Iron-substituted Beta molecular sieve: Synthesis and characterization

Microporous Materials, 2 (1994) 167-177

167

Elsevier Science B.V., Amsterdam

Iron-substituted Beta molecular sieve: synthesis and characterization Ramesh B. Borade and Abraham Clearfield* Department of Chemistry, Texas A & M University, CoUege Station, TX 77843, USA (Received 30 June 1993; aecepted 22 August 1993)

Almraet A synthesis of erystalline ferrisilicate having zeolite BEA topology (Fe-Beta) and eontaining significant quanüties of iron in the framework (22 Fe atoms per unit cell based on 192 T-atoms per unit cell) has been reported. The synthesis of Fe-Beta was earried out using tetraethylammonium hydroxide (TEAOH), 25% methanolie instead of an aqueous solution, as a souree of organic template. X-Ray powder diffracüon (XRD) and scanning eleetron micrography (SEM) were used to cheek the structural identity and phase purity of the Fe-Beta samples. A thermogravimetric analysis (TGA) study showed that the amotmt of TEAOH ions interacting with Fe-Beta framework increases with inerease in the iron content of the sample. The Fe ions in Fe-Beta were in the trivalent oxidation stare which was eonfirmed by an X-ray photoeleetron speetroseopy (XPS) study. A comparison of Si/Fe ratios determined by chemieal analysis and XPS method allowed us to eonelude that Fe 3+ speeies are distributed homogeneously from the bulk to the surfaee of Fe-Beta crystallites. However, the O1, XP speetra indicated the presence of small amounts of extraframework iron oxidic impurities in as-synthesized as weil as calcined samples. The infrared flR) spectra showed three types of hydroxyl groups in Fe-Beta loeated at 3740, 3670 and a very broad band at ~ 3540 cm- t The band at 3680 c m - : was found to be affected by pyridine adsorption and therefore was assigned to structural bridging hydroxyl groups (Si-OH-Fe) formed by subsütution of iron in the zeolite framework. The shift in the IR OH stretching frequeney towards high« ware number and the desorption of NH3 at lower temperature for Fe-Beta samples as compared to Al-Beta zeolite allowed us to conclude that the strength of Brönsted acid sites in Fe-Beta materials is lower than those Brönsted acids sites in AI-Beta zeolites.

Keywords: zeolite-Beta; Fe-Beta synthesis; acidity of Ze-Fe-Beta; infrared study of Ze-Beta; X-ray photoelectron spectroscopy of Z¢-B¢ta

Introduction One of the emerging trends in the area of zeolite synthesis is the isomorphous substitution of A! or Si in the framework by other elements such as Ge, Ga, Fe, B, Ti, etc. Isomorphously substituted zeolites are finding applications not only in the petrochemical industry but also in organic synthesis reaetions 1,1-5]. Barrer 1"6] and Tielen et al. 1'7] have reviewed possibilities of isomorphous substitution of elements other than AI and Si in the zeolite framework. Iron-substituted zeolites a r e a * Corresponding authoß 0927-6513/94/$7.00 © 1994- Elsevier Seience B.V. All rights reserved. $SDI 0927-6513(93)E0050-Q

speeial class of materials with several interests. Szostak and Thomas i"8] reported the synthesis of sodalite, a small-pore zeolite structure (eight-ring), with significant quantities of iron in the framework (SiO2/Fe203 =6-30). Several articles describe the synthesis and physicochemical properties of ferrisilicates belonging to the medium-pore (ten-ring) zeolites !"9-14]. An excellent review artiele was published by Ratnasamy and Kumar [15] deseribing the ferrisilicate analogs of zeolites. However, very little is reported about the substitution of iron in large-pore zeolites (twelve-ring), such as Beta [16] and HY [17] zeolites, in the literature. Zeolite Beta appears to be a promising catalyst for the

168 following reasons: (a) it possesses a twelvemembered ring pore opening; (b) it can be synthesized with a wide range of Si/A1 ratios as compared to the HY zeolite; and (c) it seems to have an application in the hydrodewaxing process [18]. The structure of Beta zeolite was reported by Treacy and Newsam [19] and Higgins et al. [20]. The preparation of Fe-Beta material with significant quantities of iron in the framework is of scientific interest. The highest level of iron content reported in the literature [ 16] for a Beta framework is SiO2/Fe203=37, i.e. about 9.8 Fe atoms per unit cell assuming that the framework contains 192 T-atoms per unit cell. The organic template used for the synthesis of the Fe-Beta structure was TEAOH, 20% aqueous solution. In the present work we have used a 25% methanolic instead of aqueous solution of TEAOH as an organic template which resulted in the incorporation of significant quantities of iron in the Beta framework: SiO2/Fe203= 15.0 (22.6 Fe atoms per unit cell). The synthesized products were characterized by XRD, SEM, thermal analysis, nitrogen sorption measurements, XPS and acidity by thermal desorption of NH3 and IR spectroscopy.

Experimentai Synthesis

Four different procedures were adapted for the synthesis of iron-substituted Beta zeolite. TEAOH was a 40 or 20% solution in water or a 25% methanolic solution (Alfa Products). Sodium aluminate (28.4% Na20 , 46.8% A1203, 24.8% H 2 0 ) was the source of A1 and Na for the synthesis of A1-Beta zeolite. Tetraethylorthosilicate (TEOS, Alfa Products/Aldrich) and ferric sulfate (Fluka) were the sources of Si and Fe, respectively. In the case of Fe-Beta synthesis the pH of the final gel was adjusted by adding NaOH (Mallinckrodt). The starting reaction gel composition used for various preparations is presented in Table I. Procedure A

Solution A was prepared by adding TEOS to TEAOH. Solution B was ferric sulfate in water.

R.B. Borade and A. Clearfield / Microporous Mater. 2 (1994) 167-177

The ferrisilicate gel was obtained by adding, dropwise, solution A to solution B. This mixture was stirred for 0.5 h. Finally, the NaOH solution was added and the mixture was stirred at 313 K for 4-5 h. In this way ethanol formed by hydrolysis of TEOS was allowed to evaporate. Procedure B

This method was identical to procedure A except that solution B was added to A instead of adding AtoB. Procedure C

In this method, TEOS was added directly to the ferric sulfate solution in water. Then TEAOH was added dropwise followed by addition of the NaOH solution. Procedure D

In a typical preparation (run 17 of Table 1), 20.28 g of TEOS (98%) was added slowly under stirring to 13.28g of TEAOH (25% methanolic solution). The resulting mixture was slowly added to a solution of 1.59 g of ferric sulfate in 18.9 g of water under vigorous stirring. To this a solution of 0.26 g of NaOH (Mallinckrodt) in 40.0 g of TEAOH was added. The slurry was stirred at ca. 273 K for 4-5 h and then at room temperature overnight in an open vessel. This allowed the ethanol formed during hydrolysis of TEOS to evaporate. The autoclaves were filled with this gel and heated in an air oven at different temperatures for desired lengths of time. Afterwards, to stop the crystallization process, the autoclaves were quenched in cold water. The solid product was recovered by filtration and washed with deionized hot water. The products were dried overnight in an oven at 373 K. This product was designated as as-synthesized Fe-Beta. The Na form was obtained by calcination of as-synthesized Fe-Beta at 813 K for 12 h. The sample was repeatedly ion-exchanged with 2 M NH4C1 to give the NH,~ form of FeBeta. The protonated form was obtained by calcination of the N H 2 form of Fe-Beta at 723 K for ca. 8 h. The synthesis and physicochemical properties of A1-Beta zeolite used as a reference in this study are reported elsewhere [21,22].

R.B. Borade and A. Clearfield / Microporous Mater. 2 (1994) 167-177

169

TABLE 1 Starting reaction gel composition (in molar ratios) used for the synthesis of Fe-Beta Run

N a 2 0 Fe203 SiOz

TEAOH

H20 Proce.dtlre Temperature(K)

Time(h)

Produet

1 2 3 4 5 6 7 8 9 10 I1 12 13 14 15 16 17 18 19 20

1 1 1 1 1 2.7 2.7 3.9 3.9 3.9 3.9 2.0 2.0 2.0 2.0 1 1 1 1 1

27.91 27.9a 27.9~ 27.9s 27.9a 25.9b 25.9~ 37.3a 37.3a 37.3" 37.3a 37.3b 37.3b 37.3b 37.3b 27.9¢ 27.9c 27.9~ 27.9¢ 27.9c

517 517 517 517 517 965 1835 1334 1334 1334 1334 1334 1334 1334 1334 340 340 340 340 340

90 144 258 90 216 144 267 39 II1 159 220 44 140 340 376 312 288 312 312 312/144

Amorphou8 Amorphous Amorphous Amorphous Amorphous Amorphous Amorphous Arnorphous Amorphous + Beta + U « Amorphous + Beta + U ä Amorphous + Beta + U d Amorphous Amorphous + ZSM-5 Amorphous Amorphous Fc-Bcta F¢-Beta Fc-Beta Fe-Beta Amorphous

I 1 I 1 1 1 1 1 1 I 1 1 I 1 1 1 1 I 1 1

30.9 30.9 30.9 30.9 30.9 54.3 54.3 76.7 76.7 76.7 76.7 76.7 76.7 76.6 76.7 76.7 30.0 30.0 20.0 15.0

A B A B B B C D D D D B B B B D D D D D

393 393 393 393 443 443 443 433 433 433 433 433 433 433 433 393 393 393 393 393/413

a40% TEAOH aqueous solution. b20% TEAOH aqueous solution. ~25% TEAOH methanolic solution. «Unidentified.

Characterization X R D p a t t e r n s were o b t a i n e d with a SeifertScintag P A D - I I a u t o m a t e d p o w d e r d i f f r a c t o m e t e r using CuKct r a d i a t i o n . S E M p h o t o g r a p h s were t a k e n at Texas A & M E l é c t r o n M i c r o s c o p y C e n t e r using a J E O L J S M - 2 5 S-II m i c r o s c o p e . T h e s a m p l e s were m o u n t e d in a n a l u m i n u m s a m p l e h o l d e r a n d c o a t e d with a n A u - P d e v a p o r a t e d film. T G A curves were o b t a i n e d b y using a D u P o n t M o d e l 951 t h e r m a l analyzer. T h e s a m p l e (ca. 30 mg) was p l a c e d in a q u a r t z b u c k e t a n d h e a t e d at a r a t e o f 10 K / m i n in the presence o f n i t r o g e n (flow-rate ca. 60 tal/min). Surface a r e a m e a s u r e m e n t s were c a r r i e d o u t o n a n A u t o s o r b - 6 s o r p t i o n system ( Q u a n t a C h r o m e ) at liquid n i t r o g e n t e m p e r ature. T h e s a m p l e s were d e g a s s e d at 573 K a n d 1 0 - 5 T o r r for 24 h. T o t a l analysis o f zeolite s a m p l e s was c a r r i e d o u t b y dissolving the s a m p l e in h y d r o fluoric a c i d a n d t h e n d i l u t i n g with d e m i n e r a l i z e d w a t e r to the desired volume. T h e resulting s o l u t i o n was a n a l y z e d for Si, A1, F e a n d N a b y an i n d u c -

tively c o u p l e d p l a s m a (ICP) m e t h o d using a n A R L 3510 I C P spectrometer. T h e X P S s p e e t r a of all samples were r e c o r d e d at r o o m t e m p e r a t u r e using a P e r k i n E i m e r M o d e l P H I 5500 E S C A spectrometer. T h e X - r a y source used was MgK0t (hv = 1253.6 eV). P r i o r t o the X P S s t u d y all s a m p l e s were e v a c u a t e d in a p r e t r e a t m e n t c h a m b e r (P ~ 1 0 - 3 Torr) at r o o m t e m p e r a t u r e for 1 h a n d t h e n i n t r o d u c e d into the m a i n c h a m b e r for X P S m e a s u r e m e n t s . T h e details o f the experim e n t a l p r o c e d u r e a r e d e s c r i b e d elsewhere 121,22]. The I R spectra in the h y d r o x y l a n d p y r i d i n e regions were o b t a i n e d using the following p r o cedure. T h e s a m p l e s were pressed into selfs u p p o r t i n g wafers c o n t a i n i n g ca. 7 - 9 m g of m a t e rial. T h e wafers were m o u n t e d o n stainless-steel s a m p l e h o l d e r s a n d i n t r o d u c e d into a pyrex v a c u u m cell which was designed to a c c o m m o d a t e four s a m p l e s a t a time. T h e cell was c o n n e c t e d to the v a c u u m system ( P = 1 0 -~ Torr) a n d F e - B e t a s a m p l e s were d e g a s s e d at 573 K a n d AI-Beta a t 673 K for 16 h. T h e n the cell was closed, the

170

samples were allowed to cool down to room temperature, and their IR spectra were recorded. Afterwards, pyridine vapors were admitted to the cell. Then the temperature of the sample was increased to 323 K and pyridine vapor was allowed to react overnight. Finally, excess pyridine was desorbed by evacuating the Fe-Beta samples at 423 K and the AI-Beta samples at 453 K for 16 h. The samples were cooled down to room temperature and the spectra were recorded. All the spectra were recorded in the range 4000-1000 cm-1 with a 2cm -t resolution using a Digilab FTS-40 spectrometer.

Resdts and discussion Initial attempts to synthesize Fe-Beta with starting SiO2/Fe203 ratios in the range 30-76 and using either 40 or 20% aqueous TEAOH solution were unsuccessful. Some selected runs are given in Table 1. When a 25% methanolic TEAOH solution was used, the products (runs 16-19 of Table 1) were found to be white crystalline material with XRD patterns similar to that of A1-Beta zeolite. However, the starting reaction gel composition having an SiO2/Fe203 ratio of 15 produced amorphous material. The use of methanolic TEAOH solution in the synthesis of Fe-Beta probably prevents fast precipitation of hydrated iron oxide and facilitates the formation of an ironsilicate gel that leads to the incorporation of iron in tetrahedral positions under hydrothermal conditions. For clarity, hereafter products of runs 16-19 were designated as Fe-Beta/A, Fe-Beta/B, Fe-Beta/C and FeBeta/D, respectively. The XRD pattern of an as-synthesized white Fe-Beta/D sample is shown in Fig. 1. For comparison purposes the XRD pattern of A1-Beta zeolite is also presented. It can be seen that the Fe-Beta sample is much less crystalline than the aluminum counterpart. This fact is evident in the much lower intensities of the Fe-Beta reflections which results in several weak reflections being lost in the background. FeK« fluorescence by CuKct radiation, which can cause a loss of peak intensities and an increase in the background level, is evident in the pattern but loss of intensity is so great that we

R.B. Borade and A. Clearfield / Microporous Mater. 2 (1994) 167-177

conclude that the Fe-Beta is a poorly crystalline material relaüve to the aluminum sample. The SEM photographs shown in Fig. 2 reveal that at a low level of Fe substitution the morphology of Fe-Beta samples (e.g., Fe-Beta/A and FeBeta/C) is very similar to that of Al-Beta zeolite. An appreciable effect on the shape and the size of the crystals is noüced with products having an Si/Fe ratio of 7.5 (Fe-Beta/D sample). At higher levels of Fe substitution the size of the particles varied from 0.5 to 15 ltm (Fig. 2). Each particle appears to be composed of much smaller microcrystallites. Although morphologically these microcrystals appear weil formed, the XRD patterns were of poorer quality than those obtained from the A1-Beta microcrystals. Information regarding the incorporation of iron into the Beta framework can be obtained from TGA experiments. Several workers studied [21-28] the thermal desorption of TEA species from the AI-Beta framework which allowed them to conclude that there exist two types of TEA species, occluded TEAOH and TEA ÷ interacting with the zeolite framework as charge compensating cations. The former species was found to be desorbed at lower temperatures than the charged species. Fig. 3 illustrates the TGA curves and their corresponding derivative plots for as-synthesized AI-Beta and Fe-Beta samples. TGA curves show that there are four distinct stages of weight losses which are more clearly evident in the derivative plots. These weight losses are reported in Table 2. Between room temperature and ,-323 K most of the adsorbed water is released from the zeolite pores. The amount desorbed decreases as the Fe content increases. The weight loss between 323 and 633 K can be ascribed to desorption or decomposition of occluded TEAOH species and that in the temperature range 633-823 K is due to pyrolysis of the TEA ÷ cations interacting with the zeolite framework. The weight loss related to occluded TEAOH species decreases and that of TEA ÷ species increases with increase of Fe content in the solids. The weight loss above 823 K is caused by decomposition of residual organic material occluded inside the zeolite pores. This point was confirmed by earrying out the TGA under an

171

R.B. Borade and A. Clearfield / Microporous Mater. 2 (1994) 167-177

2S~.C

AI-Beta 2153.8 1723.1 1292.3

430. g~ 0.0 o 775.0

Fe-Beta/D 646.2 516.9 387.7 ~

~

258.5 129.2 0.0

10

15

20

25

30

35

40

20(deg) Fig. 1. X-Ray powder diffractionof AI-Beta(Si/A1= 10) and Fe-Beta/D.

oxygen atmosphere. Under these conditions, the total weight loss due to organic amine was more or less the same but the difference was that peaks due to TEAOH and TEA + species were merged together and no significant weight loss above 823 K was observed. The total amount of organic material present in the zeolite pores has been calculated from the TGA results (Table 3). For the sake of comparison, we have assumed that the unit cell of Beta contains 192 T-atoms. The total number of TEA molecules per unit ceU corresponds to 13.9 + 0.1 and 15.2-16.4 for AI-Beta and Fe-Beta samples, respectively. The higher values of TEA molecules per unit cell observed for Fe-Beta samples could be due to the change in the relative

orientation of adjacent tetrahedra that modify the effective pore volume when Fe replaces AI. Now, assuming that the third step in the TGA curve is due to the pyrolysis of TEA + interacting with the zeolite framework as a charge-compensating cation, then based on the electronegativity concept, orte would expect that the interaction of TEA + caüons with the Fe-Beta framework would be weaker than in AI-Beta framework. The derivative plots shown in Fig. 3 clearly show that the peak maximum due to pyrolysis of TEA + cations occurs at 693 K for Fe-Beta samples as compared to 733 K for AI-Beta zeolite. This shift in the peak maximum towards lower temperature indicates the presence of negatively charged FeO4 tetrahedra in

172

R.B. Borade and A. Clearfield / Microporous Mater. 2 (1994) 167-177

~~

~,.~i ¸

2!'

Fig. 2. Scanning electron micrographs of as-synthesized A1-Beta (a), Fe-Beta/A (b), Fe-Beta/C (c) and Fe-Beta/D (d).

the zeolite framework. Such a shift in the peak maximum towards lower temperature was also observed in the oase of Ga-substituted Beta zeolites [28]. Fig. 4 illustrates nitrogen adsorption isotherms for the Na forms of Fe-Beta/D and AI-Beta samples obtained at 77 K. Both samples show characteristic type I isotherms which reveals an essentially microporous nature of these materials. The nitrogen BET surface area was found to be somewhat higher for Fe-Beta samples (610 ± 22 m 2 g-1) as compared to A1-Beta .zeolites [21,22-1 (500±20m 2 g-l). However, at a relative pressure P/Po=O.4, adsorbed amounts are almost the same indicating similar micropore volumes for the two materials. The presence of structural defeets or impurities can be differentiated by studying nitrogen desorption

behavior. The presence ofa hysteresis loop between the adsorption and desorption branches of A1-Beta isotherms indicates the presence of impurities which gives rise to desorption in the mesopore region (15-100 A diameter) which did not appear in the isotherm of Fe-Beta sample. The Fe2p XP spectra of as-synthesized and calcined Fe-Beta/D samples are shown in Fig. 5. Both spectra show all the characteristics of Fe 3 + at 710.5 eV and the Fe2p3/2-Fe2pl/2 splitting of 13.7eV. The presence of extraframework iron oxidic species in zeolites is weil characterized by Ols XP spectra of these materials [10,29,30]. The Ols spectra of our Fe-Beta material in various forms is presented in Fig. 6. The main peak at 532.5eV corresponds to the oxygen species involved in the formation of either S i - O - F e or

R.B. Borade and A. Clearfield / Microporous Mater. 2 (1994) 167-177

')

ITI

i

~

o

a

..~

..~ «

o

473

673

s73

~073

~27»

Temperature (K) Fig. 3. Thermat and derivative plots of as-synthesized zeolites: (a) Al-Beta; (b) Fe-Beta/A; (c) Fe-Beta/C; (d) Fe-Beta/D.

Si-O-Si linkages of the zeolitic framework. Even though as-synthesized Fe-Beta material is white, the presence of a shoulder peak at a low binding energy value at 530.5 eV is a clear indication of the presence of extra framework iron oxidic species (Table 4). This low-binding-energy peak was also present in the Fe-Beta/A, Fe-Beta/B and FeBeta/D samples. The contribution due to this lowbinding-energy Ols peak to the main O1~ peak was of the same extent in each sample. This led us to conclude that these extraframework iron oxidic

173

species are formed during the mixing and/er hydrothermal preparation of Fe-Beta samples, as a separate entity. Upon calcination of as-synthesized Fe-Beta/D at 813 K, no increase in the relative contribution due to the low-binding-energy p¢ak (extra framework iren oxidic species) to the main Ots peak was observed suggesting that the iren species associated with the zeolitic framework are stable under these conditions. This underlines the high thermal stability of the Fe-B¢ta framework. The acidity of Fe-Beta and A1-Beta zeolites was investigated via TGA on NH~-exchanged samples. The thermoanalytical curves and their derivative plots are shown in Fig. 7. All samples gare similar thermal profiles but different weight losses at •ach step. The first step between room temperature and c a . 413 K observed for Fe-Beta and up to 473 K for Al-Beta is clearly due to the removal o f adsorbed watet since, after N H 2 exchange, samples were dried at c a . 393 K and then cooled to room teml~rature (see Exp¢rimental). The amount of water desorb¢d from Fe-Beta samples was found to decrease as the Fe content increased. The weight loss between 423 and 573 K (Fig. 7b) may correspend to either desorption of NH3 from Lewis acid sites er decomposition of NH~ ions from relatively weak Brönsted acid fites er a combination of these two processes which is dit~cult to distinguish by the thermal desorption study alone. The third step, 573-823 K, is clearly due to the decomposition of NH~ ions from relatively strong Brönsted acid sites. AI-Beta zeolite showed a continuous weight loss in the 573-823 K region with a peak maximum at 650K whereas Fe-Beta showed a two-step weight loss. The assignment t h a t the weight loss in the region 413-823 K is due

TABLE 2 Thermoanalytical data for various Fe-Beta samples in comparison with A1-B¢ta Sample

Fe-B¢ta/B Fe-B¢ta/C Fe-Beta/D AI-B¢ta

Weight loss (%) I

II

III

IV

(298-323 K)

(323-633 K)

(633-823 K)

(823-1273 K)

1.80 1.80 3.20 3.96

5.60 2.18 3.79 2.65

8.23 9.18 10.60 10.91

3.63 4.74 2.85 1.64

174

R.B. Borade and A. Clearfield / Microporous Mater. 2 (1994) 167-177

TABLE 3 Physicochem/cal propertics of Fe-Beta samples

Sample

Si/M*

Organic loss (wt %)

TEA (per unit cell)

Surfaceareab (BET nitrogen)(m2/g)

Fe-Beta/A Fe-Beta/B Fe-Beta/C Fe-Beta/D AI-Beta*

30.2 13.7 12.2 7.6 10.0

16.70 17.46 16.10 17.20 15.20

15.7 16.4 15.2 16.2 14.0

586 589 631 589 518

*Si/M ratio determined by chemical analysis in their protonated form. M stands for F¢ er AI atom. bAft¢r calcination at 823 K. «From rel. 16. TABLE 4 XPS data of Fe-Beta/D sampl¢ in its various forms Sample

(Si/M)xrs"

Si2p

M2pa

Ols

AI-Beta H-Fe-Beta NH,-Fe-Beta Na-Fe-Beta As-made Beta

10.5 6.8 7.5 7.2 7.1

103.3 103.3 103.3 103.3 103.3

74.5 710.6 710.7 710.6 710.6

532.1 532.3 (529.6)b 532.6 (530.0) 532.5 (530.3) 532.6 (530.4)

"M stands for Al er Fe. bShoulder to the main Ots p¢ak.

293

AbBeta

261 228

196

,

,

I

1

0.3

0.4

)

)

t

Fe-Be/a/D 228i

131

0.0

0.1

0.2

0,5 0.6

0.7

0.8

0.9 LO

(P/Po) Fig. 4, Nitrogcn adsorption isotherms of A1-Beta and FeB¢ta/D in th¢ir ,Na form.

to the decomposifion of NH2 ions is justifi~ by IR stmctroscopy. At the activation temp¢ratur¢ of 673 K, the spectra show an absence of N - H stretching bands (see b¢low) and the presence of

an OH band at 3610 cm -1 for A1-Beta zeolites [21,22] and near 3680 cm-1 for Fe-Beta samples. The weight loss above 823 K is due to the dehydroxylation of the zeolite and is confirmed by carrying out TGA experiments with the Na and H forms of a Fe-Beta sample that was calcined at 673 K to decompose NH~- ions. Only the H form of the Fe-Beta/D sample showed a weight loss above 823 K signifying that the peak above 823 K is due to the dehydroxylaüon of zeolites. The most important observation of this study is that NH + ions associated with the Fe-Beta are decomposed (to NH3 and leaving behind protons attacbed to the framework) at a lower temperature than these in A1-Beta zeolite, suggesüng that the strength of Brönsted acid sites in Fe-Beta is lower than these in Al-Beta zeolite. The acidic nature of Fe-Beta molecular sieves can be assessed by studying the IR spectra of chemisorbed pyddine. The IR spectra in the hydroxyl stretching region before and after pyridine adsorption are shown in Fig. 8. It is important to remember that the Fe-Beta sample was activated at 573 K and AI-Beta at 673 K for 16 h before recording IR spectra at room temperature, The Fe-Beta samples were pretreated at lower temperature than that of AI-Beta mainly because Fe-Beta samples are thermally less stable and undergo dehydroxylation at 400°C [22]. Before pyddine adsorption, both Fe-Beta and A1-Beta samples show a hydroxyl band at 3740 cm -1 associated with terminal silinol groups (Fig. 8A and C). This band remains unaffected after pyridine adsorption. The sharp band at 3610cm -~ in AI-Beta and a very broad and poorly defined band at c a .

R.B. Borade and A. Clearfield / Microporous Mater. 2 (1994) 167-177

F e 2 P I/2

13,7 eV

* i

)

] ~ ~ P 3/2

,

"~_ 735

730

725

o,,/~

175

720

715

710

705

_

h

700

695

Binding Energy (er)

Fig. 5. Fe•p XP spectra of (a) as-synthesized Fe-Beta/D and ~ ) the Na form of Fe-Beta/D.

JL

\

~

a

673

873

ra b A

!

473

1073

1273

Temperature (K)

s~o

s3s

530

525

520

Binding Energy (eV) Fig. 6. Ol, XP spectra of Fe-Beta/D sample: (a) as-synthesized; (b) Na form; (c) NH + form; (d) H form.

3670 cm -1 in Fe-Beta are affected by pyridine adsorption suggesting that these OH groups are involved in the protonation of the pyridine moleeule. In other words these are the structural Brönsted acid sites in AI-Beta and Fe-Beta samples.

Fig. 7. Thermal decomposition of NH,~ ions studied by thermo8ravimetry. TGA curves and derivative plots (inset) for temperatures above 423 K are shown. (a) AI-Beta; (b) NH4-Fe-Beta/D; (c) H-Fe-Beta/D; (d) Na-Fe-Beta~.

To check the validity of the band at c a . 3670 cm- 1, the fresh Fe-Beta sample was activated at 573 and 673 K under vacuum and the IR spectra wert reeorded. The IR spectrum obtained at 573 K was very similar to spectrum a of Fig. 8. But the IR spectrum obtained at 673 K showed a reduction in intensity of the band at c a . 3670 cm-K Wc assign this reduction in intcnsity of the 3670 cm- i band to the dehydroxylaüon process wherein

176

R.B. Borade and A. Clearfield / Microporous Mater. 2 (1994) 167-177

A

A

B

1490

3670 1

I

1560

L

1520

I

L

I

1480

I

1440

W a Y e n u m b e r s , g m -1 Fig. 9. IR spectra in the region 1575-1425 cm -1 after pyridine adsorption. (A) Fe-Beta/D; activation temperature 573 K and pyridinc desorption temperature 423 K. (B) Al-Beta; activation temperature 673 K and pyridine desorption temperature 473 K.

3610

3740 I

4000

3800

I

3600

t

3400

Wavenumbers, cm« Fig. 8. IR spectra in the hydroxyl region. Fe-Beta/D: (A) before and (B) after pyridine adsorption; (C) activated at 673 K for 5 h and before pyridine adsorption. AI-Beta: (D) before and (E) after pyridine adsorption.

Brönsted acid sites are destroyed and Lewis acid sites are created. This underlines the observation that Fe-Beta samples undergo dehydroxylation at relatively low temperatures as compared to A1Beta zeolites and the shift in the hydroxyl stretching frequency towards higher wave number indicates a lower acid strength of OH groups in Fe-Beta than similar groups in AI-Beta zeolites. This is consistent with the published results of Fesubstituted materials and their corresponding A1 analogs [10,29,31]. The very broad band in the region 3600-3400 cm -1 is generally assigned to hydrogen-bonded hydroxyl groups. The IR spectra in the 1575-1425 cm -1 region show three bands at 1545, 1490 and 1445cm -1 (Fig. 9). The bands at 1545 and 1445 cm -1 are due to the formation of pyridinium ions with Brönsted acid sites and pyridine associated with Lewis acid sites, respectively. The band at 1490 cm- ~ is due to both pyridine interacting with both Brönsted

and Lewis acid sites. The presence of the band at 3670 cm-1 before pyridine adsorption and the appearance of a band at 1545 cm- 1 after pyridine adsorption in Fe-Beta is strong evidence for the isomorphous substitution of iron in the Beta framework. Fig. 9 reveals that Fe-Beta contains a lesser number of Brönsted and more Lewis acid sites. The situation is reverse for AI-Beta zeolites. The shift in the frequency of the pyridine Lewis band towards a lower value upon iron substitution again suggests that the nature of Lewis acid sites in both zeolites is different. ca.

Conclusions Ferrisilicate molecular sieves, containing significant quantities of iron in the Beta-type framework, have been synthesized and characterized. A modified synthesis procedure aUowed us to obtain FeBeta material with Si/Fe ratios as low as 7.5. SEM and a nitrogen adsorption isotherrn study did not show any indi¢ation of the presence of extra framework iron oxidic material on the surface or inside the pores of the Fe-Beta sample. Even though FeBeta samples (Si/Fe = 7.5-30) ,vere white in color the Ols XP spectra indicated the presence of small

177

R.B. Borade and A. Clearfield / Microporous Mater. 2 (1994) 167-177

amounts of extra framework iron oxidic material on the surface of these materials. Thermal desorption studies of NH~-exchanged Fe-Beta samples and IR spectra in the OH stretehing region indicated that the strength of Brönsted acid sites in Fe-Beta materials is lower than those in A1-Beta zeolites.

Acknowledgement We acknowledge with thanks financial support for this study by Regents of Texas A&M University under the Materials Science and Engineering Program.

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