [Ag+] (compare the results of Figures 8b and lob). Adsorption of Ag+ ions on the oxide surface is expected to be a fast process and autocatalysis predominates as long as free metal ions exist in the solution. On the other hand, desorption of Ag+,& from the oxide surface is expected to be slower than desorption of metal atoms and neutral clusters, which implies that growth of the Ag particles occurs mainly by agglomeration and ripening processes. Also, the rate of oxide particle dissolution will be considerable only after a significant amount of H+ ions are generated by reaction 12. Thus, the low rate of the first-order pathway as compared with the rate of the autocatalytic reaction is explained when all these factors are considered. As in the case of autocatalysis, reaction 15 is detected optically only when nonabsorbing clusters are transformed into absorbing particles. The results shown in Figure 10b are explained if the amount of Ag2O formed increases initially, at constant [OH-]R, when the concentration of silver ions increases. The net results are an increase of the oxide surface area and an increase in kapP However, the increase in the oxide surface area is limited by reaction 3. Since OH- is the limiting reagent, no further increases in the oxide surface area are possible after consumption of the OH- ions is completed. Hence, kaPDremains constant at higher concentrations of Ag+ ions. The rate of autocatalysis is dependent on the concentration of Nafion (Figure 11). Hydroxyl ions are partially neutralized in the presence of the polymer particles due to the acid dissociation of the latter. An initial concentration of 5 X 1 W M OH- ions was used in all the experiments of Figure 11. Obviously, [OH-]R and the amount of Ag2O formed via reaction 3 decrease with a raise in [Nafion]. Hence, the oxide surface area and kobdecrease as the concentration of the stabilizer increases. The reduction of Ag+ ions by 2-propanol is inhibited at 5 X 1 W N Nafion because the amount of protons generated from the dissociation of the stabilizer is sufficient to neutralize all OH- ions. However, since changes in the concentration of polymer particles may affect the rate of metal particle formation by changing the rate of the cluster agglomeration process, the weak dependence of kob on [Nafion] is probably a combined effect of the acidic nature and the stabilizing effect of the polymer particles. The reduction of Ag+ ions in basic 2-propanolis a simplemethod for the preparation of small Ag particles in solution. We have used this reaction to prepare Ag films on several supports. Interestingly, the deposition of Ag on some supports is also autocatalytic and the reaction follows a rate law that is similar

Huang et al.


to the rate law of particle formation in solution. These results will be the subject of a separate report.

in basic 2-propanol can be verified in a simple way by using optical data in conjunction with eq a7.

Acknowledgment. We wish to thank R. Zee and Y.X. Liao for their help in TEM experiments and to W. Grot (du Pont) for helpful discussions. This work was supported by the Strategic Defense Initiative Organization's Office of Innovative Science and Technology (SDIO/TNI) through Contract No. N6092191-C-0078 with the Naval Surface Warfare Center.

References and Notes

Appendix Some modifications are required in eq 10 in order to make use of the spectral data of Figure 7. Since the concentrationof metal particles needed to initiate the autocatalytic reduction ([Agnlo) is very small, it can be assumed that [Agn]~

Spontaneous formation of silver particles in basic 2-propanol - The

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