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Article Cite This: ACS Omega 2019, 4, 3198−3204
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Interaction of Pristine Hydrocalumite-Like Layered Double Hydroxides with Carbon Dioxide Anand N. Narayanappa and P. Vishnu Kamath* Department of Chemistry, Central College, Bangalore University, Bangalore 560 001, India
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S Supporting Information *
ABSTRACT: The layered double hydroxides (LDHs) of Ca2+ and trivalent cations, Al3+ and Fe3+, are single-source precursors to generate supported CaO, which picks up CO2 from the gas phase in the temperature range 350−550 °C. The supports are ternary oxides, mayenite, and Ca2Fe2O5. The uptake capacity of the Fe3+-containing LDH at 1.9 mmol g−1 is two times the capacity of the Al3+containing LDH. The product of CO2 uptake is calcite CaCO3. It is observed that the intercalated chloride ions reduce the thermal penalty by inducing the early decomposition of CaCO3. In the case of the chloride-intercalated LDHs of Ca2+ and Fe3+, the CaCO3 formed is completely decomposed at 900 °C. This is in contrast with the CaCO3 formed from bare CaO, which shows no sign of decomposition at 900 °C under similar conditions. This work shows that the hydrocalumite-like LDHs are candidate materials for CO2 mineralization.
1. INTRODUCTION
for CO2 capture when compared to the sintered LDHs for the following reasons. (i) The LDHs comprise positively charged metal hydroxide layers having the general formula [M(II)1−xM′(III)x(OH)2]x+ (x = 0.2−0.33), wherein, M(II) is commonly Mg2+ and M′(III) is commonly Al3+, Cr3+, and Fe3+. Anions are intercalated in the interlayer region. The LDHs are structurally and functionally the “inverse” of the cationic clays. While the cationic clays exhibit surface acidity, the LDHs are endowed with surface basicity.14 The interaction of pristine LDHs with CO2 has the character of an acid− base reaction whereby the LDHs are expected to have a high affinity for CO2. (ii) Structurally, the LDHs have interlayer galleries which provide space for storage. The LDHs are also known to swell and function as receptacles for increasing numbers of intercalated species by means of a hydrogen-bonded network.15 Despite these advantages, the pristine hydrotalcite-like [Mg−Al] LDH did not gain any mass when exposed to CO2 (pressure 1 atm) over the temperature range 30−400 °C,16 although the [Mg−Al] LDHs sorb carbonate anions from solution in quantities proportional to their layer charge. Another important candidate LDH for CO2 mineralization is the hydrocalumite-like LDH, in which M(II) is the Ca2+ ion and M′(III) is Al3+ or Fe3+.17 Although the conversion of
The accumulation of anthropogenic CO2 in the atmosphere is responsible for the rise in the temperature of the earth’s atmosphere. There are major efforts underway to mineralize CO2 and store it as a benign solid carbonate, an approach known as carbon capture and storage.1,2 An alternative is to use the mineralized form as a feedstock for pure CO2 for other chemical processes.3,4 In the latter approach, the temperature difference between CO2 uptake and its eventual regeneration constitutes the thermal penalty.5,6 For purposes of energy efficiency, materials with low thermal penalty are preferred. A number of materials have been used as candidates for CO2 capture and are reviewed elsewhere.7 Among them, sintered layered double hydroxides (LDHs) comprise a prominent family of materials. LDHs decompose to yield mixed metal oxides (MMOs).8 Metal oxides in general exhibit surface acidity and are not expected to mineralize CO2, as the latter is also acidic in nature. Nevertheless, a modest CO2 uptake capacity of 0.3 mmol g−1 has been reported for the MMOs of Mg2+ and Al3+ obtained by the thermal decomposition of hydrotalcite-like LDHs.9 This capacity is attributed to the adsorption of CO2 within the morphological pores of the MMO. The same MMO is also reported to exhibit a CO2 uptake capacity of 0.53 mmol g−1.10 Modifications in the morphology of the crystallites and amorphization led to only modest gains in the CO2 uptake capacity.11,12 LDHs infused with up to 30% by weight of K2CO3 were found to show CO2 uptake capacity of ∼1 mmol g−1 of LDH after calcination.13 From chemical, structural, and compositional criteria, pristine LDHs offer themselves as superior candidate materials © 2019 American Chemical Society
Received: January 10, 2019 Accepted: February 1, 2019 Published: February 13, 2019 3198
DOI: 10.1021/acsomega.9b00083 ACS Omega 2019, 4, 3198−3204
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Table 1. Composition of the Hydrocalumite-Like LDHs mass loss from TGA [%] loss of H2O 2+
3+
total loss
sample
Ca content [wt. %]a
M content [wt. %]a
anion content [wt. %]a
anion content [mol %]b
N2
CO2
N2
CO2
approximate composition
[Ca−Al−Cl] [Ca−Fe−Cl] [Ca−Al−NO3]
26.2 (28.8) 27.1 (26.1) 27.6 (27.4)
9.9 (9.7) 19.8 (18.8) 9.7 (9.2)
7.5 (7.63) 5.1 (5.1) 13.8 (12.7)
0.6 (1.0) 0.43 (1.0) 0.6 (1.0)
14.7 14.4 13.3
13.4 14.1 11.1
41.1 35.8 47.3
37.7 35.8 30
[Ca2Al(OH)6]Cl0.60(OH)0.40·2.3H2O [Ca2Fe(OH)6]Cl0.43(OH)0.57·2.45H2O [Ca2Al(OH)6](NO3)0.60(OH)0.40·2.2H2O
a Values in parenthesis are calculated based on the approximate composition given in the last column. bValues given in parenthesis are expected to be the nominal composition of the LDHs (x = 0.33).
preference to the LDH during coprecipitation.21 The [Ca−M′] LDHs are generally obtained with intercalated Cl− ions.22 The most common of these having the formula Ca2Al(OH)6Cl· 2H2O is called Friedel’s salt.17 The as-prepared [Ca−M′−Cl] (M′ = Al and Fe) LDHs are highly crystalline, and the powder X-ray diffraction (PXRD) patterns (Figure S1 in the Supporting Information) exhibit sharp 00l (l = 3n) reflections in the low-angle region (5−25° 2θ) typical of layered materials followed by hkl reflections which obey the reflection conditions −h + k + l = 3n (Table S1) in the high-angle regions (2θ > 25°), indicative of rhombohedral symmetry. The refined cell parameters are in agreement with earlier reports on these LDHs.17,22 The [Ca−Al−NO3] LDH crystallizes in the hexagonal symmetry (Figure S2, Table S2). Given the large difference in the ionic radii of Ca2+ (114 pm) and the trivalent cations (Al3+ 67.5 pm and Fe3+ 78.5 pm), the [Ca−M′] LDHs crystallize in cation-ordered structures. The a-parameter (a ≈ 5.74 Å) is a measure of the M′···M′ distance within the metal hydroxide layer and is typical of the composition x = 0.33 ([Ca2+]/[M′] = 2). The approximate formulae of the LDHs (Table 1) were arrived at by a combination of different independent estimations. As an illustration, the methodology is described in detail for [Ca−Al−Cl]. The metal contents determined independently by a combination of AAS and gravimetry correspond to the nominal composition [Ca2M′(OH)2]+. The Cl− content estimated independently by ion chromatography was found to be substoichiometric. The shortfall of negative charge required to neutralize the positive charge on the metal hydroxide layer was made up by the inclusion of hydroxyl ions in the interlayer to yield the LDH composition [Ca2Al(OH)6]Cl0.6(OH)0.4. Such an assumption is reasonable as the precipitation is carried out at high pH (11.5−12). There is evidence to suggest that LDHs comprising univalent anions are rarely single anion materials and generally contain more than one anion.23,24 The intercalated water content was estimated by thermogravimetric analysis (TGA). The TGA of the as-prepared [Ca−Al−Cl] LDH (Figure 1a) carried out in flowing N2 matches with similar data published by other authors.17 The mass loss occurs in three well-defined sigmoidal steps, the inflection temperatures being 146 °C (step I), 343 °C (step II), and 682 °C (step III) (Table 2). Assuming step I to correspond to dehydration yields a water content corresponding to 2.3 water molecules per formula unit. The total mass loss (observed 41.1%; expected 41.4%) corresponds to the complete decomposition reaction
Ca(OH)2/CaO to CaCO3 is widely studied in the context of deep well disposal of CO2 under high pressure and its subsequent mineralization to CaCO3,18 there is relatively less work on the sorption and uptake of CO2 by the [Ca−M′] (M′ = Al3+ and Fe3+) LDHs. The [Ca−M′] LDHs are of interest for many reasons in both the pristine and sintered forms. (1) Given its high natural abundance, Ca2+ is ubiquitous in soil and offers economic advantages over Mg2+. (2) Ca(OH)2 is more basic than Mg(OH)2 and is expected to take up CO2 quantitatively in a facile manner to yield CaCO3 (expected uptake 13.5 mmol g−1). (3) The LDH of Ca and Al crystallizes in a cation-ordered structure (space group R3̅), in which the Ca2+ ion has a seventh coordination with a water molecule in the interlayer.17 In the hydrocalumite structure, the Ca2+ ion is displaced away from the central plane of Al3+ ions. This displacement facilitates coordination with a water molecule in the interlayer. The puckered nature of the metal hydroxide layer helps to trap the intercalated ion. For instance, when the [Ca−Al−Cl] LDH is dehydrated, the Cl− ion moves to the seventh coordination of Ca2+ and is efficiently grafted to the layer. A similar phenomenon in the presence of CO32− could help immobilize CO2 more effectively than otherwise. Thereby, the [Ca−Al] LDH offers a different structure model for the mineralization of CO2 and its confinement in the interlayer gallery. (4) The decomposition temperature of Ca-containing LDHs is in the range of 300−400 °C. The flue gas temperature is very close to the decomposition temperature of Cacontaining LDHs facilitating the reaction between the LDH and CO2. (5) LDHs decompose with a mass loss of 30−40%. The large mass loss results in the formation of a highly porous MMO residue which is not only compositionally metastable19 but also has a large reactive surface area.20 These conditions are ideally suited for the reconstruction of the LDH with the uptake of CO2 and water vapor from the flue gas. This makes Ca-containing LDHs interesting in CO2 sorption studies. As a first step to examining the suitability of the [Ca− M′(III)] LDHs as sorbents for CO2, this work reports the uptake of CO2 by a cohort of [Ca−M′−X] LDHs (M′ = Al3+ and Fe3+; X = Cl− and NO3−).
2. RESULTS A majority of the naturally occurring as well as laboratorysynthesized LDHs have CO32− ions intercalated in the interlayer region. CO32−-intercalated [Ca−M′] LDHs are less known owing to the fact that the single cation CaCO3 forms in
[Ca 2Al(OH)6 ]Cl 0.6(OH)0.4 ·2.3H 2O → 2CaO + 0.5Al 2O3
The mass loss associated with step II corresponds to dehydration and dehydroxylation (observed 16.2%; expected 20.6%). The third mass loss (step III) corresponds to anion 3199
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Table 3. Mass Change in TGA under Flowing CO2 LDH [Ca−Al−Cl] [Ca−Al−NO3] [Ca−Fe−Cl] a
Table 2. Mass Loss Steps in TGA under Flowing N2 step I % (°C)
[Ca−Al−Cl] [Ca−Al−NO3]
14.7 (146) 13.3 (118, 161) 14.4 (161)
[Ca−Fe−Cl]
step II % (°C)
step III % (°C)
step IV % (°C)
16.2 (343) 13.2 (307)
10.2 (682) 12.5 (570)
8.3 (715)
12.1 (278)
9.4 (672)
step IIa % (°C)
step IIIb % (°C)
step IVa% (°C)
step Vc % (°C)
13.4 (132) 11.1 (123) 13.9 (147)
14.8 (380) 12.5 (351) 12.1 (322)
4.6 (497) 4.6 (467.5) 8.5 (497)
14.5 (800−869) 11.1 (511−900) 18.3 (774)
8.7 (810)
Mass loss. bMass gain in heating cycle. cMass gain in cooling cycle.
are comparable to those obtained in the control under flowing N2. However, following these two steps, unlike in the control experiment, a plateau is established from 394 to 460 °C (plateau-I). (ii) At the end of plateau-I, the LDH gains ∼4.5% mass (∼1 mmol g−1) in the temperature range 460−555 °C. Thereafter, the mass is nearly constant with a downward bias up to 750 °C (plateau-II). It is important to note that the mass gain in flowing CO2 begins after the dehydroxylation step. A prolonged stay at the plateau temperature (T = 550 °C) in a separate experiment did not lead to further mass gain (Figure S3). (iii) A mass loss of ∼4.6% is observed at 788 °C. (iv) On staying at 900 °C for up to 30 min, 37.7% total mass loss is observed. This is less than the loss expected (41.8%) for the complete decomposition of the LDH. On cooling, the residue instantaneously regains some of the mass earlier lost at T > 800 °C and reaches a constant value at 700 °C (plateau-III). However, plateau-III is nearly 5.5% mass lower than plateau-II and ∼1% mass lower than plateau-I showing considerable hysteresis. In separate experiments, the LDH was decomposed at the temperatures corresponding to the three plateaus and characterized by PXRD. The phase obtained at plateau-I (460 °C) has a reflection at ∼29.4° 2θ characteristic of 104 reflection of calcite CaCO3. The poor signal-to-noise ratio and background are indicative of a significant volume fraction of Xray amorphous content. The absence of the intense basal reflections shows that the layered structure has collapsed. On cooling back to the ambient, this phase does not regain mass, showing that dehydroxylation is not reversible. However, on further heating in flowing CO2 to the temperature of plateau-II (640 °C), CaCO3 formation is evident from the PXRD pattern. The IR spectrum of the residue at plateau-II shows an enhancement of the absorption at 1409 cm−1 because of calcite formation (see Figure S4 and the accompanying discussion of the IR spectra). At plateau-III, a mixture of CaCO3 and the ternary oxide mineral mayenite, Ca12Al14O33, is observed (Figure 2). The thermodynamically stable ternary oxide is obtained at 900 °C and does not react with CO2. Only a part of the CaO formed in excess over the ternary oxide at 900 °C gains CO2 on cooling to the temperature of plateau-III accounting for the hysteresis in the TGA. Similar observations are made in the case of the [Ca−Fe− Cl] LDH (Figure 3). The total mass loss (observed 35.8%; expected 36.5%) in flowing N2 corresponds to the decomposition reaction
Figure 1. TGA and DTG curves of [Ca−Al−Cl] LDH in (a) flowing N2 and (b) flowing CO2.
LDH
step Ia % (°C)
expulsion (observed 10.2%; expected 6%). The difference in expected and observed values at the end of the steps II and III arises because of overlapping of steps but the expected and observed total mass loss are in very good agreement. The stepwise decomposition is [Ca 2Al(OH)6 ]Cl 0.6(OH)0.4 → 2CaO + 0.6AlOCl + 0.2Al 2O3 → 2CaO + 0.5Al 2O3
A similar approach was used to arrive at the approximate formulae of the other LDHs (Table 1). It is noteworthy that in the cooling cycle, the mass of the residue is unchanged, showing that the decomposition is irreversible. The TGA measurement was repeated in flowing CO2 to examine the interaction of the LDH with CO2 in the gas phase (Figure 1b, Table 3). The following observations were made: (i) During the heating cycle, the LDH loses 28% of its mass up to 394 °C in two sigmoidal steps with inflection at 132 °C (step I) and 380 °C (step II). These two steps
[Ca 2Fe(OH)6 ]Cl 0.43(OH)0.57 ·2.45H 2O → 2CaO + 0.5Fe2O3 3200
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Figure 2. PXRD patterns of the residues obtained from the decomposition of [Ca−Al−Cl] LDH at plateaus (a) P-I, (b) P-II, and (c) P-III. Reflections marked with asterisks are due to calcite CaCO3. Unmarked reflections are due to mayenite.
Figure 4. PXRD patterns of the residues obtained from the decomposition of [Ca−Fe−Cl] LDH at (a) P-I, (b) P-II, and (c) P-III. Reflections marked with asterisk correspond to CaCO3. Unmarked reflections correspond to Ca2Fe2O5.
CaCO3 formed at plateau-II. The total mass loss corresponds to a mixed oxide residue. On cooling, this oxide residue regains mass in two minor steps (850 and 650 °C) to yield plateau-III at which the residue corresponds to a mixture of CaCO3 and Ca2Fe2O5. Evidently, the ternary oxide phase formed at 900 °C does not pick up CO2 on cooling, resulting in a much greater hysteresis between plateaus-I and III than that in the case of the [Ca−Al] LDH. As the Cl− ion is nonvolatile at moderate temperatures (T < 500 °C), similar experiments were performed on the NO3−intercalated LDH to see if the nitrate-LDH shows a higher CO2 uptake than the chloride-LDH. Under flowing N2, the [Ca−Al−NO3] LDH undergoes complete decomposition (Figure S5, observed mass loss 47.3%; expected 44.8%). Under flowing CO2, an uptake of 4.6% (∼1 mmol g−1 original LDH) is observed (Figure 5) at plateau-II (500 °C). In this
Figure 3. TGA and DTG curves of [Ca−Fe−Cl] LDH in (a) flowing N2 and (b) flowing CO2.
Figure 5. TGA and DTG curves of [Ca−Al−NO3] LDH in flowing CO2.
This mass loss is realized over three steps at 161 °C (14.4%), 278 °C (12.1%), and 672 °C (9.4%). In flowing CO2, a behavior similar to that seen in [Ca−Al−Cl] is observed. At the end of the second step of the mass loss (26.1%, 400 °C, plateau-I), the dehydroxylated residue is X-ray amorphous (Figure 4). This residue does not gain mass on cooling, but on heating in CO2, a mass gain corresponding to 8.5% of the original LDH is observed at plateau-II (680 °C) to yield CaCO3. The absence of any CaCO3 at plateau-I is responsible for the higher uptake when compared to the [Ca−Al] LDH. This mass gain corresponds to the uptake of 1.93 mmol g−1 of the original LDH. Above 700 °C, there is a steep mass loss (18.3%) until 900 °C because of the decomposition of the
case, the decomposition is incomplete at 900 °C (observed mass loss 30%) and correspondingly, there is no mass gain in the cooling cycle. The residue obtained at 900 °C was dissolved in dilute HCl and analyzed by ion chromatography. Residual NO3− ion was observed to the extent of 0.07 mol % (4.5 wt %) with respect to LDH and 0.10 mol % (6.3 wt %) with respect to residue, showing that the decomposition of the LDH is incomplete under flowing CO2 and the nitrate is retained in the residue. The nitrate-containing phase could not be identified. The phase obtained isothermally at plateau-I was a mixture of CaCO3 and Ca(OH)2 (Figure S6). The phase 3201
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obtained at plateau-II was CaCO3. At plateau-III, a mix of CaCO3, mayenite, and another ternary oxide Ca5(Al3O7)2 was observed.
3. DISCUSSION The LDH crystallizes in a rhombohedral structure. CaO crystallizes in the cubic rock salt structure. There is a wellestablished topotactic relationship between the precursor and product phases, which can be described as [00l]H/R∥[111]C (l = 1 for H: hexagonal; l = 3 for R: rhombohedral; C: cubic).25 The decomposition of the LDH takes place by the condensation of the metal hydroxide layers brought about by dehydration and dehydroxylation. This occurs just below 400 °C. CaO is therefore expected to form throughout the bulk of the LDH. The uptake of CO2 takes place thereafter at T > 460 °C. The uptake of CO2 is an addition reaction of the type CaO + CO2 → CaCO3. However, all of the Ca are not utilized in the first cycle. If indeed all of the free CaO were to be utilized, plateau-II would correspond to 86% of the mass of the [Ca− Fe] LDH (91% in the case of [Ca−Al] LDH). The observed mass gain at plateau-II corresponds to ∼26% Ca utilization (∼11% in the case of the [Ca−Al] LDH). CaCO3 formation evidently takes place on the exposed surface of the CaO crystallites. Because continued stay (up to 2 h) at various temperatures spanned by plateau-II did not lead to further mass gain (Figure S3), it is our conclusion that CaO utilization is not limited by the kinetics of CO2 uptake, but by the uptake of CO2 at the surface of the CaO crystallites. The CaCO3 formed on the surface prevents a more complete material utilization. The residue obtained at plateau-II was calcined in flowing nitrogen at 900 °C and reheated in flowing CO2. The oxide residue gained nearly 30% mass by aggressive uptake of CO2 in the temperature range of 434−685 °C (Figure S7). We conclude that intermediate calcination at high temperature exposes fresh CaO surface, facilitating increased material utilization. This is in direct contrast to silica removal by hydrotalcite-like LDHs.26,27 Silicates being nonvolatile ions accumulate in the solid phase on repeated cycling, leading to decreased efficiency. Carbonates are volatile ions and behave differently. Pure CaO is used for high-temperature CO2 cycling. However, there are two major drawbacks of CaO: (i) the thermal penalty is high.28,29 The CO2 uptake temperature is close to 650 °C for CaO and 450 °C for Ca(OH)2, whereas no decomposition of the formed CaCO3 is observed up to 900 °C (Figure 6) in both cases. (ii) The CO2 uptake falls sharply by 80% during thermal cycling;30 the fall is attributed to hightemperature sintering of the CaO crystallites.30,31 There have been major technological and scientific efforts to alleviate these problems by (i) appropriate engineering to utilize the heat generated by the exothermicity of the carbonation cycle and reduce the thermal penalty32,33 and (ii) by dispersing CaO on an inert oxide support such as mineral mayenite to prevent agglomeration and sintering of the active CaO crystallites.34,35 The use of hydrocalumite-based Ca-rich LDH as a precursor has the potential to generate in single step, active CaO, as well as mayenite (Ca2Fe2O5 in the [Ca−Fe] LDH) for effective CO2 uptake. Further, the use of the intercalated Cl− ion induces the decomposition of the formed CaCO3 at 900 °C in the case of [Ca−Al−Cl] LDH. In contrast with this, the CaCO3 formed by the uptake of CO2 by pure CaO (Figure 6)
Figure 6. TGA and DTG curves of CaO in flowing CO2.
shows no signs of decomposition at 900 °C. In the case of [Ca−Fe−Cl] LDH, the formed CaCO3 decomposes at 800 °C itself bringing about a 100 °C reduction in the thermal penalty compared to the [Ca−Al−Cl] LDH. The combined/ synergistic effect of Fe and Cl makes the [Ca−Fe−Cl] LDH interesting for practical use. In all cases isothermal treatment of the LDH at temperatures of the three plateaus did not enhance the mass gain, showing that the mass gain in the TGA is not limited by kinetics.
4. CONCLUSIONS Hydrocalumite-like LDHs are superior to the hydrotalcite-like compounds for applications related to mineralization of CO2. These Ca-containing phases are single-source precursors for both CaO and the ternary oxide supports. Among the hydrocalumites, the [Ca−Fe−Cl] phase shows the highest uptake of CO2. The product of CO2 uptake is the calcite form of CaCO3. The presence of intercalated chloride ions appears to reduce the thermal penalty when compared to CaO, whereas intercalated nitrate ions fail to induce the decomposition of the CaCO3 formed as a result of CO2 uptake. 5. EXPERIMENTAL SECTION The [Ca−M−X] LDHs (M = Al and Fe; X = Cl− and NO3−) were synthesized by coprecipitation. In separate experiments, the [Ca−Al−Cl] and [Ca−Fe−Cl] LDHs were prepared by dropwise addition of a mixed metal chloride solution in the stoichiometric ratio ([Ca]/[M] = 2, dropping rate 0.3 mL min−1) to a NaCl solution (100 mL; 0.4 mol L−1) containing five times excess Cl− ions. The pH of 11.5 was maintained for all of the syntheses by dispensing NaOH solution (0.5 mol L−1) at 60 °C using a Metrohm model 836 Titrando, operating in the pH STAT mode. N2 gas was bubbled during precipitation and ageing to avoid carbonate contamination. Once the precipitation was complete, the samples were aged in mother liquor under a N2 blanket for 18 h under vigorous stirring. In the case of [Ca−Al−NO3] LDH, rapid coprecipitation was carried out by mixing the mixed metal nitrate solution and NaOH (50 mL, 2 mol L−1) in single step in a screw cap bottle. NaNO3 (15 times in excess) was added to the NaOH solution prior to mixing. The pH of the final mixture was >12. The coprecipitated slurries were aged in mother liquor (48 h, T 85 °C). All of the precipitates were centrifuged, washed with boiled type II water (specific resistance 15 MΩ cm, Millipore 3202
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Academic water purification system), and dried in a hot air oven at 60 °C. 5.1. Wet Chemical Analysis. The compositions of the LDH were determined by a combination of techniques. The Ca and Fe contents were determined by atomic absorption spectroscopy using a Shimadzu model AA-6650 atomic absorption spectrometer. The Al content was estimated gravimetrically by dissolving a preweighed amount of the LDH in a mineral acid and precipitating Al3+ in an ammoniacal buffer. The resulting gelatinous precipitate was calcined (T ≈ 900 °C) and weighed as Al2O3. The anion contents were determined by dissolving an accurately weighed amount of the LDH in a suitable mineral acid. The anions released were estimated by ion chromatography (Metrohm model 861 advanced compact ion chromatograph with Metrosep SUP5 150 column). 5.2. Characterization. All samples were characterized by PXRD using a Bruker D8 ADVANCE powder diffractometer (source Cu Kα radiation, Ni filter, λ = 1.5418 Å) operating in reflection geometry. The PXRD patterns of the LDHs were indexed by using APPLEMAN program, part of the code PROSZKI suite of programs. The PXRD patterns of the oxide residues were indexed by “finger printing” by comparing the observed 2θ values with those reported in the powder diffraction files. TGAs were carried out using a Mettler Toledo TGA/SDTA 851e system (25−900 °C, heating rate 5 °C min−1) driven by STARe 7.01 software. The analyses were carried out in flowing N2 (control) and flowing CO2 (test) to measure the mass gain if any, as a result of CO2 uptake. In separate experiments, the LDHs were heated in the TGA balance at different temperatures determined by the appearance of plateaus in the TGA profile. The phases obtained were analyzed by PXRD. TGA measurements were also carried out on CaO and Ca(OH)2 used as controls. IR spectra were recorded by using a Bruker model Alpha-P IR spectrometer (Diamond ATR cell, 400−4000 cm−1, 4 cm−1 resolution).
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REFERENCES
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ASSOCIATED CONTENT
* Supporting Information S
The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acsomega.9b00083.
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Article
PXRD patterns of the pristine hydrocalumite-like LDHs; tables of 2θ values and indexing of the reflections of the hydrocalumite-like LDHs; infrared spectra and TGA profiles of phases obtained at the different plateau temperatures; and TGA data of the nitrate-intercalated LDH (PDF)
AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. ORCID
P. Vishnu Kamath: 0000-0002-3549-7024 Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS Authors thank the Department of Science and Technology, Government of India, for financial support. Authors also thank the Reviewers for their useful comments. 3203
DOI: 10.1021/acsomega.9b00083 ACS Omega 2019, 4, 3198−3204
ACS Omega
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DOI: 10.1021/acsomega.9b00083 ACS Omega 2019, 4, 3198−3204
