Application of calcined Mg-Al hydrotalcites for Michael additions: An investigation of catalytic activity and acid-base properties

Article abstract of DOI:10.1016/j.jcat.2005.06.004

Mg-Al hydrotalcite catalysts with different Mg/Al molar ratios (0.6, 1.4, 2.2, 3.0) were calcined and tested in liquid-phase Michael additions at room temperature. The catalysts and pure oxides (MgO and Al₂O ₃) were characterized by

Full text of DOI:10.1016/j.jcat.2005.06.004

Journal of Catalysis 234 (2005) 119–130
www.elsevier.com/locate/jcat
Application of calcined Mg–Al hydrotalcites for Michael additions:
an investigation of catalytic activity and acid–base properties
Hillary A. Prescott a, Zhi-Jian Li a, Erhard Kemnitz a,∗, Annette Trunschke b, Jens Deutsch b,
Heiner Lieske b, Aline Auroux c
a
Humboldt-Universität zu Berlin, Institut für Chemie, Brook-Taylor-Strasse 2, D-12489 Berlin, Germany
b
Institut für Angewandte Chemie Berlin-Adlershof, Richard-Willstätter-Strasse 12, D-12489 Berlin, Germany
c
Institut de Recherches sur la Catalyse, UPR CNRS 5401, 2, Av. A. Einstein, 69626 Villeurbanne cedex, France
Received 11 April 2005; revised 3 June 2005; accepted 6 June 2005
Available online 14 July 2005
Abstract
Mg–Al hydrotalcite catalysts with different Mg/Al molar ratios (0.6, 1.4, 2.2, 3.0) were calcined and tested in liquid-phase Michael addi-
tions at room temperature. The catalysts and pure oxides (MgO and Al O ) were characterized by XRD, XPS, N adsorption experiments,
2
3
2
27
Al MAS NMR, and TG-DTA. The acid–base properties of the samples were investigated by (a) FTIR spectroscopy after pyridine and CO
2
adsorption and (b) microcalorimetry with CO and benzoic acid. Both acid sites (Lewis) and base sites (Lewis and Brønsted) are present on
2
the surface of calcined hydrotalcites. The calcined hydrotalcites (CHT) catalyzed the Michael additions of CH-acid compounds to methyl
vinyl ketone to give high product yields with 100% selectivity; the Al-rich sample exhibited the highest activity.
2005 Elsevier Inc. All rights reserved.
Keywords: Hydrotalcite; Mg–Al mixed oxide; Acid–base properties; Solid base catalyst; Michael addition
1. Introduction
Choudary et al. [9] found that a rehydrated Mg–Al hydro-
talcite was an efficient, very selective catalyst for Michael
additions; the untreated and calcined hydrotalcites were in-
active.
Base-catalyzed C–C bond coupling reactions, such as the
Michael addition, are used for the production of pharmaceu-
ticals and fine chemicals. Soluble bases, such as alkali hy-
droxides, alkali alcoholates, or amines, conventionally cat-
alyze these liquid-phase reactions, yet environmental and
economical concerns are driving the replacement of soluble
bases with solid bases [1,2].
Solid base catalyst systems, such as Ba(OH)2 [3], MgO
[4,5], KF/Al2O3 [6], and Na/NaOH/Al2O3 [7], have been
used in various Michael additions. A highly active, selec-
tive base catalyst for the Michael addition of 2-methylcyclo-
hexane-1,3-dione to methyl vinyl ketone was potassium-
modified ZrO2, yet the leaching of its potassium species
made it unsuitable as a heterogeneous catalyst [8]. Recently,
Hydrotalcites are commercially available and cheap solid
bases; hydrotalcites and calcined hydrotalcites (normally re-
ferred to as mixed oxides) are highly active, selective cat-
alysts and play an important role in many base-catalyzed
reactions, such as Claisen–Schmidt condensations [10] and
Knoevenagel condensations [11]. The synthesis, the tex-
tural and acid–base properties, and the catalytic application
of hydrotalcites and calcined hydrotalcites were reviewed
in [12–15]. Their surface base sites were characterized by
temperature-programmed desorption (TPD) after adsorption
of CO2 [16,17], FTIR spectroscopy with CO2 [18–20], and
gas-phase microcalorimetry (with CO2 and SO2) [18,21–
23]. In addition to base sites, acid sites or acid–base pairs
on these materials also influence the catalytic performance.
Acid–base sites on mixed oxides are highly active sites for
*
Corresponding author. Tel: +49-30-2093-7277.
E-mail address: erhard.kemnitz@chemie.hu-berlin.de (E. Kemnitz).
0021-9517/$ – see front matter
doi:10.1016/j.jcat.2005.06.004
2005 Elsevier Inc. All rights reserved.

120
H.A. Prescott et al. / Journal of Catalysis 234 (2005) 119–130
many reactions: Meerwein–Ponndorf–Verley reactions [24],
cycloadditions of carbon dioxide to epoxides [25], and aldol
condensations to form nonenal [26]. Acid–base properties
of Mg–Al mixed oxides are governed by the Mg/Al molar
ratio [17–19], calcination temperature [21], and preparation
condition [16,20,27,28]. Although the acid–base properties
are generally known to be important for catalytic activity
and selectivity, the intricate relationship between catalytic
behavior in Michael additions, acid–base properties, and the
composition of hydrotalcite and calcined hydrotalcites is still
poorly understood.
Thus, the aim of this work was (a) to find an efficient, se-
lective catalyst that is easily acquired and cheap and whose
properties are reproducible and reliable for large-scale tech-
nical Michael additions and (b) to study the influence of
its acid–base properties and chemical composition on the
catalytic performance of the calcined hydrotalcites. The
Michael addition of 1,3-diones with different pKa values
to methyl vinyl ketone was examined on calcined com-
mercial Mg–Al hydrotalcites including an Al-rich sample.
Acid–base properties of the catalysts were investigated by
gaseous probe molecules with FTIR spectroscopy and mi-
crocalorimetry. Microcalorimetric measurements were car-
ried out with benzoic acid to study catalyst basicity under
liquid-phase conditions similar to those of the liquid-phase
test reactions.
samples were degassed at 100 ◦C until the static vacuum
was constant. The specific surface area was calculated by
the BET method. Carbon contents of hydrotalcites were de-
termined by elemental analysis (Leco CHNS-932 element
analyzer). Bulk and surface Mg/Al molar ratios were deter-
mined by ion-coupled plasma optical emission spectroscopy
(ICP-OES) (Unicam 701) and X-ray photoelectron spec-
troscopy (XPS) (VG Escalab 220 iXL spectrometer with a
Mg-Kα source and a monochromated Al-Kα source), respec-
tively.
2.2.2. 27Al MAS NMR
We performed 27Al MAS NMR experiments by accumu-
lating 64 scans on a Bruker Avance 400 spectrometer at a
resonance frequency of 104.3 MHz with an excitation of π/6
pulses and a repetition time of 5 s. A commercial Bruker
4-mm probe was used to perform the MAS experiments with
a spinning rate of 10 kHz. The reference for the chemical
shifts was an aqueous 1 M solution of aluminum chloride.
2.2.3. Thermal analysis
Thermogravimetry (TG) and differential thermal analy-
sis (DTA) were performed with a Netzsch STA409C system
equipped with a skimmer-coupled mass spectrometer in air
at a heating rate of 10 ◦C/min from room temperature to
700 ◦C (reference: Al2O3).
2.2.4. FTIR spectroscopy
2. Experimental
FTIR measurements were carried out on self-supporting
wafers (10–40 mg) in a transmission IR quartz cell with
CaF2 window. The samples were pretreated at 550 ◦C in
flowing synthetic air for 1 h and then under vacuum for
30 min. The adsorption of pyridine (Merck; freshly distilled
and stored over zeolite A) was then performed at 40 ◦C with
0–15 mbar pyridine and subsequent evacuation for 10 min
to remove physisorbed pyridine. An IR spectrum was mea-
sured every 50 ◦C during temperature-programmed desorp-
tion (TPD) from 100 to 300 ◦C (heating rate of 10 ◦C/min)
after evacuation for 10 min. In the case of CO2 adsorption
(Messer Griesheim; 99.995 vol%), the samples were pre-
treated under vacuum at 400 ◦C for 1 h. The measurement
conditions were identical to those used after pyridine adsorp-
tion, except for the TPD evacuation time of 10 min, which
was lengthened to 30 min. IR spectra (64 scans, resolution of
2 cm−1) were recorded on a Digilab FTS-60A spectrometer.
The background spectrum was measured before adsorption;
spectra were normalized by the wafer weight (10 mg).
2.1. Catalyst preparation
Hydrotalcites (Pural MG 30, 50, 61, 70 from SASOL
Germany GmbH) [29] were calcined at 550 ◦C in air for
3 h. Al2O3 was obtained by calcination of AlO(OH) (Pural
SB, SASOL Germany GmbH) [29] under the same condi-
tions. Calcination of dried Mg(OH)2 at 600 ◦C in air for
4 h produced MgO. The hydroxide was precipitated from
Mg(NO3)2 with KOH and dried at 110 ◦C. The samples are
denoted in the following way: HT0.6 or CHT0.6 signifies
the sample with a bulk Mg/Al molar ratio of 0.6; the C in
the sample code indicates that the sample has been calcined
(Table 1).
2.2. Catalyst characterization
2.2.1. XRD, N2 adsorption experiments, ICP-OES, and
XPS
X-ray powder diffraction (XRD) was performed with Cu-
Kα radiation (λ = 1.5418 Å, 40 kV, and 35 mA; RD 7,
Rich. Seifert GmbH & Co. KG, Freiberg, Germany) over
the 2θ range from 5◦ to 65◦ or 90◦. One sample, HT0.6,
was calcined further at 700 ◦C to clearly interpret the pattern
of amorphous CHT0.6 (calcination temperature: 550 ◦C).
Nitrogen adsorption experiments were carried out at 77 K
with a surface area analyzer (Micromeritics; ASAP2010);
2.2.5. Microcalorimetric measurements
2.2.5.1. Gas-phase experiments with CO2 Microcalori-
metric data were collected with a heat flow Tian–Calvet-type
calorimeter (C 80, Setaram) connected to a volumetric line
with an online injection system for pulsing reactive gases.
CO2 (> 99.9%) was pulsed from a storage vessel. After
each pulse, the equilibrium pressure was measured with
a differential pressure gauge (Barocel, Datametrics). The

H.A. Prescott et al. / Journal of Catalysis 234 (2005) 119–130
121
calorimetric and volumetric data (pressure, adsorbed vol-
ume, heats of adsorption, differential and integral enthalpies)
were stored and analyzed with a microcomputer. The sam-
ple (70–100 mg) was pretreated under vacuum overnight at
400 ◦C. The first adsorption cycle was complete after a fi-
nal equilibrium pressure of 0.6 Torr was reached at 40 ◦C;
the system was then evacuated to remove the physisorbed
CO2, and a second adsorption cycle was performed. For the
samples measured here, the level of irreversible adsorption
was almost constant above 0.2 Torr. Thus, the amount of
totally adsorbed (chemisorbed and physisorbed) CO2 was
determined from the first isotherm; the quantity of irre-
versibly adsorbed (chemisorbed) CO2 was determined by
the difference between the isotherms of the first and second
adsorption cycles.
samples (about 0.3 ml) were taken periodically and cen-
trifuged; the separated clear solutions were then concen-
trated on a rotary evaporator to remove solvent and unreacted
methyl vinyl ketone.
2.2.6.4. NMR analysis of the catalytic yield Yields of
the target product based on the 1,3-dione were quanti-
tatively determined by 1H NMR spectroscopy (solvent:
DMSO-d6 or CDCl3), with the use of the integrals of the
CH3 signals of the internal standard, dimethyl phthalate
(3.83 ppm in DMSO-d6 and 3.88 ppm in CDCl3) and those
of the following Michael adducts: (a) 2-(3-oxo-butyl)-2-
methylcyclohexane-1,3-dione (1.11 and 2.04 ppm, DMSO-
d6); (b) 2-acetyl-2-(3-oxo-butyl)-cyclopentanone (2.14 and
2.09 ppm, CDCl3); and (c) 2-acetyl-2-(3-oxo-butyl)-cyclo-
hexanone (2.08 ppm, CDCl3; only one CH3 signal was used,
because the other overlapped with that of the 1,3-dione).
2.2.5.2. Liquid-phase experiments with benzoic acid Liq-
uid-phase experiments were performed on a Titrys calorime-
ter (Setaram) with a stirring system. The samples (about
200 mg) were pretreated under vacuum overnight at 400 ◦C
and then transferred to the calorimetric cell, which con-
tained toluene (1.5 ml). The reference cell contained the
same amount of toluene before injection. The experiments
were carried out at 70 ◦C except for that with MgO, which
was done at 40 ◦C. A solution of benzoic acid in toluene
(0.0307 mol/l) was injected stepwise (0.2 ml, injection rate:
0.05 ml/min) every 2 h. The amount of unreacted benzoic
acid was measured by UV spectrofluorimetry.
2.2.6.5. Experiments on the type of catalysis These exper-
iments were carried out in the reactions with 2-methylcyclo-
hexane-1,3-dione and 2-acetylcyclohexanone for CHT0.6 in
methanol and with 2-acetylcyclopentanone in the absence of
the solvent, methanol. The catalyst was separated by cen-
trifugation from the reaction mixture after a given amount
of time: 6 h for the reaction with 2-methylcyclohexane-1,3-
dione and 4 h for the reactions with 2-acetylcyclopentanone
and 2-acetylcyclohexanone. The product yield was then
monitored with NMR after 20 h. The cation content in
the reaction solution was determined with EDTA titration
(0.01 M). In the case of Al3+, EDTA (5 ml) was added to the
solution and then the solution was back-titrated with ZnSO4
(0.01 M) at pH 5.0 (sodium acetate buffer). The Mg content
was determined by direct titration with EDTA at pH 10.0
(ammonium buffer), and subtraction of the Al content was
determined.
2.2.6. Catalytic experiments
Michael addition test reactions were carried out in a
closed 50-ml round-bottomed flask at room temperature.
2.2.6.1. Michael addition of 2-methylcyclohexane-1,3-dione
The solid, partially insoluble dione, 2-methylcyclohexane-
1,3-dione (15 mmol; Acros, 98%), recrystallized from
methanol before use, was suspended in a mixture of methyl
vinyl ketone (22.5 mmol; Aldrich, 99%) and the solvent,
methanol (10 ml; Aldrich, 99%). The mixture was stirred
for 30 min to saturate the liquid phase with the CH-acid; the
saturation concentration of the 1,3-dione was 0.15 M (10%
of its total amount in the mixture) after 30 min. The catalyst
(0.225 g HT or the respective amount of CHT after calcina-
tion of 0.225 g HT; see Table 1) was then added.
3. Results and discussion
3.1. Physical properties of the catalysts
The physical properties of the HT and CHT samples and
the reference materials, MgO and Al2O3, are listed in Ta-
ble 1. All of the HT samples except for Al-rich HT0.6 have
low specific surface areas and high average pore diameters.
The high surface area and low pore diameter of HT0.6 may
be due to a separate, minor phase of boehmit in this sample
(Fig. 1A: d for XRD pattern). Sample calcination increases
the specific surface area independently of composition and
decreases the average pore diameter (exception: CHT0.6).
The ICP-OES results verified the identical Mg/Al mo-
lar ratios of the corresponding HT and CHT samples, yet
the surface Mg/Al molar ratios (XPS measurements) of
the CHT samples deviate from the bulk ratios. The sur-
face Mg/Al ratio for CHT0.6 is slightly higher than the
bulk ratio, whereas the bulk and surface ratios of CHT1.4
2.2.6.2. Michael addition of 2-acetylcyclopentanone and
2-acetylcyclohexanone The liquid diones, 2-acetylcyclo-
pentanone (15 mmol; Aldrich, 98%) and 2-acetylcyclohex-
anone (15 mmol; Aldrich, 97%), were added to methyl vin-
yl ketone (15.0 mmol) in 10 ml methanol. The catalyst
(0.225 g) was then added to the clear solution.
2.2.6.3. Sampling The internal NMR standard, dimethyl
phthalate (3.75 mmol; Acros, 99%), was added to the re-
action mixtures described above for a quantitative determi-
nation of product yields. To monitor the product formation,

122
H.A. Prescott et al. / Journal of Catalysis 234 (2005) 119–130
Table 1
Properties of hydrotalcites (HT), calcined hydrotalcites (CHT), MgO, and Al O
2
3
a
c
b
Sample
Theor. ratio
Al/(Al + Mg)
Mg/Al molar ratio
S
V
d
Weight loss
(wt%)
C
p
p
BET
2
3
(wt%)
(Å)
(m /g)
(cm /g)
Theor.
Bulk (ICP)
Surface (XPC)
HT0.6 (MG30)
HT1.4 (MG50)
HT2.2 (MG61)
HT3.0 (MG70)
CHT0.6
CHT1.4
CHT2.2
CHT3.0
MgO
0.67
0.44
0.33
0.25
0.67
0.44
0.33
0.25
–
0.5
1.25
2.0
3.0
0.5
1.25
2.0
3.0
–
0.6
1.4
2.2
3.0
0.6
1.4
2.2
3.0
–
n.d.
n.d.
n.d.
n.d.
0.8
1.4
2.0
2.0
–
163
13
15
0.32
0.05
0.05
0.10
0.52
0.23
0.13
0.22
0.96
0.54
76
167
128
205
81
45
46
52
258
45
1.42
1.19
2.42
2.27
–
–
–
–
–
31
36
42
44
–
–
–
–
–
20
257
201
114
203
75
Al O
2
–
–
–
–
234
–
–
3
a
Average pore diameter.
Carbon content.
Based on the weight lost during calcination in air at 550 C.
b
c
◦
The XRD patterns of the HT samples show a hydrotalcite-
like structure similar to that presented in [27,28,30,31],
whose crystallinity increases with Mg content (Fig. 1A: a–d).
Additional weak peaks for boehmite are seen at 28◦ and 49◦
in the HT0.6 pattern (Fig. 1A: d, cf. Fig. 1A: e for Pural
SB, AlO(OH), PDF No. 49-133) [29]. Calcination of the HT
samples results in the destruction of the hydrotalcite-like
structure and the formation of a periclase MgO (PDF no.
45-946, Fig. 1B: a–c), which is also found for pure MgO
(Fig. 1B: f). Peaks for the periclase structure in CHT0.6
(Fig. 1B: d) were broad and low in intensity. The MgO
peaks (at about 43◦ and 63◦) of the CHT samples have
higher 2θ values than those of pure MgO; this is caused
by the incorporation of smaller Al3+ cations into the bulk
lattice of the CHT samples [25,30–32]. The MgO peaks
shift from higher to lower 2θ values as the Mg/Al ratio
increases; thus, the amount of Al3+ cations incorporated
decreases accordingly. These results agree with the XPS re-
sults. The AlO(OH) reference material forms γ -Al2O3 after
calcination (PDF no. 29-63, Fig. 1B: e). Weak, broad peaks
for MgAl2O4 (35◦, PDF no. 21-1152) and γ -Al2O3 (66◦)
are observed in the patterns of CHT0.6 and CHT1.4. The
spinel phase is formed to a higher degree after calcination of
CHT0.6 at 700 ◦C (Fig. 1B: d1).
The 27Al MAS NMR spectra of γ -Al2O3 and CHT sam-
ples (Fig. 2) have two broad signals at about 5 ppm for Al
octahedra (AlO) and 64 ppm for tetrahedrally coordinated
aluminum cations (AlT) [16,19,21,30,32]. The identical sig-
nal shapes of γ -Al2O3 and CHT0.6, particularly that of the
64-ppm signal, suggest similar Al environments. The low-
field shift of the 5-ppm signal with increasing Mg/Al ra-
tio indicates an increased number of Mg–O–AlO linkages
[19,21]. An absence of signal splitting [21] reflects indis-
tinct Al coordination sites. The relatively narrow signal (at
5 ppm) of CHT1.4 could be attributed to more ordered AlO6
octahedra and related to the sample homogeneity confirmed
by ICP-OES and XPS.
◦
Fig. 1. XRD patterns of hydrotalcites (A), hydrotalcites calcined at 550 C
(B) with Mg/Al molar ratios of 3.0 (a), 2.2 (b), 1.4 (c), and 0.6 (d), MgO
after calcination at 600 C (intensity divided by a factor of 5, B: f), and the
Al oxide before (boehmit in A: e) and after calcination (γ -Al O in B: e).
◦
2
3
◦
HT0.6 was calcined further at 700 C (B: d1).
are identical. The surface of CHT2.2 and CHT3.0 are Al-
enriched. In [16], Al-enriched surfaces were also found with
XPS (Mg/Al = 1.5 and 2.4) for calcined Mg–Al hydrotal-
cites with Mg/Al molar ratios of 1.87 and 4.57, respectively.
The AlT/AlO ratio calculated from the integrated signals
for Al2O3 is 0.38 and close to that reported in [16]. The

H.A. Prescott et al. / Journal of Catalysis 234 (2005) 119–130
123
CHT samples have identical AlT/AlO ratios (0.38), except
for that of CHT3.0 (0.59). This indicates that AlT cations
dominate in CHT3.0, possibly by the formation of inverse
spinel-like domains: Mg–O–AlT [19]. AlO4 tetrahedra are
generally formed during calcination [30,33]; the AlT/AlO
ratio can also be influenced by the preparation method [16]
and Al content [16,19]. The CHT samples were all calcined
under the same conditions; thus, the higher AlT/AlO ratio in
CHT3.0 can only result from the strong difference in bulk
and surface Al contents not observed in the other samples.
An increase in the AlT/AlO ratio with the Al surface con-
centration has been suggested before [16].
In the TG-DTA curves of HT0.6 and HT3.0 (Figs. 3a
and 3b), gradual weight loss is observed from about 60 ◦C to
approximately 600 ◦C, with two main endothermic effects at
about 250 ◦C and 416 (432) ◦C. The first large endothermic
effect at about 250 ◦C involves interlayer water loss (m/z =
18) [27,34]; the second endothermic effect, at 432 ◦C, cor-
responds to the loss of OH− groups and the decomposition
2−
of CO3 in the brucite-like layers (m/z = 44) [27,34]. The
total weight loss observed for HT0.6 and HT3.0 (ca. 33.5
and 42.5%, respectively) is similar to that calculated from
weights taken before and after calcination (Table 1).
27
Fig. 2. Al MAS NMR spectra of Al O and the calcined hydrotalcites.
2
3
Fig. 3. TG-DTA profiles of hydrotalcites with the Mg/Al molar ratios: 0.6 (a) and 3.0 (b).

124
H.A. Prescott et al. / Journal of Catalysis 234 (2005) 119–130
Fig. 4. Acid properties: pyridine adsorption on Al O (a), the calcined hy-
drotalcites with Mg/Al molar ratios of 0.6 (b), 1.4 (c), 2.2 (d), and 3.0 (e),
2
3
Fig. 5. Acid properties: temperature programmed desorption of pyridine.
Effect of the evacuation temperature on the intensity of the band of adsorbed
and MgO (f). Spectra were measured after adsorption and evacuation at
−1
◦
pyridine around 1440 cm
.
40 C.
and CHT0.6; Al2O3 is the only sample to show a band in-
tensity over 0.5 a.u. after evacuation at 300 ◦C. CHT1.4–3.0
have different amounts of acid sites, but a similar number of
stronger acid sites: the amount of weaker acid sites increases
from CHT2.2 to CHT3.0 to CHT1.4.
In summary, CHT0.6 has the highest acidity among the
CHT samples, with regard number and strength of acid sites;
the total amount of acid sites on the other CHT samples
is much lower. The acid sites on CHT0.6 and Al2O3 are
comparable in number, yet the stronger they are, the fewer
of them there are on CHT0.6 compared with Al2O3. This
explains, at least partially, the drastic activity differences be-
tween (a) CHT0.6 and Al2O3 and (b) CHT0.6 and CHT1.4–
3.0.
3.2. Acid–base properties of the catalysts
3.2.1. FTIR studies
3.2.1.1. Pyridine adsorption The spectra measured after
pyridine adsorption at 40 ◦C on Al2O3, MgO, and the CHT
samples are shown in Fig. 4. Stretching vibrations of pyri-
dine molecules coordinatively bonded to Lewis acid sites
are between 1630 and 1600 cm−1 (ν8a mode) and 1450 and
1440 cm−1 (ν19b mode). Brønsted acid sites (bands at 1640
and 1540 cm−1) [35–37] were not found on any of the sam-
ples. Weak bands at 1600 and 1440 cm−1 for the negligible
Lewis acidity of MgO (Fig. 4f) correspond to coordinatively
unsaturated Mg2+ cations [36]. With increasing Al content,
the ν8a band shifts to higher wave numbers from 1601 cm−1
(CHT3.0, Fig. 4e) to 1613 cm−1 Al2O3 (Fig. 4a), corre-
sponding to an increase in acid site strength. The band at
1613 cm−1 on Al2O3 can be attributed to pyridine bonded to
a pair of coordinatively unsaturated Al3+ ions in octahedral
(AlO) and tetrahedral (AlT) coordination [38–41]. An addi-
tional band at 1622 cm−1 (Fig. 4a) can be ascribed to pyri-
dine strongly adsorbed on coordinatively unsaturated Al3+
ions in a tetrahedral environment (AlT) [38–41]. A similar
band of much lower intensity was observed for the CHT0.6
sample (Fig. 4b), which was possibly due to the presence
of boehmite. The total amount of absorbed pyridine was
estimated by integration of the band (ν19b mode) around
1440 cm−1. The value calculated after evacuation at differ-
ent temperatures is plotted against the evacuation tempera-
ture in Fig. 5. The total number of Lewis acid sites is highest
for Al2O3 and CHT0.6 and then decreases in the following
order: CHT1.4 > CHT3.0 > CHT2.2. Few Lewis acid sites
were detected on the surface of MgO. As the desorption
temperature increased, the integrated band intensity (i.e.,
number of pyridine molecules still adsorbed) decreased. At
200 ◦C, most of the pyridine on CHT1.4–3.0 was desorbed.
Higher intensities (> 0.5 a.u.) were only observed for Al2O3
3.2.1.2. CO2 adsorption The FTIR spectra (1100–1800
cm−1) of Al2O3, MgO, and CHT samples after CO2 adsorp-
tion at 40 ◦C are shown in Fig. 6. The thermal stability of
the surface species was studied by desorption experiments
at temperatures up to 300 ◦C (spectra not shown). Analy-
sis of the spectra permits the detection of surface Brønsted
and Lewis base sites and an estimation of their respective
amounts and strengths.
3.2.1.3. Surface species The dominant species formed
with the surface OH− groups of Al2O3 (Fig. 6a) is bicarbon-
ate (HCO3−), with bands at 3609 (νOH, not shown), 1647
(asym νC=O), 1476 (sym νC=O), and 1233 cm−1 (δOH). Sur-
face carbonates are formed to a lesser extent. In the case of
surface carbonates, a splitting of the weak asymmetric νCO
vibration (ν3 mode) (at 1415 cm−1 for the free carbonate
ion in inorganic salts) is caused by a lowering of the D3h
symmetry of the carbonate ion and related to the structure of
the surface (adsorbed) carbonate (e.g., monodentate, biden-
tate, or bridged) [42]. This splitting, ꢀν3, generally reflects
the strength of the base sites: the lower the splitting, the

H.A. Prescott et al. / Journal of Catalysis 234 (2005) 119–130
125
the polydentate structures formed probably have a structure
similar to that of bulk species [42].
3.2.1.4. Strength and amount of base sites According
to [44], the formation of different types of carbonates is
related to the basicity of the surface oxygen atoms. It was
postulated that the structural nonuniformity of surface oxy-
gen ions leads to basicity differences and, thus, causes the
formation of different carbonate structures. The spectral pa-
rameter, ꢀν3, was proposed as a measure of the relative
strength of surface base sites [44]: the smaller the ꢀν3 value,
the stronger the surface basic site that is involved in the in-
teraction with CO2.
Based on ꢀν3, the following can be concluded. Al2O3
exhibits the weakest Lewis basicity, with ꢀν3 values of
552 and 439 cm−1 for the bridged and bidentate carbonate
species, respectively. The strongest basicity was observed
for MgO, which formed monodentate species with CO2 and
exhibits the smallest ꢀν3 of 134 cm−1. The calcined hy-
drotalcites are less basic than MgO but more basic than
Al2O3. Here the strength of base sites (decrease in ꢀν3)
of the CHT samples decreases with increasing Mg con-
tent: CHT0.6 (188 cm−1) > CHT1.4 (209 cm−1) > CHT2.2
(225 cm−1) ≈ CHT3.0 (226 cm−1). The relative amount of
Lewis base sites was estimated by the integration of the car-
bonate bands (MgO: 1572, 1444, 1517, 1388 cm−1; CHT:
1574–1581, 1368–1386 cm−1) after evacuation at 150 ◦C,
that is, after removal of bicarbonate species (spectra not
shown). When the amount of Lewis base sites was normal-
ized by weight, intensities decreased for the samples in the
following order: CHT3.0 ∼ CHT1.4 ∼ MgO > CHT2.2 >
CHT0.6 ꢀ Al2O3. CHT3.0 has the highest number of weak
Lewis base sites (higher splitting), whose strengths are sim-
ilar to those found on CHT2.2; Lewis base sites on CHT0.6,
the most catalytically active sample, are the strongest, but
few in number. These findings agree with [30], in which the
total amount of base sites was increased, but the proportion
of stronger base sites decreased with increasing Mg/Al mo-
lar ratios.
Fig. 6. Base properties: CO adsorption on Al O (a), calcined hydrotal-
2
2 3
cites with Mg/Al ratios of 0.6 (b), 1.4 (c), 2.2 (d), and 3.0 (e), and MgO (f).
◦
Spectra were measured after adsorption and evacuation at 40 C.
stronger the base site. Values measured here are given in the
text when applicable. In the case of Al2O3, weak bands at
1756 and 1204 cm−1 (Fig. 6a) can be assigned to νC=O vi-
brations of bridged “organic-like” complexes on Al2O3 and
indicate a ꢀν3 splitting of 552 cm−1 for a weak base site.
Weak shoulders at 1704 and 1265 cm−1 (not indicated in
Fig. 6) can be assigned to bidentate carbonate species (ꢀν3
splitting of 439 cm−1) on Al2O3. The band at 1450 cm−1
may indicate the presence of free carbonate ions.
The spectrum of MgO (Fig. 6f) after CO2 adsorp-
tion is complex, with numerous bands between 1800 and
1100 cm−1. The major bicarbonate species is indicated by
the bands at 1630, 1452, and 1217 cm−1. Weaker bands
(shoulders) at 1600, 1448, and 1200 cm−1 (not indicated in
Fig. 6) could belong to a second type of bicarbonate species.
Pairs of bands at 1686/1365 cm−1 and 1658/1304 cm−1 can
be assigned to the νC=O vibrations of two different bidentate
carbonate structures. Based on their low ꢀν3 splitting (134
and 144 cm−1, respectively), these bands are normally at-
tributed to monodentate carbonate species [42–44], yet their
corresponding high thermal stability (stable up to 300 ◦C
under vacuum, not shown) suggests they are caused by poly-
dentate carbonates [45].
The intensity of the OH− deformation band (around
1220 cm−1) for surface bicarbonates after evacuation at
40 ◦C has been used to quantitatively determine the num-
ber of basic OH− species [43]. Interestingly enough, the
decrease in the band intensities (amount of OH− species) is
CHT0.6 > CHT3.0 > CHT1.4 > CHT2.2 (Fig. 7). The shift
in the 1220 cm−1 band to lower wave number indicates an
increase in site strength with Mg content from CHT0.6 to
MgO. Thus, CHT0.6 has a high number of Brønsted base
sites; CHT3.0 has fewer but stronger sites.
A lower structural variety of carbonates was formed on
the surface of CHT samples (Fig. 6b–e). Very broad bi-
carbonate bands are seen at 1658–1650 cm−1 and 1420–
1400 cm−1. Based on the ꢀν3 splitting (see below), the max-
ima at 1591–74 cm−1 and 1386–65 cm−1 can be attributed
to monodentate surface carbonates [42–45]. The thermal sta-
bility of these species is rather high; their bands are observed
even after evacuation at 300 ◦C. Thus, as in the case of MgO,
3.2.2. Microcalorimetry
Base sites were additionally characterized by microcal-
orimetry. The probe molecule, CO2, was used for gas-phase
experiments, a common method; liquid-phase experiments
were carried out with benzoic acid in toluene to gain insights
into base site behavior under reaction conditions.

126
H.A. Prescott et al. / Journal of Catalysis 234 (2005) 119–130
−1
Fig. 7. Base properties: intensity of band around 1220 cm after evacua-
tion at 40 C.
Fig. 9. Base properties: differential heats of adsorption of benzoic acid in
toluene measured at 70 C except for that of MgO (40 C).
◦
◦
◦
performed at 40 ◦C and do not distinguish between Brøn-
sted and Lewis base sites. The FTIR measurements, on the
other hand, only characterize Lewis base sites (measurement
after evacuation at 150 ◦C).
3.2.2.2. Liquid-phase microcalorimetry with benzoic acid
Microcalorimetric experiments with a toluene solution of
benzoic acid (used to measure the strength and amount of
base sites by indicator titration [46]) were performed to char-
acterize basicity under liquid-phase conditions like those
of catalytic test reactions. The differential heats of adsorp-
tion given in Fig. 9 show a base site behavior that is much
different from that found for the experiments with CO2
(Fig. 8). Here CHT3.0 has much stronger base sites ini-
tially, but far fewer than CHT0.6; CHT1.4 and CHT2.2 have
even fewer sites, but their strengths are similar to those of
CHT0.6. The liquid-phase basicity of MgO is similar to that
of CHT0.6, but was measured at a lower experimental tem-
perature of 40 ◦C (instead of 70 ◦C); this temperature dis-
crepancy may indicate weaker and fewer base sites in MgO
than in CHT0.6.
◦
Fig. 8. Base properties: differential heats of CO adsorption at 40 C as a
function of CO uptake.
2
2
3.2.2.1. Gas-phase microcalorimetry with CO2 The dif-
ferential heats of CO2 adsorption as a function of CO2 up-
take on calcined hydrotalcites, Al2O3, and MgO (Fig. 8)
show initial values between 130 and 155 kJ/mol. The CHT
samples have more base sites than MgO and Al2O3. This
applies, in particular, to the increase in the number of weak
base sites found on the CHT samples and agrees with the
FTIR results. Again, CHT3.0 exhibits the highest number of
weak basic sites (compare Section 3.2.1 FTIR studies: CO2
adsorption). The immediate and continuous decrease in the
heats of adsorption with increasing CO2 uptakes (and the
very broad carbonate bands in the FTIR spectra after CO2
adsorption) indicates a wide distribution of site strength,
particularly for the CHT samples. The total uptake of irre-
versible (chemisorbed) CO2 decreases in the following or-
der at a differential heat value of 50 kJ/mol: CHT3.0 >
CHT1.4 > CHT2.2 > CHT0.6 > MgO > Al2O3. With re-
spect to MgO, this order differs somewhat from that ob-
served by CO2-FTIR. This could be due to the different con-
ditions applied; the microcalorimetric measurements were
3.3. Application of the catalysts in Michael additions
3.3.1. Michael addition of 2-methylcyclohexane-1,3-dione
to methyl vinyl ketone
The catalyzed Michael addition of 2-methylcyclohexane-
1,3-dione (pKa = 6.7) to methyl vinyl ketone yields 2-(3-
oxo-butyl)-2-methylcyclohexane-1,3-dione (Scheme 1). The
poorly soluble CH-acid compound partially dissolves in
methanol to slowly establish a saturation concentration of
0.15 M from a total theoretical concentration of 1.5 M in the
reaction mixture; it is then converted into the liquid reaction
product. In blank experiments (experiments without a cata-
lyst) with recrystallized 2-methylcyclohexane-1,3-dione, no
product formation was observed.
We performed catalytic studies of the HT and CHT sam-
ples by monitoring product formation; the activities of the

H.A. Prescott et al. / Journal of Catalysis 234 (2005) 119–130
127
Scheme 1. Michael addition of 2-methylcyclohexane-1,3-dione to methyl
vinyl ketone.
Scheme 2. Consecutive aldol cyclization of the Michael addition product
formed according to Scheme 1 on MgO.
less the same for CHT1.4 and CHT0.6, decreased from 69 to
49% for CHT2.2, and increased from 82 to 94% for CHT3.0
(Fig. 10B). HT2.2 and CHT2.2 exhibited the lowest catalytic
activities. The activities of the CHT samples decreased in
the same order as the HTs: CHT0.6 > CHT3.0 > CHT1.4 >
CHT2.2.
Choudary et al. found that whereas a rehydrated Mg–
Al hydrotalcite was an efficient, very selective catalyst for
Michael additions, the untreated and calcined hydrotalcites
were inactive. This finding was ascribed to their Brønsted
base sites. However, the lack of a detailed characterization
and discussion of the catalysts tested make a detailed com-
parison with the results presented here difficult. The differ-
ences in the reaction results may be related, however, to the
lower acidity of the CH-acid compounds used in [9], com-
pared with the more acidic CH-acid compounds used here.
Regarding the pure oxides, the catalytic activity of
Al2O3 was lower than that of the CHT samples (Fig. 10B)
and may be related to its Lewis acidity or weak ba-
sicity or both. Michael additions can be catalyzed by
both base and acid catalysts; acid catalysts such as acetic
acid [47], a montmorillonite-enwrapped scandium com-
plex [48], or TS-1 molecular sieve [49] have been used
before. MgO had a considerably different, less selective
behavior than that of the CHT samples or Al2O3. In the
case of MgO, a high product yield of 65% was formed
within the first 30 min, but the amount of the Michael
adduct in the reaction mixture dropped drastically during
1
the next 24 h. In the H NMR spectra of the reaction mix-
Fig. 10. Results of the Michael addition of 2-methylcyclohexane-1,3-dione
(15 mmol) to methyl vinyl ketone (22.5 mmol) in 10 ml methanol according
to Scheme 1 over the hydrotalcites (A) and over the calcined hydrotalcites,
Al O and MgO (B). Catalyst amounts were 0.225 g for the hydrotalcites,
ture, additional CH3 signals (0.97 and 1.23 ppm) were
found for a consecutive product, 6-hydroxy-1,6-dimethyl-
2,9-dioxobicyclo[3.3.1]nonane, which was formed by the in-
tramolecular aldol cyclization of 2-methyl-2-(3-oxo-butyl)-
cyclohexane-1,3-dione (Scheme 2) and identified with GC-
MS.
2
3
Al O , and MgO and the amount remaining after calcination of 0.225 g
2
3
◦
hydrotalcite at 550 C.
solid bases differed significantly (Figs. 10A and 10B, respec-
tively). The target product was formed with 100% selectiv-
ity.
Catalytic activity decreased in the following order: MgO
> CHT0.6 > CHT3.0 > CHT1.4 > CHT2.2 > Al2O3. This
correlates with the band intensities of the OH− species
(FTIR experiments), except for that of Al2O3, which has
a band intensity similar to that of CHT1.4 (Fig. 7). The
catalytic activities within the CHT sample series also re-
flect the amounts of base sites detected with liquid-phase
microcalorimetry (Fig. 9). The discrepancy between the sim-
ilar liquid-phase basicities and different catalytic behaviors
of CHT0.6 and MgO (Figs. 9 and 10B, respectively) can-
not be explained at this time. In gas-phase microcalorimetry,
The product yields obtained after 24 h on the HT samples
decreased in the order HT0.6 > HT3.0 > HT1.4 > HT2.2
and do not correlate with the Mg/Al ratio. Samples with the
lowest and highest Mg/Al ratios achieved the highest yields
after 24 h: 98% for HT0.6 and 82% for HT3.0 (Fig. 10A).
The catalytic behavior of the CHT samples formed by
calcination of the corresponding HT samples (weight loss
from dehydration and decarboxylation) remained more or

128
H.A. Prescott et al. / Journal of Catalysis 234 (2005) 119–130
Scheme 3. Michael addition of 2-acetylcyclopentanone (n = 1) and 2-ace-
tylcyclohexanone (n = 2) to methyl vinyl ketone.
ly): MgO > CHT0.6 > CHT3.0 > CHT2.2 > CHT1.4 >
Al2O3.
The Michael addition of 2-acetylcyclopentanone to meth-
yl vinyl ketone proceeded faster than that of 2-acetylcyclo-
hexanone, except in the case of MgO. This can be explained
by the higher acidity and more favored deprotonation of
2-acetylcyclopentanone by the base sites. Consequently, the
carbanion formed as the reactive species is present in higher
equilibrium concentration.
CHT0.6 exhibited an outstanding catalytic activity in
the Michael addition reaction of the less acidic CH-acid
compound, 2-acetylcyclohexanone (Fig. 11B; compare with
Fig. 11A). The activities of the other CHT samples and
Al2O3 are significantly lower. Their basicities may be
too weak to sufficiently deprotonate the CH-acid com-
pound, yet the amount of Lewis base sites determined
does not completely explain the performance of CHT0.6.
Acid sites may also facilitate the catalytic process; Al2O3
shows a considerable activity in the Michael addition of
2-acetylcyclopentanone (pKa = 7.8) (Fig. 11A). The mild
acid sites on Al2O3 and the presence of weak base sites cat-
alyze the addition with more reactive diones to some extent,
and the active sites determining catalytic activity are, in fact,
the base sites.
Fig. 11. Results of the Michael additions of 2-acetylcyclopentanone
(15 mmol) (A) and 2-acetylcyclohexanone (15 mmol) (B) to methyl vinyl
ketone (22.5 mmol) in 10 ml methanol according to Scheme 3 over calcined
hydrotalcites (0.225 g).
MgO, on the other hand, showed a similar catalytic be-
havior in the experiments with 2-acetylcyclopentanone and
2-acetylcyclohexanone and the reaction of 2-methylcyclo-
hexane-1,3-dione. Products formed rapidly initially, with
final catalytic yields drastically decreasing thereafter (cf.
Figs. 11A and 11B).
the CHT samples exhibited basicities, which do not corre-
spond to their catalytic activities in the Michael addition,
yet there is a direct relation between Mg contents and the
base site interaction with the gaseous probe molecule CO2
(Fig. 8).
A similar relationship exists between catalyst activity
and basicity (characterized by liquid-phase microcalorime-
try and the OH− band intensity at 1220 cm−1) for all three
Michael additions. The only exception is the inversed cat-
alytic activity of CHT1.4 and CHT2.2.
Solvent-free reactions with the most active samples,
CHT0.6 and CHT3.0, examined the environmentally friend-
ly implementation of the Michael additions (Scheme 3).
CHT0.6 attained the highest activities; nevertheless, yields
of the target products were lower than those found when
a solvent was used for both catalysts after 24 h. The de-
crease in yields for the solvent-free reaction with 2-ace-
tylcyclopentanone is much more drastic (Fig. 12A) than for
that with 2-acetylcyclohexanone (Fig. 12B). The low (de-
layed) product formation may be related to the absence of
carbanion stabilization by methanol. Nonetheless, the inves-
3.3.2. Michael additions of 2-acetylcyclopentanone and
2-acetylcyclohexanone to methyl vinyl ketone
To extend the catalytic application of the CHT sam-
ples, further Michael additions were performed with the
less acidic CH-acid compounds, 2-acetylcyclopentanone
(pKa = 7.8) and 2-acetylcyclohexanone (pKa = 10.1), to
give 2-acetyl-2-(3-oxo-butyl)-cyclopentanone and 2-acetyl-
2-(3-oxo-butyl)-cyclohexanone, respectively (Scheme 3). In
comparison with the solid 2-methylcyclohexane-1,3-dione,
these CH-acids are liquids and are completely soluble in
methanol. Thus, the catalytic experiments with these com-
pounds are only roughly comparable to those reported in the
previous section.
The decreasing order of catalytic activity was identi-
cal for the addition of both 2-acetylcyclopentanone and
2-acetylcyclohexanone (Figs. 11A and 11B, respective-

H.A. Prescott et al. / Journal of Catalysis 234 (2005) 119–130
129
be finely dispersed in the reaction solution. In addition, the
presence of EDTA, a good complexing agent for Mg (pKB
for Mg-EDTA complex: 8.65 [50]), allows a better dissolu-
tion of the Mg cations than that of the Al cations (pKB for
Al-EDTA complex: 16.7 [50]) during titration; this explains
the much higher Mg content found. Thus, the increase in
yield after catalyst separation is most likely caused by fine
particles of the solid catalyst still in the solution (heteroge-
neous catalysis) and not homogeneous catalysis.
The liquid-phase basicity microcalorimetry measure-
ments with benzoic acid in toluene correlated very well with
the reaction results, yet benzoic acid is not able to form a
chelate with Mg, and Mg is unlikely to leach in toluene.
Thus, leaching can be excluded from these measurements.
Based on this and the low Mg content in the solution after a
questionable separation of the catalyst, we conclude that ho-
mogeneous catalysis plays a very minor role, if any, in the
catalytic reactions studied here.
4. Conclusions
– The calcined hydrotalcites (periclase MgO stucture) ex-
hibit Lewis acid, Brønsted base, and Lewis base sites.
The calcined Al-rich sample (Mg/Al molar ratio of 0.6,
amorphous MgO) possesses Lewis acid sites similar in
strength to those found on Al2O3, but stronger than
those found on the Mg-rich hydrotalcites.
– Both the uncalcined hydrotalcites and the calcined hy-
drotalcites catalyze the Michael additions of 2-meth-
ylcyclohexane-1,3-dione, 2-acetylcyclopentanone, and
2-acetylcyclohexanone to methyl vinyl ketone with high
selectivity (no side product formation). These Michael
additions in methanol proceeded faster than the solvent-
free reactions. Experiments on the type of catalysis
involved indicated that homogeneous catalysis is less
probable, but could not be totally disregarded.
– Pure Al2O3 was the least active among the investigated
catalysts. Both the Al-rich hydrotalcite with traces of
boehmit and the calcined Al-rich sample (Mg/Al mo-
lar ratio of 0.6) gave product yields above 95% within
24 h.
– A nonlinear correlation between the Mg content and cat-
alytic activity suggested that the activity of pure MgO
surpasses that of the calcined hydrotalcites but causes
consecutive reactions of the Michael addition products,
which reduce the product selectivities and yields.
– Catalytic activity correlates with the amount of the base
sites determined by benzoic acid microcalorimetry de-
pendent on the Mg/Al molar ratio.
– An optimal balance of acid–base properties may make
the calcined Al-rich hydrotalcite an excellent catalyst in
the Michael addition of numerous 1,3-diones indepen-
dently of their pKa values.
Fig. 12. Results of the solvent-free Michael additions of 2-acetylcyclopen-
tanone (15 mmol) (A) and 2-acetylcyclohexanone (15 mmol) (B) to methyl
vinyl ketone (22.5 mmol) on CHT0.6 and CHT3.0 (0.225 g).
tigated Michael additions can be performed under solvent-
free conditions, but with longer reaction times.
3.3.3. Investigations of the type of catalysis
Experiments on the type of catalysis (homogeneous or
heterogeneous) involved in the Michael additions showed
that although the product yield continued to increase af-
ter removal of the solid catalyst (CHT0.6) in the presence
and absence of methanol, the reactions proceed consider-
ably more slowly than they do without catalyst separation
(via centrifugation). Final yields (20 h after catalyst sepa-
ration) increased from 49 to 76% and from 35 to 70% in
the reactions with cyclohexane-1,3-dione (Experiment A)
and 2-acetyl-cyclohexanone (Experiment B) in methanol.
The solvent-free reaction with 2-acetylpentanone (Experi-
ment C) had a corresponding yield of 34% up from 21%
(20 h after the catalyst was separated). The titrated reaction
solution of Experiment B had a Mg content of about 4%
of the total Mg content in the catalyst, and 0.3% of the Al
cations in the catalyst were found in this solution. However,
because the catalyst is a very fine powder, instead of being
completely separated, some of the solid catalyst may still

130
H.A. Prescott et al. / Journal of Catalysis 234 (2005) 119–130
Acknowledgments
[23] S. Casenave, H. Martinez, C. Guimon, A. Auroux, V. Hulea, E. Du-
mitriu, J. Therm. Anal. Cal. 72 (2003) 191.
[24] P.S. Kumbhar, J.S. Valente, J. Lopez, F. Figueras, Chem. Commun.
(1998) 535.
[25] K. Yamaguchi, K. Ebitani, T. Yoshida, H. Yoshida, K. Kaneda, J. Am.
Chem. Soc. 121 (1999) 4526.
[26] D. Tichit, D. Lutic, B. Coq, R. Durand, R. Teissier, J. Catal. 219 (2003)
167.
[27] D. Tichit, M.H. Lhouty, A. Guida, B.H. Chiche, F. Figueras, A. Au-
roux, D. Bartalini, E. Garrone, J. Catal. 151 (1995) 50.
[28] M.J. Climent, A. Corma, S. Iborra, K. Epping, A. Velty, J. Catal. 225
(2004) 316.
We acknowledge the Bundesministerium für Bildung und
Forschung (FKZ 03C0328) for their financial support. We
also thank Dr. M. Feist for the thermal analysis, W. Zi-
esche for adsorption experiments, Dr. J. Radnik for XPS
measurements, and Dr. D. Heidemann for MAS NMR mea-
surements.
[29] www.sasol.com, product information.
References
[30] A. Corma, V. Fornés, F. Rey, J. Catal. 148 (1994) 205.
[31] G. Fornasari, M. Gazzano, D. Matteuzi, F. Trifiro, A. Vaccari, Appl.
Clay Sci. 10 (1995) 69.
[32] K.J.D. MacKenzie, R.H. Meinhold, B.L. Sherriff, Z. Xu, J. Mater.
Chem. 3 (1993) 1263.
[33] W.T. Reichle, S.Y. Kang, D.S. Everhardt, J. Catal. 101 (1986) 352.
[34] S. Miyata, Clays Clay Miner. 28 (1980) 50.
[35] H. Knözinger, Adv. Catal. 25 (1976) 184.
[36] G. Busca, Phys. Chem. Chem. Phys. 1 (1999) 723.
[37] N. Fripiat, R. Conanec, A. Auroux, Y. Laurent, P. Grange, J. Catal. 167
(1997) 543.
[38] C. Morterra, A. Chiorino, G. Ghiotti, E. Garrone, J. Chem. Soc., Fara-
day Trans. 1 75 (1979) 271.
[39] C. Morterra, S. Coluccia, A. Chiorino, F. Boccuzzi, J. Catal. 54 (1978)
348.
[40] F. Abbattista, S. Delmastro, G. Gozzelino, D. Mazza, M. Vallino,
G. Busca, V. Lorenzelli, G. Ramis, J. Catal. 117 (1989) 42.
[41] P. Nortier, P. Fourre, A.B. Mohammed Saad, O. Saur, J.C. Lavalley,
Appl. Catal. 61 (1990) 141.
[42] J.C. Lavalley, Catal. Today 27 (1996) 377.
[43] J.A. Lercher, C. Colombier, H. Noller, J. Chem. Soc., Faraday Trans.
1 80 (1984) 949.
[44] A.A. Davydov, M.L. Shepot’ko, A.A. Budneva, Kinet. Catal. 35
(1994) 299.
[45] G. Busca, V. Lorenzelli, Mater. Chem. 7 (1982) 89.
[46] K. Tanabe, M. Misono, Y. Ono, H. Hattori, New Solid Acids and
Bases—Their Catalytic Properties, Stud. Surf. Sci. Catal., vol. 51, El-
sevier, Tokyo, 1989, p. 14.
[47] T. Bui, C.F. Barbas, Tetrahedron Lett. 41 (2000) 6951.
[48] T. Kawabata, T. Mizugaki, K. Ebitani, K. Kaneda, J. Am. Chem.
Soc. 125 (2003) 10486.
[1] H. Hattori, Stud. Surf. Sci. Catal. 78 (1993) 35.
[2] H. Hattori, Chem. Rev. 95 (1995) 537.
[3] J. Barrios, R. Rojas, A.R. Alcanrara, J.V. Sinisterra, J. Catal. 112
(1988) 528.
[4] H. Kabashima, H. Tsuji, H. Hattori, Appl. Catal. A 165 (1997) 319.
[5] H. Kabashima, H. Tsuji, T. Shibuya, H. Hattori, J. Mol. Catal. A 155
(2000) 23.
[6] H. Kabashima, H. Tsuji, S. Nakata, H. Hattori, Appl. Catal. A 194–195
(2000) 227.
[7] U. Meyer, H. Gorzawski, W.F. Hölderich, Catal. Lett. 59 (1999) 201.
[8] Z.-J. Li, H.A. Prescott, J. Deutsch, A. Trunschke, H. Lieske, E. Kem-
nitz, Catal. Lett. 92 (2004) 175.
[9] B.M. Choudary, M.L. Kantam, Ch.V. Reddy, K.K. Rao, F. Figueras,
J. Mol. Catal. A 146 (1999) 279.
[10] M.J. Climent, A. Corma, S. Iborra, J. Primo, J. Catal. 151 (1995) 60.
[11] A. Corma, V. Fornés, R.M. Martin-Aranda, F. Rey, J. Catal. 134 (1992)
58.
[12] F. Cavani, F. Trifiro, A. Vaccari, Catal. Today 11 (1991) 173.
[13] A. Vaccari, Catal. Today 41 (1998) 53.
[14] B.F. Sels, D.E. De Vos, P.A. Jacobs, Catal. Rev. Sci. Eng. 43 (2001)
443.
[15] D. Tichit, B. Coq, CATTECH 7 (2003) 206.
[16] A.L. McKenzie, C.T. Fishel, R.J. Davis, J. Catal. 138 (1992) 547.
[17] J.I. Di Cosimo, V.K. Díez, M. Xu, E. Iglesia, C.R. Apesteguía, J. Ca-
tal. 178 (1998) 499.
[18] J. Shen, J.M. Kobe, Y. Chen, J.A. Dumesic, Langmuir 10 (1994) 3902.
[19] V.K. Díez, C.R. Apesteguía, J.I. Di Cosimo, J. Catal. 215 (2003) 220.
[20] F. Prinetto, G. Ghiotti, R. Durand, D. Tichit, J. Phys. Chem. B 104
(2000) 11117.
[21] J. Shen, M. Tu, C. Hu, J. Solid State Chem. 137 (1998) 295.
[22] S. Casenave, H. Martinez, C. Guimon, A. Auroux, V. Hulea, A. Cor-
doneanu, E. Dumitriu, Thermochim. Acta 379 (2001) 85.
[49] M. Sasidharan, R. Kumar, J. Catal. 220 (2003) 326.
[50] N. Wiberg (Ed.), Lehrbuch der Anorganischen Chemie, hundred first
ed., Walter de Gruyter, Berlin, 1995, p. 1219.