642
J. Phys. Chem. C 2007, 111, 642-651
Preparation and Characterization of Sol-Gel MgAl Hydrotalcites with Nanocapsular Morphology Jaime S. Valente,*,† Manuel S. Cantu´ ,† Jose´ G. H. Cortez,† Ramo´ n Montiel,‡ Xim Bokhimi,§ and Esteban Lo´ pez-Salinas† Instituto Mexicano del Petro´ leo, Eje Central # 152, CP 07730, Me´ xico, D. F., Me´ xico, PEMEX-Petroquı´mica, I y D Tecnolo´ gico, Jacarandas 100, CP 52500, Coatz. Ver. Me´ xico, and Instituto de Fı´sica, UNAM, AP 20-364, 01000 Me´ xico, D. F., Me´ xico ReceiVed: August 15, 2006; In Final Form: October 19, 2006
A series of MgAl hydrotalcites (HTs) were synthesized by a sol-gel procedure with ethanol, 2-propanol, and 1-butanol as solvents. Particular attention was paid to controlling hydrolysis reactions by means of using a substoichiometric H2O/Al molar ratio equal to 1. The sol-gel HTs were characterized by means of thermogravimetry, mass spectrometry (MS) of the decomposition products, X-ray diffraction, diffuse reflectance infrared spectroscopy (DRIFT), electron microscopy, and N2 adsorption-desorption. Decomposition products of sol-gel HTs were identified by MS and included, among others, CxHy fragments from alkoxy groups, which were partially withdrawn between 250 and 450 °C. Alkoxy groups, from the alcohols employed as solvents, are probably intercalated within the HT layers, as indicated by XRD and DRIFT analyses. CxHy groups in these alkoxy groups were detected by DRIFT at ∼2945 and ∼2831 cm-1 in dried HTs, even after annealing at 500 °C. The HTs obtained by the sol-gel procedure described in this work showed nanocapsular morphology. Transmission electron microscopy and Rietveld refinement techniques evidenced the curvature of the HT layers. In comparison with coprecipitated MgAl-CO3 HT particles (>10 µm), sol-gel HTs showed very small particle aggregates (ca. 300∼350 nm). Calcined HTs presented specific surface area and pore volume values between 254∼332 m2/g and 0.81∼1.39 cm3/g, respectively, depending on the alcohol used in the sol-gel procedure.
Introduction Among naturally occurring anionic clays, hydrotalcite-like (HT) materials have shown the most tunable physicochemical properties.1-3 Since a wide variety of compounds with HT structure can be easily prepared,3-5 these materials are represented by the general formula [M2+(1-x)M3+x(OH)2]An-x/n‚mH2O. An HT structure is created replacing some of the M2+ cations, in the lattice of brucite, with M3+ trivalent cations, rendering the layered array positively charged. These positively charged layers are electrically compensated for by anions (An-), which are located in the interlayer region.2-4 Conventionally, HTs are prepared through the coprecipitation of metallic salts (M2+ and M3+) with a concentrated alkaline solution. The chemical composition could be controlled from the starting solution, e.g., the molar ratio M2+/M3+ and the anion’s nature. In fact, in order to produce a well-crystallized compound, a hydrothermal treatment is required.4,6 Throughout previous decades, hydrotalcites have been extensively studied as catalysts or catalyst precursors.3 Since calcination of HTs yields mixed oxides that are basic solids, HT compounds have found applications in many organic reactions considered to be catalyzed by bases such as aldol condensation,7-9 reduction of aromatic nitro compounds,10,11 Michael addition,12,13 Wadsworth-Emmons reaction,14 selective * To whom correspondence should be addressed: phone + 52 (55) 91758444; e-mail
[email protected]. † Instituto Mexicano del Petro ´ leo. ‡ PEMEX-Petroquı´mica. § Instituto de Fı´sica.
reduction of unsaturated ketones/aldehydes by hydrogen transfer from alcohols,15 cyanoethylation of alcohols,16 etc. Also, they have been used to help in environmental problems as reducing additives for SOx and NOx removal, l7,18 etc. In recent years, because of the structural anisotropy inherent to layered materials, worthwhile efforts have been reported in HT synthesis on the formation of transparent thin films.19 Considering the wide range of technological applications for these films, progress in making thin continuous films of HTs is a valuable goal. In addition, there has been an increasing interest in synthesizing inorganic nanostructures in order to use them in diverse fields such as drug delivery carriers, sensing devices, etc.20,21 In the case of HTs, several research groups worldwide have attempted to introduce biological species and organic compounds between the layers.22-25 Many of these species, including amino acids, carboxylic species, peptides, and larger proteins occur as anions at neutral to basic pH values because HTs have positively charged layers; hence, the importance of studying HTs is clear. Alternative methods have been proposed to produce hydrotalcite-like compounds, such as wave irradiation to improve crystallization,26 ultrasound assisted treatment,25,27 and the solgel method.28-32 The sol-gel method provides an attractive and convenient route since it enables accurate control over the structural and textural properties of products; for instance, materials with high specific surface area, uniform pore size distribution, and good purity can be obtained. Nevertheless, few reports have shown a pure HT structure,31,33 homogeneous pore size distribution, or high specific surfaces areas.31 In some cases, this was due to the fact that the molar
10.1021/jp065283h CCC: $37.00 © 2007 American Chemical Society Published on Web 12/14/2006
MgAl Hydrotalcites with Nanocapsular Morphology Mg2+/Al3+ ratio was too high.28,29 In order to obtain a pure HT structure, the molar ratio must be limited to 1.5 < M2+/M3+ < 4.5. Furthermore, the most important parameters to be controlled in the sol-gel process are the molar ratios and the nature of reactants.35 Hence, the polarity, reactivity, and amount of the intended solvent, generally an alcohol, are of paramount importance. Since an alkoxide exchange (RO-) takes place between the metal alkoxide and the alcohol, the reaction’s kinetics can be controlled in this manner. Also, the alkoxide dilution plays a key role: the more diluted, the higher the alcohol/metal alkoxide ratio, providing better control on the polymerization path. Water is essential to control the hydrolysis reaction, and our own experience has shown that working with substoichiometric water results in better control over the solgel process.36,37 Another important factor is the different reactivities of metal alkoxides, these reactivities being mainly related to metal electronegativity. Thus, to produce homogeneous materials when metal alkoxides with different reactivities are employed, one possibility is to complex the most reactive metal alkoxide. Acetates are suitable for this end.35 In the case of combining magnesium and aluminum alkoxides, the latter, being more reactive than the former, should be complexed. For alkoxidederived gels, condensation between surface functional groups continues to occur after the gel point. This aging process represents the time between the formation of a gel and the removal of a solvent. As long as the pore liquid remains in the matrix, a gel is not static and can undergo many transformations.36,37 This could actually be desirable because it leads to a more cross-linked network. However, extensive condensation causes the gel to shrink and is frequently detrimental to the physical properties of the material. Often, in previous studies, most of these parameters were neglected when hydrotalcites were prepared.30,32,34 The strategy employed in this study was to perform HT synthesis by the sol-gel technique with various alcohols, differing in their carbon chain length and polarity, as solvents. Particular attention was paid to (i) using a convenient alcohol/ alkoxide molar ratio, (ii) controlling hydrolysis reactions by means of a substoichiometric H2O/Al molar ratio, (iii) complexing aluminum alkoxide with acetic acid, and (iv) keeping the aging time constant. Experimental Section Synthesis Procedure. For the preparation of the sol-gel samples, aluminum tri-sec-butoxide (ATB) (Aldrich, 97%) and magnesium methoxide (MgM) (Aldrich, 10.16 wt % in methanol) were used as aluminum and magnesium sources. The hydrolysis catalyst used was HNO3 (Baker, 70%); acetic acid (AA) (Baker, 99.8%) was employed to inhibit the polymerization’s reaction. Ethanol (Baker, 99%), 2-propanol (99.9%), and 1-butanol (Aldrich 99%) were used as solvents. The synthesis procedure is as follows: the desired alcohol (ethanol, 2-propanol, or 1-butanol) was refluxed, and thereafter the aluminum tri-sec-butoxide was added and dissolved into the alcohol; the mixture was stirred for 1 h. Afterward, a 3 N HNO3 solution was added dropwise under vigorous stirring for 1 h, producing a transparent solution. The system was subsequently cooled to room temperature and acetic acid was added under vigorous stirring for 1 h. In the following step, the system was cooled to 0 °C and magnesium methoxide was slowly added. The solution was stirred at room temperature for 24 h and then deionized water was slowly added, allowing the hydrolysis to complete itself. The molar ratios of reactants were ATB:ROH
J. Phys. Chem. C, Vol. 111, No. 2, 2007 643 ) 1:60, ATB:HNO3 ) 1:0.03, ATB:MgM ) 1:3, ATB:AA ) 1:0.5, and ATB:H2O ) 1:1. The gel was poured into a glass vessel and was aged for 24 h at room temperature. The products were dried overnight at 80 °C. The samples were labeled as MgAl-E, MgAl-P, and MgAl-B where the letters E, P, and B correspond to the alcohol used. For instance, to obtain 10 g of MgAl-E the amounts of reagents were EtOH ) 108.19 mL, ATB ) 8.11 mL, HNO3(3N) ) 0.309 mL, AA ) 0.88 mL, MgM ) 78.67 mL, and H2O ) 0.276. As a comparison, a MgAl hydrotalcite was prepared by the coprecipitation method, as follows: an aqueous solution (1 M) containing 24.92 g of Mg(NO3)2‚6H2O and 11.59 g of Al(NO3)3‚ 9H2O, and a second solution (2 M) containing 10.53 g of KOH and 11.41 g of K2CO3 were simultaneously added to a glass reactor containing 100 mL of bidistilled water and mixed under vigorous mechanical stirring, with the pH maintained between 9 and 10. Then, the mixture was heated at 65 °C for 18 h. The precipitate was washed several times with deionized water and dried at 100 °C overnight. The sample was labeled as MgAlCP. Chemical Analysis. The chemical composition of solids was determined in a Perkin-Elmer model Optima 3200 Dual Vision by inductively coupled plasma atomic emission spectrometry (ICP-AES). Thermogravimetric Analysis. Thermogravimetric (TG) analyses were carried out on a Netzch STA-409EP instrument which was operated under helium flow at a heating rate of 10 °C/min from 25 to 1000 °C. In all determinations, 100 mg of finely powdered dried HT samples was used. Mass Spectrometry. The evolution of the thermal decomposition products was followed in an Autosorb-1 apparatus coupled with a Prisma-QMS 200 quadrupole spectrometer. The sample was annealed from room temperature to 900 °C in a He atmosphere at a heating rate of 10 °C/min and a He flow rate of 60 mL/min. The sample’s weight was ∼100 mg. The release of water (m/z ) 18), CO and CO2 (m/z ) 28 and 44, respectively), and organic species’ fragments (m/z ) 12, 14, 15, 29, 43, 45, 57, 59, and 73) was monitored throughout the decomposition process. Prior to the test, a scan was carried out on every sample to identify the species emitted, for the purpose of adjusting the detection parameters to each sample. In order to estimate the amount of evolved species upon annealing, total MS peak areas under the curves of all fragments, at indicated temperature intervals, were given a weight percent, corresponding to that of TG temperature interval, which represented the sum of all individual areas of the species detected by MS. Then, the weight percent contribution of a particular species or fragment was obtained by dividing the area under the curve of a particular species by the total area. All hydrocarbon-ion fragments were assumed to evolve from alkoxy anions. For the sake of clarity, only evolution of main species as a function of annealing temperature is shown in Figure 2; CO, O2, and O ionic fragments are not shown but were considered in distribution of weight percent. X-ray Diffraction. The X-ray diffraction pattern of the samples was measured in a θ-θ Bruker D-8 Advance diffractometer with Cu KR radiation, a graphite secondary-beam monochromator, and a scintillation detector. Diffraction intensity was measured between 2° and 110°, with a 2θ step of 0.02° and a counting time of 9 s/point. Hydrotalcite crystalline structure was refined via the Rietveld method using Fullprof code. Crystallite morphology was modeled with a base of spherical harmonics, which approximates well the crystalline form anisotropy. In contrast, microstrain could only be ap-
644 J. Phys. Chem. C, Vol. 111, No. 2, 2007
Valente et al.
TABLE 1: Proposed Chemical Compositions sample
solvent
chemical formulas
Mg/Al ratio
nRO-/nAl 250-450 °C
MgAl-E MgAl-P MgAl-B MgAl-CP
ethanol 2-propanol 1-butanol water
[Mg0.784Al0.216(OH)2](OC2H5)0.193‚(CO3)0.027‚0.57H2O [Mg0.776Al0.224(OH)2](OCH(CH3)2)0.158(CO3)0.087‚0.95H2O [Mg0.763Al0.237(OH)2](OC4H9)0.076(CO3)0.089‚0.83H2O [Mg0.751Al0.249(OH)2](CO3)0.145‚0.86H2O
3.63 3.47 3.22 3.02
0.89 0.70 0.32 0
proximated with an isotropic model because of the lack of information about the relative positions in unit cell of the atoms in the intercalated alcohol molecules. The background was modeled with a polynomial function that, in addition to the constant, linear, quadratic, and cubic terms in 2θ, also contained the terms (1/2θ) and (1/2θ)2. The standard deviations given in text and in tables, which shows the last figure variation of a number, are given in parentheses. When they correspond to Rietveld refined parameters, their values are not estimates of the probable error in the analysis as a whole but only of the minimum possible errors based on their normal distribution.38 Infrared Spectroscopy. Diffuse reflectance infrared (DRIFT) spectra were recorded on a Bruker Equinox 55 spectrometer equipped with a Harrick diffuse reflection attachment (H-DRPBR-3) and a Harrick reaction chamber (HVC-DR2). Spectra were taken with a deuterated L-alanine truglycine sulfate (DLATGS) high-sensitivity detector over the ranges of 40002500 and 1750-1000 cm-1 with a resolution of 4 cm-1, and 300 scans were collected. Powdered samples were heat-treated in situ under vacuum at 50, 250, and 500 °C for 30 min, and then the spectra were recorded at these same temperatures. Note: The infrared spectra were taken both under vacuum and under helium atmosphere. However, better resolution was obtained under vacuum; thus these results are here presented and discussed. Electron Microscopy. Samples were analyzed through transmission electron microscopy (TEM) in a Jeol 200 Kv JEM2200FS. The microscope is equipped with a Schottky-type field emission gun in ultra-high-resolution (UHR) configuration (Cs ) 0.5 mm; Cc ) 1.1 mm; point-to-point resolution, 0.19 nm) and in-column omega-type energy filter. Fresh or calcined (at 500 °C/4 h in air) samples were dispersed in ethanol before they were placed on the copper grid with Formvar support. Scanning electron microscopy (SEM) analysis was carried out in a Philips XL30 ESEM with an acceleration voltage of 25 kV. Prior to analysis, the samples were covered with gold and mounted over a carbon film. Textural Analysis. The texture of the calcined samples (at 500 °C/4 h in air) was analyzed by N2 adsorption-desorption isotherms at -196 °C on an Autosorb-I apparatus. Prior to the analysis, the samples were outgassed in a vacuum (10-5 Torr) at 400 °C for 5 h. The surface areas were calculated by the Brunauer-Emmett-Teller (BET) method, and the pore size distribution and total pore volume were determined by the Brunauer-Joyner-Hallenda (BJH) method.
A nominal ratio of 3 was fixed for all HTs synthesis, but as shown in Table 1, the observed molar Mg/Al ratio varied from 3.02 for MgAl-CP to 3.63 for MgAl-E. Assuming that the ROspecies are responsible for the Al3+ charge balance, a nROH/nAl ) 1 ratio means that all the net positive charge is compensated for by RO- anions. Thus, a simple way to corroborate the intercalation of the RO- anions in the interlayer is to analyze the nROH/nAl molar ratio. In sol-gel HTs, the nROH/nAl molar ratio decreased as the alcohol chain length increased, being 0.89, 0.70, and 0.32 for ethanol, 2-propanol, and 1-butanol, respectively (see Table 1). This behavior may be related to steric hindrance of the RO- anion to accommodate itself within the layers as the R group increases, thus giving rise to partial CO32occupancy (see Mass Spectrometry and Infrared Spectroscopy sections below). The weight loss curves of HTs (see Figure 1) were used to set the limits of temperature intervals at which a particular weight loss event occurred as well as its magnitude. The curves shapes were very similar among them, differing mainly in weight loss magnitude, which depends on the amount and nature of decomposable species. Hence, similar temperature intervals were set to represent all events in sol-gel HTs as well as in the coprecipitated one. The MS analysis was divided into three temperature regions: (i) between 25 and 250 °C from weakly bound species such as physisorbed water, hydroxyls, carbon species such as CO2, and ion fragments from the organic precursors; (ii) between 250 and 450 °C from the withdrawal of species located in the interlaminar region and from the partial decomposition of the layers; and (iii) above 450 °C from the total decomposition of the layers and the decomposition of the remaining species. In order to better understand the decomposition of sol-gel HTs, it is appropriate to start with the decomposition of a typical coprecipitated HT (MgAl-CP in Figure 2), since that of similar ones has already been widely studied.39-41 A weight loss of 18.7 wt % was obtained in the range between 25 and 250 °C, which has been reported as corresponding mainly to physisorbed water and CO2.4 MS analysis indicates that this loss derives
Results and Discussion Thermal Decomposition of Hydrotalcites. Commonly, water is the solvent of choice when HTs are prepared by coprecipitation methods, but when aliphatic alcohols are used as solvents, alkoxide anions form in the reaction medium and their incorporation into HTs is feasible.19 Approximate chemical formulas of sol-gel HTs and the coprecipitated one (as a comparison) are shown in Table 1, where Mg2+ and Al3+ cations were quantified by emission spectroscopy, while H2O and alkoxy anions were estimated from TG weight loss and MS spectra, as explained in the Experimental Section.
Figure 1. TG profiles for hydrotalcites.
MgAl Hydrotalcites with Nanocapsular Morphology
J. Phys. Chem. C, Vol. 111, No. 2, 2007 645
Figure 2. Profiles of products from thermally decomposing hydrotalcites, followed by mass spectrometry.
mostly from H2O + OH and a small contribution from CO2, assigned mainly to crystallization water and/or weakly bound hydroxyls and physisorbed CO2. Accordingly, the relative amount of each species will vary depending on the chemical composition, particle size, and texture of the sample,42,43 the layered structure of hydrotalcite being kept in this temperature range, as indicated elsewhere.44 Between 250 and 450 °C, a weight loss of 22.8% was recorded, attributed to H2O + OH and CO2. Water produced by dehydroxylation of the interlayer OH groups is still the main decomposition product, notwithstanding, and carbonate (CO2) is released at 418 °C. In the case of MgAl-CP, removal of carbonates from the interlayer region causes the structure to collapse, resulting in poorly crystalline MgO.4 In the range of 450-650 °C the sample lost ∼3.5% weight, which we ascribed to H2O + OH and CO2. This weight loss is obtained from the remaining CO32- anions strongly bonded and from the remaining hydroxyls of the laminae; only a negligible weight loss was obtained beyond this temperature range. The thermal decomposition of MgAl-E by TG and MS (see Figures 1 and 2) shows that between 25 and 250 °C, the corresponding weight loss is 16.2 wt % from the following fragment ions: H2O + OH, CO2, and CxHy. The thermal decomposition of the species from the sol-gel samples is more complex in comparison with the one prepared by coprecipitation. Here, physisorbed ethanol and adsorbed CO2 (captured from the environment) are released as ionic fragments of the ethoxy groups and CO2; however, the organic fraction amount in MgAl-E is fairly small below 250 °C. It can be observed that the ion fractions evolved in the 250-450 °C range (21.4 wt %) consisted of H2O + OH, CO2, and CxHy; therefore, within this temperature interval the majority of the organic species contained in the solid were released, thus indicating that the anions located in the interlaminar region are mostly organic in nature. Noteworthy, in the MS decomposition pattern, water is released
within the 250-450 °C range in two stages: (i) at ∼350 and (ii) at ∼415 °C. This behavior can be rationalized when it is considered that structural OH groups near the edges of the layers are the first ones to be released, followed by the ones located toward the bulk of the layers, which demand higher temperatures. In fact, between these two stages, emission of CxHy fragments is maximum (∼370 °C), suggesting that ethoxy group decomposition/release is not totally concomitant with that of OH groups. The weight loss (2.3 wt %) occurring between 450 and 650 °C is made up of H2O + OH, CO2, and CxHy. The TG analysis of MgAl-P showed a 22.6 wt % loss between 25 and 250 °C (see Figures 1 and 2) mainly from H2O + OH, a small contribution of CO2, and CxHy. In the 250-450 °C interval, a weight loss of 37.8% can be ascribed to H2O + OH, CO2, and CxHy. In this interval, and similar to the ethanol samples, OH groups are withdrawn in two steps at (i) ∼350 and (ii) ∼410 °C; CO2 and CxHy species are released at 380390 °C. The release of CxHy fragments at temperatures well above the boiling points of the corresponding alcohols is a strong indication that the organic species are mainly located inside the interlaminar region, that is, requiring a high energy for their decomposition. The weight loss above 450 °C (4.2 wt %) corresponds to H2O + OH, CO2, and CxHy. The material prepared with 1-butanol, MgAl-B, lost 26.8 wt % from 25 to 250 °C (see Figure 1 and 2) and corresponds to H2O + OH, CO2, and CxHy ion fractions. As opposed to the above sol-gel materials, the detection of CxHy fragments in MgAl-B at ∼150 °C indicates physisorbed 1-butanol or weakly bounded butoxy species, probably located on the external surface. Between 250 and 450 °C, a 23.2% weight loss from H2O + OH, CO2, and CxHy, was observed. From chemical analysis, a nROH/nAl ) 0.32 value was attained, which implies that there is a considerable shortage of alkoxy anions balancing the positive charge between the layers; hence, hydroxyls and/ or carbonate anions may be compensating the electronic charge
646 J. Phys. Chem. C, Vol. 111, No. 2, 2007
Valente et al.
TABLE 2: Hydrotalcite, Space Group R-3m: Lattice Parameters sample
a (nm)
c (nm)
MgAl-E MgAl-P MgAl-B MgAl-CP
0.308 91(5) 0.309 17(5) 0.308 39(7) 0.305 66(1)
2.811(2) 2.697(2) 2.603(2) 2.3192(3)
as well. In the 450-650 °C interval (2.5 wt % loss, from TG), the ion fractions determined by MS analysis were H2O + OH, CO2, and CxHy. The decomposition MS pattern of MgAl-CP (see Figure 2) is very different from those of sol-gel materials, but there is one common aspect in all hydrotalcite-like materials, regardless of being sol-gel or coprecipitated: dehydroxylation of the layers is partially interrupted when anions (carbonates or alkoxy groups) are withdrawn from the interlayer region. Carbonates in MgAl-CP are released at higher temperatures, ∼425 °C, than the corresponding fragments of alkoxy-containing sol-gel materials, that is, ∼370 °C. The reason for this may lie on the fact that MgAl-CP hydrotalcite is more crystalline than the solgel hydrotalcites (see Figure 3). Furthermore, the temperature interval at which ion fragments from the alkoxy groups are emitted is far from the boiling points of the parent alcohols (being 78.5, 82.4, and 117.2 °C, for ethanol, 2-propanol, and 1-butanol, respectively), indicating that the interaction of alkoxy species and hosting HT layers is a strong one. Alkoxy Intercalation. In a first approximation, capsule shells (see the Morphology Examination section) had the atomic local order of crystalline hydrotalcites, as was observed from the X-ray diffraction patterns (see Figure 3). The diffraction peak broadening may have three causes: The first one was linked to the small shells’ thickness, which produced size broadening; the second one was associated with the walls’ curvature, causing deformation of the octahedral sheets, which in turn distorted the unit cells and generated microstrains; and the last one was produced by the stacked alcohol molecules, which were disordered, causing additional microstrain. In spite of these difficulties, the diffraction patterns were modeled in a first approximation by using the Rietveld refining
Figure 4. Rietveld refinement plot of the sample prepared with 2-propanol (RF ) 2.42). In the upper curves the points correspond to the experimental diffraction pattern while the continuous line corresponds to the calculated one. The lower curve shows the difference between calculated and experimental data. The marks correspond to hydrotalcite.
method (Figure 4) to obtain the unit cell parameters (see Table 2). Since the Rietveld method is a full pattern analysis, the obtained lattice parameter c is more reliable than the one derived from only the (003) reflection, whose precise position depends on the background (Figures 3 and 4). The quasi-hydrotalcite atom distribution in the particles was modeled with a hexagonal unit cell, with the assumption of the R-3m space group in rhombohedral symmetry and the atom positions indicated in Table 3. Since almost no information is available about the corresponding alcohol’s atom distribution in the unit cell, they were simply modeled with the positions of the carbon and oxygen atoms in the carbonate group, associated with an ideal hydrotalcite without intercalated molecules. This causes a bad model of the (006) reflection (Figure 4), because the alcohol molecules are in the planes that generate this reflection. This rough approximation of the atom positions associated with the intercalated molecules also causes small variations in the other diffraction peaks, which are evident in the curve that describes the difference between experimental and calculated diffraction patterns (Figure 4). In the model used for the refinement, magnesium atoms occupied randomly 3/4 of the cation sites, while aluminum atoms occupied randomly 1/4 of them. After the Rietveld refinement analysis was performed, it was observed that the lattice parameter a was almost independent of the alcohol used in the synthesis (see Table 2); it was, however, a little larger than the a parameter obtained in the coprecipitated sample. This difference was probably caused by the curved surfaces associated with the nanocapsules (see the Morphology Examination section and Figure 7). In contrast, the TABLE 3: Hydrotalcite, Space Group R-3m: Fractional Atom Positionsa
Figure 3. X-ray powder diffraction patterns of the sol-gel hydrotalcites with different alcohols, and that of the coprecipitated HT.
atom
site
x
y
z
Mg Al O1 O2 C
3a 3a 6c 18h 6c
0.0 0.0 0.0 u 1 /3
0.0 0.0 0.0 -u 2 /3
0.0 0.0 w w1 w2
a Site 3a was 2/3 occupied by the magnesium atom and 1/3 by the aluminum atom. Site 18h had an occupancy of 1/6 as well as the site occupied by the carbon atom. The zero approximations for the variable parameters were w ) 0.38, u ) 0.13, w1 ) 0.51, and w2 ) 0.53.
MgAl Hydrotalcites with Nanocapsular Morphology
J. Phys. Chem. C, Vol. 111, No. 2, 2007 647
Figure 5. Proposed arrangement of alcohol molecules within the HT interlayer: (a) ethanol, (b) 2-propanol, and (c) 1-butanol.
effect of the alcohol on the lattice parameter c was much larger and depended on the alcohol type (see Table 2). Intercalation of alcohol molecules into the HT interlayer resulted in increased basal spacing, as shown in Table 2, where the values of the c parameter ranged from 28.11 Å for MgAl-E to 23.192 Å for MgAl-CP. If it is assumed that the brucite layer has a thickness of 4.8 Å45 and the c parameters reported in Table 2 are used, the gallery heights are determined as 4.57, 4.19, 3.87, and 2.93 Å for MgAl-E, MgAl-P, MgAl-B, and MgAl-CP respectively. These interlayer distances were in agreement with the alkoxy anions’ size in various orientations (see Figure 5). For instance, ethoxy anions are thought to be intercalated perpendicular to the HT layers, while 1-butoxy anions could be oriented with the alkyl chain along the HT layers (Figure 5). This fact explains why the molar nROH/nAl ratios of ethoxy-intercalated HT showed the highest value, for example, close-packed orientation, while the butoxy-intercalated HT showed the lowest ratio (see Table 1). Given their tetrahedral symmetry around the secondary carbon, 2-propoxy anions could be accommodated with the oxygen pointing toward the layers with the two methyl groups antipode to the oxygen, as illustrated in Figure 5b. 2-Propoxy anions, being bulkier, showed a lower nROH/nAl ratio than the ethoxy ones (see Table 1). These findings agree with previous studies, where different alcohols were used as solvents or intercalated in HT materials.19,46 Evidence for Intercalated Alkoxy Groups. The DRIFT spectra (4000-2500 cm-1 region) of the hydrotalcite-like compounds are shown in Figure 6 at the indicated temperatures. The MgAl-E material heat-treated at 50 °C shows (i) a peak at 3707 cm-1 from isolated Mg-OH groups;47 (ii) a broad tailing peak at about 3550 cm-1, typical of hydrogen-bonded OH groups forming MgAl HT layers48 (including those of interstitial
Figure 6. DRIFT spectra of hydrotalcites after heat treatment at (a) 50 °C, (b) 250 °C, and (c) 500 °C.
or adsorbed water); and (iii) two peaks at ∼2945 and ∼2831 cm-1 from -CH2 and/or -CH3 stretching modes.49 These groups are likely to be associated with ethoxy groups from ethanol. Accordingly, Prinetto et al.31 detected similar bands in sol-gel MgAl hydrotalcites prepared with magnesium ethoxide, aluminum acetylacetonate, and ethanol. Moreover, Ranjit and Klabunde50 observed IR bands (2920, 2846, and 2788 cm-1) associated with methoxy groups in the hydrolysis of Mg(OCH3)2. When MgAl-E was heat-treated at 250 or 500 °C, practically all these bands can be observed, although with lower intensity, especially those of OH groups (i.e., 3705 and 3550 cm-1), indicating a partial dehydroxylation event. In spite of lower intensities, the -CHx group bands can be observed even after annealing at 500 °C, suggesting high stability in remaining
648 J. Phys. Chem. C, Vol. 111, No. 2, 2007 ethoxy groups. The DRIFT behavior of both MgAl-P and MgAl-B (see Figure 6) is very similar to that of the one prepared with ethanol, showing bands clearly associated with -CHx groups, even at 500 °C, again indicating high stability in remaining propoxy and butoxy groups, as discussed above. The remaining -OH and -CHx groups at 500 °C are very strongly bound, and as indicated in the MS profiles (see discussion above), they require temperatures above 600 °C to be fully withdrawn. As a comparison, the DRIFT spectra of MgAl-CP at 50, 250, and 500 °C are shown in Figure 6. The spectra at 50-500 °C are mainly characterized by broad tailing peaks between 3600 and 2800 cm-1 (from -OH groups as discussed above), but bands associated with -CH3 or -CH2 groups are absent. Notice in all spectra in Figure 6 that removal of -OH groups is incomplete at 500 °C, suggesting that remaining -OH groups are very strongly fixed to HT materials. The removal of these refractory -OH groups becomes difficult; as the dehydroxylation advances, the remaining pseudolayers of HT material acquire a more oxidic character because the population of strong basic sites increases, interacting more strongly with isolated OH groups. Additionally, an OH needs vicinal -OH groups to first form water and then be withdrawn from the layers, but this is hindered by the OH isolation conditions as dehydroxylation advances. In the IR region between 1750 and 1000 cm-1, a characteristic 1085 cm-1 band from CO in alkoxy groups can be observed in all samples prepared by the sol-gel method. This vibration, however, is absent in the sample prepared by a coprecipitation procedure where alcohols were not employed. A band at 1650 cm-1, typical of water bending mode, can be noted in all spectra heated at 50 °C, which disappeared upon calcining at 250 °C (compare all spectra a and b in Figure 6), indicating full removal of water molecules below 250 °C. Other vibrations can be observed in most sol-gel HTs, indicating that intercalated alkoxy anions are accompanied by carbonate and/or bicarbonate anions. The various bands observed between 1350 and 1650 cm-1 are mainly due to CO32- anions in different symmetries. For instance, the bands at 1581 and 1360-1385 cm-1 in MgAl-E, -P, and -B arise from CO32- in D3h planar symmetry. The band at 1420 cm-1 is very likely to arise from bicarbonate anions. However, upon annealing at 250 and at 500 °C, the 1350-1650 cm-1 region becomes overlapped with new bands. The individual assignment of these bands is complicated by the fact that (i) the ν3 vibration mode of CO32anions generally splits into two other bands when HTs are annealed51,52 and (ii) CO2, resulting from the thermal decomposition of alkoxy anions, may be adsorbing on basic sites and contributing to yield other carbonate species with a mixture of different symmetries. Commonly carbonate anions in coprecipitated HTs show an intense band at ca. 1360-1370 cm-1 from ν3 stretching mode,53,54 but in MgAl-CP the band at 1415 cm-1 indicates the presence of bicarbonates, whose formation has been claimed in some cases.4 Interlayered HCO3- species, form because of pH lowering during washing of the samples (CO32- + H2O f HCO3- + OH-).4 As the annealing temperature increases, this band splits into several bands at 250 °C, probably from variations in HCO3- symmetry, and at 500 °C a band at 1363 cm-1 indicates the presence of planar CO32- anions.52 The small band at 1168 cm-1 is due to Mg-O bonds.55 Morphology Examination. The transmission electron micrograph of the samples prepared by the sol-gel method shows that the samples were made up of rings (see Figure 7). Because
Valente et al.
Figure 7. TEM micrographs of the samples prepared with ethanol (MgAl-E), 2-propanol (MgAl-P), or 1-butanol (MgAl-B) and by coprecipitation (MgAl-CP).
the image was formed with transmitted electrons, these apparent rings correspond to nanocapsules with diameters between 10 and 40 nm (Figure 7); a similar morphology was observed in boehmite obtained by the sol-gel method.37 In contrast, the sample prepared by coprecipitation was made up of ribbons (see Figure 7, MgAl-CP), probably formed by the interaction among them when capsules were opened since the aggregation or coalescence of such capsules is favored by the water’s presence, which is the case in conventional HT synthesis. However, it should be recalled that the capsules, which are colloidal particles of the sol made during the synthesis, could aggregate even in the absence of water. This aggregation depends on several synthesis parameters such as the reaction temperature, the solute concentration, and the addition of external molecules like surfactants to the solution during the synthesis, etc. Thus, the capsule’s size and homogeneity could be adjusted by controlling such parameters. As far as we know, this is the first time that a capsular morphology of HT compounds has been reported, regardless of the technique employed.3,19,28-34 The importance of this result is based on the increasing interest in producing hollow inorganic nanostructures in order to use them as drug-delivery carriers, catalysts, optical and sensing devices, etc.20 Accordingly, the mild synthesis conditions developed in the present work open new application issues. It must be remarked that high-resolution microscopy, performed on the HT sol-gel samples, shows that the capsule’s shell is lamellar (see inset in Figure 7, MgAl-B). As discussed in the X-ray Diffraction section, the interlayer distance depended on the nature of the alcohol used as a solvent. Thus, the interlayer distance observed in the inset in Figure 7 agrees with the c value determined by XRD (see Table 2). Hence, the crystalline platelets in the spheroids are oriented perpendicular to this crystallographic axis. It is important to notice that after calcination lamellar HT structure collapses and then a solid solution is produced. Figure 8 shows the TEM image of calcined hydrotalcite MgAl-P; the nanocapsular morphology allows the formation of primary particles of spheroidal shape, with sizes ranging from 7 to 9
MgAl Hydrotalcites with Nanocapsular Morphology
Figure 8. TEM micrograph of a calcined sample prepared with 2-propanol (MgAl-P).
Figure 9. SEM micrographs of the samples prepared with ethanol (MgAl-E), 2-propanol (MgAl-P), or 1-butanol (MgAl-B) and by coprecipitation (MgAl-CP).
nm. Similar results were obtained over the other sol-gel calcined hydrotalcites. On the other hand, the sample synthesized by coprecipitation (MgAl-CP), after calcination, showed the characteristic platelet-like morphology of hydrotalcites, which was in good agreement with findings already stated elsewhere.39 SEM analysis of MgAl-CP (see Figure 9) revealed that the sample was made up of irregularly shaped platelets with sizes larger than 10 µm, while MgAl-E, MgAl-P, and MgAl-B consisted of rounded aggregates with average sizes of 338, 354, and 293 nm respectively (see Figure 9). The morphology in the sample obtained by coprecipitation (MgAl-CP) agrees well with that already reported in the literature, while the SEM micrographs of the sol-gel samples show that their morphology is quite different. Even if this difference was noted in some works, no explanation was given for it.28,31 The origin of this difference most likely lies within the capsular morphology, which is a consequence of the method used here. The meaningfulness of this result can be appreciated with the following example: Pinnavaia’s group19 has reported a successful synthesis of intercalated hydrotalcites to obtain thin films. Accordingly, SEM images of films cast from an intercalated HT with methoxy anions and a conventional coprecipitated HT were compared. The intercalated HT provided a transparent and smooth continuous film while the conventional
J. Phys. Chem. C, Vol. 111, No. 2, 2007 649 HT afforded an opaque, rough film.19 Hence, we are certain that the procedure proposed here should lead to similar results because the small rounded aggregates could be easily deposited on a surface. Textural Properties. The adsorption-desorption isotherms of the calcined samples showed that all are of type IV, which characterizes mesoporous solids, as shown in Figure 10. A multilayer nitrogen adsorption process takes place from the beginning due to the low relative pressure (P/P0) where the adsorption start indicates the weak adsorption enthalpy of adsorbate-adsorbent. However, the sol-gel samples showed clearly two steps in their adsorption path, the first one at approximately P/P0 ) 0.6 and the second one at approximately 0.85, while the path in the coprecipitated sample was a continuous one (see Figure 10). Taking into account that the adsorption process is intimately related to the solid’s morphology,56 it is obvious that morphological variations should occur when the sol-gel method is employed instead of coprecipitation. This difference could be understood by analyzing the hysteresis loops since these manifest a capillary condensation of adsorbate in meso and macropores of solids. The hysteresis loop for the calcined coprecipitated HT sample, MgAl-CP, is type H3,56 which, in agreement with the morphology of coprecipitated HT materials, is usually given by adsorbents containing slit-shaped pores.56 The two steps observed in the hysteresis loops of the sol-gel HT samples indicate a bimodal pore size distribution (see Figure 10). The first narrow hysteresis loop in the range 0.6 < P/P0 < 0.85 evidence pores with a regular geometry, while the second one, in the range 0.85 < P/P0 < 1.0, denotes meso- and macropores with a nonuniform distribution, as shown in Figure 10 insets, where a bimodal pore size distribution was observed, with the first one being homogeneous and the second one presenting a wide pore size distribution. The coprecipitated sample showed a heterogeneous pore size distribution (see corresponding inset in Figure 10). In the case of sol-gel samples, the total pore volume, reported in Table 4, is mainly due to the first homogeneous pore size distribution. The pore size population observed on these solids is meaningful because smaller pores are predominant over larger ones. In contrast, during the coprecipitation procedure the pore volume was originated by the interparticle stacking created during the calcination treatment. In this sense, the bimodal distribution of the sol-gel samples can be explained as follows: the first distribution was produced by intraparticle porosity related to the collapsed capsules, since spheroidal particles ranging from 7 to 9 nm are formed after thermal treatment of the sol-gel samples (see Figure 8), while the second one was likely created by interparticle arrangement during the calcination process associated with the agglomerates observed by SEM (see Figure 9). Specific surface areas, pore volume, and average pore size of calcined solids are reported in Table 4. The surface areas were 289, 274, 332, and 254 m2/g for MgAl-E, MgAl-P, MgAlB, and MgAl-CP, respectively. It is important to note the high pore volume found on the calcined compounds because pore volume ranges from 0.812 cm3/g for MgAl-E to 1.390 cm3/g for MgAl-B. In summary, the sol-gel samples showed the highest specific surface areas, reaching the maximum on the calcined MgAl-B with 332 m2/g. In Table 4, the maximum pore size distributions are also presented. Because the sol-gel samples showed a well-defined bimodal distribution, two average pore diameters are reported, I and II. The average pore diameters of sol-gel solids are between 50 and 90 Å, and those in the coprecipitated HT are between 200 and 500 Å.
650 J. Phys. Chem. C, Vol. 111, No. 2, 2007
Valente et al.
Figure 10. N2 adsorption-desorption isotherms of HTs, calcined at 500 °C: (2) MgAl-E, (() MgAl-P, (b) MgAl-B, and (9) MgAl-CP. (Inset) Corresponding pore size distributions.
Conclusions The hydrotalcites obtained by the sol-gel procedure described in this work show nanometric hollow spheroidal morphology. To our knowledge, this is the first time that such morphology has been reported in the case of HT materials. In contrast to other sol-gel methods used to produce hydrotalcites, our method focused on restricting the addition of water to control hydrolysis reactions. A direct benefit of such nanocapsular shape is reflected in the textural properties of these sol-gel HTs; that is, specific surface area and pore volume values ranged between 289∼332 m2/g and 0.81∼1.39 cm3/g, respectively, depending on the alcohol used in the sol-gel procedure. Alkoxy groups, from the alcohols employed as solvents, are very likely to be located within the interlayer region of sol-gel hydrotalcites, as indicated by XRD and DRIFT analyses. In comparison with coprecipitated MgAl-CO3 HT particles (>10 µm), sol-gel HTs showed very small particle aggregates (ca. 300∼350 nm). Future applications of nanocapsular HTs are very attractive, given the increasing interest in producing hollow inorganic nanostructures that can use them as drug delivery carriers, catalysts, optical (thin films) and sensing devices, etc. TABLE 4: Textural Analysis of Calcined HTs pore diameter (Å) samplea
BET (m2/g)
pore volume (cm3/g)
MgAl-E MgAl-P MgAl-B MgAl-CP
289 274 332 254
0.812 0.814 1.390 1.148
a
Calcined at 500 °C for 4 h.
I
II
67 50 80 124
415 500 388 397
Acknowledgment. We thank the Mexican Institute of Petroleum for the financial support and Jaime F. Jaramillo for providing technical support. References and Notes (1) Braterman, P. S.; Xu, Z. P.; Yarberry, F. Handbook of Layered Materials; Marcel Dekker: New York, 2004. (2) De Roy, A.; Forano, C.; El Malki, K.; Besse, J. P. Synthesis of Microporous Materials; Van Nostrand-Reinhold: New York, 1992. (3) Figueras, F. Top. Catal. 2004, 29, 189. (4) Cavani, F.; Trifiro, F.; Vaccari, A. Catal. Today 1991, 11, 173. (5) Miyata, S. Clays Clay Miner. 1975, 23, 369. (6) Miyata, S. Clays Clay Miner. 1980, 28, 50. (7) Reichle, W. T. U.S. Patent 4 086 188, 1978. (8) Figueras, F.; Tichit, D.; Bennani N. M.; Ruiz, R. Catalysis of Organic Reactions; Marcel Dekker Inc.: New York, 1998. (9) Rao, K. K.; Gravelle, M.; Valente, J. S.; Figueras, F. J. Catal. 1998, 173, 115. (10) Kumbhar, P. S.; Valente, J. S.; Figueras, F. Tetrahedron Lett. 1998, 39, 2573. (11) Kumbhar, P. S.; Valente, J. S.; Millet, J. M.; Figueras, F. J. Catal. 2000, 191 (2), 467. (12) Choudary, B. M.; Kantam, M. L.; Reddy, Ch.V.; Rao, K. K.; Figueras, F. J. Mol. Catal. A 1999, 146, 279. (13) Prescott, H. A.; Li, Z-J.; Kemnitz, E.; Trunschke, A.; Deutsch, J.; Lieske, H.; Auroux, A. J. Catal. 2005, 234, 124. (14) Choudary, B. M.; Kantam, M. L.; Reddy, C.; Reddy, V.; Bharathi, B.; Figueras, F. J. Catal. 2003, 218, 191. (15) Kumbhar, P. S.; Valente, J. S.; Lopez, J.; Figueras, F. J. Chem. Soc., Chem. Commun. 1998, 7 (5), 535. (16) Kumbhar, P. S.; Valente, J. S.; Figueras, F. J. Chem. Soc., Chem. Commun. 1998, 10, 1091. (17) Obalova´, L.; Jira´tova´, K.; Kovanda, F.; Pacultova´, K.; Lacny´, Z.; Mikulova, Z. Appl. Catal. B 2005, 60, 289. (18) Cantu´, M.; Lo´pez-Salinas, E.; Montiel, R.; Valente, J. S. EnViron. Sci. Technol. 2005, 39, 9715.
MgAl Hydrotalcites with Nanocapsular Morphology (19) Gardner, E.; Huntoon, K. M.; Pinnavaia, T. J. AdV. Mater. 2001, 16, 13. (20) Sun, Y.; Xia, Y. Science 2002, 298, 2176. (21) Itaya, K.; Chang, H.-C.; Uchida, I. Inorg. Chem. 1987, 26, 624. (22) Darder, M.; Lo´pez-Blanco, M.; Aranda, P.; Leroux, F.; Ruiz-Hitzky, E. Chem. Mater. 2005, 17, 1969. (23) Reinholdt, M. X.; Kirkpatrick, R. J. Chem. Mater. 2006, 18, 2567. (24) Choy, J.-H.; Kwak, S.-Y.; Park, J.-S.; Jeong, Y.-J.; Portier, J. J. Am. Chem. Soc. 1999, 121, 1399. (25) Jaubertie, C.; Holgado, M. J.; San, Roma´n, M. S.; Rives, V. Chem. Mater. 2006, 18, 3114. (26) Tichit, D.; Rolland, A.; Prinetto, F.; Fetter, G.; Martinez-Ortiz, M. J.; Valenzuela, M. A.; Bosch, P. J. Mater. Chem. 2002, 12, 3832. (27) Climent, M. J.; Corma, A.; Iborra, S.; Epping, K.; Velty, A. J. Catal. 2004, 225, 316. (28) Lopez, T.; Bosch, P.; Ramos, E.; Gomez, R.; Novaro, O.; Acosta, D.; Figueras, F. Langmuir 1996, 12, 189. (29) Lopez, T.; Bosch, P.; Asomoza, M.; Gomez, R.; Ramos, E. Mater. Lett. 1997, 31, 311. (30) Wang, J. A.; Morales, A.; Bokhimi, X.; Novaro, O.; Lo´pez, T.; Go´mez, R. Chem. Mater. 1999, 11, 308. (31) Prinetto, F.; Ghiotti, G.; Graffin, P.; Tichit, D. Microporous Mesoporous Mater. 2000, 39, 229. (32) Bolognini, M.; Cavani, F.; Scagliarini, D.; Flego, C.; Perego, C.; Saba, M. Microporous Mesoporous Mater. 2003, 66, 77. (33) Tichit, D.; Lorret, O.; Coq, B.; Prinetto, F.; Ghiotti, G. Microporous Mesoporous. Mater. 2005, 80, 213. (34) Othman, M. R.; Rasid, N. M.; Fernando, W. J. N. Microporous Mesoporous Mater. 2006, 93, 23. (35) Brinker, C. J.; Scherer, G. W. Sol-Gel Science; Academic Press: New York, 1989. (36) Lima, E.; Valente, J. S.; Bosch, P.; Lara, V. J. Phys. Chem. B 2005, 109, 17435. (37) Bokhimi, X.; Lima, E.; Valente, J. S. J. Phys. Chem. B 2005, 109, 22222.
J. Phys. Chem. C, Vol. 111, No. 2, 2007 651 (38) Prince, E. J. Appl. Crystallogr. 1981, 14, 157. (39) Reichle, W. T.; Kang, S. Y.; Everhardt, D. S. J. Catal. 1986, 101, 352. (40) Sato, T.; Kato, K.; Endo, T.; Shimada, M. React. Solids 1986, 2, 253. (41) Roelofs, J. C. A. A.; van Bokhoven, J. A., van Dillen, A. J.; Geus, J. W.; de Jong, K. P. Chem. Eur. J. 2002, 8, 5571. (42) Yun, S. K.; Pinnavaia, T. J. Chem. Mater. 1995, 7, 348. (43) Valente, J. S.; Figueras, F.; Gravelle, M.; Khumbar, P.; Lo´pez, J.; Besse, J.-P. J. Catal. 2000, 189, 370. (44) Yang, W., Kim, Y., Liu, P. K. T., Sahimi, M., Tsotsis, T. T. Chem. Eng. Sci. 2002, 57, 2945. (45) Kanezaki, E.; Kunugawa, K.; Ishikawa, Y. Chem. Phys. Lett. 1994, 226, 325. (46) Costantino, U.; Marmottini, F.; Nocchetti, M.; Vivani, R. Eur. J. Inorg. Chem. 1998, 1439. (47) Coluccia, S.; Lavagnino, S.; Marchese, L. Mater. Chem. Phys. 1988, 18, 445. (48) Utamapanya, S.; Klabunde, K. J.; Schlup, J. R. Chem. Mater. 1991, 3, 175. (49) Harris, D. C.; Bertolucci, M. D. Symmetry and Spectroscopy; Dover: New York, 1989. (50) Ranjit, K. T.; Klabunde, K. J. Chem. Mater. 2005, 17, 65. (51) Brindley, G. W.; Kikkawa, S. Am. Mineral. 1979, 64, 836. (52) Rey, F.; Fornes, V.; Rojo, J. M. J. Chem. Soc., Faraday Trans. 1992, 88, 2233. (53) Nakamoto, K. Infrared and Raman spectra of inorganic and coordination compounds, 3rd ed.; Wiley: New York, 1978; p 243. (54) Yang, W.; Kim, Y.; Liu, P. K. T.; Sahimi, M.; Tsotsis, T. T. Chem. Eng. Sci. 2002, 57, 2945. (55) Yamaguchi, Y.; Taguchi, H.; Miyata, Y.; Yoshinaka, M.; Shimizu, K. Polyhedron 1987, 6, 1587. (56) Rouquerol, F.; Rouquerol, J.; Sing, K. Adsorption by Powder & Porous Solids; Academic Press: New York, 1999.
