Reaction and Crystallization Mechanism of Potassium Dititanate

CRYSTAL GROWTH & DESIGN

Reaction and Crystallization Mechanism of Potassium Dititanate Fibers Synthesized by Low-Temperature Calcination

2005 VOL. 5, NO. 4 1399-1404

Chang Liu,# Ming He,§ Xiaohua Lu,*,§ Qitu Zhang,# and Zhongzi Xu# College of Chemical Engineering and College of Materials Science and Engineering, Nanjing University of Technology, Nanjing 210009, P. R. China Received November 27, 2004;

Revised Manuscript Received March 3, 2005

ABSTRACT: The reaction and crystallization mechanism of K2Ti2O5 fibers from amorphous titania and K2CO3 by a calcination method at low temperatures are investigated in this study. K2Ti2O5 nuclei are formed in the solidstate reaction stage at 640 °C due to the suppression of TiO2 crystal growth in the presence of K2CO3. K2Ti2O5 nanofibers are obtained in the crystallization stage at 820 °C due to the appearance of K2O melts. The separation between the nucleation and the fiber growth process is critical to control the morphologies and microstructures of K2Ti2O5 fibers at low temperatures. Introduction Potassium titanate fibers of the chemical formula K2O‚nTiO2 (n ) 2, 4, 6, 8) are a class of one-dimensional inorganic materials whose diameters range from 1 to 10 microns and lengths can be up to several hundred microns. Fibers of potassium titanate with n ) 6 and 8, i.e., K2Ti6O13 and K2Ti8O17, have been used as reinforcement materials for plastics and metals and frictional materials for brakes due to their stable tunnel structures and chemical stability.1 Fibers of potassium titanate with n ) 4, i.e., K2Ti4O9, possess a zigzag layered structure and have been used as inorganic ion exchangers to remove and immobilize radioactive nuclides such as 137Cs and 90Sr from high-level liquid wastes.2,3 Fibers of potassium titanate with n ) 2, i.e., K2Ti2O5, also possess a layered structure. However, its layered structure is not stable in water or acid solution. Thus, there have been few research reports about K2Ti2O5 fibers since they were synthesized first in 1960.4 Recently, it has been found that K2Ti2O5 exhibits the highest catalytic activity and photoluminescence, even at room temperature, compared to that of other layered titanates.5,6 Our previous works have shown that mesoporous fibrous titania with a large specific surface area7 and H2Ti8O17 nanorods with a high photocatalytic activity8 can be prepared using K2Ti2O5 fibers as precursors. These properties and applications have a direct relationship to morphologies, sizes, water contents, and crystallinities of K2Ti2O5. K2Ti2O5 fibers are usually prepared by calcination of potassium carbonate and anatase or rutile. Calcination is a traditional method for preparation of one-dimensional inorganic material that is simple and can be used on a large scale. Usually, a relatively high temperature is required for the formation of fibers, and particles but no fiber products are synthesized below 1000 °C.9 In our recent work, the synthesis temperature for K2Ti2O5 * To whom correspondence should be addressed. [email protected]. # College of Materials Science and Engineering. § College of Chemical Engineering.

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fibers was decreased sharply by using an amorphous titania instead of anatase or rutile as the reactant.10-13 Mesoporous fibrous titania with a large special area had been prepared by the calcination of the amorphous titania and potassium carbonate at about 800 °C followed by a hydrolytic and a thermal decomposition process.7 However, K2Ti2O5 fibers prepared by lowtemperature calcination have poor crystallinity and imperfect fibrous morphology, compared with that prepared by conventional calcination. Therefore, it is necessary to study the formation mechanism of K2Ti2O5 fibers prepared by the low-temperature calcination. In the present work, the reaction and crystallization process of K2Ti2O5 fibers using an amorphous titania as the precursor are investigated, and the formation mechanism of K2Ti2O5 fibers is discussed. Theory Figure 1 shows an idealized crystal structure of K2Ti2O5.14 The titanium atom is bonded to five oxygen atoms in a distorted trigonal bipyrimid coordination. These trigonal bipyrimids form an endless string by edge sharing, and the strings are linked by having one oxygen atom in common. When combined in this way, they form an infinite two-dimensional sheet with a composition of (Ti2O5)2-. The layered sheets are subsequently held together by ionic interactions with potassium ions. Figure 2 shows a canonical X-ray diffraction pattern with the input structural parameters for K2Ti2O5. It can be seen that the intensities of the (001) peak is only one-third of that of the (111) peak in the calculated diffraction pattern. However, it often exhibits some preferred orientation due to the layered structure, and the relative peak intensities in the powder diffraction pattern deviate from the intrinsic relative diffraction intensities. The TiO5 unit layers are parallel to the (001) face. Therefore, the relative intensities of the (001) peak will be increased, and the relative intensities of other peaks will be depressed. To explore the preferred orientation of samples, a method that was proposed by Leventouri15 is used in this work. With this method, the coefficient R0 of the March-Dollase function16 is used

10.1021/cg049602a CCC: $30.25 © 2005 American Chemical Society Published on Web 05/12/2005

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Figure 1. Idealized crystal structure of K2Ti2O5.

Figure 3. The calculated intensity ratios I(001)/I(111) as a function of the March coefficient R0 for K2Ti2O5.

Experimental Methods

Figure 2. Calculated diffraction patterns for K2Ti2O5, assuming a completely random orientation of the crystallites in the powdered sample.

to generate a calibration curve for estimating the degree of preferred orientation in samples from conventional X-ray diffraction patterns. The form of the MarchDollase equation is shown in eq 1:

[

Icorr ) Iobs R02 cos2 γk +

]

sin2 γk R0

-(3/2)

(1)

where Icorr is the intensity as corrected for the preferred orientation with the measured intensity, Iobs is the measured intensity, γk is the acute angle between the scattering vector of reflection k and the orientation direction, and R0 is the March coefficient that ranges from 0 to 1, where zero refers to a 100% preferentially oriented crystallites in a specimen, and 1 refers to a completely random orientation of its crystallites. For textured samples of K2Ti2O5, the relative intensities of (001) peak is the strongest among the reflections. Thus, the ratios of the intensities I(001) to the intensities I(111) of the characteristic Bragg peak (111) are calculated in a software “materials studio” (version 2.0) and plotted in Figure 3. A least-squares fit to the data yields eq 2:

I(001) I(111)

) 0.762R0-0.2497

(2)

Equation 2 can be considered as a relationship between the relative intensity of the (111) Bragg peak and the degree of preferred orientation. This relation will be used to estimate the preferred orientation in samples heated at different temperatures.

General Procedure. The starting material, K2CO3, was commercially available with purities of 99.5%. TiO2‚nH2O (hydrous titanium dioxide) was prepared by hydrolyzing TiOSO4 in hot water with vigorous stirring. Deionized water was added to the mixture of TiO2‚nH2O and K2CO3. The chemical composition (TiO2/K2O molar ratio) was controlled at 2. The mixing processes were preformed by ball milling with water. Final mixtures were dried in an oven at 90 °C for 10 h. Calcinations were performed in a muffle furnace at a heating rate of 5 °C/min. The heating temperatures were 640 and 820 °C. The heating time was prolonged up to 8 h. The heated samples were quenched by quickly removing them from the furnace and immediately cooling them to room temperature; they were then ground into a fine powder to perform the X-ray diffraction (XRD) analysis. Characterizations. Thermal analysis was carried out on a Perkin-Elmer TGA7 under CO2 atmosphere from 25 to 1000 °C at a heating rate of 10 °C/min. X-ray powder diffraction patterns were obtained using a D8 advance (Bruker AXS, Karlsruhe, Germany). Cu KR radiation with a nickel filter was used, operating at 40 kV and 30 mA. All samples were measured in the continuous scan mode at 5-40° (2θ) with a scanning rate of 0.05°/s. Raman spectra were collected using a JY Horiba LabRam HR800 microRaman spectrograph at room temperature. A 200 mW laser emitting at 488 nm was used as the excitation source. The laser power is 1-2 mW on the sample. The samples used for TEM observation were prepared by dispersing some products in distilled water followed by mild ultrasonic vibration for half an hour, then dripping a drop of the dispersion onto a copper grid coated with a layer of amorphous carbon. A JEM-200CX TEM was used to study the morphologic characteristics and structure.

Results and Discussions Figures 4 and 5 show the results of X-ray diffraction patterns for anatase-K2CO3 and TiO2‚nH2O-K2CO3 samples quenched at different heating times during isothermal heating at 640 °C. When the anatase-K2CO3 sample was quenched at 640 °C for 1 h, the main phases were the raw materials, anatase and K2CO3, and a weak (111) peak of product K2Ti2O5 appeared in the XRD pattern, indicating that the solid-state reaction between anatase and K2CO3 is so slow in this condition that more reaction time is needed. As the heating time was prolonged to 8 h, the main phase was

Reaction and Crystallization of Potassium Dititanate

Figure 4. X-ray powder diffraction patterns of anataseK2CO3 heated at 640 °C for different times: (9) K2Ti2O5; (0) K2CO3; (4) anatase; (2) intermediate.

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Figure 6. X-ray powder diffraction patterns of TiO2‚nH2OK2CO3 heated at 820 °C for different times: (9) K2Ti2O5.

Figure 7. TEM image of TiO2‚nH2O-K2CO3 heated at 640 °C for 8 h. Table 1. Effect of Precursor and Thermal Treatment on the March Coefficient R0 of K2Ti2O5

Figure 5. X-ray powder diffraction patterns of TiO2‚nH2OK2CO3 heated at 640 °C for different times: (9) K2Ti2O5.

K2Ti2O5, and K2CO3 was still found in the XRD pattern. However, when the TiO2‚nH2O-K2CO3 sample was quenched at 640 °C for 1 h, only K2Ti2O5 was observed and the (111) peak was the strongest in the XRD pattern. As the heating time was prolonged to 8 h, the K2Ti2O5 diffraction lines became sharper and stronger, and the intensity of the (001) peak became as strong as that of the (111) peak. This indicates that it is difficult for anatase-K2CO3 samples to react completely at 640 °C; however, for the TiO2‚nH2O-K2CO3 sample, the reaction progresses quickly. Figure 6 shows the results of XRD patterns for TiO2‚ nH2O-K2CO3 samples quenched at different heating times during isothermal heating at 820 °C. When the sample was quenched at 820 °C for 1 h, only K2Ti2O5 was observed in the XRD pattern, and it could be noticed that the diffraction lines were similar with that sample quenched at 640 °C for 8 h. As the heating time was prolonged to 8 h, the (001) peak became the strongest in the XRD pattern, and the intensity of the (111) peak was only about one second that of the (001) peak. It

thermal treatment precursor anatase-K2CO3 TiO2‚nH2O-K2CO3 TiO2‚nH2O-K2CO3 TiO2‚nH2O-K2CO3 TiO2‚nH2O-K2CO3

temperature (°C) soaking time (h) 640 640 640 820 820

8 1 8 1 8

R0 0.932 0.933 0.730 0.749 0.610

means that the process at 820 °C for TiO2‚nH2O-K2CO3 samples is only a growth process of K2Ti2O5 fibers, and the process can be studied with the extent of preferred orientation. The results of the extent of preferred orientation for different samples are listed in Table 1. It can be seen that the March coefficient R0 of the anatase-K2CO3 sample heated at 640 °C for 8 h is close to 1, indicating that the orientation directions of crystallites in the sample are random. After calcination for 8 h, the March coefficient R0 of TiO2‚nH2O-K2CO3 samples heated at 640 °C is decreased to 0.73, while R0 of TiO2‚nH2O-K2CO3 samples heated at 820 °C is decreased to 0.61. The magnified pictures of samples heated at 640 and 820 °C for 8 h, taken after the removal of the soluble matters with HCl solution, are shown in Figures 7-9. It can be seen that TiO2‚nH2O-K2CO3 samples heated

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Figure 8. TEM image of TiO2‚nH2O-K2CO3 heated at 820 °C for 8 h.

at 640 °C for 8 h were composed of nanocrystals, and TiO2‚nH2O-K2CO3 samples heated at 820 °C for 8 h were composed of nanofibers with diameters ranging from 30 to 40 nm and lengths up to 400 nm. It should be noted that the samples in XRD measurements are calcination products without washing. However, the samples in TEM images are calcination products after washing in water. Our previous work has shown that the K2Ti2O5 fibers obtained by calcination of K2CO3 and TiO2‚nH2O are a kind of single crystal.7,12 Therefore, these tiny fibers will show a stronger preferred orientation in X-ray diffraction patterns than those crystal grains which have been confirmed partly by the observations from the XRD patterns. However, the anataseK2CO3 samples heated at 820 °C for 8 h showed some irregular aggregated crystals. This indicates that for TiO2‚nH2O-K2CO3 samples the reaction between TiO2‚ nH2O and K2CO3 below 640 °C is a solid-state reaction in which the nuclei of K2Ti2O5 are formed, and the reaction process around 820 °C is a growth process for K2Ti2O5 fibers. These results are consistent with the observations from the XRD patterns. It can be seen that the coarsening of K2Ti2O5 nanocrystals is negligible in Figure 7, which can be ascribed to the existence of K2CO3 in TiO2‚nH2O-K2CO3 samples. TiO2‚nH2O is gel-like and exhibits a long-range order like crystalline anatase. TiO2‚nH2O has a specific surface area of approximately 300 m2/g,17 and the size of the crystallite particles in TiO2‚nH2O is about 510 nm. In thermal treatment, the primary particles are organized to larger aggregates, and the growth of TiO2 nano- and microcrystals proceeds by coalescence with neighboring crystals, i.e., in the same way that several drops of a liquid coagulate to one larger drop. However, when K2CO3 is added to TiO2‚nH2O, the number of contacts between primary particles are increased by necks of salt, and, in consequence, the number of direct contacts between crystals is reduced. In addition, the aggregated particles in TiO2‚nH2O can be dissolved to primary particles due to the higher pH values.18 Figure 10 and Figure 11 show the TG-DTA curves for anatase-K2CO3 and TiO2‚nH2O-K2CO3 mixtures under N2 atmosphere, respectively. In the TG trace of TiO2‚nH2O and K2CO3 mixtures, a continuous weight loss with various slopes was observed at 100-840 °C. This means that the dehydration of TiO2‚nH2O and the solid-state reaction between K2CO3 and TiO2 primary particles occurred during calcination, and the crystal growth of

Liu et al.

Figure 9. TEM image of anatase-K2CO3 heated at 820 °C for 8 h.

Figure 10. TG-DTA curves for anatase and K2CO3 mixtures under N2 atmosphere.

Figure 11. TG-DTA curves for TiO2‚nH2O and K2CO3 mixtures under N2 atmosphere.

TiO2 primary particles are suppressed. Therefore, when TiO2‚nH2O-K2CO3 samples are heated at 640 °C, nanoparticles of K2Ti2O5 are obtained. However, in the TG trace of anatase and K2CO3 mixtures, the TGA trace after the first weight loss step attributed to the decrease of free water was horizontal at 220-500 °C, and no endothermic peak was observed on the DTA trace between 220 and 500 °C. This indicates that no reaction occurred between anatase and K2CO3 at T < 500 °C. Therefore, when anatase-K2CO3 samples are heated at 640 °C, the particle aggregates will be obtained. Our recent work has shown that the starting reaction temperature between TiO2‚nH2O and K2CO3 is lower

Reaction and Crystallization of Potassium Dititanate

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Figure 12. Raman spectra of TiO2‚nH2O and K2CO3 mixtures heated at 820 °C for 8 h.

Figure 13. TG-DTG curves for TiO2‚nH2O and K2CO3 mixtures under CO2 atmosphere.

than that between anatase and K2CO3.10-13 It can be seen from Figures 10 and 11 that for the anatase and K2CO3 mixtures the weight loss due to the release of CO2 occurred above 500 °C, indicating that the starting reaction temperature is above 500 °C. However, in the TG trace of TiO2‚nH2O and K2CO3 mixtures, a continuous weight loss with various slopes was observed between 200 and 840 °C, and most of the weigh loss occurred below 640 °C.

that K2Ti2O5 was the major component, and peaks at 190, 284, 459, 533, 648, and 898 cm-1 are assigned to K2Ti2O5.19 However, one can find a weak peak at 1061 cm-1 assigned to K2CO3.20 This indicates that a small amount of residual K2CO3 exists after calcination at 820 °C for 8 h. Considering the melting point of K2CO3 is 891 °C, the endothermic peak at 820 °C in the DTA curve for TiO2‚nH2O and K2CO3 mixtures seems not to be due to the melt of residual K2CO3. To understand the growth process at 820 °C for TiO2‚nH2O and K2CO3 mixtures, the TG-DTG curves for TiO2‚nH2O and K2CO3 mixtures under CO2 atmosphere was investigated, which can be seen in Figure 13. The weight loss processes under CO2 atmosphere occurred in three stages. The first weight loss stage below 158 °C was ascribed to the decrease of free water and hydrate water. The second weight loss stage at 158-608 °C was due to the solid-state reaction between TiO2 and K2CO3, and the temperature range was similar to that of the weight loss stage at 200-640 °C under N2 atmosphere, indicating that the solid-state reaction between TiO2 and K2CO3 is not affected by CO2 partial pressure in the gaseous phase. However, the third weight loss stage at 640-820 °C under N2 atmosphere shifted to 866 °C under CO2 atmosphere, indicating that it is a decomposition process for K2CO3 and the product K2O will be melted when the temperature is over its melting point 790 °C. Therefore, the endothermic peak at 820 °C in DTA curve for TiO2‚nH2O and K2CO3 mixtures under N2 atmosphere is due to the melt of K2O. Now it is clear that when amorphous titania is used as a precursor to prepare K2Ti2O5 fibers, the reaction between potassium carbonate and amorphous titania below 640 °C is ascribed to solid-state reactions and the nuclei of K2Ti2O5 are formed, and the reaction process around 820 °C is ascribed to the K2Ti2O5 fiber growth process due to the appearance of the K2O melt. The separation between the nucleation and the fiber growth process is critical to control the morphologies and microstructures of K2Ti2O5 fibers. However, when anatase is used as the precursor to obtain K2Ti2O5, the fiber growth process is accompanied by a solid-state reaction process. The coarsening of crystals is favorable,21 and the morphology of the product is irregular crystal aggregates, which can be seen in Figure 9.

K2CO3 + 2TiO2 (anatase) f K2Ti2O5 + CO2 (3) K2CO3 + 2TiO2‚nH2O f K2Ti2O5 + CO2 + 2nH2O (4) In our recent studies, The Gibbs free energy changes for eqs 3 and 4 were calculated to study the influence of titania reactants on the lowest generation temperature of potassium dititanate by calcination. The Gibbs free energy changes are lower than zero at T > 508 °C for eq 3 and at T > 295 °C for eq 4, respectively, indicating that estimated starting generation temperatures for preparing potassium dititanate from anatase and TiO2‚nH2O are 508 and 295 °C, respectively.12 Thus, the solid-state reaction at 640 °C for TiO2‚nH2O and K2CO3 mixtures is easy to perform nearly completely, which is in accordance with the results from the XRD patterns. Until now, the detail of the growth process at 820 °C for TiO2‚nH2O and K2CO3 mixtures was not very clear. A weak endothermic peak appeared at 820 °C on the DTA curve of Figure 11, and for the anatase and K2CO3 mixtures the endothermic peak appeared at 860 °C on the DTA curve of Figure 12. Generally, the growth mechanisms for fibers include vapor-solid (VS), liquid-solid (LS), and vapor-liquid-solid (VLS) mechanisms, etc., and a liquid or vapor phase is usually necessary for fiber growth. Therefore, the endothermic peak would be due to the melt of a certain material. From the XRD pattern for TiO2‚nH2O-K2CO3 samples at 820 °C only K2Ti2O5 was observed; however, the melting point of K2Ti2O5 is about 930 °C, which can be seen on the DTA curve of Figure 11. Thus, the endothermic peak at 820 °C is not due to the melt of K2Ti2O5. Raman spectra of TiO2‚nH2O-K2CO3 samples heated at 20 °C for 8 h are shown in Figure 12. It can be seen

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Conclusion The formation mechanism of K2Ti2O5 fibers prepared with amorphous titania and K2CO3 at low temperatures is investigated in this study. The growth and aggregation of K2Ti2O5 nucleus are prevented due to the existence of K2CO3 in TiO2‚nH2O-K2CO3 samples, and K2Ti2O5 nanofibers can be obtained at 820 °C. The reaction between potassium carbonate and amorphous titania below 640 °C is ascribed to the solid-state reaction process, and the reaction process around 820 °C is ascribed to the K2Ti2O5 fiber growth process due to the appearance of the K2O melt. The separation between the reaction and the crystallization process is critical to control the morphologies and microstructures of K2Ti2O5 fibers at low temperatures. Acknowledgment. Financial support from Chinese National Science Foundation for Outstanding Young Scholars (No. 29925616) and Chinese National Natural Science Foundation (No. 20236010, No. 20246002) are gratefully acknowledged. References (1) Lee, C. T.; Um, M. H.; Kumazawa, H. J. Am. Ceram. Soc. 2000, 83, 1098. (2) Sasaki, T.; Watanabe, M.; Komatsu, Y.; Fujiki, Y. Inorg. Chem. 1985, 24, 2265.

Liu et al. (3) Sasaki, T.; Komatsu, Y.; Fujiki, Y. Inorg. Chem. 1989, 28, 2776. (4) Andersson, S.; Wadsley, A. D. Nature 1960, 187, 499. (5) Kudo, A.; Sakata, T. J. Mater. Chem. 1993, 3, 1081. (6) Kudo, A.; Kaneko, E. Chem. Commun. 1997, 4, 349. (7) He, M.; Lu, X. H.; Feng, X.; Yu, L.; Yang, Z. H. Chem. Commun. 2004, 19, 2202. (8) Bao, N. Z.; Feng, X.; Yang, Z. H.; Shen, L. M.; Lu, X. H. Environ. Sci. Technol. 2004, 38, 2729. (9) Fujiki, Y.; Ohsaka, T. J. Ceram. Soc. Jpn. 1982, 90, 19. (10) Bao, N. H.; Feng, X.; Shen, L. M.; Lu, X. H. Cryst. Growth Des. 2002, 2, 437. (11) Bao, N. Z.; Feng, X.; Lu, X. H.; Yang, Z. H. J. Mater. Sci. 2002, 37, 3035. (12) Bao, N. Z.; Feng, X.; Lu, X. H.; Shen, L. M.; Yanagisawa, K. AIChE J. 2004, 50, 1568. (13) Bao, N. Z.; Shen, L. M.; Feng, X.; Lu, X. H. J. Am. Ceram. Soc. 2004, 87, 326. (14) Andersson, S.; Wadsley, A. D. Acta Chem. Scand. 1961, 15, 663. (15) Leventouri, T. Physica C 1997, 277, 82. (16) Dollase, W. A. J. Appl. Crystallogr. 1986, 19, 267. (17) Gesenhues, U. Chem. Eng. Technol. 2001, 24, 685. (18) Jalava, J.-P.; Hiltunen, E.; Kahkonen, H.; Erkkila, H.; Harma, H.; Taavitsainen, V.-M. Ind. Eng. Chem. Res. 2000, 39, 349. (19) Bamberger, C. E.; Begun, G. M.; Macdougall, C. S. Appl. Spectrosc. 1990, 44, 30. (20) Bates, J. B.; Boyd, G. E.; Brooker, M. H.; Quist, A. S. J. Phys. Chem. 1972, 76, 1565. (21) Penn, R. L. J. Phys. Chem. B 2004, 108, 12707.

CG049602A