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Homogeneous Precipitation by Formamide Hydrolysis: Synthesis, Reversible Hydration, and Aqueous Exfoliation of the Layered Double Hydroxide (LDH) of Ni and Al G. V. Manohara,† Daniel A. Kunz,‡ P. Vishnu Kamath,*,† Wolfgang Milius,‡ and Josef Breu*,‡ †
Department of Chemistry, Central College, Bangalore University, Bangalore 560 001, India, and ‡ Department of Inorganic Chemistry I, University of Bayreuth, Bayreuth, Germany Received June 2, 2010. Revised Manuscript Received August 28, 2010
Homogenous precipitation by formamide hydrolysis results in the formation of a formate-intercalated layered double hydroxide (LDH) of Ni(II) and Al(III). The formate-LDH is sensitive to the atmospheric humidity and reversibly exchanges its intercalated water with atmospheric moisture. The hydration/dehydration cycle is complete within a narrow range of 0-30% relative humidity with significant hysteresis and involves a randomly interstratified intermediate phase. When immersed in water, the formate ion grows its hydration sphere (osmotic swelling), eventually leading to the exfoliation of the metal hydroxide layers into lamellar particles having in-plane dimensions of 100-200 nm and a thickness of 9-12 nm. These nanoplatelets restack to thicker tactoids again upon evaporation of the dispersion. The intercalated formate ion can be exchanged with nitrate ions in solution but not with iodide ions. These observations have implications for many applications of LDHs in the area of carbon dioxide sorption and catalysis.
1. Introduction The layered double hydroxides (LDHs) are a class of inorganic layered materials with applications in catalysis,1 environmental amelioration,2 carbon dioxide adsorption,3 fire retardants,4 magnetic inorganic-organic hybrids,5 and electrodes for alkaline secondary batteries.6 For all of these applications, it is necessary to enhance the external surface area of the material. The LDHs comprise positively charged layers having the composition [M1 - x2þM0 x3þ(OH)2]xþ (M = Mg, Co, Ni; M0 = Al, Cr, Fe; 0.2 e x e 0.33) and thereby possess large internal surfaces that are generally not accessible because the interlayer region is occupied by carbonate ions.7 The carbonates not only restore charge neutrality but also bond with the metal hydroxide slabs by strong H bonds and facilitate a compact packing of the metal hydroxide slabs. Of particular interest is the property of surface basicity, which makes the LDHs potent materials for base-catalyzed reactions8 of industrial importance. Even more important than the bulk basicity is the site-specific basicity, which is based on the local structure of the solid, and the anion and its mode of incorporation. By replacing the carbonate with other anions, the surface basicity can be tuned.9 The LDHs are low-temperature phases and completely decompose to yield oxide residues at 450 °C, which is far below typical sintering temperatures. They undergo a 30-42% mass loss during decomposition to yield *Corresponding authors. E-mail:
[email protected] (P.V.K.),
[email protected] (J.B.).
(1) Cavani, F.; Trifiro, F.; Vaccari, A. Catal. Today 1991, 11, 173. (2) Prasanna, S. V.; Kamath, P. V. Ind. Eng. Chem. Res. 2009, 48, 6315. (3) Hutson, N. D.; Speakman, S. A.; Payzant, A. E. Chem. Mater. 2004, 16, 4135. (4) Costache, M. C.; Wang, D. Y.; Heidecker, M. J.; Manias, E.; Wilkie, C. A. Polym. Adv. Technol. 2006, 17, 272. (5) Kurmoo, M. Chem. Mater. 1999, 11, 3370. (6) Kamath, P. V.; Dixit, M.; Indira, L.; Shukla, A. K.; Kumar, V. G.; Munichandraiah, N. J. Electrochem. Soc. 1994, 141, 2956. (7) Taylor, H. F. W. Miner. Mag. 1973, 39, 377. (8) Reichle, W. T. J. Catal. 1985, 94, 574. (9) Constantino, V. R. L.; Pinnavaia, T. J. Inorg. Chem. 1995, 34, 883. (10) Behrens, M.; Kasatkin, I.; K€uhl, S.; Weinberg, G. Chem. Mater. 2010, 22, 386.
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nanoparticulate oxide materials that are useful solid acid catalysts.10 To exploit the full potential of LDHs for all of these applications, it is necessary to go beyond the morphological outer surface of the material and make the structural inner surface available for sorption and catalysis. To achieve this objective, the metal hydroxide layers have to be pried apart by exfoliating them into thinner tactoids having a thickness of nanometer dimensions.11 The limiting surface area is generated when the thickness of the tactoids approaches that of a single layer, a phenomenon known as delamination.11 The delamination of LDHs is of interest in the emerging area of nanoelectronics because it generates positively charged nanosheets, in contrast with other oxide nanosheets that carry a negative charge.12 Any candidate material for exfoliation must have the ability to exchange its interlayer species for others in solution. Furthermore, any increase in the basal spacing, for instance, by incorporating water molecules into the interlamellar space (swelling), weakens the Coulombic attraction between the layer and interlayer and renders the tactoid more labile to shearing, thus fostering exfoliation.13 The LDHs exhibit a wide range of anion-exchange reactions. 14 A typical case is the family of carboxylate-exchanged LDHs, the chief interest being in the ability to tune the interlayer spacing from 7.6 A˚ in the case of the carbonate to 44 A˚ in the case of an interdigitated dodecylcarboxylate bilayer included between the metal hydroxide sheets.15 However, exfoliation is difficult to achieve because the layer charge is high compared to that of the smectite-type clays16 and the solvation enthalpy of interlayer anions tends to be lower than that of interlayer cations in smectites. Although the exfoliation of smectite-type clays is facile in water, the exfoliation of (11) Lagaly, G.; Gardolinsky, J. E. F. C. Clay Miner. 2005, 547. (12) Osada, M.; Sasaki, T. J. Mater. Chem. 2009, 19, 2503. (13) Moller, M. W.; Handge, U. A.; Kunz, D. A.; Lunkenbein, T.; Altst€adt, V.; Breu, J. ACS Nano 2010, 4, 717. (14) Meyn, M.; Beneke, K.; Lagaly, G. Inorg. Chem. 1990, 29, 5201. (15) Newmann, S. P.; Jones, W. New. J. Chem. 1998, 105. (16) Pinnavaia, T. J. Science 1983, 220, 4595.
Published on Web 09/14/2010
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LDHs so far could be achieved only by the interplay of hydrophilic/hydrophobic forces. Soaking long-chain alkylsulfonate/ phosfonate/carboxylate-LDHs in a solvent such as n-alkanol (n = 4-6) causes the n-alkanol molecules to enter the hydrophobic interlayer region and the metal hydroxide layers to be exfoliated.17,18 We decided to attempt the exfoliation of the LDHs in a single-step procedure in water by employing a different strategy based on growing the hydrogen-bonded network of water molecules in the coordination sphere of the intercalated anion. For this purpose, it is crucial to choose an anion with a high hydration enthalpy. MD simulations show that the hydration of citrate ions has no potential energy minimum and approaches that of bulk water.19 Hibino et al.20 have demonstrated the exfoliation of lactate-intercalated LDH in water. However, smallest among the carboxylates is the formate ion, which is expected to have the highest hydration enthalpy. Efforts to make formateintercalated LDHs by conventional coprecipitation techniques have generally been unsuccessful.21 There could be many reasons for this. (1) Navrotsky and co-workers22 have shown that the thermodynamic stability of the LDH decreases if the solubility of the salt of the divalent metal increases. Divalent metal formates are generally more soluble than the corresponding carbonates, and this adversely impacts the stability of formate-LDHs. (2) Carbonates are ubiquitous in the aqueous reaction media and successfully compete for inclusion in the interlayer region of the LDHs in preference to the formate ion on account of their higher charge. Iyi and co-workers23 reported the synthesis of the formateintercalated LDH among other carboxylate-LDHs in both the hydrated (basal spacing 11 A˚) and dehydrated (basal spacing 7.8 A˚) forms. These materials were prepared by ion exchange starting from a Cl--LDH precursor. The exchange was carried out in two steps, each involving a 380-fold stoichiometric excess of the formate ion. Gordijo and co-workers21 reported the formation of a formate-LDH by the decarbonation of the carbonate-LDH in a formamide-ethanol mixed solvent. They report a basal spacing of 7.6 A˚ similar to that of the carbonate precursor. In this article, we report a facile, direct synthesis of the Ni-Al formate LDH by homogeneous precipitation by formamide hydrolysis and report on its exchange reactions, reversible hydration, and exfoliation in water.
2. Experimental Section 2.1. Synthesis. Ni(NO3)2 3 6H2O, Al(NO3)3 3 9H2O, and formamide were procured from Merck (India) and used without further purification. All solutions were prepared using type-II water (specific resistance 15 MΩ cm, Millipore Elix-3 water purification system). Metal nitrate solutions of Ni2þ and Al3þ of known strengths were placed in Teflon-lined stainless steel autoclaves (80 mL capacity) along with formamide. The resultant reaction mixture was made up to 40 mL with decarbonated water to get final concentrations of 0.335 M (Ni), 0.165 M (Al), and 5 M (formamide). The reaction mixture (pH 2.6) was hydrothermally (17) Khaldi, M.; De Roy, A.; Chaouch, M.; Besse, J. P. J. Solid State Chem. 1997, 130, 66. (18) Liu, Z.; Ma, R.; Ebina, Y.; Iyi, N.; Takada, K.; Sasaki, T. Langmuir 2007, 23, 861. (19) Padmakumar, P.; Kalinichev, A. G.; Kirkpatrick, R. J. J. Phys. Chem. B 2006, 110, 3841. (20) Hibino, T.; Kobayashi, M. J. Mater. Chem. 2005, 15, 653. (21) Gordijo, C. R.; Constantino, V. R. L.; Silva, D. O. J. Solid State Chem. 2007, 180, 1967. (22) Allada, R. K.; Navrotsky, A.; Berbeco, H. T.; Casey, W. H. Science 2002, 296, 721. (23) Iyi, N.; Ebina, Y.; Sasaki, T. Langmuir 2008, 24, 5591.
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treated at 150 °C for 24 h, and the resultant product was recovered by centrifugation, washed, and dried at 65 °C for 24 h. For water decarbonation, dissolved CO2 was expelled by heating water to 65 °C and then bubbling nitrogen gas. 2.2. Anion Exchange. The Ni/Al-HCOO LDH (0.3 g) was placed in a 100 mL screw-capped bottle. To this, 50 mL of a decarbonated solution of NaX (X = NO3, I) containing 3 times the stoichiometric requirement of the incoming anion was added. The suspension was stirred for 24 h, and the product was separated by filtration, washing, and drying. In the case of Iexchange, the solid after the exchange reaction was recovered by fast filtration to minimize the effects of washing. 2.3. Exfoliation. The Ni/Al-HCOO LDH (0.1 g) was placed in a 250 mL screw-capped bottle. To this, 100 mL of decarbonated water was added, and the bottle was tightly sealed. This suspension was stirred for 96 h, diluted two times to about 200 mL in total, and ultrasonicated for 20 min. 2.4. Characterization. The Ni and Al contents in the LDH were estimated using an inductively coupled plasma optical emission spectroscopy (ICP-OES) technique (Perkin-Elmer Model Optima-2100DV). The formate content was determined by ion chromatography using a Metrohm 861 Advanced Compact ion chromatograph with a Metrosep A SUPP5 250 anion column and a conductivity detector. The formate content was less than what was expected from the Al content. To obtain a charge-balanced formula, this shortfall was made up by the inclusion of the requisite number of carbonate ions in the molecular formula. The unaccounted for weight was attributed to the intercalated water content to arrive at the complete formula for the LDH. The net mass loss expected for this composition was verified by TGA studies (Mettler Toledo TG/SDTA 851e system driven by stare software, heating rate 5 °C min-1, flowing N2). All LDH samples were characterized by powder X-ray diffraction (PXRD) using a Bruker aXS model D8 Advance diffractometer (θ-2θ scan, Cu KR radiation, λ = 1.541 A˚). Data were collected at a continuous scan rate of 2° 2θ min-1. In situ measurements of the PXRD patterns under controlled humidity condition were carried out using a PANanalytical X’pert pro X-ray diffractometer (θ-2θ scan, step size of 0.017° 2θ, Cu KR radiation, λ = 1.541 A˚, Bragg-Brentano geometry) equipped with an X’celerator scientific RTMS and an Anton Parr temperature humidity chamber driven by a VTI Corp. RH-200 humidity generator. Multiple measurements were made at each value of the relative humidity to monitor the equilibration. Infrared spectra of all of the samples were recorded using a Bruker Alpha-P FTIR spectrometer (ATR mode, diamond crystal, 400-4000 cm-1, 4 cm-1 resolution). Scanning electron microscopy (SEM) images were obtained using a JEOL JSM 6490LV scanning electron microscope operated at 15 kV. Powder samples were spread over conducting carbon tape and sputter coated with Pt to improve the conductivity. Topographical images of the exfoliated LDHs were recorded using an MFP3D atomic force microscope (Asylum Research, Santa Barbara, CA) equipped with silicon cantilevers (silicon tip, type NSC15/A1BS, μmash, Tallin, Estonia). A few drops of the LDH suspension were dropped onto a silicon wafer previously cleaned with a CO2 snow jet and spin coated at 2000 rpm. The scan rate was 1 Hz.
3. Results Because the formate-containing mixed-metal solution is acidic (pH 2.6), hydrothermal treatment at elevated temperatures results in acid-catalyzed formamide hydrolysis24 HCONH2 þ H2 O f HCOOH þ NH3 This is an ideal route to synthesizing formate-containing LDHs because the generated ammonium hydroxide helps in precipitating (24) Krug, J. P.; Popelier, P. L. A.; Bader, R. F. W. J. Phys. Chem. 1992, 96, 7604.
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Figure 1. PXRD patterns of the Ni/Al-HCOO LDH: (a) as prepared, (b) thermally dehydrated at 100 °C, and (c) (b) on rehydration at RT and 100% RH.
metal hydroxides and formic acid serves as a source of formate ions. On the basis of this expectation, the synthesis of a number of formate-containing LDHs was attempted. The only successful synthesis was that of the Ni-Al system. Reactions in other systems did not yield any solid products. The pH (observed, 4.7; calculated by buffer-maker freeware (www.chembuddy.com), 4.63) of the reaction medium at the end of the hydrothermal treatment corresponds to that of decomposed formamide (a 1:1 mixture of formic acid and ammonium ions) and was not high enough to nucleate other LDHs. The chemical analysis of the product yields a composition of Ni0.57Al0.400.03(OH)2[(CO3)0.02(HCOO)0.3] 3 (H2O)0.43 [(observed wt %) Ni, 33.34; Al, 10.71; HCOO-, 13.5; (assumed wt %) CO32-, 1.2; OH-, 34.0; (unaccounted for wt %) H2O, 7.25; 0, cation vacancy). TG data (Supporting Information SI.1) show a total mass loss of 42% (expected, 37.4%) up to 800 °C. This can be fit to the following reactions: Ni0:57 Al0:4 00:03 ðOHÞ2 ½ðCO3 Þ0:02 ðHCOOÞ0:3 3 ðH2 OÞ0:43 f ½Ni0:57 Al0:4 ðOHÞ2 ½ðCO3 Þ0:02 ðHCOOÞ0:3 f 0:37NiO þ 0:2NiAl2 O4 The PXRD pattern of the Ni/Al-HCOO LDH (Figure 1a) exhibits three basal reflections, 00l (l = 1, 2, 3), below 25° 2θ and two weak “sawtooth” reflections characteristic of a sample with turbostratic disorder.25 A d spacing of 11 A˚ (2θ = 8°) corresponds to a hydrated HCOO- in a bilayer arrangement obtained in the absence of any control over the humidity. This basal spacing agrees with that reported by Iyi and co-workers.23 Additional evidence of HCOO- ions is given by the IR spectra (Figure 2a), which shows absorptions at 1560 and 1347 cm-1 due to the antisymmetric and symmetric stretching vibrations of the carboxylate group. The TGA data (SI.1) show a prominent mass loss (12.5%) below 100 °C. This low-temperature mass loss is due to the loss of exchangeable interlayer water. In contrast, the carbonate-containing LDH shows negligible mass loss below 100 °C. HCOO-LDH obtained after equilibration at 100 °C exhibits a smaller d spacing of 8.7 A˚ (2θ = 10.2°) (Figure 1b). The IR spectrum (Figure 2b) of this phase shows the continued presence of formate ions. When this phase was exposed to ∼100% relative (25) Warren, B. E.; Bodenstein, P. Acta Crystallogr. 1966, 20, 602.
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Figure 2. Infrared spectra of the Ni/Al-HCOO LDH: (a) as prepared and (b) thermally dehydrated at 100 °C.
Figure 3. PXRD patterns of the Ni/Al-HCOO LDH hydrated at (a) 20% RH and (b) 25% RH. Vertical lines mark the positions of Bragg peaks of the dehydrated (-) and hydrated (---) LDHs.
humidity (RH) in a humidity chamber, the 11.2 A˚ d spacing together with all of the features of the pristine LDH was restored (Figure 1c). The exchangeable water content estimated from isothermal drying studies is ∼25 mass %. To examine the dehydration/hydration phenomenon in detail, in situ measurements of the basal spacings were carried out as a function of the atmospheric humidity at ambient temperature. The 11.2 A˚ phase was found to dehydrate at ambient temperature in dry (∼2% RH) flowing N2. On increasing the humidity in a stepwise manner, the 11.2 A˚ phase was restored at 30% RH with a short equilibration time (∼20 min). (See Supporting Information SI.2). Interesting intermediate phases were found on equilibration at lower RH values. (1) At 20% RH, a broad peak is observed at 9.5° 2θ (d = 9.4 A˚), lying between the 001 reflections of the dehydrated and hydrated phases (Figure 3a). (2) The higher harmonic of the 9.4 A˚ reflection is nearly extinguished, but a broad hump is seen in the 20-23° 2θ range. This “peak” lies between the 002 reflection of the dehydrated phase and the 003 reflection of the hydrated phase. (3) The 002 reflection of the hydrated phase does not manifest itself in the intermediate phase. (4) The fwhm of the 9.4 A˚ reflection (1.83° 2θ) is also very large in comparison to the fwmh values of the end members (dehydrated phase, 0.79° 2θ; hydrated phase, 0.49° 2θ). Langmuir 2010, 26(19), 15586–15591
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Figure 4. Variation of the basal spacing of the Ni/Al-HCOO LDH in one complete hydration/dehydration cycle.
All of these observations are in line with Mering’s principles26 and point to a disordered interstratification of the end members arising from the insertion of water molecules in randomly chosen galleries of the dehydrated 8.8 A˚ phase. The large fwhm value of the 9.4 A˚ phase is indicative of this disorder. On equilibration at 25% RH, two peaks are seen at 10.92 A˚ (8.1° 2θ) and 10.0 A˚ (8.8° 2θ) (Figure 3b). Higher harmonics of the 10.92 A˚ reflection are seen at 5.56 A˚ (15.9° 2θ) and 3.73 A˚ (23.8° 2θ). The rationality of these reflections shows the growth of the hydrated phase at 25% RH. The 10 A˚ peak corresponds to the 002 reflection of a phase with a superstructure obtained by the ordered interstratification of the hydrated and dehydrated phase (cs = 11.2 A˚ þ 8.8 A˚ = 20.0 A˚; s, superlattice reflection). However, the 001 superstructure reflection could not be observed, indicating the possible surplus of hydrated interlayers beyond the 50:50 ratio required for a perfectly ordered interstratification. The ratio of the hydrated phase in the ordered interstratification increases with increasing RH, and finally the material is fully hydrated at 30% RH. These observations suggest that on increasing the RH, the hydrated phase coexists with an ordered interstratified phase. The fact that two hydration states coexist at the same RH indicates that the intracrystalline reactivity of different tactoids of the material varies slightly. This is indicative of minor variations in layer charge densities. Dehydration follows a different route (Supporting Information SI.3): (1) There is a large amount of hysteresis, and dehydration does not set in until the humidity is brought down to 15%, at which time the 9.4 A˚ phase is observed. Complete dehydration leading to the 8.8 A˚ phase is seen only at 10% RH. (2) There is no evidence of the formation of the ordered interstratified phase. To conclude, the hydration/dehydration behavior of HCOO-LDH is complete within a narrow range of relative humidity values of 0-30% (Figure 4). Under laboratory conditions (24-27 °C, g60% RH), the 11.2 A˚ hydrated phase is stable and was studied further. The hydration energy of the formate ion has no distinct minima as a function of water content and in fact approaches the energy of free bulk liquid water.27 The formate ion draws a large amount of water into the interlayer region by means of a hydrogen bonding network. This is responsible for the reversible hydration behavior of the LDH. An extreme manifestation of this behavior is first in the swelling and then in the exfoliation of the layers when the LDH is stirred in water. The osmotic swelling results in a phase with a low-angle reflection at 3-4° 2θ (Figure 5a). After 48 h of (26) Mering, J. Acta Crystallogr. 1949, 2, 371. (27) Padmakumar, P.; Kalinichev, A. G.; Kirkpatrick, R. J. J. Phys. Chem. C 2007, 111, 13517.
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Figure 5. PXRD pattern of Ni/Al-HCOO (a) on swelling and (b) after restacking.
stirring, a translucent colloidal dispersion is obtained. The colloidal state of the LDH dispersion is illustrated by light scattering (Supporting Information SI.4). In a series of experiments, up to 130 mg of the LDH could be dispersed in 100 mL of water. These dispersions were stable for up to 3 weeks on standing. The PXRD pattern of the sol does not contain any of the Bragg peaks of the LDH; only a broad feature due to diffuse scattering is observed. On evaporation of the solvent, restacking of the layers was observed as shown by the reappearance of reflections due to the basal planes in the PXRD pattern (Figure 5b). AFM images of the colloid reveal the platelet morphology (Figure 6). The stepwise magnification of typical particles shows in-plane dimensions of 100-200 nm and a thickness of 9-12 nm. Exfoliation is evident from the fact that the thickness of the platelets corresponds to tactoids of 8-11 layers and suggests a high degree of separation and dispersion of the layers, which is reversible. The dry LDH has the typical hexagonal platelet morphology of these materials. The SEM images (Figure 7) show that these platelets have aggregated by surface-to-edge interactions to yield a sand-rose morphology. Crystal growth takes place by the accretion of atoms over several length scales. The accretion pattern is dictated by energy minimization. In layered materials, anisotropic bonding favors in-plane growth rather than out-ofplane growth. Therefore, crystallites acquire a lamellar morphology. Urea hydrolysis reactions yield large platelets with in-plane dimensions exceeding 10 μm.18 Formamide hydrolysis is less facile than urea hydrolysis; therefore, crystal growth is much slower. The crystallite edge comprises high-energy sites, and in the absence of adequate growth, the edges bond to neighboring crystallites to yield sand-rose aggregates. The preference of LDHs for different anions decreases in the order of CO32- > SO42- > Cl-, Br- > NO3- > I-.28 To determine the position of the formate ion in this series, exchange reactions were performed with HCOO-LDH. The intercalated formate ion is readily exchangeable. HCOO-LDH was suspended in a NaNO3 solution to affect the formate exchange for NO3ions. The exchange took place readily as shown by the decrease in the basal spacing from 11.2 to 8.8 A˚ in the product LDH (Figure 8a). The displacement of formate by nitrate ions was verified by the chemical analysis of the exchanged product. This reduced basal spacing corresponds to NO3- ions intercalated with their plane perpendicular to the plane of the metal hydroxide layers. (28) Miyata, S. Clays Clay Miner. 1983, 31, 305.
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Figure 6. (a) AFM topographical image of representative platelets of the Ni/Al-HCOO LDH. (b) Height profiles corresponding to the platelets in (a).
Figure 7. SEM images of the pristine Ni/Al-HCOO LDH exhibiting the sand-rose morphology.
exchange of formate ions for I- taken in solution as shown by PXRD (Figure 8b) and IR (Supporting Information SI.5b). Iuptake from solution was below the limits of detection of chemical analysis. This shows that the preference for these anions decreases in the order NO3- > HCOO- > I-. This observation is in keeping with those of Meyn and co-workers,14 who could not insert the formate ion into a variety of LDHs by anion exchange reactions carried out over several days even at elevated temperatures (65 °C).
4. Discussion
Figure 8. PXRD pattern of (a) Ni/Al-NO3 obtained by anion exchange by soaking the Ni/Al-HCOO LDH in a NaNO3 solution and (b) the Ni/Al-HCOO LDH soaked in a NaI solution.
IR spectra (Supporting Information SI.5a) show that the carboxylate-related absorptions are greatly reduced and the strong absorption due to the intercalated nitrate appears. Chemical analysis of the LDH after exchange yields the formula [Ni0.57Al0.400.03(OH)2](NO3)0.24(HCOO)0.04(CO3)0.03 3 (H2O)0.18 [(observed wt %) NO3-, 14.9; HCOO-, 1.8), showing that 87% of the formate has been exchanged. However, there was no 15590 DOI: 10.1021/la103108f
The most common anion found in laboratory-prepared as well as mineral LDHs is the carbonate ion. The structure of the CO32(molecular symmetry D3h) ion is compatible with the local symmetry of the trigonal prismatic interlayer site. Consequently, the LDH has a very high affinity for CO32- ions. The layerinterlayer bonding is reinforced by H bonding between the oxygen of the intercalated CO32- and water with OH- groups of the metal hydroxide layers. CO32- ions thereby mediate an ordered stacking of the metal hydroxide sheets. From a thermodynamic point of view, the carbonate-LDHs are also stable22 and the carbonates are not exchangeable. We surmised that the formate ion, also having planar geometry (molecular symmetry C2v) and a lower charge would substitute for CO32- and yield an ordered phase and yet be exchangeable. Many of our attempts to obtain HCOO-LDH by conventional synthesis techniques failed. This is because HCOO-LDH is unstable on account of the very high hydration energy of the formate ion.19 Langmuir 2010, 26(19), 15586–15591
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This also explains the inability to prepare HCOO-LDHs by exchange reactions using LDH precursors containing other anions. The interlayer region restricts the growth of the hydration sphere, thereby making formate intercalation enthalpically expensive. Furthermore, the lower charge diminishes the Coulombic forces compared to the carbonate ion. The number of oxygen atoms is one fewer than in carbonate and nitrate ions, limiting the capacity to hydrogen bond with an ordered stacking of the metal hydroxide layers. Therefore, formate intercalation is enthalpically unfavorable in every way. Exchange reactions are generally enthalpydriven. Therefore, we used the homogeneous precipitation from solution method by urea hydrolysis.29 However, the hydrolysis of urea generates CO32- ions, and the products are CO32- LDHs. Formamide, HCONH2, undergoes hydrolysis by the generation of HCOO- ions, but the extent of hydrolysis is insignificant at ambient temperature. In the present study, formamide hydrolysis was carried out at 150 °C. Under carbonate-starved, formate-rich conditions, HCOO-LDH was precipitated but only in the Ni/Al system. The product, however, was turbostratically disordered because of the excessive hydration of the formate ion in the interlayer. Increased inclusion of water molecules in the interlayer weakens the direct layer-anion interaction and introduces turbostratic disorder due to loss of registry between successive metal hydroxide layers. The large hydration enthalpy of the formate ion is responsible for the humidity-induced reversible hydration/dehydration behavior of formate-LDH. This is reminiscent of similar behavior of the cationic clays. There are some differences: the natural clays exhibit this behavior over the 0-100% RH range because of layer charge inhomogeneities. The narrower RH range observed with LDH would thus indicate a comparatively homogeneous charge density similar to that of synthetic clays.30 Furthermore, the natural clays do not exhibit either a step profile or hysteresis in the hydration/dehydration cycle. This behavior is attributed to the charge inhomogeneity of the aluminosilicate layers, which provide hydration sites with a wide and continuous range of binding energies for water molecules. The laboratory-synthesized clays obtained from melt precursors exhibit a stepwise hydration and a large amount of hysteresis in the dehydration cycle, a behavior characteristic of materials with charge homogeneity.31 This behavior is unique to the melt-derived clays. The hydration behavior of the interlamellar space perhaps is the most sensitive sensor for charge inhomogeneity. Although the double peak observed at 20% RH for the LDH reported here indicates charge-density variations, these are minute as compared to what is observed for natural clays. Given the close relationship between the aluminosilicate cationic clays and the LDH anionic clays, we ask the following question: are the LDH layers charge-homogeneous? Although this question has not been answered directly, a closely related question has engaged the attention of many authors: are the LDHs cations ordered? The results of different investigations are ambiguous. X-ray diffraction studies have generally answered in the negative, not because the LDHs are cation-disordered but more so because of the limitation of the technique: (1) XRD is insensitive to the short-range structure and (2) the superlattice 100 reflection corresponding to the cationordered LDH structure has a very low intensity. EXAFS32 and (29) Costantino, U.; Marmottini, F.; Nocchetti, M.; Vivani, R. Eur. J. Inorg. Chem. 1998, 10, 1439. (30) Malikova, N.; Cadene, A.; Dubois, E.; Marry, V.; Durand-Vidal, S.; Turq, P.; Breu, J.; Longeville, S.; Zanotti, J. M. J. Phys. Chem. C 2007, 111, 17603. (31) Breu, J.; Wolfgang, S.; Stoll, A. J.; Lange, K. G.; Probst, T. U. Chem. Mater. 2001, 13, 4213. (32) Vucelic, M.; Jones, W.; Moggridge, G. D. Clays Clay Miner. 1997, 6, 803.
Langmuir 2010, 26(19), 15586–15591
Article
solid-state NMR studies33 have predicted cation order (and thereby charge inhomogeneity), but this is in some sense not surprising. Both of these techniques are sensitive to the shortrange structure, and lattice energy considerations preclude trivalent cations from occupying neighboring cation sites.34 On a smaller length scale, divalent cation sites are clearly distinguishable from trivalent cation sites. Trivalent cations offer sites of greater binding energy to the interlayer atoms. This distinction is obliterated on a larger length scale. In the context of LDHs, the more important question is the following: do the intercalated water molecules bond strongly with each other, or do they bond to the metal hydroxide layers? If the former were true, then the water molecules would be insensitive to the positive charge distribution of the layer and behave cooperatively. This cooperativity would typically result in a barrier to dehydration and hysteresis in the decreasing-RH part of the cycle. This is indeed observed in the hydration/dehydration cycle of the HCOO-LDH (Figure 4), especially with reference to the 9.4 A˚ phase. The current experimental evidence suggests that the 11.2 A˚ hydrated phase promotes cooperative behavior, reminiscent of a material with charge homogeneity. In contrast, the 8.8 A˚ phase behaves in a manner reminiscent of charge inhomogeneity by generating the randomly interstratified 9.4 A˚ phase. Within this picture, we are unable to explain the observation of the 10 A˚ ordered interstratified phase. Evidently, the activity of water, in the decreasing versus increasing RH parts of the cycle, also has a role to play. The large energy of hydration of the formate ion can in principle be exploited for the exfoliation and eventual delamination of the HCOO-LDH. Complete delamination should be possible if the sand-rose morphology of the pristine LDH is suppressed. The observation of exfoliation is itself significant because it is carried out entirely in the aqueous medium. Most other delamination studies utilize the interplay of hydrophilic-hydrophobic forces and involve the use of organic solvents. This work shows the possibility of employing purely hydrophilic interactions for exfoliation and possible delamination of the LDH materials.
5. Conclusions The intercalation of anions with a high enthalpy of hydration presents a new strategy for realizing the aqueous exfoliation of layered double hydroxides. This is a green alternative to the traditional method of exfoliation with the use of organic solvents. However, the incorporation of anions with a high hydration enthalpy in the interlayer of the LDH is a major challenge because the intercalation of such anions involves an enthalpy loss that has to be compensated for by other means. Existing models used for the evaluation of the stability of LDHs do not account for the loss of the hydration enthalpy of the anion. Acknowledgment. G.V.M. and P.V.K. thank the Department of Science and Technology (DST), Government of India for financial support. P.V.K. is a recipient of the Ramanna Fellowship of the DST. D.A.K., W.M., and J.B. are financially supported by the Deutsche Forschungsgemeinschaft (SFB 840). We thank Prof. A. Fery for making AFM equipment available. Supporting Information Available: TGA data of HCOOLDH. In situ PXRD patterns of HCOO-LDH in a hydrationdehydration cycle. Tyndall effect in an exfoliated HCOOLDH suspension. IR spectra of anion-exchanged LDH. This material is available free of charge via the Internet at http:// pubs.acs.org. (33) Sideris, P. J.; Nielson, U. G.; Gan, Z.; Grey, C. P. Science 2008, 321, 113. (34) Brindley, G. W.; Kikkawa, S. Am. Mineral. 1979, 64, 836.
DOI: 10.1021/la103108f
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