Langmuir 1994,10, 1207-1212
1207
Intercalation of Macrocyclic Compounds (Crown Ethers and Cryptands) into 2:l Phyllosilicates. Stability and Calorimetric Study P. Aranda,? B. Casal,?J. J. Fripiat,j and E. Ruiz-Hitzky*pt Instituto d e Ciencia de Materiales, CSIC, CISerrano 115-bis, 28006 Madrid, Spain, a n d Department of Chemistry and Laboratory for Surface Studies, P.O.Box 413, University of Wisconsin-Milwaukee, Milwaukee Wisconsin 53201 Received May 3, 1993. I n F i n a l Form: February 3, 1994" Crown ether and cryptand intercalation processes into 2:l phyllosilicates are exothermic. The nature of ligand and the nature of the interlayer cations are the most important factors influencing the net heat of intercalation. The enthalpy of intercalation is more exothermic than that of the corresponding complexation reaction of the same macrocycle with the same cation in homogeneous medium. The characteristics of the host silicate as well as the solvent used in the intercalation process also influence the enthalpy. The thermal treatment of the intercalated materials shows that the complexes remain unaltered up to 300 "C under nitrogen.
Introduction
The structural features of crown ether and cryptand cationiccomplexes, and the nature of macrocyclic ligandcation interactions have been extensively studied in homogeneous medium.l-9 In particular, thermodynamic and kinetic studies have been reported for a vast number of complexes.1*2Thermodynamic parameters have been correlated with different factors such as the ratio of t h e size of t h e cation t o the size of the macrocycle, t h e nature of the anion and/or of t h e solvent, the nature and geometry of binding sites, etc.13 On the other hand, t h e formation of crown ethers and cryptand cationic complexes in heterogeneous media by intercalation into t h e interlayer spaceof swelling clay minerals has been reported earlier.The present contribution aims t o study the energetics of a selected number of these processes by adsorption microcalorimetry. Indeed, i t seems interesting t o correlate the energy associated with t h e interlayer complex formation with the ion-transport properties of these materials, which could be used as solid electrolytes or as ion-selective membranes.Sl1 Along this line of thought it is also important to study t h e thermal stability in order to
establish t h e temperature range over which the intercalates could be used as solid electrolytes.
Experimental Section
Materials and Methods. Wyoming montmorillonite (Upton) and hectorite (Hector, CA) were purified and prepared as homoionic samples (particle size 600 2.6 42 (endo.) DB24C8/Na-mont. 54 333 9.4 341 (endo.) >600 4.2 626 (endo.) C222/Na-mont. 51 1.2 51 (endo.) 227 (endo.) 320 14.8 376 (endo.) 446 (endo.) 559 4.2 sample 12C4/Na-mont.
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35
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20
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(endo.) = endothermic process.
for the Na-, K-, and Ba-montmorillonites preheated at 550 "C (24 h). The barium-smectite can intercalate 18C6 even after a thermal pretreatment at 650 "C, although in this case the collapsed phase of Ba-montmorillonite ( d ~ = 9.7 A) is also observed. It was confirmed that pure methanol is not able to produce the swelling of any of these calcined samples. As discussed later, the intercalation enthalpy of the thoroughly dehydrated samples ( A H 3 in Table 1)is always lower than (i) that obtained for the uncalcined samples (MI)and even than (ii) that of the complexation reaction carried out in homogeneous conditions (A&). The intercalation of macrocyclic compounds into clays preheated above 450 "C has never been shown before. It is well-known19 indeed, that in similar experimental conditions,water and ethylene glycol do not swellsmectites (like montmorillonite and hectorite) calcined above 500
"C. Thermal Stability. The thermal stability of macrocyclic intercalation compounds is inferred from the TGA, DTG, and DTA results, which are summarized in Table 5. These thermal analyses were carried out on Namontmorillonite intercalation compounds under NZ atmosphere. In general,the weight loss between 25 and 700 "C occurs in three steps. The first step (2' < 100 "C) is related with the loss of water readsorbed upon exposing the samples to a relative humidity of 55% prior to the thermal analysis. The second step appears at 250-350 "C and correspondsto an endothermicprocess associated with the loss of the organic ligands. Differences in the temperature of these endothermic processes depend on (19) Brindley, G. W. Clay Miner. 1966,6, 237.
b
% org % org comp comp (chem (TGA) anal.)
Figure 4. X-ray diffraction patterns and b a d spacings (dL)of 12C4/Na-montmorillonite thermally treated for 30 min (under vacuum) at 25 "C (a), 280 "C (b), 330 O C (c), and 360 "C (d).
the nature of the macrocyclic compound. The third step above 550 "C is the endothermic process related to the dehydroxylation of the host lattice.20721 Special features are observed for the 12C4and DB24C8 intercalated into Na-montmorillonite. In the first case, the macrocyclic compound is lost in two steps corresponding to the change from a 2:l complex to a 1:l complex. This assertion is supported by the XRD data of samples heated at different temperatures (Figure 4), which show that the 2 1 stoichiometry of the 12C4 intercalation compound is preserved up to about 300 "C (A~L = 17.5 A at 280 "C). When the samples are heated above 300 "C, the stoichiometry corresponds to a 1:l ligand-cation complex (A& = 13.6 A at 330 "C). The difference in temperature at which the release of 12C4 shows up in the thermal analysis (>258 "C), compared to that observed in the XRD study (>280 "C), can be explained by the different experimental conditions in which the measurements have been carried out (Nz flow us vacuum and different heating rates). On the other hand, the DB24C8/Na-montmorillonite compound exhibita a weight loss much lower than the organic content determined by elemental microanalysis. The IR spectrum of this compound in the 1800-1200 cm-l region shows significant changes when the sample is heated. The interpretation of these modifications up to 350 "C is difficult, but they could be related to a polycondensation of benzene rings. This hypothesis is supported by the analogy with the IR spectra reported by Soma and coworkerszz of polyphenylene species intercalated into montmorillonite. These polymeric species remain intercalated even at high temperatures and, consequently,the mass balance deduced from TGA and from elemental microanalysis does not coincide. (20) Green-Kelly, R. Cluy Miner.Bull. 1952,1, 221. (21) McKenzie, R. C., Ed. Differential Thermal Analysis;Academic Press: London & New York, 1970. (22) Soma, Y.;Soma, M.;Harada, I. J.Phys. Chem. 1984,88, 3034.
Intercalation of Macrocyclic Compounds
Discussion In homogeneous medium, the complex formation requires a complete rearrangement of the cation solvation shell, the solvent being replaced partially by the macrocyclic ligand. For a specific solvent, the complexation enthalpy is usually deduced from the energy balance between different terms involving the ion desolvation energy, the energy associated to conformational changes of the macrocyclic ligands, and the energy evolved in the direct cation-ligand binding.l*23 In heterogenous conditions, i.e. in the intracrystalline environment of a phyllosilicate,additionalfactors related with structural features of the host solid must be taken into account. Thus, the measured intercalation enthalpy (AHl) can be expressed as indicated in eq 2
where AHWmpler is the enthalpy of the cation-ligand coordination which is always exothermic since it corresponds to a process involvingbond formation. The AHsheU term is ascribedto the energy required to remove a fraction of the solvent in the coordination shell around the interlayer cations. This term must be endothermic. The last term, AHothe,, accounts for several factors such as (i) change of the interlayer spacing (interlayer expansion), (ii) ligand silicate surface interactions, (iii) lateral interactions between neighboring complexes imposed by a constrained geometry existing in the silicate interlayer region, (iv) interaction between the ligand and uncomplexed cations, and (v) redistribution of the electrical charge in the octahedral and tetrahedral sheets of the host lattice due to the screening of the cationic charge by the macrocyclic compound. The sign of Mothe cannot ls be predicted a priori. Analysis of the enthalpy values presented in Table 1 arises to several conclusions: 1. Except for the C(222) macrocycle, the enthalpy of complexation per macrocycle and cationic equivalent in the swollen interlayer is not very different from the enthalpy of complexationof the correspondingmetal salt in the solvent used for swelling the clay. This suggests that the macrocycle is efficientin screeningthe interaction between the cation and the solvent. 2. In the two cases where the clay interlayer was indicating collapsed, the - A H 3 is much less than that a fraction of the energy of complexation is used for swelling and solvatingthe clay interlayer. By subtracting A H 1 from AH3, this energy is about 31 kJ/mol per 18C6 macrocycle. The cross-section area of 18C6 is 95 Az per molecule.24 According to Giese2sthe energy associated with cleaving 1mol of mica is in the order of 120 kJ. This corresponds to the separation of two 45 A2 surfaces from the equilibrium distance (with both the faces in contact) up to a few A. The experimental value (AH1 - AH31 represents only about l/4 of the value calculated by Giese. A relatively large discrepancy is expected since (i) a mica would not swell in the presence of a macrocycle and (ii) the cleavage energy is not known for smectites because stacking disorder does not allow computation. The conclusionis that the assignment of the value (AH1- A H 3 1 to the swelling energy is reasonable. (23) Lindenbaum, S.; Rytting,J. H.; Stemson, L. A. In Progress in Macrocyclic Chemistry;Izatt, R. M., Christensen,J. J., Eds.;John Wiley & Sone: New York, 1979; Vol. 1, p 219. (24) Izatt, R. M., Christensen, J. J., Eds. Progress In Macrocyclic Chemistry Vol. 1 & 2; John Wiley & Sons: New York, 1979 and 1981. (25) Giese, R. Clays Clay Miner. 1978,26, 55.
Langmuir, Vol. 10, No. 4, 1994 1211
3. The C(222)macrocycle behaves very differently, AH1 being almost twice as large as those reported for the other macrocycles. While as stated in item 1,the 15C5,18C6, and DB24C8 macrocycles are efficiently screening the cation from its solvation shell ( A H 1 = A&), this does not appear to be the case of C(222). The relative uncertainty on the stoichiometry of the (3222) complexes does not permit further speculation. The influence of the nature of the solvent appears to be somewhat different from that observed for reactions in solution. Indeed, the intercalation enthalpy of 15C5 in Na-montmorilloniteis lower in acetonitrile than in methanol (Table 31, which is contrary to the observations reported in solution.scJ8 This can be explained in assuming that acetonitrile does not produce a completereplacement of the water molecules of the interlayer cation solvation shell and, consequently,an additional energy is needed to remove the solvationshell. That the intercalation of 18C6 from acetonitrile presents an adsorption enthalpy close to zero can be justified by taking into account the strong interaction of 18C6 with acetonitrile molecules.3c Consequently, a significant ligand desolvation enthalpy (included in the Mothe term) rs must be considered in this case. Concerning the influence of, the host silicate nature (Table 2), two facts could be related to the observed differences: (i) location of the charge in the silicate sheets and (ii) difference in the CEC. Hectorite presents lower CEC and then the available interlayer area per cation is larger than in montmorillonite. In general, this fact produces a larger percentage of complexed interlayer cations in hectorite than in montmorillonite. So, the interactions between other neighboring complexes and uncomplexed cations are different. Although enthalpy differences between hectorite and montmorillonite are in general small, the contribution of these interactions can give differences in the value of the AH,,then term. The effect of the location of the charge in the latter might, perhaps, explain the large differences in adsorption enthalpies of C(222) in both silicates. Montmorillonite presents a partial “beidellitic character” consisting in localization of negative charge in the tetrahedral sheets (about 10% in this sample). When the C(222)compound is intercalated in a montmorillonite where not all of the water molecules are exchanged by the solvent (in our case methanol) an acid-base reaction could take place. This type of reaction is strongly exothermic and could not be negligible. However, IR spectroscopy does not reveal protonation of the macrocycle. In spite of lack of spectroscopic confirmation the much less exothermic intercalation of C(222) in preheated samples where interlayer water is completely removed could be indirect evidence. The difference between the enthalpy values obtained for the C(222) intercalation into homoionic montmorillonites and the enthalpy of complexationin homogeneous conditions is quite remarkable. The so-called “cryptate effect”,namely the ability of these bicyclic compounds to isolate the cations from the environment providing an exceptionally efficient envelope as solvation shell, is invoked for complexes formed in solution26and in the interlayer space of smectites.6 The effectiveness of the cryptand-cation interaction could be enhanced in the complexes formed in the intracrystalline region of the smectites, because the characteristic diffuse charge of the (26) Lehn, J. M.; Sauvage,J. P. J. Chem. Soc., Chem. Commun. 1971, 440.
1212 Langmuir, Vol. 10, No. 4, 1994
silicatelayers acta as “counterions”of the interlayer cations and can prevent the strong anion-complexed-cation interactions which operate in the complexesin homogeneous solution. In addition, it can be expected that lateral interactions in a closely packed arrangement of interlayer alkali-metal cryptates and those with the uncomplexed cations give a significant exothermic contribution. Compare, for instance, the C(222) intercalation in montmorillonite with divalent cations (Ba2+) with those of monovalent cations, the former being more spaced than the latter. In fact, in monovalent homoionic smectites, due to steric hindrance, about 50% of the cations remain uncomplexed, probably in side interaction with the cryptand (exothermiceffect),whereas in divalentsaturated smectites all the cations are involved in the complex formation. Assuming that the energetic balance involved in the overall process described by eq 2 is mainly related with complexation and, consequently, with the stability of the intercalated complex, the stability of the Na-montmorillonite interlayer complexesshould rank as follows: C(222) >>> 18C6 > 15C5 > DB24C8 > 12C4. This sequence is similar to the one known in solution.13 The interactions between the interlayer cations (Na+ ions) and the crown ethers have been clearly established by 2sNaNMR spectroscopystudy! and the 2sNachemical shift uncorrected for the second order quadrupole effect can be correlated with the number of oxygen atoms in the macrocyclic ligand (Figure 5). Thus, the stronger interactions between cations and macrocycleswith five and six oxygen atoms, as indicated by the more exothermic AH1 value, can be correlated with upfield chemical shifts, that is an increase in the shielding of the sodium nucleus. Finally,the results from the adsorptionmicrocalorimetry can be related to the ion-transport properties of the complexed interlayer cations. The ionic conductivity measured in different crown ether/Na-montmorillonite complexescan be correlated with the adsorption enthalpy. In fact, the Na+ ion conductivity increases in complexes which have been obtained involving the less exothermic proce~ses.g~~~ Thus, in the DB24C8 complex, where the macrocyclic cavity size is too large to accommodate the
Aranda et al. chemlcal shlft (ppm) 0 wIr..xpond (c SO% R.H.) hoctorlt. (1)
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oxygen atoms in the macrocycle Figure 5. Chemical shift of =Na NMR signal for different macrocycle/Na-hectoriteintercalation compounds against the number of oxyethyleneunite of the macrocyclic compound. Note 1: Signal correspondingto the startingNa-hectoritein a sample exposed to air about 50% relative humidity.
cation, the conductivity is greater than that determined for the other crown ether complexes of greater affmity toward the interlayer Na+ cations. This behavior is of great interest in the design of controlled ion-conducting materials based on the intercalation of oxyethylene compounds into s m e c t i t e ~ . ~ ~ * ~ ~ * %
Acknowledgment. Financial support from the Ram611 Areces Foundation and from the CICYT, Spain (Projects: MAT 90-725 and MAT-91-0952-C04-02), is gratefully acknowledged. We are indebted to Dr. J. Laynez and Dr. J. Sanz for their helpful assistance in the microcalorimetry and NMR techniques, respectively. ~~
(27) Aranda, P.; C d ,B.; Galvan, J. C.; Ruiz-Hitzky, E. In Chemical Physics of Intercalation ZI; Bernier, P., Fisher, J. E., Roth, S.,Soh, 5. A., Ede.; Plenum Prw: New York, 1993; p 397. (28) (a) Ruiz-Hitzky, E.; Aranda, P. Adu. Mater. ISSO, 2, 646. (b) Aranda, P.; Ruiz-Hitzky, E. Chem. Mater. 1992,4,1397. (29) Lamb,J. D.; Izatt, R. M.; Swain,C. S.;Christensen, J. J. J. Am. Chem. SOC.1980,102,475.
