Similarities between Zinc Hydroxide Chloride Monohydrate and Its

Article pubs.acs.org/JPCC

Similarities between Zinc Hydroxide Chloride Monohydrate and Its Dehydrated Form: A Theoretical Study of Their Structures and Anionic Exchange Properties Sérgio R. Tavares,† Viviane S. Vaiss,† Fernando Wypych,‡ and Alexandre A. Leitaõ *,† †

Departamento de Química, Universidade Federal de Juiz de Fora, 36036-330 Juiz de Fora-MG, Brazil CEPESQ - Research Center in Applied Chemistry, Departamento de Química, Universidade Federal do Paraná, P.O. Box 19081, 81531-980 Curitiba-PR, Brazil

‡

ABSTRACT: The electronic structures of zinc hydroxide chloride monohydrate (ZHC) and its dehydrated form were calculated by ab initio density functional theory methodology in order to study some of their properties, such as anionic exchange. It was observed that the dehydrated structure showed no major differences in the unit cell parameters compared to hydrated one, except for the basal parameter c, which was smaller in the dehydrated structure. Charge difference plots and projected densities of states were computed to study the interactions between water molecules and the grafted chlorides. Differences between this kind of layered hydroxide salt and hydrotalcite-type compounds could be noticed and understood by the calculations. It could be observed that there were no drastic changes due to the dehydration. The calculated dehydration potentials were used to predict the water loss temperature and the thermodynamic features of the chloride anionic exchange reactions of Zn5(OH)8Cl2·H2O using different halide anions, and it was found that only the exchange with iodine can be performed spontaneously at 298.15 K.

can be seen in Figure 1. Its structure was first discovered in 1961 with crystals synthesized by Nowacki and Silverman.2 However, Hawthorne and Sokolova3 reconducted the structural refinement of simonkolleite with polycrystals. As mentioned above, although less investigated, the layered hydroxide salts share many similarities with the layered double hydroxides, and consequently, they have similar potential

1. INTRODUCTION Natural or synthetic layered compounds are studied because of their industrial applications as catalysts, anion exchangers, nanostructured oxide precursors, matrixes for slow delivery of different drugs, functional fillers in polymers, and so on. Layered double hydroxides (LDHs) and layered hydroxide salts (LHSs) are good examples of those materials with layers similar to the mineral brucite, Mg(OH)2. Brucite’s structure is obtained by stacking of two-dimensional layers built from edge-sharing octahedra whose centers are occupied by Mg2+ cations with six hydroxyls at the vertices. With only divalent cations, the neutral trioctahedral layers are connected by hydrogen bonds and the interlayer regions are empty. The brucite layer is very important because it occurs as part of the structural building blocks (layers) of many different clay minerals, including talc, different micas, and chloride mineral groups. During the synthesis, modifications can be induced in order to obtain the LHS structure. If only divalent cations are maintained, some hydroxyls can be exchanged by monovalent and divalent anions (type I). Part of the octahedral sites can also be removed and split into two tetrahedral sites occupying the upper and lower parts of the vacancy (type II). Those tetrahedrons may contain anionic species (type IIa) or neutral species (type IIb) that are bound to the central atom.1 The typical formulation of LHSs can be represented by M2+(OH)2−x(An−)x/n·yH2O. Zinc hydroxide chloride monohydrate (ZHC), Zn5(OH)8Cl2·H2O, has a rhombohedral structure and is found as a mineral called simonkolleite. Its space group is R3m, and the chlorides are coordinated to the tetrahedrons, as © 2014 American Chemical Society

Figure 1. Side views of (A) the conventional unit cell and (B) the reduced unit cell of a simonkolleite structure fragment. Received: April 25, 2014 Revised: July 29, 2014 Published: July 29, 2014 19106

dx.doi.org/10.1021/jp504051z | J. Phys. Chem. C 2014, 118, 19106−19113

The Journal of Physical Chemistry C

Article

molecular graphics were generated using the XCRYSDEN24,25 graphical package, and the simulated X-ray powder diffractograms were obtained using Mercury software.26 In order to reduce computational costs, we constructed a reduced unit cell based on the original ZHC structure (Figure 1A), whose unit cell has 77 atoms, belongs to the space group R3m (polytype 3R2), and has the following parameters: a = b = 6.3412 Å, c = 23.650 Å, α = β = 90°, γ = 120°. This cell can also be described by the vectors v1 = (a, 0, 0) and v2 = (−a/2, 31/2a/2, 0) and the stacking vector v3 = (0, 0, c). In the reduced unit cell (Figure 1 B), we constructed a new stacking vector based on previous work of our group.11 This procedure substitutes the point symmetry for the translational one, but it does not alter the crystalline structure. Thus, the reduced unit cell can be described by the vectors u1 = v1 = (a, 0, 0), u2 = v2 = (−a/2, 31/2a/2, 0), and u3 = v1/3 + 2v2/3 + (0, 0, c/3) and has 26 atoms. For comparison with experimental data, the parameter c was recalculated after the total geometry optimization had been done using the expression c = 3u3z, where u3z is the z component of u3. We consider the following reaction for the dehydration process:

applications. Many intercalation and immobilization studies have been performed using LHSs and LDHs as matrixes for many reasons, since they provide protection for those intercalated species and their toxicities are decreased.4,5 Another interesting property is the selectivity of the anion exchanges, which shows great potential for the removal of pollutants and analytical separation.6 An improvement in the yields of intercalated catalysts is also observed, along with the advantage of separarating the catalyst from the reaction medium and the possibility of reusing the material after a simple process of washing and drying.7 Those layered materials are good precursors of oxides, which are largely used as industrial catalysts, by means of thermal treatment.8 More recently, LHSs were also proposed as functional fillers in polymer nanocomposites, replacing clay minerals with many advantages.9,10 We have been exploring density functional theory (DFT) combined with periodic boundary conditions to propose a structural model for LDHs with the 3R1 polytype11 and to examine several properties of these materials, including changes occurring in the dehydration of hydrotalcites containing Cl− and CO32− counteranions;12 layer−anion and intermolecular interactions in Zn−Al−An− hydrotalcite-like compounds (An− = Cl−, F−, Br−, OH−, CO32−, NO3−) obtained from anion exchange on the LDH Zn2/3Al1/3(OH)2Cl1/3·2/3H2O;13 and the reaction pathways connected to the initial steps of thermal decomposition of LDH compounds.14 We have also been investigating brucite-like compounds and their properties, such as the formation energy of mixed neutral layered hydroxides,12 the reaction mechanisms of fluoride incorporation into brucite layers [Mg(OH)2−xFx compound] originating from the reaction between HF and brucite,15 the degradation reaction mechanism of the Sarin molecule using brucite,16 and the formation reaction mechanisms of hydroxide anions from Mg(OH)2 layers.17 However, in contrast to brucite-like and hydrotalcitelike compounds, there have been no theoretical studies involving layered hydroxide salts. Therefore, it would be very interesting to study the existing interactions of those compounds and some phenomena such as dehydration and anion exchange. Obviously, this work should be able to provide guidance to existing and future experiments. This report is divided in three parts: a structural comparison of Zn5(OH)8Cl2·H2O and Zn5(OH)8Cl2, a thermodynamic study of the dehydration process, and potential anion exchange of Zn5(OH)8Cl2·H2O with different halides such as F−, Br−, and I−.

Zn5(OH)8 Cl 2·H 2O(s) → Zn5(OH)8 Cl 2(s) + H 2O(g) (R1)

This reaction occurs experimentally in the temperature range from 100 to 110 °C.27,28 The dehydrated compound was built by simply removing the water in the interlayer space of the optimized structure. The anion exchange energies for the studied systems were calculated considering the following reaction: Zn5(OH)8 Cl 2· H 2O(s) + 2A−(aq) → Zn5(OH)8 A 2 ·H 2O(s) + 2Cl−(aq)

(R2)

where A− represents an anionic species (F−, Br−, I−) that can replace the Cl− in the struture. We carried out vibrational calculations to characterize optimized structures as either minimal or maximal. The vibrational modes were obtained from phonon calculations, which were based on the harmonic approximation by density functional perturbation theory (DFPT),29,30 at the Γ q point, and the convergence threshold was 10−14. The infrared absorption spectra were simulated using a Lorentzian distribution curve with a full width at half-maximum of 10 cm−1. In order to analyze the thermodynamics, we computed the variation of entropy, enthalpy, and Gibbs free energy related to the dehydration and exchange reactions, which were obtained from the calculated thermodynamic property of each solid or isolated ionic species. The enthalpies H(T) and entropies S(T) of minerals in the solid state were calculated through the following approximation:

2. METHODOLOGY All of the ab initio calculations were performed using the codes available in the Quantum Espresso package,18 which implements the DFT19,20 framework with periodic boundary conditions. We used the generalized gradient approximation (GGA/PBE)21 for the exchange−correlation functional, and the ion cores were described by the Vanderbilt22 ultrasoft pseudopotential. The Kohn−Sham one-electron states were expanded in a plane-wave basis set until the kinetic cutoff energy of 60 Ry (480 Ry for the density). Monkhorst−Pack23 meshes of 5 × 5 × 5 k-point sampling in the first Brillouin zone were used for all cells. Equilibrium lattice parameters and atomic positions for all structures were found by minimizing the total energy gradient. For each set of lattice parameters, the relative ion positions were relaxed until all of the force components were smaller than 0.001 Ry/bohr. All of the

H(T ) = E elec + EZPE + Evib(T )

(1)

S(T ) = S vib(T )

(2)

where Eelec, EZPE, Evib(T), and Svib(T) are the electronic energy, the zero-point energy (ZPE), and the vibrational contributions to the enthalpy and entropy, respectively. The vibrational data were also used to calculate both the contribution of the lattice thermal vibration to the total energy and the ZPE as shown in our previous work.11,12,15,16 19107



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The ions and H2O molecule in the gas state were treated according to the formalism described in our previous work.11,12,15,16 In this case, the enthalpy contribution was calculated as

Table 2. Main Bond Distances in ZHC

H(T ) = E elec + EZPE + Evib(T ) + Etrans(T ) + Erot(T ) (3)

+ RT

a

where Etrans(T) and Erot(T) are the translational and rotational contributions to the enthalpy, respectively, both of which are equal to 3RT/2, where R is the gas constant. The RT term is equal to pV. The entropy S(p,T) for a gas can be calculated by the expression: S(p , T ) = Strans(p , T ) + S rot(T ) + S vib(T ) trans

(4)

where S (p, T) and S (T) are the translational and rotational contributions to the entropy, respectively. We also considered an additional contribution to the electronic energy in order to remove the long-range Coulomb interaction between the periodic images of charged molecules. The Makov−Payne correction31 was used to compensate for this effect. In order to obtain the contributions of H, S, and G of (aq) (aq) the anions in the aqueous solution [H(aq) A , SA , and GA , respectively], we considered the formation of a hydrated ion from the gas state, as shown in the following equation: ΔEA◦(hyd)

(R3)

Thus, we considered the following expressions: ΔHA(aq) = HA(g) + ΔHA◦(hyd)

(5)

ΔSA(aq) = SA(g) + ΔSA◦(hyd)

(6)

ΔGA(aq) = GA(g) + ΔGA◦(hyd)

(7)

H(g) A ,

S(g) A ,

G(g) A

where and are the calculated enthalpy, entropy, and Gibbs free energy of the anion in the gas phase, respectively, and ΔHA°(hyd), ΔSA°(hyd), and ΔGA°(hyd) are the standard enthalpy, entropy, and Gibbs free energy of formation of the hydrated ion A− from its gaseous state, respectively, which were taken from experimental measurements reported in the literature.32

3. RESULTS AND DISCUSSION 3.1. Structural Analysis. For the sake of simplicity, Zn5(OH)8Cl2·H2O(s) and Zn5(OH)8Cl2 will be called ZHC and d-ZHC, respectively. The simulated cell parameters and other main interatomic distances of ZHC can be found in Tables 1 and 2, respectively. As can be seen, good agreement between the simulated and experimental values was obtained, and the highest relative Table 1. Main Geometrical Parameters of ZHC a (Å) b (Å) c (Å) α (deg) β (deg) γ (deg) V (Å3) a

simulated

experimentala

relative error (%)

6.45 6.44 23.67 90.09 89.90 120.03 852.0

6.34 6.34 23.65 90.00 90.00 120.00 823.4

1.76 1.61 0.07 0.10 0.22 0.02 3.47

experimentala

relative error (%)

H···Cl (Å) Zn−Cl (Å) ZnOh−O (Å)

2.24 2.35 2.16

2.26 2.31 2.13

0.61 1.57 1.48

ZnTd−O (Å)

1.97

1.95

1.30

Taken from ref 3.

errors found were for the a and b axes (1.76 and 1.61%, respectively). The relative error for the c parameter was extremely low (0.07%) compared with those. The main bond distance errors show that the hydrogen bonds between the chlorides and the water molecules were better described than the covalent bonds (0.61 and >1.30%, respectively). It has been reported experimentally that the chlorides interact with the water molecules and with the upper and lower hydroxyls by hydrogen bonds. All of those hydrogen bonds in the structure were predicted as described in the experimental work.3 The tetrahedral and octahedral sites were also maintained during optimization, and it is also good to point out that the reduced unit cell constructed has approximately 33% of the volume of the original ZHC structure.11 Kozawa et al.8 conducted X-ray diffraction (XRD) experiments at higher temperatures for ZHC and noticed no radical changes in the diffraction patterns. Consequently, there was no need to include corrections for thermal expansion in the structure optimization. The polytype of the compound was maintained, as can be seen from the comparison of the simulated and experimental diffraction patterns in Figure 2A. The simulated and experimental (003) peaks overlap completely, and the simulated and experimental (006) peaks differ slightly. Those features also corroborate the simulated unit cell parameters presented in Table 1. This simulated XRD pattern is quite similar to several other ones reported in the literature, including our own data.33−36 As can be seen in Table 3, the simulated infrared absorption spectrum (Figure 3) also showed a good agreement with the experimental data.36−38 The absorption band at 744 cm−1 corresponds to the angular deformation of the hydroxyls. The bands located between 800 and 1200 cm−1 correspond to angular deformations of the Zn−O. The wagging vibration of the water molecule is also located in this region. The absorption bands corresponding to the symmetric and asymmetric stretching of the water molecule and to the stretching of hydroxyls appear between 3400 and 3600 cm−1. Those assignments were done with the eigenvectors of phonon calculations. For d-ZHC, the basal parameter c was 0.48% lower than the hydrated one and the octahedral sites were also maintained, as can be seen in Figure 4 and Table 4. There was no major modification of the other unit cell parameters due to the water removal. The number of hydrogen bonds between the chlorides and the hydroxyls on both sides of the layers are the same. This behavior is not observed for the Zn−Al−Cl HDL compounds, since the presence of intercalated chloride anions changes the number of bonds with the hydroxyls during the dehydration process.12 Figure 2B shows the simulated XRD pattern of dZHC compared with that of ZHC. One can notice that there are no differences between them, especially for the peaks (003) and (006). This can be explained by the small variation of the c parameter in relation to that of the hydrated compound. As

rot

A−(g) → A−(aq)

simulated

Experimental data were taken from ref 3. 19108



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Figure 3. Comparison of the simulated IR spectra of ZHF, ZHC, ZHB, and ZHI.

Figure 2. XRD patterns of (A) ZHC and (B) d-ZHC (Cu Kα).

Table 3. Wavenumber Values of the Main Absorption Modes of ZHC

a

assignment

simulated (cm−1)

exptla (cm−1)

exptlb (cm−1)

νOH νH2O

3506 3464

3520 3480

3500 3454

Figure 4. d-ZHC structure.

δ H2 O

1567

1612

1622

δOH δOH δOH

1025 909 744

1035 895 715

1045 906 725

Table 4. Comparison of the Optimized ZHC and d-ZHC Structures a (Å) b (Å) c (Å) α (deg) β (deg) γ (deg)

Taken from ref 37. bTaken from ref 36.

mentioned previously, the experimental diffraction peaks do not change with temperature below 150 °C, which is in good agreement with our simulated diffraction pattern.8 Charge density difference plots were also made for ZHC and d-ZHC in order to study the interactions between the layers, the chloride anions, and the hydration water molecules within the layers. Two 1 × 1 × 2 supercells were constructed in order to compute the total charge density, and the layers of this resulting supercell were removed individually in order to subtract their charge densities. The charge density differences Δρ(r) at a point r in Figures 5 and 6 are defined in eqs 8 and 9, respectively:

ZHC

d-ZHC

6.45 6.44 23.66 90.09 89.90 120.03

6.44 6.44 23.55 89.83 90.04 119.98

Δρ(r) = ΔρZHC (r) − ΔρH O (r) − ΔρZHC without H O (r) 2

2

(8)

Δρ(r) = Δρ1 × 1 × 2 (r) − Δρupper layer (r) − Δρlower layer (r) (9)

Figure 5B indicates that the water molecules interact with both chlorides equally. There is a charge transfer from the water hydrogens to both chlorides. Figure 6A,B shows the 19109



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Figure 7. Simulated IR spectrum of d-ZHC.

band at 3725 cm−1, corresponding to hydroxyl stretching modes, that could be hidden by other bands in the case of an experimental spectrum. The most obvious change in this spectrum is the loss of the water deformation band. Despite that, it is very clear that the removal of the water molecule does not lead to any remarkable change in the vibrational spectrum. The projected density of states (pDOS) shows that the chlorides’ basicity does not change as a result of dehydration (Figure 8). A different behavior was observed when the Zn− Al−Cl HDLs were submitted to the dehydration process.12 Also, Figure 6 shows that the density difference analyses indicated much more interaction between the chlorides and the layer than between the chlorides and the water interlayer molecules. Comparing the ZHC and d-ZHC results in Figure 8, one can practically observe no perturbation in the electronic transfer with the chlorides when the water molecules are absent. For hydrotalcites (LDHs), the chlorides are more connected with the interlayer water molecules, and after the dehydration they change to the center of the prismatic sites formed by the hydroxyls. Also, the pDOS and HOMO analyses revealed that the chlorides are as basic as the hydroxyls in the HDL dehydrated form. Those calculations also showed that the tetrahedral and octahedral hydroxyls from both sides of the layers are equivalent. These hydroxyls are represented in Figure 9 3.2. Thermodynamic Study of the Dehydration Process. Figure 10 shows the computed Gibbs free energy change for the dehydration process as a function of temperature from 10 to 160 °C at 1 atm. When the temperature reaches above 104 °C, the process becomes spontaneous. This result is in good agreement with the experimental data, which show that this process becomes spontaneous in the range 100−110 °C.27,28 This calculation validates the simulated structure of dehydrated simonkolleite. The entropic contribution of this process tends to increase with increasing temperature from TΔS = 38.42 kJ mol−1 at 10 °C to 56.75 kJ mol−1 at 150 °C, whereas the enthalpy change decreases very slightly from ΔH = 51.11 kJ mol−1 at 10 °C to 50.57 kJ mol−1 at 150 °C. This also shows that the dehydration process is endothermic. We can understand the lower dehydration temperature compared with that of Zn−Al−Cl HDLs (around 125 °C) on the basis of the difference in the interactions as we previously commented. In the hydrotalcite-type compounds, the water molecules are more connected with the chlorides, whereas in

Figure 5. Charge density differences for (A) ZHF [Δρ(ZHF) − Δρ(H2O) − Δρ(ZHF without H2O)], (B) ZHC [Δρ(ZHC) − Δρ(H2O) − Δρ(ZHC without H2O)], (C) ZHB [Δρ(ZHB) − Δρ(H2O) − Δρ(ZHB without H2O)], and (D) ZHI [Δρ(ZHI) − Δρ(H2O) − Δρ(ZHI without H2O)]. The red region indicates an increase in charge density, and the blue one indicates a depletion of charge density. The contour spacing is 0.003 electrons/bohr3.

Figure 6. Charge density differences for (A) ZHC [Δρ(two layers of ZHC) − Δρ(first layer of ZHC) − Δρ(second layer of ZHC)] and (B) d-ZHC [Δρ(two layers of d-ZHC) − Δρ(first layer of d-ZHC) − Δρ(second layer of d-ZHC)]. The red region indicates an increase in charge density, and the blue one indicates a depletion of charge density. The contour spacing is 0.003 electrons/bohr3.

interactions between the layers of ZHC and d-ZHC, respectively, which are merely charge interactions. It can be noticed that those charge interactions occur between the chlorides and the hydroxyls (ZHC and d-ZHC) and between the water molecule and a hydroxyl (ZHC). Figure 7 shows the infrared absorption spectrum of d-ZHC. The absorption bands between 505 and 1050 cm−1 correspond to the angular deformation of the layer hydroxyls. The bands located in the region between 3415 and 3730 cm−1 correspond to hydroxyl stretching modes. Those bands are narrow compared with those of the hydrated ZHC because there is a loss of vibrational modes due to the removal of water. One can also observe in Figure 7 that there is a narrow, low-intensity 19110



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Figure 10. Gibbs free energy of the dehydration process as a function of temperature. Experimental data were taken from refs 27 and 28.

similar to that of ZHC, but the numbers of hydrogen bonds and the basal parameter c were reduced and increased, respectively, for the compounds with bromide (24.45 Å) and iodide (25.52 Å), while the structure containing fluoride presented more hydrogen bonds and a smaller basal spacing. Unfortunately, attemps to synthesize these compunds using a procedure similar to that used for ZHC were unfruitful. Their unit cell volumes also showed a variation, since there are more interactions in the interlayer region with increasing electronegativity of the anion; hence, ZHF has the lowest unit cell volume, whereas ZHI has the highest one. However, the angle values changed very slightly. Their main cell parameters can be found in Table 5. Table 5. Simulated Main Geometrical Parameters and Properties of the Halogenated Compounds Figure 8. DOS and pDOS of both chlorides from (A) ZHC and (B) d-ZHC. The Fermi energy was considered as the zero of energy and is represented by the vertical dotted lines.

a (Å) b (Å) c (Å) α (deg) β (deg) γ (deg) V (Å3)

ZHF

ZHB

ZHI

6.40 6.40 21.55 89.72 90.12 120.27 762.8

6.48 6.48 24.44 89.97 89.86 120.02 889.6

6.53 6.49 25.51 89.72 89.99 119.77 932.8

Charge density difference plots were made in order to compare the interactions between the water molecules and the layers (Figure 5). Those plots show that interactions between the water molecules and the halogens tend to decrease with the diminishing of electronegativity. Thus, the fluorides interact the most with the water molecule in comparison with the other ones. These observations are also in agreement with their geometries. Their hydrogen-bond distances (1.82, 2.36, 2.53, and 2.71 Å for ZHF, ZHC, ZHB, and ZHI, respectively) tend to decrease with the electronegativity, indicating that the interaction strength in the interlayer region is driven by the anion. The simulated IR absorption spectra for the halogenated compounds (Figure 3) were also constructed for the sake of comparison. The bands between 1520 and 1640 cm−1 were ascribed to the scissor-type deformation of the water molecule in the ZHF, ZHC, ZHB, and ZHI. There is a shift of the water deformation wavenumber values due to the interaction between

Figure 9. Representation of the hydroxyls: colored hydroxyls represent (A) octahedral and (B) tetrahedral hydroxyls.

the case of ZHC the chlorides are much more connected with the layers and the hydrogen bonds are weaker. The final effect is that the water molecule is less imprisoned in the ZHC structure and can leave at a lower temperature. 3.3. Anion Exchange Reactions. In this study, we used the ZHC structure as the precursor. The exchange capability was evaluated by changing the Cl− anions of the precursor compound to F − , Br − , or I − anions, producing the Zn5(OH)8A2·H2O compounds, where A denotes the anion. For the sake of simplicity, the compounds containing fluoride, bromide, and iodide will be called ZHF, ZHB, and ZHI, respectively. The halogenated compounds presented structures 19111


The Journal of Physical Chemistry C the hydrogens and the halogens. ZHF presents the highest wavenumber value for its water deformation because the fluorine interacts with the water molecules. The absorption bands in the range of 3210 to 3620 cm−1 correspond to the symmetric and asymmetric stretching modes of the water molecule and to the hydroxyl stretchings. The computed Gibbs free energy changes for the topotactic anionic exchange reactions are shown in Table 6. The

ΔG (kcal mol−1)

ΔH (kcal mol−1)

TΔS (kcal mol−1)

ZHC → ZHF ZHC → ZHB ZHC → ZHI

1.91 1.99 −1.62

3.35 −0.27 −4.61

1.44 −2.26 −2.99

■

REFERENCES

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thermodynamic properties were divided by the number of the zinc cations located at the unit cell. A negative value of the Gibbs free energy change indicates that a certain reaction is spontaneous, and therefore, the values presented in Table 6 show that the only spontaneous exchange at 298.15 K is the exchange between chloride and iodide. It can also be observed that this process is more exothermic (ΔH < 0) than the exchange for bromide (i.e., this anionic exchange leads to a decrease of energy). There is also a higher contribution from the enthalpy than from the entropy for this exchange (ZHC → ZHI). Calculations such as these were previously performed for the anion exchange reactions of hydrotalcites containing Zn and Al cations with Cl−, F−, Br−, OH−, CO32−, and NO3−,13 and good agreement with experimental data collected by Myaita39 was obtained.

4. CONCLUSIONS DFT calculations were employed in order to understand the main interactions in the layered compound ZHC and the thermodynamics of its possible anion exchanges. The structural model was validated by comparison with the experimental unit cell parameters and by the agreement between the theoretical and experimental bands from infrared absorption spectroscopy. Density difference analyses indicated much more interaction between the chlorides and the layers than between the chlorides and the interlayer water molecules. Therefore, the dehydration process leads to minimal changes in the layered structure. The simulated dehydrated structure was proved to be the true dehydrated structure by computation of the Gibbs free energy as a function of temperature. We can understand the smaller dehydration temperature compared with those of Zn− Al−Cl HDLs on the basis of the difference in the interactions. Thermodynamic calculations showed that the only spontaneous exchange at 298 K is between chlorides and iodides. However, it must be noted that the energy barrier of this process was not computed, and therefore, its kinetics could highly influence its duration.

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ACKNOWLEDGMENTS

The authors thank the Brazilian agencies CAPES (Ph.D. fellowship for S.R.T.), CNPq (research grant for A.A.L. and F.W. and process 477706/2013-4), and FAPEMIG (CEX BDP00256/13 and CEX PPM-00262/13) and Vale S.A./FAPEMIG (CEX RDP-00138/10) for financial support. We also thank the CENAPAD-SP computational center for the use of its facilities.

Table 6. Thermodynamic Potentials for the Anion Exchange Rections of ZHC at 298.15 K Anion

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