Reinvestigation of Dehydration and Dehydroxylation of Hydrotalcite

Jun 1, 2010 - ... i.e. OH−(M(II)3) and OH−(M(II)2M(III)). The thermal stability of these two hydroxyl groups is principally determined by the natu...
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J. Phys. Chem. C 2010, 114, 10768–10774

Reinvestigation of Dehydration and Dehydroxylation of Hydrotalcite-like Compounds through Combined TG-DTA-MS Analyses Jia Zhang, Yun Feng Xu, Guangren Qian,,* Zhi Ping Xu,*,* Chun Chen, and Qiang Liu School of EnVironmental and Chemical Engineering, Shanghai UniVersity, Shanghai 200072, PR China, and Australian Research Council Centre of Excellence for Functional Nanomaterials, Australian Institute for Bioengineering and Nanotechnology, The UniVersity of Queensland, Brisbane, QLD 4072, Australia ReceiVed: April 7, 2010; ReVised Manuscript ReceiVed: May 12, 2010

We reinvestigated dehydration and dehydroxylation of hydrotalcite and four other hydrotalcite-like compounds (HTlcs) with the combined thermogravimetry/differential thermal analysis/mass spectrometry (TG-DTA-MS) technique. The observations indicate that HTlcs undergo dehydration first and then dehydroxylation, followed by or overlapping with decomposition of interlayer anions. The detail deconvolution of MS curves of m/z ) 18 reveals that both dehydration and dehydroxylation occurs in two consecutive steps. Dehydration of crystalline water molecules starts on the surface and edge in the first step, and then continues to the interlayer water. Both steps are influenced by the nature of the interlayer anions and layer cations. For the first time, we have assigned two dehydroxylation steps to two types of hydroxyl groups in the HT lattice, i.e. OH-(M(II)3) and OH-(M(II)2M(III)). The thermal stability of these two hydroxyl groups is principally determined by the nature of the cations. According to this research, the thermal stability of hydroxyl groups in HTlcs is OH(Ca3) (∼480 °C) > OH-(Mg2Al) (∼410 °C) > OH-(Mg2Fe) (∼350 °C) ≈ OH-(Mg3) (300-370 °C) ≈ OH-(Ca2Al) (∼330 °C) > OH-(Ca2Fe) (∼290 °C) at a ramping rate of 2 °C/min. 1. Introduction Hydrotalcite-like compounds (HTlcs), also known as layered double hydroxides (LDHs), have received extensive research in recent decades, as they find many potential applications, such as catalysis, anion adsorbents, gene delivery vectors, and polymer stabilizers, due to their versatile composition and tailorable physicochemical properties.1–4 Most HTlcs can be chemically expressed as a general formula of [M1-x2+Mx3+(OH)2](An-)x/n · yH2O,5,6 where M2+ and M3+ are any divalent and trivalent cations, or their combination, with x ) 0.2 to 0.4;7 An- is an exchangeable interlayer anion (organic and inorganic); and water molecules are either crystalline in the lattice or physicosorbed on the surface/edge. Different combination of M2+, M3+, and An- results in a large family of mixed hydroxides, the physicochemical properties of which or derived mixed oxides can thus be finely tuned for various catalytic reactions.8,9 To make mixed oxide catalysts with tailored activity that lies in the control at the atomic level rather than at the particle level, scientists often face a challenge in that how HTlcs are transferred into the mixed oxides and in what controllable way is very much a concern, which is still not completely understood. The transformation from HTlcs to mixed oxides is normally carried out through thermal decomposition, involving a number of competing reactions such as dehydration, decomposition of interlayer anions, and dehydroxylation of layer hydroxide groups.10 Many thermal decomposition investigations demonstrate that dehydration occurs first at a lower temperature, while deanation and dehydroxylation follow at a higher temperature in competition or sequence. For example, Yang et al. found * Corresponding author. (G.Q.) Tel: +86-21-56338094; fax: +86-2156333052; e-mail: [email protected]. (Z.P.X.) Tel: 61-7-33463809; fax: 61-7-33463973; e-mail: [email protected]. † Shanghai University. ‡ Australian Institute for Bioengineering and Nanotechnology.

that dehydroxylation is complete before decarbonation of MgAl-CO3-HT11 and Kameda et al. noted similar decomposition behavior for MgAl-Cl-HT,12 while Xu and Zeng observed that decomposition of NO3- in MgAl-NO3-HTlcs simultaneously takes place with dehydroxylation in some cases.13 In particular, dehydroxylation collapses the layer structure and leads to the formation of mixed oxides, which seems to undergo two steps.11,14 However, the assignment is unclear or incorrect in terms of the HT structure, as clearly revealed recently.15 In addition, it is commonly known that the thermal stability of HTlcs is strongly dependent on the nature of the layer cations and interlayer anions, while the relationship is not well understood. Therefore, in this research we re-examined thermal decomposition pathways of hydrotalcite (Mg3Al-CO3-HT) and four other HTlcs (Table 1) through a combined thermogravimetry/ differential thermal analysis/mass spectrometry (TG-DTA-MS) method. In particular, we deconvoluted their MS curves of evolved water vapor and analyzed the details of water vapor evolution during the thermal decomposition, together with TGDTA profiles. So the objectives of this research were to (i) assign two dehydroxylation steps in terms of the HT structure and (ii) understand the effects of the interlayer anions and layer cations on the dehydration and dehydroxylation of HTlcs. 2. Experimental Section Material Preparation. Hydrotalcite (Mg3Al-CO3-HT) was bought from Shin Woun Chemical Co., Ltd. (Jeungwang-Dong Shiheung-City, South Korea). In some cases, the commercial HT was particularly calcined for 4 h at 190, 245, or 355 °C, and then stored in a desiccator for further investigation. Friedel’s salts (Ca2Al-Cl-HT and Ca2Fe-Cl-HT), Mg3Al-Cl-HT, and iowaite (Mg3Fe-Cl-HT) were prepared by the coprecipitation method.16 For example, Ca2Al-Cl-LDH

10.1021/jp103115q  2010 American Chemical Society Published on Web 06/01/2010

Dehydration/Dehydroxylation of HTlcs through TG-DTA-MS

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TABLE 1: Physical Parameters and Peak Temperatures of Deconvoluted MS Curves of Water Release from As-Prepared HTlcsa sample

a (nm)

c (nm)

t (nm)

De-H2O1 (°C)

De-H2O2 (°C)

De-OH1 (°C)

De-OH2 (°C)

Mg3Al-CO3-HT Mg3Al-Cl-HT Mg3Fe-Cl-HT Ca2Al-Cl-HT Ca2Fe-Cl-HT

0.305 0.305 0.310 0.573 0.585

2.275 2.360 2.394 1.555 1.552

18.7 11.3 11.6 30.8 24.1

194 159 133 156 143

259 218 193 181,244 192,267

338 370 298

414 424 350 331 307

480

a

Note: De-H2O1 and De-H2O2 mean the first and second dehydration peaks; De-OH1 and De-OH2 the dehydroxylation peaks of OH-(M(II)3) and OH-(M(II)2M(III)), respectively.

was synthesized by mixing a salt solution containing 50 mmol of CaCl2 and 25 mmol of AlCl3 in 50 mL of water with a basic solution containing 150 mmol of NaOH in 100 mL of water under vigorous magnetic stirring in the nitrogen stream and then aging for 24 h at room temperature. The resultant slurry was collected and then washed by distilled water through filtration, and dried at 105 °C for 24 h. The dried sample was ground and stored in a desiccator. Ca2Fe-Cl-LDH, Mg2Fe-Cl-LDH, and Mg3Al-Cl-LDH were prepared similarly. Material Characterization. The XRD patterns of all HTlcs were recorded in an XRD DLMAX-2550 (Rigaku Co.) using Cu KR radiation (λ ) 0.15418 nm) from 2θ ) 5 to 80° at a scanning rate of 3° per minute. The dried sample was finely ground with KBr and dried under an infrared light (JC18-1) before being pressed (YP-2) into a disk using 18 MPa of pressure for 1 min. The infrared spectrum of the KBr pellet was then recorded by accumulating 32 scans at 4 cm-1 resolution between 400 and 4000 cm-1 in a Nicolet 380 Fourier transform infrared spectrometer (Thermo Scientific). The TG and DTA profile of each sample and the MS profile of the evolved gaseous water (m/z ) 18) and carbon dioxide (m/z ) 44) were collected by means of NETZSCH Simultaneous TG-DTA/DSC (Apparatus STA 449C/4/G Jupiter-QMS 403C Aeolos) in N2 atmosphere at a ramping rate of 2 °C/min from 30 to 700 °C. For each measurement, about 50-60 mg of dried sample was used. Deconvolution of MS. The mathematical manipulation of the MS profile, such as spectrum smoothing and baseline sunstracting, were performed using the software package Proteus. Then MS profile in 30-700 °C was deconvoluted using the software package Peakfit by selecting and adjusting the peak position and width. Peak fitting was done using a combined Gauss and Lorentz function with minimum deconvoluted peaks in the fitting process. The fitting was undertaken until reproducible results were obtained with a squared correlation of R2 > 0.98.

Figure 1. XRD patterns of examined LDHs: (A) Mg3Al-CO3-LDH, (B) Mg3Al-Cl-LDH, (C) Mg3Fe-Cl-LDH, (D) Ca2Al-Cl-LDH, (E) Ca2Fe-Cl-LDH.

3. Results and Discussion Physicochemical Features of HTlcs. Curve A of Figure 1A represents the typical X-ray diffraction (XRD) pattern of hydrotalcite (Mg3Al-CO3-HT), characteristic of peaks (003), (006), (012), (015), (018), (110) and (113), with 3R1 polytype and cell parameters a ) 0.305 nm and c ) 2.275 nm (Table 1). A similar pattern was recorded for Mg3Al-Cl-HT (curve B of Figure 1) and Mg3Fe-Cl-HT (curve C of Figure 1), with the interlayer a bit expanded (Table 1) On the other hand, curves D and E of Figure 1 (Ca2Al-Cl-HT and Ca2Fe-Cl-HT, respectively) show the typical pattern of Friedel’s salts. For example, cell parameters a ) 0.573 nm and c ) 1.555 nm, identical to those reported elsewhere for Ca2Al-Cl-HT17 and on PDF 78-1219. The nominal particle thickness (Table 1) was estimated by the instrument software from the measured widths

Figure 2. TG-DTA-MS curves of hydrotalcite (Mg3Al-CO3-LDH).

of peaks (003) and (006) (or (002) and (004)) using the Debye-Scherrer equation without any corrections: D ) 0.9λ/B cos θ,18 where B is the full-width-half-maximum in radian, λ the X-ray wavelength (0.15418 nm for Cu KR radiation) and θ the diffraction angle. Thermal Dehydration and Dehydroxylation of Hydrotalcite (Mg3Al-CO3-HT). Figure 2 shows the TG-DTA-MS profiles of hydrotalcite. The smooth TG profile indicates two major weight losses at around 250 and 390 °C, and the MS

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profile of released water (m/z ) 18) displays two similar events at 256 and 415 °C. However, the DTA contour shows three major endothermic peaks at 244, 401, and 550-600 °C. It is well understood that the two events at temperature below 500 °C are attributed to dehydration of various water molecules (physically adsorbed, crystalline, and H-bound water in the interlayer) and dehydroxylation of OH groups as well as decomposition of interlayer anions in the HT lattice.19–21 The endothermic event above 500 °C is due to crystallization of rock salt (MgO) as well as spinel (MgAl2O4), without significant weight loss. The MS profile of CO2 release (Figure 2) indicates that decarbonation is mainly overlapping with the dehydroxylation process. Detailed investigation of the MS profile of evolved water vapor suggests a minor dehydroxylation between 256 and 415 °C. This takes place at 338 °C, as clearly reflected by the deconvolution fitting (Figure 2). In a more detail examination, a minor endothermic event at 186 °C in the DTA curve and a minor MS peak at 194 °C (Figure 2) can be observed, probably due to dehydration of crystalline water near the surface and edge. Therefore, water is released in four stages during thermal decomposition, demonstrating four types of water origins in the dried hydrotalcite sample. It is interesting to note that the peak temperature in the DTA profile is normally 10-20 °C lower than that in the MS profile (also in Figures 3 and 5). This is probably because the MS detection of water molecules is much more delayed due to the slow release of water during heating and the transport to the MS detection chamber. These water origins have different thermal stability as they can be consecutively removed at temperatures of 194, 256, 338, and 415 °C. As shown in Figure 3B, the small peak at 186 °C in the DTA profile and 194 °C in the MS profile is gone after heating at 190 °C for 4 h. Also note that the peak area of the second event at around 240 °C is very much reduced (Table 2), revealing that crystalline interlayer water molecules start dehydration even at 190 °C. Figure 3C indicates that calcination at 245 °C for 4 h dehydrates all crystalline interlayer water molecules,13 leaving a major peak at 420 °C and a shoulder peak at 350 °C in the MS profile, which are due to hydroxylation.14 Calcination at 355 °C further removes the shoulder peak, leaving the only event at 430 °C. All these observations are in good agreement with the public reports elsewhere for HT, and it is no doubt that crystalline water molecules are dehydrated in the first two steps.11,13,22 However, dehydroxylation in two steps seems not to be well understood. Some researchers assigned the first dehydroxylation to OH groups in Mg-(OH)-Al and the second one to those in Mg-(OH)-Mg.11,14 In our opinion, this assignment is not correct, at least not exact, and thus not meaningful. As well discussed by Sideris et al.,15 there are two types of hydroxyl groups in Mg3Al-CO3-HT. As outlined in Figure 4, each OH group is bound with three adjacent metal cations. In an ideal distribution of metal cations in the brucite-like layer, there are 25% OH-(Mg3) and 75% OH-(AlMg2). Sideris et al.15 have found from 1H NMR that there are about 20% OH-(Mg3) and 80% OH-(AlMg2) in their prepared HT. Although there are some OH-(MgAl2) and OH-(Al3) existing in a random case, their portion is too small to be recognized.15 In this connection, we have thus assigned the event at around 340 °C to hydroxylation of OH-(Mg3) and that at ∼420 °C to OH-(AlMg2). This assignment predicts a water loss ratio of 1:3 (or 1:4 in reality15), which is in quantitative agreement with the observed value (28.3:71.7 and 18.0:82.0 in Table 2). This assignment seemingly contradicts with the relative thermal stability of Mg(OH)2 and Al(OH)3. It has been reported that

Zhang et al.

Figure 3. Band component analysis of the MS curves of Mg3Al-CO3-LDH: (a) origin form, (b) calcited at 190 °C, (c) calcited at 245 °C, (d) calcited at 355 °C.

dehydroxylation of Al(OH)3 occurs at 290 °C, while that of Mg(OH)2 occurs at 350 °C.11 However, the higher thermal stability of OH-(AlMg2) than OH-Mg3) is still understandable. First of all, the thermal stability of OH-(Mg3) should be similar to that of OH in Mg(OH)2, which is supported by the similar dehydroxylation temperature (340-350 °C). Second, the coordination of Al in Al(OH)3 is much different from that in Mg3Al-CO3-HT. In Al(OH)3 Al3+ is tetrahedronally coordinated with 4 OH groups while in hydrotalcite with 6 OH groups. Thus the thermal stability of OH in Al(OH)3 is not comparable with that in OH-(Mg2Al). Third, interactions between Al3+ and OH- in OH-(Mg2Al) are much stronger than those between Mg2+ and OH- in OH-(Mg3) because Al3+ has one more positive charge than Mg2+. This probably stabilizes OH-(Mg2Al) and delays its dehydroxylation. In addition, we have also found that dehydroxylation analyses of Mg3Al-Cl-LDH discussed below qualitatively and quantitatively support this assignment. It should be noted that the fitting at around 500 °C seems not so good, which is also seen in other cases (except for Ca2Fe-Cl-HT) in Figure 5. Therefore, here we fitted an extra minor peak at 450-500 °C for all these samples. This weak peak (1-4% area of dehydroxylation peaks) could be assigned to the tailing of interior hydroxide group dehydroxylation.

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TABLE 2: Peak Temperatures and Relative Areas (in Parentheses) of Deconvoluted MS Curves of Hydrotalcites (Mg3Al-CO3-LDH) Treated under a Certain Conditiona heated at 105 190 245 355 a

°C °C °C °C

De-H2O1 (°C)

De-H2O2 (°C)

De-OH1 (°C)

De-OH2 (°C)

196 (1.1)

259 (42.1) 253 (17.6)

336 (28.3) 359 (18.0) 349 (11.1)

415 423 421 432

(71.7) (82.0) (88.9) (100)

The values in parentheses were the relative peak area by setting the total area of De-OH1 and De-OH2 to be 100.

Figure 4. Schematic representation of OH types in the lattice with ideal cation distribution of cations in the brucite-like layer of hydrotaclite (Mg3Al(OH)8(CO3)0.5 · 2H2O). Each octahedron contains one cation in the center. The shaded triangle means an Al3+ in the center; otherwise Mg2+ is in the center.

Thermal Decomposition Behaviors of Other HTlcs. Figure 5 presents MS, TG, and DTA curves of other four HTlcs. TG and DTA profiles are very similar to reported elsewhere. For example, Mg3Al-Cl-LDH dehydrated the surface and interlayer water below 250-300 °C, and then dehydroxylated (and de-Cl) mainly in 350-450 °C.23 As expected, the MS curves can be well fitted with four water release peaks. For example, the MS curve of Mg3Al-Cl-HT is similarly deconvoluted into four peaks: 159, 218, 370, and 429 °C (Figure 5A). The latter two peaks with an area ratio of 1:3.29 (De-OH1:De-OH2 ) 3.3:76.7 in Table 3) are very similar to those in Mg3Al-CO3-HT in the peak temperature and the relative peak area, further supporting our hypothesis that two dehydroxylation steps are attributed to OH-(Mg3) and OH-(AlMg2), respectively. This similarity also indicates that the anion type may not obviously affect dehydroxylation of OH-(AlMg2) and OH-(Mg3). However, the former two peaks due to dehydration are shifted down by 20-30 °C when compared with Mg3Al-CO3-HT, reflecting the relatively weaker interactions of Cl- with crystalline water molecules. In sharp contrast, it is very different in the case of Mg3Fe-Cl-HT. As shown in Figure 4B, two steps of dehydration (133 and 193 °C) and dehydroxylation (298 and 350 °C) are overlapping to some extent, and are not as well separated as in Mg3Al-CO3-HT and Mg3Al-Cl-HT. Moreover, dehydration and dehydroxylation occur at much lower temperatures in comparison with those in Mg3Al-CO3-HT. Therefore, incorporation of Fe3+ into the hydroxide layer severely weakens the H-bond interactions of lattice OH groups with crystalline water molecules, and meanwhile destabilizes the thermal stability of OH groups associated with FeMg2. It seems a bit difficult to assign the two dehydroxylation steps. Tentatively, we assigned the first peak at 298 °C in Mg3Fe-Cl-HT to OH-(Mg3) dehydroxylation, which is a bit lower than that in Mg3Al-CO3-HT (338 °C) and Mg3Al-Cl-HT (370 °C) (Table 1), which could be attributed to the destabilization effect of adjacent Fe3+. The second peak at 350 °C is assigned to OH-(FeMg2) dehydroxylation due to the destabilization of Fe(III) in the structure, about 70-80 °C lower than that for OH-(AlMg2) dehydroxylation. This seems to also be supported by the peak ratio of 26.9:73.1 (Table 3).

For the two Friedel’s salts (Ca2Al-Cl-HT and Ca2Fe-Cl-HT), the major dehydration takes place at 185-190 °C. The difference is that there is a peak at 245 °C in Ca2Al-Cl-HT, while a peak at 147 °C in Ca2Fe-Cl-HT, probably due to the stronger H-bond interaction with OH-(Ca2Al) than that with OH-(Ca2Fe). Also note that the temperature of major dehydroxylation in Ca2Al-Cl-HT is 330 °C, 70 °C lower than that in Mg3Al-Cl-HT. This event is assigned to dehydroxylation of OH-(Ca2Al) because there is 97% OH-(Ca2Al) in Ca2Al-Cl-HT.15 This observation reflects that replacing Mg2+ with Ca2+ leads to a lower thermal stability of OH-(M(II)2Al). In contrast, the temperature of major dehydroxylation in Ca2Fe-Cl-HT is further lowered to 307 °C, again reflecting the destabilization of Fe3+ for the thermal stability of OH-(Ca2M(III)), similar to the comparison of Mg3Al-Cl-HT with Mg3Fe-Cl-HT, as discussed previously. The peak at 267 °C in Figure 5D is tentatively assigned to dehydration of H-bond water in the interlayer, as addressed in the following section. A more interesting observation is that there is a weak peak at 480 °C in the Ca2Fe-Cl-HT MS curve (1.9 wt % in total dehydroxylation), which can be assigned to dehydroxylation of OH-(Ca3). This assignment is supported by two facts: (i) Ca(OH)2 decomposes in 450-500 °C,24 and (ii) OH-(Ca3) exists in a small portion in as-synthesized Ca2Fe-Cl-HT. As reported by Sideris et al.,15 there are 3% OH-(Mg3) and 97% OH-(Mg2Al) in the synthesized Mg2Al-HT. Thus the released water (De-OH1 Table 3) is quantitatively agreed with the prediction. It is worth mentioning that the total dehydration of crystalline water in Mg3Al-CO3-HT, Mg3Al-Cl-HT, and Mg3Fe-Cl-HT (De-H2O1+De-H2O2 in Table 3) is nearly 50% of the water released from dehydroxylations (De-OH1+DeOH2), in good agreement with the prediction (Table 3). This is also nearly true for Ca2Al-Cl-HT and Ca2Fe-Cl-HT (Table 3). Effects of Cations and Anions. As shown in Figure 6A, water molecules interact directly with interlayer anions and lattice OH groups and indirectly with metal cations mainly via electrostatic forces, so dehydration is mainly determined by these interactions, i.e., the nature of the interlayer anions and layer cations. We observed that the temperatures for dehydration of Mg3Al-CO3-HT are about 30-40 °C higher than those of Mg3Al-Cl-HT (Table 1), just because interactions between H2O and CO32- (H-bond and dipole-dipole interaction) are much stronger than those between H2O and Cl- (mainly dipole-dipole interaction). The stronger effect of carbonate on dehydration has been also reflected in samples Ca2Al-Cl-HT and Ca2Fe-Cl-HT. As shown in Figure 7, the IR peaks are characteristic of these HTlcs, such as the broad band at 3400-3500 cm-1 (νOH), a weak peak at 1620 cm-1 (H2O bending vibration), and bands below 1000 cm-1 (due to M-O vibrations and M-O-H bending).24,25 In particular, carbonate is clearly indicated by the strong peak at 1366 cm-1 [due to ν(CO32-)] in sample Mg2Al-CO3-HT.24,25 Similarly, there is a intermediate band at 1440 cm-1 in Ca2Fe-Cl-HT (Figure

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Figure 5. MS profile of Mg3Al-Cl-HT (A), Mg3Fe-Cl-HT (B), Ca2Al-Cl-HT (C), and Ca2Fe-Cl-HT (D).

TABLE 3: Relative Peak Area of Deconvoluted MS Curves of Water Release from Examined HTlcsa sample

De-H2O1

De-H2O2

De-OH1

De-OH2

De-OH1 predicted

Mg3Al-CO3-HT Mg3Al-Cl-HT Mg3Fe-Cl-HT Ca2Al-Cl-HT Ca2Fe-Cl-HT

1.1 8.9 5.7 12.4 7.7

42.1 55.0 46.3 49.8 91.9

28.3 23.3 26.9

71.7 76.7 73.1 100 98.1

25 25 25 3 3

1.9

a

De-H2O1 and De-H2O2 mean the first and second dehydration peaks; De-OH1 and De-OH2 the dehydroxylation peaks of OH-(M(II)3) and OH-(M(II)2M(III)), respectively. The addition of De-OH1 and De-OH2 areas was set at 100. So in theory, the total area of crystalline water dehydration (De-H2O1+De-H2O2) was 50 for Mg3Al(OH)8(CO3)0.5 · 2H2O, Mg3Al(OH)8Cl · 2H2O, and Mg3Fe(OH)8Cl · 2H2O, and 66.7 for Ca2Al(OH)6Cl · 2H2O and Ca2Fe(OH)6Cl · 2H2O. De-OH1 predicted was cited from ref 15 for a practical case. In all cases, the area of the peak at 450-500 °C (1-4% of De-OH1+De-OH2) has not been counted in this calculation.

7E), which clearly indicates that the sample is contaminated by CO32-. The CO32- contamination is likely to be responsible for the MS peak at 267 °C (Figure 5D). In comparison, the weak MS peak at 244 °C in Figure 5C is probably related to the weak IR band at 1470 cm-1 (Figure 7D).25,26 The effect of cations on dehydration also seems obvious. Comparing hydration profiles of Mg3Al-Cl-HT and Mg3Fe-Cl-HT (Figure 5A,B) demonstrates incorporation of Fe3+ eases hydration, lowering the temperature by 25 °C. As Al-OH is more acidic than Fe-OH, the partial positive charge on H in Fe-OH is smaller than that in Al-OH. Therefore the H-bond of H2O with OH-(Mg2Fe) is weaker than that with OH-(Mg2Al), and dehydration in Mg3Fe-Cl-HT is easier than in Mg3Al-Cl-HT. Similarly, Ca-OH is more basic than Mg-OH, so the partial negative charge of OH group in Ca-OH is more than that in Mg-OH, i.e., less proton positive charge in Ca-OH, and thus a weaker H-bond of H2O with

OH-(Ca2Al) than with OH-(Mg2Al). Therefore, replacement of Mg2+ with Ca2+ likely eases hydration to some extent (Table 1). It is worth mentioning that the crystallite size also affects the dehydration. If crystalline water molecules are restricted in the interlayer, their evaporation is prohibited in the space. However, the evaporation of surface and edge water molecules is not sterically hindered and takes place relatively easily. Therefore, larger crystallites have a smaller portion of surface, and thus less water dehydrated from surface and edge. As shown in Tables 1 and 3, the relatively larger thickness of Mg3Al-CO3-HT compared to that of Mg3Al-Cl-HT and Mg3Fe-Cl-HT results in a smaller portion of surface and edge, and thus less water in the first dehydration step. On the other hand, dehydroxylation is less affected by the nature of the interlayer anion, but severely influenced by the nature of the metal cations in the layer. This is because hydroxyl

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J. Phys. Chem. C, Vol. 114, No. 24, 2010 10773 that of OH-(Mg2Al). This seemingly contradicts with the fact that Ca(OH)2 is more thermally stable than Mg(OH)2.11 It is well known that Ca(OH)2 is a strong base, while Mg(OH)2 is a weak base, and Al(OH)3 is a weak base/acid (amphiprotic), so it is our belief that the larger difference between Al(OH)3 and Ca(OH)2 in this property may lead to a less thermally stable OH-(Ca2Al) in comparison with OH-(Mg2Al). In the Ca/Fe combined HT (Ca2Fe-Cl-HT), the thermal stability of OH-(Ca2Fe) is further reduced to 290 °C, again due to the thermal instability of Fe-OH. However, the thermal decomposition of OH-(Ca3) in Ca2Fe-Cl-HT occurs at 480 °C, as in the case of Ca(OH)2 (420-510 °C at ramping rate of 20 °C/ min).24 In this connection, Ca2Al-Cl-HT seemingly does not contain OH-(Ca3).

Figure 6. Schematic representation of interactions among intercalated anions, water, lattice OH groups, and cations before hydration (A) and after hydration (B). Each octahedron contains one cation in the center. The shaded triangle means an M(III) in the center; otherwise M(II) is in the center.

4. Conclusions In conclusion, we observed that HTlcs undergo two-step dehydration and consequently two-step dehydroxylation. Dehydration of crystalline water molecules starts on the edge and surface in the first step, and then continues to the interlayer. Both steps are influenced by the nature of the anions and cations, more severely by the former. We have for the first time ascribed two dehydroxylation steps to two types of hydroxyl groups in the HT lattice, i.e., OH-(M(II)3) and OH-(M(II)2M(III)). The thermal stability of these two hydroxyl groups is determined by the nature of the cations, but rarely by anions. According to the current research, we found that the thermal stability of these hydroxyl groups in HTlcs is OH-(Ca3) (∼480 °C) > OH-(Mg2Al) (∼410 °C) > OH-(Mg2Fe) (∼350 °C) ≈ OH-(Mg3) (300-370 °C) ≈ OH-(Ca2Al) (∼330 °C) > OH-(Ca2Fe) (∼290 °C) at a heating rate of 2 °C/min. Acknowledgment. This project is supported by the National Nature Science Foundation of China, Nos. 20477024, 20677037, and 20877053, and the key subject of Shanghai Municipality (S30109). This work was accomplished with the assistance of the Institute of Solid Waste Recycle and Safety Disposal, Shanghai University, China. The support from the ARC Centre of Excellence for Functional Nanomaterials, funded by the Australia Research Council under its Centre of Excellence Scheme, is also appreciated. References and Notes

Figure 7. FTIR spectra of HTlcs: (A) Mg-Al-CO32- LDH, (B) Mg-Al-Cl LDH, (C) Mg-Fe-Cl LDH, (D) Ca-Al-Cl LDH, (E) Ca-Fe-Cl LDH.

groups are covalently bound with metal cations, while the possible H-bond of OH with interlayer anions would be broken after dehydration (as presented in the dashed arrows in Figure 6B). In fact, we observed that there is no significant difference in terms of peak temperatures for two-stage dehydroxylations of Mg3Al-CO3-HT and Mg3Al-Cl-HT (Figure 3A and 4A, Table 1). As shown in Figure 5A and 5B, when Fe3+ is incorporated into the lattice to replace Al3+, dehydroxylation of OH-(Mg2Fe) occurs at a temperature by about 130 °C lower than that of OH-(Mg2Al). This can be attributed to the lower stability of Fe-OH because thermal decomposition of Fe(OH)3 mainly occurs at 240 °C, while Al(OH)3 decomposes at around 290 °C at a ramping rate of 2 °C/min.11 As also discussed previously, OH-(Mg2Al) is more thermally stable than OH-Al in Al(OH)3, and is thus much more stable than OH-(Mg2Fe). Data in Table 1 show that dehydroxylation of OH-(Ca2Al) in Ca2Al-Cl-HT takes place at a much lower temperature than

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