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Synthesis and characterization of near-monodisperse yttria particles by homogeneous precipitation method
Powder Technology 142 (2004) 136 – 153 www.elsevier.com/locate/powtec
Synthesis and characterization of near-monodisperse yttria particles by $ homogeneous precipitation method Seungman Sohn a,*, Youngshik Kwon b, Yeunsik Kim c, Dongsu Kim d a
Chemical Engineering, Virginia Tech., Blacksburg, VA 24061, USA Environmental Engineering, Suwon Science College, San 9 – 10, Hwasung, Kyounggido 445-742, South Korea c Materials Science and Engineering, Seoul National University, Shinlim, Kwan-ak, Seoul 151-172, South Korea d Environmental Science and Engineering, Ewha Womans University, Daehyun, Seodaemun, Seoul 120-750, South Korea b
Received 19 October 2003; received in revised form 10 March 2004; accepted 20 March 2004 Available online
Abstract A wide application of yttria particles for advanced ceramics has attracted a number of studies concerning the preparation of colloidal yttria particles. In the present work, near-monodisperse yttria particles are synthesized in urea aqueous solution by a homogeneous precipitation method. The effect of four experimental variables, such as concentrations of yttrium and urea, reaction temperature, and solution pH were characterized in terms of morphology and reaction kinetics. It was found that yttrium concentration, varying between 0.005 and 0.04 M, has a profound impact on the average size of particles, which systematically increases from 65 nm to over 220 nm. We also found that as yttrium concentration increases, not only does the size distribution broaden, but particles also start to agglomerate above 0.025 M of concentration, mainly due to the reduction of zeta potential. The rate of precipitation reaction, however, is shown to be independent of yttrium concentration. In contrast, as urea concentration increases from 0.04 to 4.0 M, the average particle size exhibits a gradual decrease from ca. 220 to ca. 100 nm. At extremely high urea concentration such as 7.0 M, a significant level of inter-particle agglomeration is observed. The rate of precipitation is found to increase with urea concentration up to 3.0 M. Above 3.0 M, the concentration dependence is weakened. Temperature mainly affects the kinetics of precipitation, not thermodynamic quantities such as average particle size, which is a strong function of the concentrations of reactants (both yttrium and urea). Assuming Arrhenius-type reaction kinetics, the activation energy for precipitation is obtained as 29 kcal/mol. This value is quite comparable with the activation energy of urea decomposition (28 – 32 kcal/mol). Based on this, we propose that urea decomposition may be a rate-determining step in the formation of yttria particles by homogeneous precipitation method. Particle morphology and the reaction kinetics are also sensitive to solution pH. While at low pH ( < 2.0) particles experience severe agglomeration, and the rate of precipitation is slow. As pH increases above 3.0, near-monodisperse yttria particles are obtained. Lastly, based on two observations that, first, pH does not change during precipitation (maintaining pH 5.5 – 6.0), and second, the equilibrium constant of yttrium ion hydrolysis at this pH (i.e., K=[Y(OH)2 +]/[Y3 +]) is too small to serve as a reasonable source of hydroxyl ions in the final yttria compound, we propose, along with some experimental supports, a new reaction mechanism that is more plausible in yttria particle precipitation from urea aqueous solution. D 2004 Elsevier B.V. All rights reserved. Keywords: Yttria particles; Near-monodisperse homogeneous precipitation; Urea; Particles size; Reaction kinetics
1. Introduction Rare earth metals and their compounds (mostly oxides) were found to be versatile as key materials for high $
This work is a part of Dr. Youngshik Kwon’s Ph.D. dissertation. * Corresponding author. Current address: E-Ink Corporation, 733 Concord Avenue, Cambridge, MA 02138, USA. Tel.: +1-617-499-6180; fax: +1-617-499-6200. E-mail address: [email protected] (S. Sohn). 0032-5910/$ - see front matter D 2004 Elsevier B.V. All rights reserved. doi:10.1016/j.powtec.2004.03.013
temperature applications, catalysts, fluorescence, laser applications, and electronic applications such as memory chips and superconductors [1 – 10]. Among many rare earth oxides, the most widely used are zirconium, yttrium, titanium, and cerium. Zirconia (ZrO2) has an important role due to its excellent chemical resistance, refractory character, ionic conductivity and polymorphic nature [3 –5]. Yttria (Y2O3) has a high melting point (2698 K) with excellent chemical stability and low volatility in vacuum. Because of these outstanding properties, yttria has been of interest for
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high temperature applications such as alloying elements and additives for advanced ceramics such as Si3N4 and AlN [6 – 8]. Cerium oxide (CeO2) has also attracted important applications such as catalysts and grinding materials for optical glasses [9,10]. To improve the efficiency during the sintering process, which significantly affects the final material’s properties, ultra fine size (in the order of submicron) and a narrow size distribution are required. Various methods have been proposed to prepare fine and near-monodisperse metal oxide particles [11 – 21]. These methods may largely be classified into three categories, such as solid, liquid and vapor methods, depending on the phase of materials involved. Solid-phase methods include thermal dissociation [11,12] and mechanical milling [13,14]. Liquid-phase methods are based on homogeneous precipitation (HP) [15,16], hydrolysis [17,18], and sol – gel techniques [19]. Lastly, vapor condensation, vapor– vapor, vapor – liquid and vapor –solid reactions are typical examples of vapor-phase methods [20,21]. Among these methods, it is generally agreed that liquid-phase methods are suitable to precisely control chemical compositions. Furthermore, the HP method has an advantage in preparation of fine particles with relatively narrow size distribution. Some fundamentals of the HP method follow. When a certain type of organic compound, such as urea, slowly dissolves in a solution containing metal cations, the generated anions, with metal cations, form precipitation nuclei above a critical point of super-saturation. This decreases the anion concentration in the solution, and thereafter, the generated anions are only consumed by the growth of nuclei. This leads to a controlled separation of nucleation and growth of precipitates. By adjusting the concentration of reactants, solution pH, and reaction temperature, we can control the concentration of nuclei (i.e., control the point of super-saturation), which eventually leads to a control of shape, size and size distribution of precipitates. For instance, under a constant concentration of metal ions, a greater number of nuclei or higher critical point of super-saturation will lead to the smaller precipitates. The first use of the HP method in the preparation of metal oxide particles was reported in 1937 by Williard and Tang [22]. They synthesized near monodisperse aluminum basic sulfate in reaction with aluminum and urea utilizing the slow decomposition of the latter into ammonia and carbon dioxide. They mostly studied the effect of pH and the types of anions on the precipitates. In 1950, rare earth metal oxide was first synthesized by Salutsky and Quill who obtained carbonates of Lanthanum, Neodymium and Samarium through the hydrolysis of rare earth trichloroacetates (e.g., Nd(C2Cl3O2)3) in a homogeneous phase reaction yielding a pure crystalline carbonate (e.g., Nd2(CO3)3) [23]. Later on, the importance of HP in rare earth oxide particle preparation further attracted more in-depth studies; among them, the contributions from Matijevic [15,17,21,24,25] and Akinc [26] may be distinguishable. They studied the HP
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method based on urea decomposition to prepare various rare earth metal oxides such as zirconia and yttria. As results, some important properties of particles such as chemical composition, isoelectric point (IEP), and chemical steps in thermal dissociation were characterized. In addition, they exercised their efforts in understanding the effects of urea concentration, the type and concentration of anions, reaction time on the particle shape, size and its distribution. Although, these pioneering studies unveiled some of the characteristics of yttria particles precipitation using the HP method, a more systematic study is necessary to understand the complete picture of the effects of several processing variables. In the present work, we will show an extensive study on the synthesis and characterization of yttria particles prepared from the HP method. Our primary goals are to understand the effects of yttrium and urea concentrations, reaction temperature, and initial solution pH on particle morphology (e.g., shape, size and size distribution) and reaction kinetics. Based on these understandings, we propose a new reaction mechanism, which is more consistent with the experimental observations over previously proposed mechanism by Matijevic et al. [25]. As a last part of this study, we investigated the effect of calcination on the change of morphology of yttria particles.
2. Experiments 2.1. Materials Yttrium oxide (Y2O3; purity>99.9%) was purchased from Strem Chemicals (USA), and both nitric acid (HPLC grade) and urea ((NH2)2CO) were provided by Dongyang Chemicals (Seoul, Korea). Y(NO3)3 standard solution was prepared as follows. First, to remove any trace of moisture and other organic impurities, yttrium oxide was calcinated at 850 jC for several hours. Subsequently, it was dissolved in nitric acid, and the remaining solvent was removed by prolonged heating above the solvent boiling point. Y(NO3)3 thus obtained was further diluted using distilled water adjusting the solid concentration to 0.5 M with pH 6.0, and was used as a standard solution. Standard aqueous urea solutions 2.0 and 5.0 M were prepared and stored in a refrigerator (below 5 jC). Although urea is known to decompose in an aqueous solution, the rate of decomposition at this relatively low temperature was so slow that the temporal change of initial urea concentration could be neglected (the measured decomposition amount of urea after 7 days at 5 jC was about 7 10 4%). Urea standard solutions were either used or discarded within a week. Urea solutions with higher concentration than 5.0 M were always prepared right before the experiment to minimize the concentration variation.
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2.2. Preparation of yttria particles The previous two solutions having appropriate concentrations, Y(NO3)3 and urea, were mixed and then filtered through a filter paper (#41, Whatman). About 25 ml of mixed solution was placed in a glass tube, and this was inserted in an isothermal bath maintaining a desired reaction temperature. In these experiments, the effects of concentration of reactants, reaction temperature, solution pH on the morphology and kinetics of yttria particle formation were studied. After a complete reaction, the precipitates were separated, thoroughly washed with ethanol, and further dried at 100 jC overnight. Whenever necessary, a fraction of solution was extracted and the unreacted yttrium concentration in the solution was determined by titration method using 0.001 M EDTA (ethylenediaminetetraacetic acid). 2.3. Characterization of yttria particles The morphology of yttria particles was examined by transmission electron microscopy (TEM: JEM-200CX, JEOL) and scanning electron microscopy (SEM: JSM-35, JEOL). Ethanol was used to disperse the particles, and either carbon or gold was surface coated for TEM and SEM measurements, respectively. In this study, the main purpose of using TEM is to check the shape and the size of precipitates. As will be detailed in Results and discussion, particle shape and size are strong functions of the concentrations of yttrium and urea, temperature, and solution pH. To avoid any complications from kinetic effects, all the micrographs were taken at the terminal stage of each reaction. Thus we can safely assume that the precipitates were full grown in each set of experiment. The only exception would be the experiment showing the effect of time on the growth of precipitates. Particle size and distribution estimated from TEM micrographs were quantitatively confirmed using a particle size analyzer (Mastersizer, Malvern). The Zeta potential of the particles was monitored by Autosizer-III (Malvern). To understand the thermal behavior of yttria particles, differential scanning calorimetry (DSC: DSC7, Perkin Elmer), and thermogravimetric analysis (TGA: TGA-7, Perkin Elmer) were utilized. Carbon content in the particles was determined by a carbon –hydrogen – nitrogen analyzer (MT2, Yanaco).
3. Results and discussion Yttria particles are generated through a well-known reaction that yttrium cations (Y3 +) precipitate with anions such as CO32 or OH from the dissolution of urea [25]. In brief, yttrium cations precipitate as follows: Y3 + + OH + CO32 + 1.5 H2O = Y(OH)CO31.5 H2O. See Eqs. (9) – (15) for complete details of reaction schemes. Accordingly, the concentrations of yttrium ions and urea will significantly
affect the shape and size distribution of yttria particles. In addition, reaction temperature and initial solution pH will influence on the final products. In this section we will evaluate the effects of these variables. 3.1. The effect of yttrium concentration A series of TEM micrographs in Fig. 1 presents the effect of yttrium concentration on the shape and size of yttria particles. All other experimental conditions, such as the concentration of urea, solution pH, and reaction temperature, were all fixed as given in the figure caption. Two important observations can be made. First, as anticipated, the effect of yttrium concentration on particle size is significant, varying from 65 to 140 nm in the concentration range between 0.005 and 0.025 M. Fig. 2 shows the average size variation as a function of yttrium concentration. The error bar inserted in each data point serves as a standard deviation of particle size distribution. Below 0.025 M particles do not agglomerate, the particles size was nearmonodisperse (10 – 18% of standard deviation in size distribution). Severe agglomeration, however, at higher concentrations above 0.03 M decreased the reliability of measurement so the two data points denoted as open circles are only approximations with much broader size distributions. Second, within the range of yttrium concentration studied, the shape of precipitates is a well-rounded sphere. This may indicate that particles precipitate through a nucleation and growth mechanism, in which nucleation density is a strong function of yttrium concentration. Thus below a critical concentration not to invoke an impingement of growth front, we anticipate spherical-shaped precipitates. At higher yttrium concentration, however, we may observe non-spherical type particles due to frequent impingements. Indeed, a close comparison of particle shapes between Fig. 1c and g reveals that the shape of precipitates at higher concentrations is more irregular. Figs. 3 and 4 present time-dependent morphology changes in two different yttrium concentrations, 0.02 and 0.04 M, which represent below and above a critical concentration of 0.025 M for agglomeration. TEM micrographs were all taken in 10-min intervals in each case. At initial stage of precipitation (10 min), well-rounded spheres of size on the order of 60 – 100 nm evolve in both concentrations. As time proceeds, the precipitates in 0.02 M solution grow maintaining the shape of spheres, yet the particles in 0.04 M solution start agglomerate and form irregularly shaped conglomerates. Temporal average particle size variation in Figs. 3 and 4 are shown in Fig. 5. The open square symbol in Fig. 5 denotes the high uncertainty in size measurement due to severe agglomeration (in case of 0.04 M). In conjunction with Fig. 2, Fig. 5 suggests that the average particle size in the final stage of reaction matches well with the previous results within experimental uncertainty (e.g., compare Figs. 1d and 3c). This strongly indicates that not only is the
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control of particle size reproducible, but also, time can be an effective factor to manipulate the size of precipitates. The observation that, for a given time of experiment, higher yttrium concentration leads to bigger particles is in full accordance with our notion: under a fixed concentration of urea and constant level of supersaturation for nucleation, more yttrium ions are available for the growth of precipitates. The existence of agglomeration, however, is not expected due to a particle size effect. Experimentally, we could not observe particle agglomeration in the case of 0.02 M concentration even at much longer times. To address this issue properly, we measured the zeta potential of precipitates as a function of yttrium concentration, and the result is shown in Fig. 6. At lower concentrations up to 0.02 M, the zeta potential is above 13 mV, however, as concentration increases, potential undergoes a rapid drop and reaches a plateau value around 2 – 3 mV. This is most likely because as [Y(NO3)3] increases, the higher concen-
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tration of anions, such as NO3, lead to a decrease of zeta potential, which, in turn, will facilitate the agglomeration of precipitates. From a study of homogeneous yttria precipitation, Matijevic et al. reported without specifying the mechanism that there exists an upper limit of yttrium concentration to obtain mono-disperse yttria precipitates [25]. In the present work, we observed a similar phenomenon; furthermore, we suggest that the decrease of zeta potential with yttrium concentration causes particle agglomeration, and this is perhaps one of the reasons for the existence of the upper [Y(NO3)3] concentration to obtain near-monodisperse yttria precipitation. These findings suggest that any forced dispersions using techniques such as ultrasonic or mechanical disturbance may be of help to increase the critical upper yttrium concentration by avoiding particle – particle agglomeration. Indeed, we found that the particle shape and size distribution became more spherical and narrow when the precipitates were made under the influence of ultrasonic dispersion even in
Fig. 1. A schematics of experimental set up for the production of yttria oxides particles. 1. Water bath; 2. Heating element; 3. Isothermal temperature controller; 4. Thermocouple; 5. Reaction tube; 6. Mechanical stirrer; 7. Water. TEM micrographs of yttrium compound prepared under different Y(NO3)3 concentrations as (a) 0.005 M, (b) 0.01 M, (c) 0.015 M, (d) 0.02 M, (e) 0.025 M, (f) 0.03 M, and (g) 0.04 M (in this case, urea concentration and pH were fixed at 1.0 and 6.0, respectively, and reaction temperature was 95 F 1 jC).
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Fig. 1 (continued).
Fig. 2. Variation of mean particle size as a function of yttrium concentration. The error bars are based on standard deviation of size distribution in each concentration. The open circle represents approximated average particle size in the presence of agglomeration. Note that, upon agglomeration, not only the size distribution broadens, but also the mean value deviates from the linear extrapolation (short-dashed line).
0.03 M yttrium solution. Fig. 7 shows SEM micrographs of homogeneous yttria precipitation with and without ultrasonic dispersion. It is apparent that ultrasonic dispersion effectively reduces the degree and the size of agglomeration. A similar result was previously reported [27]. The ultrasonic dispersion effect, however, was less distinctive as the yttrium concentration further increased to 0.04 M. It should be also noted that ultrasonic dispersion does not affect the average size and size distribution below 0.02 M concentration. In Fig. 8, the effect of yttrium concentration on the rate of precipitation is presented. A fraction of solution was extracted at appropriate time intervals, and the unreacted yttrium concentration was analyzed by titration method using 0.001 M EDTA (i.e., the amount of precipitation in Fig. 8 is given as the difference between the initial and the unreacted yttrium concentration). Solutions with three different concentrations (0.01, 0.02 and 0.03 M) were used, and other experimental conditions were controlled. Fig. 8 clearly tells us that, within the concentration range between 0.01 and 0.03 M, the rate of precipitation is virtually a constant, being equal to 6.56 10 4 mol/ l min. In addition, as shown in Fig. 9, we can define the
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Fig. 3. TEM micrographs showing the morphology change of yttrium compound prepared for different reaction time as (a) 10 min, (b) 20 min, (c) 30 min, and (d) 40 min (in this case, [Y(NO3)3] and [urea] were, respectively, 0.02 and 1.0 M. Solution pH was 6.0, and the reaction temperature was 95 F 1 jC).
necessary time for a complete reaction from the plateau region in Fig. 8. The 100% reaction time, ranging between ca. 20 and 60 min, linearly increases with the yttrium concentration. These results, combined with Fig. 2 showing the variation of average size with concentration, are of practical importance to control the particle size and its distribution. 3.2. The effect of urea concentration To study the effect of urea on shape, size, and the distribution of yttria precipitates, the urea concentration, [urea], was systematically varied between 0.04 and 7.0 M, while other conditions were controlled. A series of TEM micrographs in Figs. 10 and 11 presenting the variation of average particle size clearly show that an increase of urea concentration decreases the average particle size, from ca. 220 nm ([urea] = 0.04 M) to 105 nm ([urea] = 4.0 M). Particle size almost linearly decreases with the logarithm of urea concentration (see the insert in Fig. 11). At the highest concentration studied, particles experience such an excessive degree of agglomeration that the morphology of
precipitates appears to be a network structure, and individual particles cannot be distinguished (thus 7.0 M data point is omitted in Fig. 11). The shape of precipitates was spherical between 0.04 and 4.0 M of urea concentration. The inverse relationship between the urea concentration and the average particle size may be explained as follows. The higher the urea concentration, the greater the amount of anions generated, such as CO32 and OH, which will lead to a higher degree of supersaturation before the onset of precipitation. The amount of urea decomposition with urea concentration linearly increased, in which the rate of decomposition was 1.3 10 3/min. In this case, we expect an increase of nucleation density; therefore, under a given yttrium concentration, the average particle size will decrease. At very high urea concentration, such as 7.0 M, an excessively high level of anions will lead to a high nucleation density and a fast growth of precipitates, which contribute to form (even continuous) aggregates. Once again, these observations reinforce the theorem in homogeneous precipitation that particle shape and size control must be understood in terms of the balance between nucleation and growth.
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Fig. 4. TEM micrographs showing the morphology change of yttrium compound prepared for different reaction time as (a) 10 min, (b) 20 min, (c) 30 min, and (d) 40 min (in this case, [Y(NO3)3] and [urea] were, respectively, 0.04 and 1.0 M. Solution pH was 6.0, and the reaction temperature was 95 F 1 jC).
Fig. 5. Variation of mean particle size with reaction time for two different yttrium concentrations below and above a critical concentration for agglomeration. The data point represented by an open square is only an approximated value under severe agglomeration.
Fig. 6. Variation of zeta potential of the precipitates formed under different yttrium concentration.
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Fig.9. The time for complete reaction as a function of yttrium concentration.
Fig. 7. SEM micrographs of yttrium compound prepared without (a) and with (b) ultrasonic disturbance (in this case, [Y(NO3)3] and [urea] were, respectively, 0.03 and 1.0 M. Solution pH was 6.0, and the reaction temperature was 95 F 1 jC).
Fig. 8. Temporal variation of the precipitated amount of yttrium for three different yttrium concentrations (in this case, urea concentration was 1.0 M. Solution pH was 6.0, and the reaction temperature was 95 F 1 jC).
The effect of urea concentration on the rate of precipitation is shown in Fig. 12, in which the temporal variation of the amount of precipitation is obtained similarly as shown in Fig. 8. One immediate observation is that as urea concentration increases, precipitation rate accelerates. This is in good contrast with the results in Fig. 8, where yttrium concentration did not increase the rate of precipitation but only affected the time for complete reaction. Urea concentration strongly influences both the rate of precipitation (Fig. 13) and the time for complete reaction (Fig. 14). The rate of precipitation, as defined by the slope of each curve in Fig. 12, increases fast with urea concentration up to 3.0 M (Regime I), and reaches a plateau value (Regime II). The behavior in Regime I may be readily understood from the fact that as the concentration of anions from urea decomposition increases, the precipitation of yttria particles accelerates. However, the existence of Regime II deserves further explanation. For this, first we need to understand the equilibrium reactions involved with anions such as CO32 and OH[28]. CO2 3 þ H2 O ¼ HCO3 þ OH
ð1Þ
HCO 3 þ H2 O ¼ H2 CO3 þ OH
ð2Þ
H2 CO3 ¼ CO2 þ H2 O
ð3Þ
In a system where the concentration of yttrium cation (Y3 +) is fixed, the excessive amount of (CO32 ) over the balance of the reaction stoichiometry will be consumed through the sequential reactions as shown above. If this hypothesis were true we would expect an increase of solution pH at high urea concentrations. An increase in solution pH may be equally possible at long reaction times at which the precipitation reaction is completed and thus all the anions generated must be consumed via above reactions.
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The variation of solution pH for selected systems is inserted in Fig. 12. At two lower concentrations, solution pH was invariant (5.7 F 0.1) until the precipitation reaction is completed. In case of [urea] = 1.0 M, the later stage pH increase around 35 –40 min must be ascribed to the consumption of anions after the reaction is completed (see also Fig. 14 showing that the time for complete reaction in this case occurs around 38 min). On the other hand, at high urea concentration such as 4.0 M, even at the first measurement done at 5 min, the time well before the reaction is complete (18 minutes in Fig. 14), the solution pH is already higher than 6.0, and keeps increasing afterwards. These observations are in good agreement with the above hypothesis, and may explain the existence of Regime II in that excessive supply of carbonyl anions from urea decomposition will be consumed through 1) reactions (2) reactions (3) and do not contribute to form yttria precipitates. The time for complete reaction in Fig. 14 has another noticeable point. At low concentrations up to 1.0 M, the reaction time, t, reduces following an exact inverse relationship between t and [urea] (i.e., t~1/[urea]). This
matches well with Fig. 13 showing a linear increase of the rate of precipitation in low urea concentrations. At high concentrations above 3.0 M, the reaction time does not reduce appreciably because the concentrations of anions such as (CO32 ) are already saturated. Combined with Fig. 11, Fig. 14 will be an important map to control the yttria precipitation process. 3.3. The effect of reaction temperature Similar to other phase transition behaviors, such as crystallization, involving nucleation and growth mechanism, the rate of yttria precipitation will critically depend on reaction temperature. To quantify this effect, reaction temperature was varied between 70 and 95 jC. Other experimental conditions were kept constant ([Y(NO3)3]:0.02 M; [urea]:1.0 M; pH: 6.0). TEM micrographs in Fig. 15 show the morphology of yttria particles formed under different reaction temperatures. The variation of average particle size as a function of temperature is presented in Fig. 16. Within the temperature range studied, precipitates maintained well-
Fig. 10. TEM micrographs of yttrium compound prepared under different urea concentrations as (a) 0.04 M, (b) 0.1 M, (c) 0.5 M, (d) 1.0 M, (e) 2.0 M, (f) 4.0 M, and (g) 7.0 M (in this case, yttrium concentration and pH were fixed at 0.02 and 6.0 M, respectively, and reaction temperature was 95 F 1 jC).
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Fig. 10 (continued).
rounded spherical shape, and the average size and the size distribution did not vary significantly. Particle size showed a slight decrease with temperature, yet the overall variation, from 135 nm (70 jC) to 126 nm (95 jC), was within the resolution of measurements. In contrast to relatively insignificant effect on particle size, we found that temperature has a profound impact on the kinetics of precipitation. In Fig. 17, the amount of precipitation is given with reaction temperature. At 70 jC, precipitates form very slowly, and even after 3 h the total amount of precipitation is about 25% of theoretical limit (i.e., in this case, 0.02 M). As temperature increases, the rate of precipitation increases significantly; at 95 jC, the reaction finishes around 40 min. Fig. 17 in combination with Fig. 16 confirms that reaction temperature predominantly affects the growth kinetics of precipitates not the thermodynamic factors such as an average particle size, which is governed by the concentrations of yttrium and urea. Fig. 18 presents the rate of precipitation, defined by the slope of each line in Fig. 17, as a function of reaction temperature. The rate exponentially increases more than an
order of magnitude within 25 jC temperature window. The decomposition rate of urea is known to exhibit similar exponential dependency in the same temperature range, and the literature data has been reproduced in Fig. 18 [29]. This indicates that a strong temperature dependence of the rate of precipitation mainly comes from an increased dissolution rate of urea at high temperature. From the slope of each line in Fig. 17, we can also deduce an activation energy for yttria precipitation by plotting lnK versus 1/T, in which K is a rate constant, and T is an absolute temperature (i.e., use of Arrhenius-type temperature dependence: K = Koexp( Ea/RT), where Ea is an activation energy). The result is inserted in Fig. 18. The calculated activation energy is ca. 29 kcal/mol. It is interesting to note that, according to Shaw and Bordbaux [29], the activation energy for urea decomposition in aqueous solution is in the range of 28 – 32 kcal/mol. The close match between the two activation energies may suggest that urea decomposition step is a rate-determining step in the formation of yttria particles from a homogeneous precipitation. In Fig. 19, the reaction time for complete reaction as a function of temperature is presented. The complete reaction
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Fig. 11. Variation of mean particle size as a function of urea concentration. The insert shows a linear decrease of the mean particle size with the logarithm of urea concentration.
time at each temperature is defined by the intercept of corresponding line to 0.02 M as shown in Fig. 17. Over 800 min of reaction time at 70 jC dramatically decreases to 40 min at 95 jC. Recall that the wide range of reaction rate at different reaction temperature, however, does not affect the particle size obtained at the terminal stage.
3.4. The effect of initial solution pH It is well known that above 7.0 pH, yttrium ions (Y3 +) precipitates with OHanion in the form of Y(OH)3 when [Y3 +] is in the range of 0.1 – 0.01 M (reaction constant, K=[Y3 +][OH]3 = 6.0 10 24 at 25 jC) [28].
Fig. 12. Temporal variation of the precipitated amount of yttrium for six different urea concentrations (in this case, yttrium concentration was 0.02 M. Solution pH was 6.0, and the reaction temperature was 95 F 1 jC). The insert shows time-dependent variation of solution pH for three urea concentrations. For details of the correlation between the precipitation amount and solution pH, see text.
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of yttrium ions (Y3 +) (Eq.(6)), and (3) precipitation reaction (Eq. (7)). Urea decomposition ðNH2 Þ2 CO ¼ NHþ 4 þ OCN
ð4Þ
OCN þ 2Hþ þ H2 O ¼ CO2 þ NHþ 4
ð5Þ
Hydrolysis Y3þ þ H2 O ¼ ½YðOHÞ 2þ þ Hþ
ð6Þ
Precipitation ½YðOHÞ 2þ þ CO2 þ 2H2 O ¼ YðOHÞCO3 H2 O þ 2Hþ ð7Þ Overall, Fig. 13. Precipitation rate change as a function of urea concentration. For the explanation of Regime-I and -II, see text.
Y3þ þ ðNH2 Þ2 CO þ 4H2 O ¼ YðOHÞCO3 H2 O þ þ 2NHþ 4 þH
Therefore, in the present study, pH = 7.0 is will be an upper limit to obtain yttria particles through homogeneous precipitation. In this section we investigated the effect of solution pH below 7.0. TEM micrographs in Fig. 20 show the morphology of yttria particles under varying pH’s between 1.3 and 6.0. It is clear that below pH 2.0, particles experience severe agglomeration, and only above pH 3.0, spherically shaped near-monodisperse yttria particles are obtained. The effect of pH on the size and distribution of particles was not significant when pH is in the range of 3.0 and 6.0. Although we did not measure the zeta potential of precipitates, we suspect that the origin of agglomeration at low pH may be ascribed to the decrease of zeta potential in the presence of a higher concentration of counter anions (NO3) which compress the electrical double layer and facilitate the agglomeration of particles. Similar to the previous sections, we also measured the amount of precipitation with time, from which we could deduce the time for complete reaction and the rate of precipitation, as shown in Fig. 21a and b, respectively. These two additional data confirm that the effect of pH on the kinetics of precipitation is not distinguishable at pH 3.0 or higher. As a reason for the decreased reaction kinetics at low pH, we suggest that at significantly lower pH than ca. 5.5 –6.0, which is required for the onset of precipitation (see Fig. 12, for instance), it takes longer time to reach this pH through the decomposition of urea.
ð8Þ
These chemical equilibriums lead to an overall reaction that one mole of urea reacts with one mole of yttrium ions and forms an equi-molar yttria precipitates. In addition, as reaction proceeds, the solution pH will decrease due to the generation of protons. Contrary to the previous report, we found that solution pH remains constant during the precipitation (see, for instance, Fig. 12) performed under wide different experimental conditions, varying reactant concentrations (both yttrium and urea), and temperature. In addition to this apparent mismatch, we also found that the equilibrium constant, K, of the hydrolysis reaction in Eq. (6) is too small to believe that it is a main source supplying the hydroxyl groups in the overall reaction in Eq. (8) (e.g., at
3.5. Reaction mechanism According to Matijevic et al., the reaction mechanism of the formation of yttria particles through homogeneous precipitation is claimed to include three distinct steps [25]: (1) decomposition of urea (Eqs. (4) and (5)), (2) hydrolysis
Fig. 14. Time for complete reaction as a function of urea concentration.
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Fig. 15. TEM micrographs of yttrium compound prepared under different reaction temperatures as (a) 70 jC, (b) 80 jC, (c) 90 jC and (d) 95 jC (in this case, [Y(NO3)3] and [urea] were, respectively, 0.02 and 1.0 M. Solution pH was 6.0).
pH 6.0, [Y(OH)]2 +/[Y3 +] = K 106 = 10 8.04 106 c0.01, thus only about 1% of yttrium ions exist in the form of yttrium hydroxide ions) [30].
To resolve these contradictions, we suggest new routes of chemical reactions that are compliant to the experimental observations. Urea Decomposition ðNH2 Þ2 CO ¼ NHþ 4 þ OCN
ð9Þ
þ OCN þ 2H2 O ¼ CO2 3 þ NH4
ð10Þ
CO2 3 þ H2 O ¼ HCO3 þ OH
ð11Þ
HCO 3 þ H2 O ¼ H2 CO3 þ OH
ð12Þ
H2 CO3 ¼ CO2 ðgÞ þ H2 O
ð13Þ
Precipitation Y3þ þ OH þ CO2 3 þ 1:5H2 O ¼ YðOHÞCO3 1:5H2 O ð14Þ Overall, Fig. 16. Variation of mean particle size with reaction temperature. The large error bars show approximately 10 – 13% of standard deviation involved in multiple measurements.
Y
3þ
þ 1:5ðNH2 Þ2 CO þ 5H2 O
¼ YðOHÞCO3 1:5H2 O þ 3NHþ 4 þ 1=2CO2 ðgÞ
ð15Þ
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Fig. 17. Temporal variation of the precipitated amount of yttrium for five different reaction temperatures (in this case, [Y(NO3)3] and [urea] were, respectively, 0.02 and 1.0 M. Solution pH was 6.0).
According to newly proposed mechanism, 1.5 mol of urea reacts with 1 mol of yttrium ions to form 1 mol of yttrium compound (Y(OH)CO31.5H2O). This contrasts with the previous mechanism by Matijevic et al. [25] in that the amount of consumed urea per mole of yttrium ions is 50% larger, and accordingly, the reaction will produce, as a by-product, 3 mol of NH4+ions rather than 2 mol (compare Eqs. (8) and (15)). To compare the validity of the reaction mechanisms proposed in Eq. (15) over Eq. (8), we measured the amount of NH4+as a function of reaction time and compared it with
149
Fig. 19. Time for complete reaction as a function of reaction temperature.
the precipitated amount of yttrium ions. This result will support, at least indirectly, which reaction mechanism is more feasible in the explanation of experimental observations. A brief experimental procedure follows. Sixteen glass tubes with 25 ml of solution ([(Y(NO 3 ) 3 ) = 0.02 M, [urea] = 1.0 M, and pH = 6.0 ) were placed in a water bath at 95 jC. The amounts of NH4+ and yttrium ions were measured in every 5 min. Two samples were used for each measurement. It is well known that, at room temperature, ammonium ions (NH4+) are stable up to pH = 7.0, and as pH increases they gradually transform to ammonia (NH3) [30]. At pH 12.0 or higher only ammonium species are stable.
Fig. 18. Precipitation rate change as a function of reaction temperature (closed circles). Open squares show the variation of rate constant in urea decomposition from the literature [25]. Note that how these two different quantities exhibit similar temperature dependence. The insert presents Arrhenius-type plot, from which activation enrgy for yttria precipitation can be calculated.
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Fig. 20. TEM micrographs of yttrium compound prepared under different solution pHs as (a) 1.3, (b) 2.0, (c) 3.0, (d) 4.0 and (e) 6.0 (in this case, [Y(NO3)3] and [urea] were, respectively, 0.02 and 1.0 M and reaction temperature was 95 F 1 jC).
Based on this thermodynamic equilibrium, ammonium ions were first changed to ammonia at pH = 12.0 and then determined using ammonia gas sensing electrode (NH31501, Phoenix, USA). Fig. 22 compares the temporal variation of the precipitated amounts between yttrium ions and ammonium ions. As proposed in the present study, the temporal variation of ammonium ion concentration quantitatively matches with three times of yttrium ions. The
inserted line showing twice the amount of [Y3 +] is based on the previous study by Matijevic et al. [25]. Another difference between the Eqs. (8) and (15) lies in the number of water molecules contained in yttrium compounds. It is largely accepted that general chemical composition of yttria precipitates is Y(OH)CO3xH2O, which further changes to Y(OH)CO3 at 200 jC, and finally to Y2O3 above 640 jC [25,26]. In this work we carefully followed
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tion of precipitates. Upon calcination, the average size decreases about 20% with slight broadening. The average size of as-prepared particles was ca. 170 nm, which reduces to 135 nm after calcination. A similar result was reported previously [25]. From the known densities of yttrium compound [31] and yttrium oxide [32], which are, respectively, ca. 3.9 and ca. 5.0 g/cm3, we can expect about 23% size reduction after calcination. This theoretical size decrease is in good match with the actual observation.
4. Summary An emerging interest of utilizing rare earth metals and their compounds in advanced ceramics has driven many recent studies. As a part of these efforts, we studied a homogeneous precipitation method in urea aqueous solution as a route to synthesize near-monodisperse yttria particles. Among others, the concentrations of yttrium and urea, reaction temperature, and solution pH were chosen for the most significant processing variables affecting both morphology of particles and reaction kinetics. Based on this, we conclude as follows.
Fig. 21. Precipitation rate change (a) and time for complete reaction (b) as a function of solution pH.
(1) Yttrium concentration, within the range studied (0.005 M and 0.04 M), increased the average size of particles from 65 nm to over 220 nm. Yttrium concentration is also involved with the broadening of size distribution. Above 0.025 M particles start to coagulate, which is ascribed to the reduction of zeta potential with yttrium concentration. The rate of precipitation reaction is shown to be independent of yttrium concentration.
the mass changes upon thermal decomposition at two different temperatures of 250 and 850 jC using TGA (oxygen purging rate: 1.0 l/min), and we confirmed that, indeed, the exact chemical composition of yttrium precipitates is Y(OH)CO31.5H2O. These two independent experimental evidences support the validity of the newly proposed mechanism in homogeneous precipitation of yttria particles in the presence of urea. 3.6. Calcination of yttrium compound As noted earlier, the precipitated yttrium compound undergoes chemical change to yttrium oxide (Y2O3) above 640 jC. It is important to understand the morphological changes upon calcinations, since one of the target applications of yttrium particles includes the utility as an alloying element for advanced ceramics through sintering. TEM micrographs in Fig. 23 contrast the morphology before (a and b) and after (c and d) calcinations performed at 850 jC for 2 h. One immediate observation is that, surface roughness increases after calcination. This is mainly due to the evaporation of chemically bonded water and CO2 during heating. Calcination also affects the size and size distribu-
Fig. 22. Quantitative comparison between the amounts of yttrium ions (open squares) involved in precipitation reaction and ammonium ions (open circle) formed as a by-product. An excellent match between these two quantities strongly supports the validity of newly proposed reaction mechanism (see Eq. (15) and text for details). Triangles show twice the amount of yttrium ions used in precipitation based on the reaction mechanism from the literature [25].
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Fig. 23. Morphological change of yttria compound before (a and b) and after (c and d) calcination performed at 850 jC for 2 h (reaction conditions were [Y(NO3)3] = 0.02 M, [urea] = 1.0 M, pH = 6.0, and temperature of 95 F 1 jC).
(2) From the study of urea concentration ranging from 0.04 to 4.0 M, we found that the average particle size exhibits a gradual decrease from ca. 220 to ca. 100 nm. As similarly observed previously, extremely high urea concentration such as 7.0 M, caused inter-particle coagulations. The rate of precipitation, which increases with concentration up to 3.0 M of urea, starts to level off at higher concentrations. (3) Temperature is shown to significantly increase the rate of precipitation; however, thermodynamic quantities, such as average particle size, is only a function of the concentrations of reactants (both yttrium and urea). Based on Arrhenius-type reaction kinetics, the activation energy for precipitation is found to be 29 kcal/mol. This value is in good match with the activation energy of urea decomposition (28 –32 kcal/mol). We suggest that this similarity may be an indication that urea decomposition is a rate-determining step in the formation of yttria particles in homogeneous precipitation method.
(4) Solution pH is also proven to be a controlling factor affecting the particle morphology and the reaction kinetics. Severe coagulation that occurs at low pH ( < 2.0) deters to obtain well-separated individual yttria particles, yet as pH increases above 3.0, near-monodisperse yttria particles can be achieved. The rate of precipitation showed a similar trend as coagulation: at low pH the rate is slow, and as pH increases the rate enhances reaching a plateau value above pH 3.0. From the study of reaction mechanism, we observed two important features. First, pH does not change during precipitation (maintaining pH 5.5– 6.0), and second, the equilibrium constant of yttrium ion hydrolysis (i.e., K= [Y(OH)2 +]/[Y3 +]) at this pH range is too small to serve as a reasonable source of hydroxyl ions in the final yttrium compound. Based on these two findings, we propose a new reaction mechanism that is more plausible in yttria particle precipitation from urea aqueous solution. As supportive experimental evidence, we found a quantitative match
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between the amount of ammonium ions generated and the precipitated yttrium ions proposed in the new reaction mechanism. [16] [17]
Acknowledgements Editorial work by Mr. Derek Klinedinst is appreciated.
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Synthesis and characterization of near-monodisperse yttria particles by $ homogeneous precipitation method Seungman Sohn a,*, Youngshik Kwon b, Yeunsik Kim c, Dongsu Kim d a
Chemical Engineering, Virginia Tech., Blacksburg, VA 24061, USA Environmental Engineering, Suwon Science College, San 9 – 10, Hwasung, Kyounggido 445-742, South Korea c Materials Science and Engineering, Seoul National University, Shinlim, Kwan-ak, Seoul 151-172, South Korea d Environmental Science and Engineering, Ewha Womans University, Daehyun, Seodaemun, Seoul 120-750, South Korea b
Received 19 October 2003; received in revised form 10 March 2004; accepted 20 March 2004 Available online
Abstract A wide application of yttria particles for advanced ceramics has attracted a number of studies concerning the preparation of colloidal yttria particles. In the present work, near-monodisperse yttria particles are synthesized in urea aqueous solution by a homogeneous precipitation method. The effect of four experimental variables, such as concentrations of yttrium and urea, reaction temperature, and solution pH were characterized in terms of morphology and reaction kinetics. It was found that yttrium concentration, varying between 0.005 and 0.04 M, has a profound impact on the average size of particles, which systematically increases from 65 nm to over 220 nm. We also found that as yttrium concentration increases, not only does the size distribution broaden, but particles also start to agglomerate above 0.025 M of concentration, mainly due to the reduction of zeta potential. The rate of precipitation reaction, however, is shown to be independent of yttrium concentration. In contrast, as urea concentration increases from 0.04 to 4.0 M, the average particle size exhibits a gradual decrease from ca. 220 to ca. 100 nm. At extremely high urea concentration such as 7.0 M, a significant level of inter-particle agglomeration is observed. The rate of precipitation is found to increase with urea concentration up to 3.0 M. Above 3.0 M, the concentration dependence is weakened. Temperature mainly affects the kinetics of precipitation, not thermodynamic quantities such as average particle size, which is a strong function of the concentrations of reactants (both yttrium and urea). Assuming Arrhenius-type reaction kinetics, the activation energy for precipitation is obtained as 29 kcal/mol. This value is quite comparable with the activation energy of urea decomposition (28 – 32 kcal/mol). Based on this, we propose that urea decomposition may be a rate-determining step in the formation of yttria particles by homogeneous precipitation method. Particle morphology and the reaction kinetics are also sensitive to solution pH. While at low pH ( < 2.0) particles experience severe agglomeration, and the rate of precipitation is slow. As pH increases above 3.0, near-monodisperse yttria particles are obtained. Lastly, based on two observations that, first, pH does not change during precipitation (maintaining pH 5.5 – 6.0), and second, the equilibrium constant of yttrium ion hydrolysis at this pH (i.e., K=[Y(OH)2 +]/[Y3 +]) is too small to serve as a reasonable source of hydroxyl ions in the final yttria compound, we propose, along with some experimental supports, a new reaction mechanism that is more plausible in yttria particle precipitation from urea aqueous solution. D 2004 Elsevier B.V. All rights reserved. Keywords: Yttria particles; Near-monodisperse homogeneous precipitation; Urea; Particles size; Reaction kinetics
1. Introduction Rare earth metals and their compounds (mostly oxides) were found to be versatile as key materials for high $
This work is a part of Dr. Youngshik Kwon’s Ph.D. dissertation. * Corresponding author. Current address: E-Ink Corporation, 733 Concord Avenue, Cambridge, MA 02138, USA. Tel.: +1-617-499-6180; fax: +1-617-499-6200. E-mail address: [email protected] (S. Sohn). 0032-5910/$ - see front matter D 2004 Elsevier B.V. All rights reserved. doi:10.1016/j.powtec.2004.03.013
temperature applications, catalysts, fluorescence, laser applications, and electronic applications such as memory chips and superconductors [1 – 10]. Among many rare earth oxides, the most widely used are zirconium, yttrium, titanium, and cerium. Zirconia (ZrO2) has an important role due to its excellent chemical resistance, refractory character, ionic conductivity and polymorphic nature [3 –5]. Yttria (Y2O3) has a high melting point (2698 K) with excellent chemical stability and low volatility in vacuum. Because of these outstanding properties, yttria has been of interest for
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high temperature applications such as alloying elements and additives for advanced ceramics such as Si3N4 and AlN [6 – 8]. Cerium oxide (CeO2) has also attracted important applications such as catalysts and grinding materials for optical glasses [9,10]. To improve the efficiency during the sintering process, which significantly affects the final material’s properties, ultra fine size (in the order of submicron) and a narrow size distribution are required. Various methods have been proposed to prepare fine and near-monodisperse metal oxide particles [11 – 21]. These methods may largely be classified into three categories, such as solid, liquid and vapor methods, depending on the phase of materials involved. Solid-phase methods include thermal dissociation [11,12] and mechanical milling [13,14]. Liquid-phase methods are based on homogeneous precipitation (HP) [15,16], hydrolysis [17,18], and sol – gel techniques [19]. Lastly, vapor condensation, vapor– vapor, vapor – liquid and vapor –solid reactions are typical examples of vapor-phase methods [20,21]. Among these methods, it is generally agreed that liquid-phase methods are suitable to precisely control chemical compositions. Furthermore, the HP method has an advantage in preparation of fine particles with relatively narrow size distribution. Some fundamentals of the HP method follow. When a certain type of organic compound, such as urea, slowly dissolves in a solution containing metal cations, the generated anions, with metal cations, form precipitation nuclei above a critical point of super-saturation. This decreases the anion concentration in the solution, and thereafter, the generated anions are only consumed by the growth of nuclei. This leads to a controlled separation of nucleation and growth of precipitates. By adjusting the concentration of reactants, solution pH, and reaction temperature, we can control the concentration of nuclei (i.e., control the point of super-saturation), which eventually leads to a control of shape, size and size distribution of precipitates. For instance, under a constant concentration of metal ions, a greater number of nuclei or higher critical point of super-saturation will lead to the smaller precipitates. The first use of the HP method in the preparation of metal oxide particles was reported in 1937 by Williard and Tang [22]. They synthesized near monodisperse aluminum basic sulfate in reaction with aluminum and urea utilizing the slow decomposition of the latter into ammonia and carbon dioxide. They mostly studied the effect of pH and the types of anions on the precipitates. In 1950, rare earth metal oxide was first synthesized by Salutsky and Quill who obtained carbonates of Lanthanum, Neodymium and Samarium through the hydrolysis of rare earth trichloroacetates (e.g., Nd(C2Cl3O2)3) in a homogeneous phase reaction yielding a pure crystalline carbonate (e.g., Nd2(CO3)3) [23]. Later on, the importance of HP in rare earth oxide particle preparation further attracted more in-depth studies; among them, the contributions from Matijevic [15,17,21,24,25] and Akinc [26] may be distinguishable. They studied the HP
137
method based on urea decomposition to prepare various rare earth metal oxides such as zirconia and yttria. As results, some important properties of particles such as chemical composition, isoelectric point (IEP), and chemical steps in thermal dissociation were characterized. In addition, they exercised their efforts in understanding the effects of urea concentration, the type and concentration of anions, reaction time on the particle shape, size and its distribution. Although, these pioneering studies unveiled some of the characteristics of yttria particles precipitation using the HP method, a more systematic study is necessary to understand the complete picture of the effects of several processing variables. In the present work, we will show an extensive study on the synthesis and characterization of yttria particles prepared from the HP method. Our primary goals are to understand the effects of yttrium and urea concentrations, reaction temperature, and initial solution pH on particle morphology (e.g., shape, size and size distribution) and reaction kinetics. Based on these understandings, we propose a new reaction mechanism, which is more consistent with the experimental observations over previously proposed mechanism by Matijevic et al. [25]. As a last part of this study, we investigated the effect of calcination on the change of morphology of yttria particles.
2. Experiments 2.1. Materials Yttrium oxide (Y2O3; purity>99.9%) was purchased from Strem Chemicals (USA), and both nitric acid (HPLC grade) and urea ((NH2)2CO) were provided by Dongyang Chemicals (Seoul, Korea). Y(NO3)3 standard solution was prepared as follows. First, to remove any trace of moisture and other organic impurities, yttrium oxide was calcinated at 850 jC for several hours. Subsequently, it was dissolved in nitric acid, and the remaining solvent was removed by prolonged heating above the solvent boiling point. Y(NO3)3 thus obtained was further diluted using distilled water adjusting the solid concentration to 0.5 M with pH 6.0, and was used as a standard solution. Standard aqueous urea solutions 2.0 and 5.0 M were prepared and stored in a refrigerator (below 5 jC). Although urea is known to decompose in an aqueous solution, the rate of decomposition at this relatively low temperature was so slow that the temporal change of initial urea concentration could be neglected (the measured decomposition amount of urea after 7 days at 5 jC was about 7 10 4%). Urea standard solutions were either used or discarded within a week. Urea solutions with higher concentration than 5.0 M were always prepared right before the experiment to minimize the concentration variation.
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2.2. Preparation of yttria particles The previous two solutions having appropriate concentrations, Y(NO3)3 and urea, were mixed and then filtered through a filter paper (#41, Whatman). About 25 ml of mixed solution was placed in a glass tube, and this was inserted in an isothermal bath maintaining a desired reaction temperature. In these experiments, the effects of concentration of reactants, reaction temperature, solution pH on the morphology and kinetics of yttria particle formation were studied. After a complete reaction, the precipitates were separated, thoroughly washed with ethanol, and further dried at 100 jC overnight. Whenever necessary, a fraction of solution was extracted and the unreacted yttrium concentration in the solution was determined by titration method using 0.001 M EDTA (ethylenediaminetetraacetic acid). 2.3. Characterization of yttria particles The morphology of yttria particles was examined by transmission electron microscopy (TEM: JEM-200CX, JEOL) and scanning electron microscopy (SEM: JSM-35, JEOL). Ethanol was used to disperse the particles, and either carbon or gold was surface coated for TEM and SEM measurements, respectively. In this study, the main purpose of using TEM is to check the shape and the size of precipitates. As will be detailed in Results and discussion, particle shape and size are strong functions of the concentrations of yttrium and urea, temperature, and solution pH. To avoid any complications from kinetic effects, all the micrographs were taken at the terminal stage of each reaction. Thus we can safely assume that the precipitates were full grown in each set of experiment. The only exception would be the experiment showing the effect of time on the growth of precipitates. Particle size and distribution estimated from TEM micrographs were quantitatively confirmed using a particle size analyzer (Mastersizer, Malvern). The Zeta potential of the particles was monitored by Autosizer-III (Malvern). To understand the thermal behavior of yttria particles, differential scanning calorimetry (DSC: DSC7, Perkin Elmer), and thermogravimetric analysis (TGA: TGA-7, Perkin Elmer) were utilized. Carbon content in the particles was determined by a carbon –hydrogen – nitrogen analyzer (MT2, Yanaco).
3. Results and discussion Yttria particles are generated through a well-known reaction that yttrium cations (Y3 +) precipitate with anions such as CO32 or OH from the dissolution of urea [25]. In brief, yttrium cations precipitate as follows: Y3 + + OH + CO32 + 1.5 H2O = Y(OH)CO31.5 H2O. See Eqs. (9) – (15) for complete details of reaction schemes. Accordingly, the concentrations of yttrium ions and urea will significantly
affect the shape and size distribution of yttria particles. In addition, reaction temperature and initial solution pH will influence on the final products. In this section we will evaluate the effects of these variables. 3.1. The effect of yttrium concentration A series of TEM micrographs in Fig. 1 presents the effect of yttrium concentration on the shape and size of yttria particles. All other experimental conditions, such as the concentration of urea, solution pH, and reaction temperature, were all fixed as given in the figure caption. Two important observations can be made. First, as anticipated, the effect of yttrium concentration on particle size is significant, varying from 65 to 140 nm in the concentration range between 0.005 and 0.025 M. Fig. 2 shows the average size variation as a function of yttrium concentration. The error bar inserted in each data point serves as a standard deviation of particle size distribution. Below 0.025 M particles do not agglomerate, the particles size was nearmonodisperse (10 – 18% of standard deviation in size distribution). Severe agglomeration, however, at higher concentrations above 0.03 M decreased the reliability of measurement so the two data points denoted as open circles are only approximations with much broader size distributions. Second, within the range of yttrium concentration studied, the shape of precipitates is a well-rounded sphere. This may indicate that particles precipitate through a nucleation and growth mechanism, in which nucleation density is a strong function of yttrium concentration. Thus below a critical concentration not to invoke an impingement of growth front, we anticipate spherical-shaped precipitates. At higher yttrium concentration, however, we may observe non-spherical type particles due to frequent impingements. Indeed, a close comparison of particle shapes between Fig. 1c and g reveals that the shape of precipitates at higher concentrations is more irregular. Figs. 3 and 4 present time-dependent morphology changes in two different yttrium concentrations, 0.02 and 0.04 M, which represent below and above a critical concentration of 0.025 M for agglomeration. TEM micrographs were all taken in 10-min intervals in each case. At initial stage of precipitation (10 min), well-rounded spheres of size on the order of 60 – 100 nm evolve in both concentrations. As time proceeds, the precipitates in 0.02 M solution grow maintaining the shape of spheres, yet the particles in 0.04 M solution start agglomerate and form irregularly shaped conglomerates. Temporal average particle size variation in Figs. 3 and 4 are shown in Fig. 5. The open square symbol in Fig. 5 denotes the high uncertainty in size measurement due to severe agglomeration (in case of 0.04 M). In conjunction with Fig. 2, Fig. 5 suggests that the average particle size in the final stage of reaction matches well with the previous results within experimental uncertainty (e.g., compare Figs. 1d and 3c). This strongly indicates that not only is the
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control of particle size reproducible, but also, time can be an effective factor to manipulate the size of precipitates. The observation that, for a given time of experiment, higher yttrium concentration leads to bigger particles is in full accordance with our notion: under a fixed concentration of urea and constant level of supersaturation for nucleation, more yttrium ions are available for the growth of precipitates. The existence of agglomeration, however, is not expected due to a particle size effect. Experimentally, we could not observe particle agglomeration in the case of 0.02 M concentration even at much longer times. To address this issue properly, we measured the zeta potential of precipitates as a function of yttrium concentration, and the result is shown in Fig. 6. At lower concentrations up to 0.02 M, the zeta potential is above 13 mV, however, as concentration increases, potential undergoes a rapid drop and reaches a plateau value around 2 – 3 mV. This is most likely because as [Y(NO3)3] increases, the higher concen-
139
tration of anions, such as NO3, lead to a decrease of zeta potential, which, in turn, will facilitate the agglomeration of precipitates. From a study of homogeneous yttria precipitation, Matijevic et al. reported without specifying the mechanism that there exists an upper limit of yttrium concentration to obtain mono-disperse yttria precipitates [25]. In the present work, we observed a similar phenomenon; furthermore, we suggest that the decrease of zeta potential with yttrium concentration causes particle agglomeration, and this is perhaps one of the reasons for the existence of the upper [Y(NO3)3] concentration to obtain near-monodisperse yttria precipitation. These findings suggest that any forced dispersions using techniques such as ultrasonic or mechanical disturbance may be of help to increase the critical upper yttrium concentration by avoiding particle – particle agglomeration. Indeed, we found that the particle shape and size distribution became more spherical and narrow when the precipitates were made under the influence of ultrasonic dispersion even in
Fig. 1. A schematics of experimental set up for the production of yttria oxides particles. 1. Water bath; 2. Heating element; 3. Isothermal temperature controller; 4. Thermocouple; 5. Reaction tube; 6. Mechanical stirrer; 7. Water. TEM micrographs of yttrium compound prepared under different Y(NO3)3 concentrations as (a) 0.005 M, (b) 0.01 M, (c) 0.015 M, (d) 0.02 M, (e) 0.025 M, (f) 0.03 M, and (g) 0.04 M (in this case, urea concentration and pH were fixed at 1.0 and 6.0, respectively, and reaction temperature was 95 F 1 jC).
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Fig. 1 (continued).
Fig. 2. Variation of mean particle size as a function of yttrium concentration. The error bars are based on standard deviation of size distribution in each concentration. The open circle represents approximated average particle size in the presence of agglomeration. Note that, upon agglomeration, not only the size distribution broadens, but also the mean value deviates from the linear extrapolation (short-dashed line).
0.03 M yttrium solution. Fig. 7 shows SEM micrographs of homogeneous yttria precipitation with and without ultrasonic dispersion. It is apparent that ultrasonic dispersion effectively reduces the degree and the size of agglomeration. A similar result was previously reported [27]. The ultrasonic dispersion effect, however, was less distinctive as the yttrium concentration further increased to 0.04 M. It should be also noted that ultrasonic dispersion does not affect the average size and size distribution below 0.02 M concentration. In Fig. 8, the effect of yttrium concentration on the rate of precipitation is presented. A fraction of solution was extracted at appropriate time intervals, and the unreacted yttrium concentration was analyzed by titration method using 0.001 M EDTA (i.e., the amount of precipitation in Fig. 8 is given as the difference between the initial and the unreacted yttrium concentration). Solutions with three different concentrations (0.01, 0.02 and 0.03 M) were used, and other experimental conditions were controlled. Fig. 8 clearly tells us that, within the concentration range between 0.01 and 0.03 M, the rate of precipitation is virtually a constant, being equal to 6.56 10 4 mol/ l min. In addition, as shown in Fig. 9, we can define the
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Fig. 3. TEM micrographs showing the morphology change of yttrium compound prepared for different reaction time as (a) 10 min, (b) 20 min, (c) 30 min, and (d) 40 min (in this case, [Y(NO3)3] and [urea] were, respectively, 0.02 and 1.0 M. Solution pH was 6.0, and the reaction temperature was 95 F 1 jC).
necessary time for a complete reaction from the plateau region in Fig. 8. The 100% reaction time, ranging between ca. 20 and 60 min, linearly increases with the yttrium concentration. These results, combined with Fig. 2 showing the variation of average size with concentration, are of practical importance to control the particle size and its distribution. 3.2. The effect of urea concentration To study the effect of urea on shape, size, and the distribution of yttria precipitates, the urea concentration, [urea], was systematically varied between 0.04 and 7.0 M, while other conditions were controlled. A series of TEM micrographs in Figs. 10 and 11 presenting the variation of average particle size clearly show that an increase of urea concentration decreases the average particle size, from ca. 220 nm ([urea] = 0.04 M) to 105 nm ([urea] = 4.0 M). Particle size almost linearly decreases with the logarithm of urea concentration (see the insert in Fig. 11). At the highest concentration studied, particles experience such an excessive degree of agglomeration that the morphology of
precipitates appears to be a network structure, and individual particles cannot be distinguished (thus 7.0 M data point is omitted in Fig. 11). The shape of precipitates was spherical between 0.04 and 4.0 M of urea concentration. The inverse relationship between the urea concentration and the average particle size may be explained as follows. The higher the urea concentration, the greater the amount of anions generated, such as CO32 and OH, which will lead to a higher degree of supersaturation before the onset of precipitation. The amount of urea decomposition with urea concentration linearly increased, in which the rate of decomposition was 1.3 10 3/min. In this case, we expect an increase of nucleation density; therefore, under a given yttrium concentration, the average particle size will decrease. At very high urea concentration, such as 7.0 M, an excessively high level of anions will lead to a high nucleation density and a fast growth of precipitates, which contribute to form (even continuous) aggregates. Once again, these observations reinforce the theorem in homogeneous precipitation that particle shape and size control must be understood in terms of the balance between nucleation and growth.
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Fig. 4. TEM micrographs showing the morphology change of yttrium compound prepared for different reaction time as (a) 10 min, (b) 20 min, (c) 30 min, and (d) 40 min (in this case, [Y(NO3)3] and [urea] were, respectively, 0.04 and 1.0 M. Solution pH was 6.0, and the reaction temperature was 95 F 1 jC).
Fig. 5. Variation of mean particle size with reaction time for two different yttrium concentrations below and above a critical concentration for agglomeration. The data point represented by an open square is only an approximated value under severe agglomeration.
Fig. 6. Variation of zeta potential of the precipitates formed under different yttrium concentration.
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Fig.9. The time for complete reaction as a function of yttrium concentration.
Fig. 7. SEM micrographs of yttrium compound prepared without (a) and with (b) ultrasonic disturbance (in this case, [Y(NO3)3] and [urea] were, respectively, 0.03 and 1.0 M. Solution pH was 6.0, and the reaction temperature was 95 F 1 jC).
Fig. 8. Temporal variation of the precipitated amount of yttrium for three different yttrium concentrations (in this case, urea concentration was 1.0 M. Solution pH was 6.0, and the reaction temperature was 95 F 1 jC).
The effect of urea concentration on the rate of precipitation is shown in Fig. 12, in which the temporal variation of the amount of precipitation is obtained similarly as shown in Fig. 8. One immediate observation is that as urea concentration increases, precipitation rate accelerates. This is in good contrast with the results in Fig. 8, where yttrium concentration did not increase the rate of precipitation but only affected the time for complete reaction. Urea concentration strongly influences both the rate of precipitation (Fig. 13) and the time for complete reaction (Fig. 14). The rate of precipitation, as defined by the slope of each curve in Fig. 12, increases fast with urea concentration up to 3.0 M (Regime I), and reaches a plateau value (Regime II). The behavior in Regime I may be readily understood from the fact that as the concentration of anions from urea decomposition increases, the precipitation of yttria particles accelerates. However, the existence of Regime II deserves further explanation. For this, first we need to understand the equilibrium reactions involved with anions such as CO32 and OH[28]. CO2 3 þ H2 O ¼ HCO3 þ OH
ð1Þ
HCO 3 þ H2 O ¼ H2 CO3 þ OH
ð2Þ
H2 CO3 ¼ CO2 þ H2 O
ð3Þ
In a system where the concentration of yttrium cation (Y3 +) is fixed, the excessive amount of (CO32 ) over the balance of the reaction stoichiometry will be consumed through the sequential reactions as shown above. If this hypothesis were true we would expect an increase of solution pH at high urea concentrations. An increase in solution pH may be equally possible at long reaction times at which the precipitation reaction is completed and thus all the anions generated must be consumed via above reactions.
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The variation of solution pH for selected systems is inserted in Fig. 12. At two lower concentrations, solution pH was invariant (5.7 F 0.1) until the precipitation reaction is completed. In case of [urea] = 1.0 M, the later stage pH increase around 35 –40 min must be ascribed to the consumption of anions after the reaction is completed (see also Fig. 14 showing that the time for complete reaction in this case occurs around 38 min). On the other hand, at high urea concentration such as 4.0 M, even at the first measurement done at 5 min, the time well before the reaction is complete (18 minutes in Fig. 14), the solution pH is already higher than 6.0, and keeps increasing afterwards. These observations are in good agreement with the above hypothesis, and may explain the existence of Regime II in that excessive supply of carbonyl anions from urea decomposition will be consumed through 1) reactions (2) reactions (3) and do not contribute to form yttria precipitates. The time for complete reaction in Fig. 14 has another noticeable point. At low concentrations up to 1.0 M, the reaction time, t, reduces following an exact inverse relationship between t and [urea] (i.e., t~1/[urea]). This
matches well with Fig. 13 showing a linear increase of the rate of precipitation in low urea concentrations. At high concentrations above 3.0 M, the reaction time does not reduce appreciably because the concentrations of anions such as (CO32 ) are already saturated. Combined with Fig. 11, Fig. 14 will be an important map to control the yttria precipitation process. 3.3. The effect of reaction temperature Similar to other phase transition behaviors, such as crystallization, involving nucleation and growth mechanism, the rate of yttria precipitation will critically depend on reaction temperature. To quantify this effect, reaction temperature was varied between 70 and 95 jC. Other experimental conditions were kept constant ([Y(NO3)3]:0.02 M; [urea]:1.0 M; pH: 6.0). TEM micrographs in Fig. 15 show the morphology of yttria particles formed under different reaction temperatures. The variation of average particle size as a function of temperature is presented in Fig. 16. Within the temperature range studied, precipitates maintained well-
Fig. 10. TEM micrographs of yttrium compound prepared under different urea concentrations as (a) 0.04 M, (b) 0.1 M, (c) 0.5 M, (d) 1.0 M, (e) 2.0 M, (f) 4.0 M, and (g) 7.0 M (in this case, yttrium concentration and pH were fixed at 0.02 and 6.0 M, respectively, and reaction temperature was 95 F 1 jC).
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Fig. 10 (continued).
rounded spherical shape, and the average size and the size distribution did not vary significantly. Particle size showed a slight decrease with temperature, yet the overall variation, from 135 nm (70 jC) to 126 nm (95 jC), was within the resolution of measurements. In contrast to relatively insignificant effect on particle size, we found that temperature has a profound impact on the kinetics of precipitation. In Fig. 17, the amount of precipitation is given with reaction temperature. At 70 jC, precipitates form very slowly, and even after 3 h the total amount of precipitation is about 25% of theoretical limit (i.e., in this case, 0.02 M). As temperature increases, the rate of precipitation increases significantly; at 95 jC, the reaction finishes around 40 min. Fig. 17 in combination with Fig. 16 confirms that reaction temperature predominantly affects the growth kinetics of precipitates not the thermodynamic factors such as an average particle size, which is governed by the concentrations of yttrium and urea. Fig. 18 presents the rate of precipitation, defined by the slope of each line in Fig. 17, as a function of reaction temperature. The rate exponentially increases more than an
order of magnitude within 25 jC temperature window. The decomposition rate of urea is known to exhibit similar exponential dependency in the same temperature range, and the literature data has been reproduced in Fig. 18 [29]. This indicates that a strong temperature dependence of the rate of precipitation mainly comes from an increased dissolution rate of urea at high temperature. From the slope of each line in Fig. 17, we can also deduce an activation energy for yttria precipitation by plotting lnK versus 1/T, in which K is a rate constant, and T is an absolute temperature (i.e., use of Arrhenius-type temperature dependence: K = Koexp( Ea/RT), where Ea is an activation energy). The result is inserted in Fig. 18. The calculated activation energy is ca. 29 kcal/mol. It is interesting to note that, according to Shaw and Bordbaux [29], the activation energy for urea decomposition in aqueous solution is in the range of 28 – 32 kcal/mol. The close match between the two activation energies may suggest that urea decomposition step is a rate-determining step in the formation of yttria particles from a homogeneous precipitation. In Fig. 19, the reaction time for complete reaction as a function of temperature is presented. The complete reaction
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Fig. 11. Variation of mean particle size as a function of urea concentration. The insert shows a linear decrease of the mean particle size with the logarithm of urea concentration.
time at each temperature is defined by the intercept of corresponding line to 0.02 M as shown in Fig. 17. Over 800 min of reaction time at 70 jC dramatically decreases to 40 min at 95 jC. Recall that the wide range of reaction rate at different reaction temperature, however, does not affect the particle size obtained at the terminal stage.
3.4. The effect of initial solution pH It is well known that above 7.0 pH, yttrium ions (Y3 +) precipitates with OHanion in the form of Y(OH)3 when [Y3 +] is in the range of 0.1 – 0.01 M (reaction constant, K=[Y3 +][OH]3 = 6.0 10 24 at 25 jC) [28].
Fig. 12. Temporal variation of the precipitated amount of yttrium for six different urea concentrations (in this case, yttrium concentration was 0.02 M. Solution pH was 6.0, and the reaction temperature was 95 F 1 jC). The insert shows time-dependent variation of solution pH for three urea concentrations. For details of the correlation between the precipitation amount and solution pH, see text.
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of yttrium ions (Y3 +) (Eq.(6)), and (3) precipitation reaction (Eq. (7)). Urea decomposition ðNH2 Þ2 CO ¼ NHþ 4 þ OCN
ð4Þ
OCN þ 2Hþ þ H2 O ¼ CO2 þ NHþ 4
ð5Þ
Hydrolysis Y3þ þ H2 O ¼ ½YðOHÞ 2þ þ Hþ
ð6Þ
Precipitation ½YðOHÞ 2þ þ CO2 þ 2H2 O ¼ YðOHÞCO3 H2 O þ 2Hþ ð7Þ Overall, Fig. 13. Precipitation rate change as a function of urea concentration. For the explanation of Regime-I and -II, see text.
Y3þ þ ðNH2 Þ2 CO þ 4H2 O ¼ YðOHÞCO3 H2 O þ þ 2NHþ 4 þH
Therefore, in the present study, pH = 7.0 is will be an upper limit to obtain yttria particles through homogeneous precipitation. In this section we investigated the effect of solution pH below 7.0. TEM micrographs in Fig. 20 show the morphology of yttria particles under varying pH’s between 1.3 and 6.0. It is clear that below pH 2.0, particles experience severe agglomeration, and only above pH 3.0, spherically shaped near-monodisperse yttria particles are obtained. The effect of pH on the size and distribution of particles was not significant when pH is in the range of 3.0 and 6.0. Although we did not measure the zeta potential of precipitates, we suspect that the origin of agglomeration at low pH may be ascribed to the decrease of zeta potential in the presence of a higher concentration of counter anions (NO3) which compress the electrical double layer and facilitate the agglomeration of particles. Similar to the previous sections, we also measured the amount of precipitation with time, from which we could deduce the time for complete reaction and the rate of precipitation, as shown in Fig. 21a and b, respectively. These two additional data confirm that the effect of pH on the kinetics of precipitation is not distinguishable at pH 3.0 or higher. As a reason for the decreased reaction kinetics at low pH, we suggest that at significantly lower pH than ca. 5.5 –6.0, which is required for the onset of precipitation (see Fig. 12, for instance), it takes longer time to reach this pH through the decomposition of urea.
ð8Þ
These chemical equilibriums lead to an overall reaction that one mole of urea reacts with one mole of yttrium ions and forms an equi-molar yttria precipitates. In addition, as reaction proceeds, the solution pH will decrease due to the generation of protons. Contrary to the previous report, we found that solution pH remains constant during the precipitation (see, for instance, Fig. 12) performed under wide different experimental conditions, varying reactant concentrations (both yttrium and urea), and temperature. In addition to this apparent mismatch, we also found that the equilibrium constant, K, of the hydrolysis reaction in Eq. (6) is too small to believe that it is a main source supplying the hydroxyl groups in the overall reaction in Eq. (8) (e.g., at
3.5. Reaction mechanism According to Matijevic et al., the reaction mechanism of the formation of yttria particles through homogeneous precipitation is claimed to include three distinct steps [25]: (1) decomposition of urea (Eqs. (4) and (5)), (2) hydrolysis
Fig. 14. Time for complete reaction as a function of urea concentration.
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Fig. 15. TEM micrographs of yttrium compound prepared under different reaction temperatures as (a) 70 jC, (b) 80 jC, (c) 90 jC and (d) 95 jC (in this case, [Y(NO3)3] and [urea] were, respectively, 0.02 and 1.0 M. Solution pH was 6.0).
pH 6.0, [Y(OH)]2 +/[Y3 +] = K 106 = 10 8.04 106 c0.01, thus only about 1% of yttrium ions exist in the form of yttrium hydroxide ions) [30].
To resolve these contradictions, we suggest new routes of chemical reactions that are compliant to the experimental observations. Urea Decomposition ðNH2 Þ2 CO ¼ NHþ 4 þ OCN
ð9Þ
þ OCN þ 2H2 O ¼ CO2 3 þ NH4
ð10Þ
CO2 3 þ H2 O ¼ HCO3 þ OH
ð11Þ
HCO 3 þ H2 O ¼ H2 CO3 þ OH
ð12Þ
H2 CO3 ¼ CO2 ðgÞ þ H2 O
ð13Þ
Precipitation Y3þ þ OH þ CO2 3 þ 1:5H2 O ¼ YðOHÞCO3 1:5H2 O ð14Þ Overall, Fig. 16. Variation of mean particle size with reaction temperature. The large error bars show approximately 10 – 13% of standard deviation involved in multiple measurements.
Y
3þ
þ 1:5ðNH2 Þ2 CO þ 5H2 O
¼ YðOHÞCO3 1:5H2 O þ 3NHþ 4 þ 1=2CO2 ðgÞ
ð15Þ
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Fig. 17. Temporal variation of the precipitated amount of yttrium for five different reaction temperatures (in this case, [Y(NO3)3] and [urea] were, respectively, 0.02 and 1.0 M. Solution pH was 6.0).
According to newly proposed mechanism, 1.5 mol of urea reacts with 1 mol of yttrium ions to form 1 mol of yttrium compound (Y(OH)CO31.5H2O). This contrasts with the previous mechanism by Matijevic et al. [25] in that the amount of consumed urea per mole of yttrium ions is 50% larger, and accordingly, the reaction will produce, as a by-product, 3 mol of NH4+ions rather than 2 mol (compare Eqs. (8) and (15)). To compare the validity of the reaction mechanisms proposed in Eq. (15) over Eq. (8), we measured the amount of NH4+as a function of reaction time and compared it with
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Fig. 19. Time for complete reaction as a function of reaction temperature.
the precipitated amount of yttrium ions. This result will support, at least indirectly, which reaction mechanism is more feasible in the explanation of experimental observations. A brief experimental procedure follows. Sixteen glass tubes with 25 ml of solution ([(Y(NO 3 ) 3 ) = 0.02 M, [urea] = 1.0 M, and pH = 6.0 ) were placed in a water bath at 95 jC. The amounts of NH4+ and yttrium ions were measured in every 5 min. Two samples were used for each measurement. It is well known that, at room temperature, ammonium ions (NH4+) are stable up to pH = 7.0, and as pH increases they gradually transform to ammonia (NH3) [30]. At pH 12.0 or higher only ammonium species are stable.
Fig. 18. Precipitation rate change as a function of reaction temperature (closed circles). Open squares show the variation of rate constant in urea decomposition from the literature [25]. Note that how these two different quantities exhibit similar temperature dependence. The insert presents Arrhenius-type plot, from which activation enrgy for yttria precipitation can be calculated.
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Fig. 20. TEM micrographs of yttrium compound prepared under different solution pHs as (a) 1.3, (b) 2.0, (c) 3.0, (d) 4.0 and (e) 6.0 (in this case, [Y(NO3)3] and [urea] were, respectively, 0.02 and 1.0 M and reaction temperature was 95 F 1 jC).
Based on this thermodynamic equilibrium, ammonium ions were first changed to ammonia at pH = 12.0 and then determined using ammonia gas sensing electrode (NH31501, Phoenix, USA). Fig. 22 compares the temporal variation of the precipitated amounts between yttrium ions and ammonium ions. As proposed in the present study, the temporal variation of ammonium ion concentration quantitatively matches with three times of yttrium ions. The
inserted line showing twice the amount of [Y3 +] is based on the previous study by Matijevic et al. [25]. Another difference between the Eqs. (8) and (15) lies in the number of water molecules contained in yttrium compounds. It is largely accepted that general chemical composition of yttria precipitates is Y(OH)CO3xH2O, which further changes to Y(OH)CO3 at 200 jC, and finally to Y2O3 above 640 jC [25,26]. In this work we carefully followed
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tion of precipitates. Upon calcination, the average size decreases about 20% with slight broadening. The average size of as-prepared particles was ca. 170 nm, which reduces to 135 nm after calcination. A similar result was reported previously [25]. From the known densities of yttrium compound [31] and yttrium oxide [32], which are, respectively, ca. 3.9 and ca. 5.0 g/cm3, we can expect about 23% size reduction after calcination. This theoretical size decrease is in good match with the actual observation.
4. Summary An emerging interest of utilizing rare earth metals and their compounds in advanced ceramics has driven many recent studies. As a part of these efforts, we studied a homogeneous precipitation method in urea aqueous solution as a route to synthesize near-monodisperse yttria particles. Among others, the concentrations of yttrium and urea, reaction temperature, and solution pH were chosen for the most significant processing variables affecting both morphology of particles and reaction kinetics. Based on this, we conclude as follows.
Fig. 21. Precipitation rate change (a) and time for complete reaction (b) as a function of solution pH.
(1) Yttrium concentration, within the range studied (0.005 M and 0.04 M), increased the average size of particles from 65 nm to over 220 nm. Yttrium concentration is also involved with the broadening of size distribution. Above 0.025 M particles start to coagulate, which is ascribed to the reduction of zeta potential with yttrium concentration. The rate of precipitation reaction is shown to be independent of yttrium concentration.
the mass changes upon thermal decomposition at two different temperatures of 250 and 850 jC using TGA (oxygen purging rate: 1.0 l/min), and we confirmed that, indeed, the exact chemical composition of yttrium precipitates is Y(OH)CO31.5H2O. These two independent experimental evidences support the validity of the newly proposed mechanism in homogeneous precipitation of yttria particles in the presence of urea. 3.6. Calcination of yttrium compound As noted earlier, the precipitated yttrium compound undergoes chemical change to yttrium oxide (Y2O3) above 640 jC. It is important to understand the morphological changes upon calcinations, since one of the target applications of yttrium particles includes the utility as an alloying element for advanced ceramics through sintering. TEM micrographs in Fig. 23 contrast the morphology before (a and b) and after (c and d) calcinations performed at 850 jC for 2 h. One immediate observation is that, surface roughness increases after calcination. This is mainly due to the evaporation of chemically bonded water and CO2 during heating. Calcination also affects the size and size distribu-
Fig. 22. Quantitative comparison between the amounts of yttrium ions (open squares) involved in precipitation reaction and ammonium ions (open circle) formed as a by-product. An excellent match between these two quantities strongly supports the validity of newly proposed reaction mechanism (see Eq. (15) and text for details). Triangles show twice the amount of yttrium ions used in precipitation based on the reaction mechanism from the literature [25].
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Fig. 23. Morphological change of yttria compound before (a and b) and after (c and d) calcination performed at 850 jC for 2 h (reaction conditions were [Y(NO3)3] = 0.02 M, [urea] = 1.0 M, pH = 6.0, and temperature of 95 F 1 jC).
(2) From the study of urea concentration ranging from 0.04 to 4.0 M, we found that the average particle size exhibits a gradual decrease from ca. 220 to ca. 100 nm. As similarly observed previously, extremely high urea concentration such as 7.0 M, caused inter-particle coagulations. The rate of precipitation, which increases with concentration up to 3.0 M of urea, starts to level off at higher concentrations. (3) Temperature is shown to significantly increase the rate of precipitation; however, thermodynamic quantities, such as average particle size, is only a function of the concentrations of reactants (both yttrium and urea). Based on Arrhenius-type reaction kinetics, the activation energy for precipitation is found to be 29 kcal/mol. This value is in good match with the activation energy of urea decomposition (28 –32 kcal/mol). We suggest that this similarity may be an indication that urea decomposition is a rate-determining step in the formation of yttria particles in homogeneous precipitation method.
(4) Solution pH is also proven to be a controlling factor affecting the particle morphology and the reaction kinetics. Severe coagulation that occurs at low pH ( < 2.0) deters to obtain well-separated individual yttria particles, yet as pH increases above 3.0, near-monodisperse yttria particles can be achieved. The rate of precipitation showed a similar trend as coagulation: at low pH the rate is slow, and as pH increases the rate enhances reaching a plateau value above pH 3.0. From the study of reaction mechanism, we observed two important features. First, pH does not change during precipitation (maintaining pH 5.5– 6.0), and second, the equilibrium constant of yttrium ion hydrolysis (i.e., K= [Y(OH)2 +]/[Y3 +]) at this pH range is too small to serve as a reasonable source of hydroxyl ions in the final yttrium compound. Based on these two findings, we propose a new reaction mechanism that is more plausible in yttria particle precipitation from urea aqueous solution. As supportive experimental evidence, we found a quantitative match
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between the amount of ammonium ions generated and the precipitated yttrium ions proposed in the new reaction mechanism. [16] [17]
Acknowledgements Editorial work by Mr. Derek Klinedinst is appreciated.
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