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Amorphous Metal Fluorides with Extraordinary High Surface Areas.

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Angewandte
Chemie
High-Surface-Area Aluminum Fluoride
Amorphous Metal Fluorides with Extraordinary
High Surface Areas**
Erhard Kemnitz,* Udo Groß, Stephan Rdiger, and
Chandra S. Shekar
Inorganic materials having high surface areas are of interest,
for example, in heterogeneous catalysis where the activity of a
catalyst depends largely on its surface.[1] Metal oxides with
very high surface areas can be prepared by sol–gel routes, by
freeze-drying, or under supercritical conditions as xerogels or
aerogels.[2] However, no such synthetic routes exist to date for
metal fluorides, which because of their high lattice energy
tend to form regular crystalline structures. Thus, aluminum
trifluoride prepared conventionally by precipitation in aqueous media exhibits surface areas in the range from 10 to
60 m2 g1 only.[3] In spite of this, AlF3 is widely used as
catalyst,[4] and other metal fluorides, such as magnesium
fluoride, are used as supports for catalytically active materials.[5] By a sophisticated NF3 plasma-etching treatment of
aluminum-rich zeolites it could be shown that AlF3 with
surface areas of up to 190 m2 g1 is in principal accessible.[6]
Herein we report a novel facile non-aqueous synthesis
route to X-ray amorphous metal fluorides that have high
surface areas. First a solid, fluoridated metal compound is
prepared, which is structurally highly distorted because of the
presence of bulky groups in stoicheiometrical quantities. This
precursor is then fully fluorinated under soft conditions, that
is, below crystallization temperature to preserve the distortion.
As in case of the sol–gel syntheses of metal oxides where
the final removal of the alcohol from the metal–alkoxide
precursor (aerogels vs. xerogels) is crucial for the development of their specific morphological properties, for highsurface-area metal fluorides the final fluorinating treatment
of the partially fluorinated precursor is the determining step
of the synthesis. For metal oxides, high-surface materials have
unusual properties, particularly when it comes to catalysis.
The same is true for metal fluorides. Thus, aluminum fluoride
prepared as described below, has not only a very high surface
area but is also an extremely strong Lewis acid, comparable
with some of the strongest known Lewis acids SbF5[7] and
aluminum chlorofluoride (ACF).[8]
The route, which will be detailed for AlF3, consists of two
consecutive steps. First, aluminum triisopropoxide
(Al(OiPr)3) was treated in water-free isopropanol with a
moderate excess of anhydrous hydrogen fluoride dissolved in
[*] Prof. Dr. E. Kemnitz, Dr. U. Groß, Dr. S. R#diger, Dr. C. S. Shekar
Institute of Chemistry
Humboldt University
Brook-Taylor-Strasse 2, 12489 Berlin (Germany)
Fax: (+ 49) 30-2093-7277
E-mail: erhard.kemnitz@chemie.hu-berlin.de
[**] We would like to thank Dr. M.-M. Pohl, ACA Berlin, for recording the
TEM micrographs.
Angew. Chem. Int. Ed. 2003, 42, 4251 –4254
isopropanol or diethyl ether. After removing all volatile
material at elevated temperature in vacuum a white powdery
solid remained; the precursor of the target fluoride. The
precursor had a ratio of fluorine to aluminum of about 2:1,
and a carbon content of about 30 %. Thus, despite having
treated the Al(OiPr)3 with an excess of HF, there was
obviously some isopropylate together with some coordinatively bound isopropanol in the compound (Scheme 1).
Scheme 1. Route used to prepare HS-AlF3.
The second step comprises treatment of the precursor
with a fluorinating agent, such as CCl2F2 at 350 8C, to remove
the organic component by replacing it fully by fluorine. The
final product obtained had a ratio of fluorine to aluminum of
3:1 and the carbon content was less than 0.5 %. There is also a
small amount of chlorine present, which becomes zero if
CHF3 is used instead of CCl2F2 (analysis (%) calcd for AlF3 :
Al 32.13, F 67.87; found: Al 31.2, F 67.2, C 0.4, Cl 0.3).
Strikingly, the aluminum fluoride prepared this way has a
surface area of about 200 m2 g1 and showed a very high Lewis
acidity and corresponding catalytic activity. Because of its
high specific surface area we denote it HS-AlF3. In a similar
way, other metal fluorides with very high surface areas can
also be prepared. Thus, for example, magnesium fluoride
prepared accordingly had a BET(N2) surface area of
190 m2 g1.
In case of aluminum fluoride, both the precursor and the
final product, HS-AlF3, were X-ray amorphous, but thermal
analysis of both compounds revealed the exothermic formation of a crystalline phase at 540–570 8C identified by X-ray
powder diffraction as a-AlF3. The presence of iPrO units in
the precursor was confirmed by (in addition to carbon
analysis) the IR spectra of the solid, and also by 1H NMR
spectrum of a suspension in CD3CN. Attempts to remove the
organic from the precursor with supercritical CO2 failed;
more than two thirds of the organic remained (Scheme 1).
Heating the precursor to 350 8C in a nitrogen flow gave an
amorphous product with only 1 % carbon content and a BET
surface area of 47 m2 g1, but which was not catalytically
active. These results together with the elemental analysis
indicate that the precursor has a composition corresponding
to the formula AlF3x(OiPr)x·(iPrOH)y with x 1 and y < 1.
Only the loosely bound alcohol can be removed by super-
DOI: 10.1002/anie.200351278
2003 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
4251
Communications
critical CO2, whereas the alkoxy groups can be gently
removed by a soft fluorination under preservation of the
high morphological disorder. Hence, the presence of alkoxy
groups in the precursor is decisive for obtaining high-surfacearea metal fluorides.
The BET(N2) surface properties were compared for HSAlF3, its precursor, and b-AlF3, a common AlF3-type, which is
often used as a Lewis acid catalyst in fluorine chemistry. The
specific surface area (SBET) of b-AlF3 is only 33 m2 g1,
treatment with CCl2F2 at 350 8C did not improve that value.
The HS-AlF3 has a BET surface area of 206 m2 g1 (BJH
adsorption cumulative surface area 251 m2 g1, BJH desorption cumulative surface area 281 m2 g1; all data obtained with
N2 using an ASAP 2010, Micrometrics) and the values for the
precursor are even higher (Table 1). In addition, HS-AlF3 had
a much narrower pore size distribution than b-AlF3
(Figure 1).
Table 1: Properties of selected aluminum fluorides.
Sample
Crystallinity
SBET [m2 g1]
F:Al
precursor
HS-AlF3
b-AlF3
AlF3[b]
amorphous
amorphous
crystalline
crystalline
430
206
33
113
2:1
3:1
3:1
3:1
C Content [%]
31
0.4
n.a.[a]
0.15
Figure 2. SEM micrographs of HS-AlF3, view: 2 F 2 mm2 (upper) and
100 F 100 mm2 (lower section).
[a] n.a. = not applicable; [b] Obtained conventionally by an aqueous
route; treated with CCl2F2 at 350 8C.
Figure 1. HS-AlF3 adsorption/desorption isotherms and pore size
distribution, obtained by N2 adsorption technique. * adsorption,
* desorption; p/p0 = relative pressure, Vads = adsorbed volume,
D = pore diameter, dV/dlgD = pore volume.
Figure 3. TEM micrograph of HS-AlF3.
SEM micrographs (Figure 2) show the amorphous HSAlF3 as well as its precursor to be composed of particles,
which according to BET results are mesoporous, and which
are probably agglomerates of much smaller particles of
irregular shape. In TEM micrographs well-defined domains
of about 10 nm diameter can be seen (Figure 3), some of
which are regular, that is, lattice planes could be detected in
some regions. However these areas are too small to be
detected by X-ray diffraction. For both the precursor and HSAlF3, interlattice-plane separations of approximately 3.5 B
were estimated, which are typical for aluminum fluorides.
4252
2003 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
Energy-dispersive X-ray spectroscopy (EDX) confirms
the elemental analysis of the HS-AlF3 (after the fluorination/
activation treatment) that it consists of Al, F, and some O. The
oxygen content might be a result of surface hydration when
the solid was handled in open atmosphere. IR- and 27Al-magic
angle spinning (MAS)-NMR-spectra of the solid HS-AlF3
were recorded (Figure 4). The 27Al-MAS-NMR-spectra were
collected at 132.2 MHz at a spinning rate of 10 kHz and allow
the short-range order at Al atoms to be described and enable
discrimination of different Al species. The chemical shift of
www.angewandte.org
Angew. Chem. Int. Ed. 2003, 42, 4251 –4254
Angewandte
Chemie
ence of catalytic CCl2F2 dismutation [Eq. (1)] was studied in a
cat:
5 CCl2 F2 ƒ! CCl3 F þ 3 CClF3 þ CCl4
Figure 4.
27
Al-MAS-NMR spectrum of HS-AlF3.
ð1Þ
flow reactor under conditions that equilibrium is reached (ca.
0.5 g of catalyst in a Ni tube of 0.5 cm i.d., CCl2F2 flow 2
mL min1, contact time ca. 2 s). Upon varying the contact time
no change in product composition could be observed above
0.5 sec.
Further, the catalytic isomerization of 1,2-dibromohexafluoropropane to 2,2-dibromohexafluoropropane was studied.[10] According to the data in Table 2, the dismutation of
CCl2F2 proceeds over HS-AlF3 even at very low temperatures
with much higher conversion rates than over the amorphous
AlF3 from the aqueous synthesis route, which also has a high
Table 2: Catalytic activity of selected aluminum fluorides.
d = 13.8 ppm is indicative to an octahedral
CCl2F2 Dismutation [%]
C3Br2F6 Isomerization [%]
environment around the Al sites of a poly- Sample
morphous solid. Compared with crystalline T [8C]
300
250
200
150
100
25
0
0
0
0
aluminum fluoride, for example, a-AlF3, the precursor
96
96
87
59
22
> 90
resonance is shifted down field, that is, it is HS-AlF3
93
47
0
indicative of a more coordinatively unsatu- b-AlF[a]3
96
90
55
20
8
5
AlF3
rated species.
[a]
Obtained
conventionally
by
an
aqueous
route;
treated
with
CCl
F
at
350
8C.
2 2
The large sideband manifold of the
satellite transitions of the quadrupolar aluminum nucleus (Figure 4) indicates large
quadrupole coupling constants (QCCs) for aluminum, consurface area compared to other metal fluorides. This differsistent with a stronger distortion of the octahedral environence in reactivity is clearly founded in the presence of
ments. Clear discontinuities because of first-order QCCs were
substantially more acidic sites in HS-AlF3, as demonstrated by
not found because of the wide distribution of the QCCs.
the NH3-TPD curves shown in Figure 5. Although their
Sideband simulations indicate QCCs of about 1.5 MHz which
profile is rather similar, the NH3 desorption from the
is much larger than that of the highly symmetric a-AlF3, but
comparable with the estimated 1.4 MHz of the nearly
amorphous material.[6]
The IR spectra of HS-AlF3 support these findings. In the
region from ~n = 700–4000 cm1, there should be no absorption
bands for AlF3. In fact, only one band at 667 cm1 (Al-F
valence vibration n3) could be seen, which was very broad
owing to the amorphous state of the compound. Laser-Raman
analysis was not successful because of fluorescence phenomena, clearly a result of the carbon present. The chemical
properties of the surface of HS-AlF3 are demonstrated by its
pyridine absorption, by temperature programmed desorption
of NH3 (NH3-TPD), and more importantly by its catalytic
properties. After treatment with pyridine and subsequent
heating in an N2 stream at 150 8C the photoacoustic IR spectra
of the solid showed intense bands at ~n = 1454 and 1492 cm1
Figure 5. Temperature programmed desorption of NH3 (NH3-TPD)
which indicate Lewis acid sites. The NH3 TPD profile
from HS-AlF3 (a) and amorphous AlF3 (c).
resembles that of b-AlF3, which shows two distinct maxima
at about 150 and 375 8C, and the end of the desorption at
about 500 8C.[9] HS-AlF3, however, shows maxima at about
amorphous AlF3 is completed at around 500 8C, as found for
200 and 450 8C and desorption is complete at about 600 8C
b-AlF3 (see above)[9] whereas with HS-AlF3 completion is
which gives evidence for much stronger Lewis acid sites than
around 600 8C. The heterogeneity of the acidic strength of the
in b-AlF3. The catalytic activity was compared to the
surface sites is the reason for the different temperature
dependencies of this reaction in which the energy of
precursor, to b-AlF3, and to an amorphous AlF3 prepared
activation for HS-AlF3 is apparently higher than for amorconventionally by an aqueous route from AlF3·3 H2O desiccated at 200 8C in vacuum followed by subsequent CCl2F2
phous AlF3. The extraordinarily strongly acidic sites in HSAlF3 are clearly those at which the isomerization of 1,2treatment at 350 8C. As test reaction the temperature dependAngew. Chem. Int. Ed. 2003, 42, 4251 –4254
www.angewandte.org
2003 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
4253
Communications
dibromohexafluoropropane, which is only catalyzed by the
strongest known Lewis acids (SbF5 and ACF), is activated.
Consequently, with amorphous AlF3 nearly no conversion
occurs.
However, and in striking contrast to ACF, HS-AlF3 is not
destroyed by contact with moisture, it completely regains its
original catalytic activity after removing the adsorbed water
at higher temperatures. Thus, it is suited for industrial
application.
To conclude, the new two-step fluorination method is a
powerful tool with which to prepare metal fluorides that have
extraordinarily high specific surface areas and, depending on
the metal, exhibit very high Lewis acidity.
Received: February 25, 2003
Revised: June 25, 2003 [Z51278]
.
Keywords: aluminum fluoride · heterogeneous catalysis ·
materials science · super-acidic systems
[1] M. Boudart in Handbook of Heterogeneous Catalysis (Eds.: G.
Ertl, H. KnIzinger, J. Weitkamp), VCH, Weinheim, 1997.
[2] J. B. Miller, E. I. Ko in Advanced Catalysts and Nanostructered
Materials (Ed.: W. R. Moser), Academic Press, San Diego, 1996.
[3] E. Kemnitz, D.-H. Menz, Prog. Solid State Chem. 1998, 26, 97.
[4] L. E. Manzer, V. N. M. Rao, Adv. Catal. 1993, 39, 329.
[5] M. Wojciechowska, J. Goslar, W. Kania, M. Pietrowski, J.
Fluorine Chem. 1998, 91, 141.
[6] J. L. Delattre, P. J. Chupas, C. P. Grey, A. M. Stacey, J. Am.
Chem. Soc. 2001, 123, 5364.
[7] K. O. Christe, D. A. Dixon, D. McLemore, W. W. Wilson, J. A
Sheehy, J. A. Boatz, J. Fluorine Chem. 2000, 101, 151.
[8] V. A. Petrov, C. G. Krespan, B. E. Smart, J. Fluorine Chem. 1996,
77, 139.
[9] A. Hess, E. Kemnitz, J. Catal. 1994, 149, 449.
[10] V. A. Petrov, C. G. Krespan, B. E. Smart, J. Fluorine Chem. 1998,
89, 125.
4254
2003 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
www.angewandte.org
Angew. Chem. Int. Ed. 2003, 42, 4251 –4254
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