Shape Control of Silver Nanoparticles by Stepwise Citrate Reduction

Mar 27, 2009 - Under low pH, the product was mainly dominated by triangle or polygon silver nanoparticles due to the slow reduction rate of the precur...
5 downloads 4 Views 199KB Size
Download PDF
J. Phys. Chem. C 2009, 113, 6573–6576

6573

Shape Control of Silver Nanoparticles by Stepwise Citrate Reduction Xinyi Dong, Xiaohui Ji, Hongli Wu, Lili Zhao, Jun Li, and Wensheng Yang* State Key Laboratory of Surpramolecular Structures and Materials, College of Chemistry, Jilin UniVersity, Changchun, 130012, People’s Republic of China ReceiVed: January 26, 2009; ReVised Manuscript ReceiVed: March 1, 2009

Growth of silver nanoparticles by the citrate reduction of silver nitrate under the range of pH from 5.7 to 11.1 was investigated systematically and quantitatively. Reduction of the silver precursor (Ag+) was promoted with increased pH, attributed to the higher activity of the citrate reductant under high pH value. Under high pH, the product was composed of both spherical and rod-like silver nanoparticles as a result of the fast reduction rate of the precursor. Under low pH, the product was mainly dominated by triangle or polygon silver nanoparticles due to the slow reduction rate of the precursor. The product that is dominated by spherical silver nanoparticles cannot be acquired by the one-step citrate reduction method in the range of pH investigated, indicating the poor balance between the nucleation and growth processes in the reactions. On the basis of the results of quantitative analyses, a stepwise reduction method, in which the nucleation and growth processes were carried out at high and low pH, respectively, was proposed for the syntheses of spherical silver nanoparticles. Introduction Gold and silver nanoparticles have been widely investigated during the past decades due to their unique electronic and optical properties and application potentials in catalysis,1-3 biological and chemical sensing,4,5 nonlinear optics,6 surface-enhanced Raman spectroscopy (SERS),7 and electronics.8 During the past decades, great efforts have been contributed to studying control over the shape and size and related optical and electric properties of gold and silver nanoparticles.9-19 Reduction of gold precursors (HAuCl4) by citrate, which was first invented by Turkevich in 1951,20 has become one of the most used methods for the preparation of gold nanoparticles. In 1982, Lee’s group extended the citrate reduction method to the preparation of silver nanoparticles.21 However, the citrate reduction method is less successful in control over the size and shape of silver nanoparticles than gold nanoparticles. In the citrate reduction method, a simple strategy for control over the size of the nanoparticles is to change the molar ratio of citrate and the gold or silver precursors.16,18,22 Our previous work proved that the pH value of the system will increase with the increased molar ratio of citrate and HAuCl4 used due to the basic character of citrate.19 The increased pH will induce the variation of coordination situation of Au(III) ions as well as the reducing power of citrate, thus mediating the nucleation and growth processes of the gold nanoparticles. With control over the pH value of the reaction system, nearly monodispersed spherical gold nanoparticles can be prepared successfully.19 This encourages us to investigate systematically the growth kinetics and temporal size/shape evolution of silver nanoparticles under different pH values while fixing the molar ratio of citrate and silver precursor (AgNO3). Quantitative analysis showed that the reduction rate of the precursor was promoted with increasing pH value due to the higher activity of the citrate reductant under high pH. The product dominated by spherical silver nanoparticles cannot be acquired in the range of pH investigated due to * To whom correspondence should be addressed. Phone: +86-43185168185. Fax: +86-431-85168086. E-mail: [email protected].

the poor balance between the nucleation and growth processes. According to the well-known LaMer model,23 a stepwise reduction method, in which the nucleation and growth stages were carried out at high and low pH, respectively, is proposed. The shape control over the spherical silver nanoparticles is improved greatly due to the improved balance between nucleation and growth stages in the stepwise method. Experimental Section Materials. Silver nitrate (AgNO3), trisodium citrate (Na3C6H5O7), sodium hydroxide (NaOH), nitric acid, 1,10phenanthroline (PHEN), 2,4,5,7-tetrabromofluorescein (TBF), sodium acetate, and acetic acid were of analytical grade and were not purified further before use. All water was distilled and subsequently purified to Millipore Milli-Q quality. All glassware used was cleaned in a bath of freshly prepared aqua regia solution (HCl/HNO3, 3:1), then rinsed thoroughly with H2O before use. Synthesis of Silver Nanoparticles. A typical reaction is as follows, which is similar to the well-known Turkevich method.20 A 100-mL sample of aqueous trisodium citrate (7.0 mM) was prepared in a 250-mL flask, containing controlled amounts of citric acid or NaOH at room temperature. The pH value of various reaction solutions was adjusted by further addition of nitric acid or NaOH. The solution with different pH values (from 5.7 to 11.1) was first brought to a boil while being stirred, and then 1.0 mL of 0.1 M aqueous silver nitrate was added to start the reaction. The color of the reaction solution was observed to change generally from colorless to yellow, then turbid. At last no more change took place, indicating the reaction was completed. The stepwise method was carried out by fast nucleation at higher pH value (shown in the text), and the solution was suddenly adjusted to pH 6.1 by addition of HNO3 at the end of the nucleation stage. UV-Visible Absorption Spectra. UV-vis absorption spectroscopy of silver nanoparticles and Ag ion concentration determination was recorded on a SHIMADZU UV-2450 UV-vis spectrophotometer. Aliquots of the solution were taken out at

10.1021/jp900775b CCC: $40.75  2009 American Chemical Society Published on Web 03/27/2009

6574 J. Phys. Chem. C, Vol. 113, No. 16, 2009

Dong et al.

Figure 1. UV-visible spectra for the silver nanoparticles synthesized at different pH conditions. The insert is the magnification for the 700-900-nm window.

Figure 3. Temporal evolution of the UV-visible absorption spectra (a), the absorption peak position (b), the absorption peak intensity (c), and the precursor concentration (d) during the syntheses of silver nanoparticles at pH 8.3. Figure 2. Transmission electron images for the silver nanoparticles synthesized under pH values of 11.1, 8.3, 6.1, and 5.7.

different times during the reaction, and cooled in ice water to quench the reaction. Such samples were diluted 8 times before being measured. The determination of Ag+ concentration was based on the classical spectrophotometric method established by Mohamed and Roland (details are shown in Figure S1 in the Supporting Information).24 Transmission Electron Microscopy. Transmission electron microscopic (TEM) observations were carried out under a Hitachi H-8100 IV electron microscope at 200 kV after the diluted dispersions of the silver nanoparticles were dropped onto carbon-coated copper grids. Results and Discussion In all the reactions, the concentration of the silver precursor (AgNO3) was set as 1.0 × 10-3 M, and the molar ratio of citrate and AgNO3 was 7:1. The pH value of the reaction solutions was adjusted by addition of nitric acid or NaOH. Figure 1 shows the UV-vis absorption spectra of the silver nanoparticles synthesized under different pH values. As the pH was mediated from 11.1 to 5.7, the absorption peak of the nanoparticles shifted from 402 to 466 nm, and the full width at half-maximum (fwhm) of the spectra increased from 84 to 328 nm, possibly due to the increased size and polydispersity of the particles.25-27 TEM observations were carried out to observe the morphology of the silver nanoparticles synthesized under different pH values. Figure 2 gives the TEM images of the silver nanoparticles synthesized under pH values of 11.1, 8.3, 6.1, and 5.7, respectively. At pH 11.1 and 8.3, the product obtained was a mixture of spherical and rod-like particles. At pH 6.1 and 5.7, almost no rod-like particles were observable in the product. However, in addition to the spherical particles, silver nanoparticles with triangle or polygon shape occurred. Especially, the case became very serious at pH 5.7. The broadened fwhm of absorption peaks (shown in Figure 1) at low pH should be

Figure 4. Temporal evolution of Ag+ concentration for the reactions carried out at different pH values.

attributed to the particles with triangle or polygon shape. If the rod-like particles were not considered, the average sizes and standard deviations of the particles prepared at pH 11.1, 8.3, 6.1, and 5.7 were determined to be 30 nm ((13%), 59 nm ((18%), 69 nm ((16%), and 96 nm ((25%), respectively, as shown by the TEM observations. The average particle size increased evidently while the pH was decreased from 11.1 to 5.7, consistent with the red-shift of the absorption peaks. Moreover, the integrated absorbance in the 700-900-nm window (see Figure 1), should be mainly ascribed to the rodlike ones for the particles prepared at high pH,10,28 and to the scattering effect of the large particles prepared at the lowest pH (5.7).17,29,30 Temporal evolutions of UV-vis spectra were recorded to further understand the citrate reduction reaction. A representative temporal evolution of UV-vis spectra for the reaction carried out at pH 8.3 was shown in Figure 3a. The corresponding evolution of the position (λmax) and intensity of the absorption peaks are summarized in Figure 3, parts b and c. It is likely that the whole reaction could be divided into three stages. At the initial stage (I) of the reaction (0-3 min), the solution was colorless, and there was no absorbance observed from the UV-vis spectra. At the end of stage I, the solution turned slightly yellow and a weak absorption peak was observed at 406 nm. So stage I is assumed to be the nucleation stage.20,31

Shape Control of Silver Nanoparticles

J. Phys. Chem. C, Vol. 113, No. 16, 2009 6575

Figure 5. Comparisons of the typical one-step syntheses at pH 8.1 and 6.1 with the stepwise syntheses initiated at pH 8.3 and then adjusted to pH 6.1 at the end of stage I: (a) temporal evolution of Ag+ concentration, (b) the UV-visible absorption spectra for the as-prepared silver nanoparticles, and (c) the TEM image of the silver nanoparticles synthesized by the stepwise method.

Figure 6. TEM images of the silver nanoparticles synthesized by the stepwise method with initial pH values of (a) 9.0, (b) 7.7, and (c) 6.9. All three reactions were adjusted to pH 6.1 at the end of stage I. (d) UV-visible absorption spectra for the samples of a, b, and c.

In stage II of the reaction (3-10 min), the absorption peak shifted from 406 to 430 nm gradually, accompanied by the continuous increase in the absorbance intensity, meaning the continuous growth of the silver nanoparticles.26,27,32,33 In stage III (10-20 min), the absorption peak began to experience a slight blue-shift and then remained steady. At the same time, the absorbance intensity remained almost unchanged during this stage. Temporal evolution of the silver precursor (Ag+) concentration was determined as shown in Figure 3d, which is approximately a mirror image of the corresponding temporal evolution of the absorption peak intensity in the UV-vis absorption spectra. Only a few of the precursor molecules were consumed during stage I, consistent with the assumption that stage I represents the nucleation process. The remaining precursor was consumed almost completely during stage II, consistent with the occurrence of the growth process. Since no precursor was left in stage III, the slight blue-shift of the absorption peak should be attributed to the intraparticle ripening.34,35 The temporal evolution of the precursor concentration under other pH values was also investigated as shown in Figure 4. Although three similar stages as described above were also observable for each temporal evolution, the reduction rate of the precursor showed great difference. For example, stages I and II lasted only 3 and 9 min for the reaction at pH 8.3. However, they were prolonged to 50 and 80 min respectively for the reaction carried out at pH 5.7. It is obvious that the reduction rate of Ag+ was decelerated with the decrease of the pH value, which is consistent with the pH dependent reduction activity of citrate (see the Supporting Information). Under higher pH, such as pH 11.1 and 8.3, both the nucleation and growth processes were fulfilled rapidly due to the fast reduction of the

precursor. Combined with the TEM results shown in Figure 2a,b, the rod-like particles formed under the higher pH conditions should be related to the high instantaneous concentration of monomer produced in the reaction.35 With the decrease of pH, the nucleation and growth processes were slowed down synchronously due to the slow reduction rate of the precursor. As a result, formation of the rod-like particles was suppressed almost completely under low pH value (Figure 2c,d). However, some anisotropic silver nanoparticles, such as triangle or polygon ones, were formed. Especially, the case became very serious at pH 5.7. Therefore, it is difficult to have control over the shape of the silver nanoparticles under either low or high pH attributed to the poor balance between nucleation and growth processes in the typical citrate reduction reaction. According to the well-known LaMer model,23 a fast nucleation process followed by a slow growth should be beneficial to the shape control of the nanoparticles.18-20 Thus, a stepwise reduction method is proposed by adopting different pH values at the nucleation and growth stages, for example, the fast nucleation stage is acquired at high pH and the following growth stage is slowed down by decreasing the pH. In a representative reaction, stage I was carried out under pH of 8.3, and then the reaction solution was adjusted to pH 6.1 by addition of HNO3 at the end of stage I. Once the pH value was adjusted to 6.1, the reduction rate was decelerated as expected, as shown in the temporal evolution of Ag+ concentration (Figure 5a). The UV-visible absorption spectra (Figure 5b) reveal the difference between the resulting silver nanoparticles obtained by the stepwise reduction method and the particles obtained by the typical one-step reduction method. The integrated absorbance in the 700-900-nm window almost disappeared and the fwhm become very narrow for the particles prepared by the stepwise method. TEM images show that the product of the stepwise reaction is mainly dominated by the spherical silver nanoparticles (Figure 5c). The average size of the particles was determined to be 56 nm, with a standard deviation of (8%. Both the rod-like particles in the product synthesized under pH 8.3 (Figure 2b) and the anisotropic silver nanoparticles with triangle or polygon shape synthesized at pH 6.1 (Figure 2c) are suppressed greatly in the stepwise reaction. Control experiments in which stage I was initiated under different pH values (6.9, 7.7, 9.0) and stage II carried out at pH 6.1 were further conducted to check the stepwise method. TEM observations and UV-visible absorption spectra indicate that all three stepwise reactions show improved shape control over the particles compared with the typical one-step citrate reduction reactions carried out under the corresponding initial pH (Figure 6), indicating the improved balance between nucleation and growth processes in the stepwise reactions. It should be mentioned that there are still some nonspherical particles (such as ellipsoids or some other irregular shapes)

6576 J. Phys. Chem. C, Vol. 113, No. 16, 2009 synthesized by the stepwise method, and the relative amounts of the nonspherical particles in parts a, b, and c of Figure 6 were estimated to be about 35%, 12%, and 21%, respectively. The average particle sizes and standard deviations of the three samples were 71 nm ((15%), 58 nm ((11%), and 56 nm ((26%) respectively for the reactions carried out at initial pH 6.9, 7.7, and 9.0. It is noted that the stepwise reaction carried out at the initial pH value of 7.7 presents better shape control ability than the ones carried at the initial pH values of 9.0 and 6.9, meaning a better balance between the nucleation and growth processes in the reaction. Compared to the reported methods for the synthesis of spherical silver nanoparticles, such as the laser irradiation method,16 such a stepwise method represents a facile way to have control over the shape of the silver nanoparticles. Conclusion In summary, in the citrate reduction reaction, the pH of the reaction solution affects the size/shape evolution of the resulting silver nanoparticles due to the pH-dependent reduction activity of citrate and thus the pH dependent reduction rate of the precursor (Ag+). The product of the reaction carried out under high pH is composed of both spherical and rod-like silver nanoparticles as a result of fast reduction of the Ag+ precursor. Anisotropic silver nanoparticles with triangle or polygon shape tend to occur in the reaction carried out under low pH conditions due to the slow nucleation and growth processes. With this understanding, a stepwise reduction method is developed to improve the balance between the nucleation and growth, in which the fast nucleation stage is carried out at high pH value and the following growth stage is slowed down by decreasing the pH value. The shape control of spherical silver nanoparticles is improved greatly due to the improved balance between the nucleation and growth processes. It is expected that more precise adjustment on the balance between nucleation and growth will allow better control over the shape of silver nanoparticles. Acknowledgment. This work was supported by the National Natural Science Foundation of China (Nos. 20773053, 20803029, 50825202), the National Research Fund for Fundamental Key Project (2009CB939701), and the Program for NCET in the University of Chinese Ministry of Education. We thank Prof. X. Peng for his helpful suggestions and discussion. Supporting Information Available: The standard curve for the determination of silver ion concentration, and a discussion on the effect of pH value on the activity of the citrate. This material is available free of charge via the Internet at http:// pubs.acs.org.

Dong et al. References and Notes (1) Daniel, M. C.; Astruc, D. Chem. ReV. 2004, 104, 293. (2) Xiong, Y. J.; Wiley, B.; Xia, Y. N. Angew. Chem., Int. Ed. 2007, 46, 7157. (3) Lewis, L. N. Chem. ReV. 1993, 93, 2693. (4) Cao, Y. W.; Jin, R.; Mirkin, C. A. J. Am. Chem. Soc. 2001, 123, 7961. (5) Taton, T. A.; Mirkin, C. A.; Letsinger, R. L. Science 2000, 289, 1757. (6) Yang, Y.; Nogami, M.; Shi, J. L.; Chen, H. R.; Ma, G. H.; Tang, S. H. Appl. Phys. Lett. 2006, 88, 081110. (7) Munro, C. H.; Smith, W. E.; Garner, M.; Clarkson, J.; White, P. C. Langmuir 1995, 11, 3712. (8) Tominaga, J.; Mihalcea, C.; Bu¨chel, D.; Fukuda, H.; Nakano, T.; Atoda, N.; Fuji, H.; Kikukawa, T. Appl. Phys. Lett. 2001, 78, 2417. (9) Tao, A. R.; Habas, S.; Yang, P. D. Small 2008, 4, 310. (10) Wiley, B.; Sun, Y. G.; Xia, Y. N. Acc. Chem. Res. 2007, 40, 1067. (11) Jin, R. C.; Cao, Y. W.; Mirkin, C. A.; Kelly, K. L.; Schatz, G. C.; Zheng, J. G. Science 2001, 294, 1901. (12) Wu, X. M.; Redmond, P. L.; Liu, H. T.; Chen, Y. H.; Steigerwald, M.; Brus, L. J. Am. Chem. Soc. 2008, 130, 9500. (13) Murphy, C. J.; Jana, N. R. AdV. Mater. 2002, 14, 80. (14) Khanal, B. P.; Zubarev, E. R. J. Am. Chem. Soc. 2008, 130, 12634. (15) Pastoriza-Santos, I.; Liz-Marza´n, L. M. J. Mater. Chem. 2008, 18, 1724. (16) Pyatenko, A.; Yamaguchi, M.; Suzuki, M. J. Phys. Chem. C 2007, 111, 7910. (17) Suber, L.; Sondi, I.; Matijevic´, E.; Goia, D. V. J. Colloid Interface Sci. 2005, 288, 489. (18) Frens, G. Nature (London), Phys. Sci. 1973, 241, 20. (19) Ji, X. H.; Song, X. N.; Li, J.; Bai, Y. B.; Yang, W. S.; Peng, X. G. J. Am. Chem. Soc. 2007, 129, 13939. (20) Turkevich, J.; Stevenson, P. C.; Hillier, J. Discuss. Faraday Soc. 1951, 11, 55. (21) Lee, P. C.; Meisel, D. J. Phys. Chem. 1982, 86, 3391. (22) Henglein, A.; Giersig, M. J. Phys. Chem. B 1999, 103, 9533. (23) LaMer, V. K.; Dinegar, R. H. J. Am. Chem. Soc. 1950, 72, 4847. (24) El-Ghamry, M. T.; Frei, R. W. Anal. Chem. 1968, 40, 1986. (25) Kreibig, U.; Vollmer, M. Optical Properties of Metal Clusters; Springer Series in Materials Science 25; Springer: Berlin, Germany, 1995. (26) Kelly, K. L.; Coronado, E.; Zhao, L. L.; Schatz, G. C. J. Phys. Chem. B 2003, 107, 668. (27) Noguez, C. J. Phys. Chem. C 2007, 111, 3806. (28) Gonza´lez, A. L.; Reyes-Esqueda, J. A.; Noguez, C. J. Phys. Chem. C 2008, 112, 7356. (29) Tao, A.; Sinsermsuksakul, P.; Yang, P. D. Angew. Chem., Int. Ed. 2006, 45, 4597. (30) Velikov, K. P.; Zegers, G. E.; van Blaaderen, A. Langmuir 2003, 19, 1384. (31) Kumar, S.; Gandhi, K. S.; Kumar, R. Ind. Eng. Chem. Res. 2007, 46, 3128. (32) Kumbhar, A. S.; Kinnan, M. K.; Chumanov, G. J. Am. Chem. Soc. 2005, 127, 12444. (33) Evanoff, D. D.; Chumanov, G. J. Phys. Chem. B 2004, 108, 13948. (34) Hoang, T. K. N.; La, V. B.; Deriemaeker, L.; Finsy, R. Langmuir 2002, 18, 10086. (35) Peng, Z. A.; Peng, X. G. J. Am. Chem. Soc. 2001, 123, 1389.

JP900775B