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Journal of Natural Gas Chemistry 15(2006)247-252
Journalof Natural Gas Chemstry
SCIENCE PRESS
www.elsevier.m%catdjomnoeate/jngc
Article
The Effect of Sulfate Ion on the Isomerization of
n-Butane to iso-Butane
Sugeng Triwahyonol*,
Zalizawati Abdullah2,
Aishah Abdul Jali12
1 . Ibnu Sina Instatute for Fundamental Science Studies, 2. Faculty of Chemical and Natural
Resources Engineering, Universiti Teknologi Malaysia, 8131 0 Skudai, Johor, Malaysia
[Manuscript received October 18, 20061
Abstract:
The effect of sulfate ion (SO%-) loading on the properties of Pt/SO:--ZrOz and on the
catalytic isomerization of n-butane to iso-butane was studied. The catalyst was prepared by impregnation
of Zr(0H)A with H2S04 and platinum solution followed by calcination at 600 "C. Ammonia T P D and
FT-IR were used to confirm the distribution of acid sites and the structure of the sulfate species. Nitrogen
physisorption and X-ray diffraction were used to confirm the physical structures of Pt/SO:--ZrO2. XRD
pattern showed that the presence of sulfate ion stabilized the metastable tetragonal phase of zirconia
and hindered the transition of amorphous phase to monoclinic phase of zirconia. Ammonia TPD profiles
indicated the distributions of weak and medium acid sites observed on 0.1 N and 1.0 N sulfate in the
loaded catalysts. The addition of 2.0 N and 4.0 N sulfate ion generated strong acid site and decreased
the weak and medium acid sites. However, the XRD results and the specific surface area of the catalysts
indicated that the excessive amount of sulfate ion collapsed the structure of the catalyst. The catalysts
showed high activity and stability for isomerization of n-butane to iso-butane at 200 "C under hydrogen
atmosphere. The conversion of n-butane to iso-butane per specific surface area of the catalyst increased
with the increasing amount of sulfate ion owing to the existence of the bidentate sulfate and/or polynucleic
sulfate species ((ZrO)zSOz), which acts as an active site for the isomerization.
Key words: sulfate ion; strong acid site; isomerization; n-butane; Zr; Pt
1. Introduction
Catalyzed isomerization of alkane is one of the
important processes in petroleum refining to produce high quality gasoline because of the capability
to modify the octane number of gasoline. In industrial processes, acid catalyst is known as a media for
the conversion of alkane into iso-alkane. However, the
catalysts such as HF, and catalysts containing halides
have many disadvantages and are not suitable for the
isomerization of alkanes. HF is particularly dangerous while catalysts containing halides such as AlC13 or
sulphuric acid are corrosive and pose significant environmental challenges including the disposal of waste
[l-3).
* Corresponding author. E-mail: sugeng8ibnusina.utm.my
Considerable interest has been focused on the use
of strong solid acids based on anion-modified zirconium oxide catalyst. Recently, many investigations
have been focused on Pt/SO;--ZrOz because it was
reported t o exhibit higher activity and selectivity in
the isomerization of c 4 - C ~ [4,5]. In addition, the
sulfated zirconia showed catalytic activities for diversified acid-catalyzed reactions at low temperature.
This catalytic performance is unique when compared
to typical solid acid catalysts, such as zeolites, which
are mesoporous materials showing no activity for the
reaction a t such low temperature.
Zirconia possesses weak acid and basic properties
and has no capacity for alkane isomerization. It has
been realized that the catalytic activity depends on
248
Sugeng Tkiwahyono et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
the acid-based properties of ZrO2. The sulfate ion
(SO;-) content has a significant effect on the catalyst performance. The presence of SO;- promotes
the acidity and activity towards the isomerization of
alkane. Yori et al. proposed that the addition of
the SO;- ion, an electron-rich anion, produces a very
strong Lewis acid-based pair [6]. This is because of
the inductive effect of the S=O groups, which produces an electronic deficiency markedly increasing the
Lewis acidity of the zirconia cations.
The acid sites are classified into two groups,
Broensted and Lewis acid sites. The Broensted acid
site also known as the protonic acid site usually exists
in the form of a hydroxyl group, whereas the Lewis
acid site usually appears as an unsaturated metal.
The Broensted acid site is involved in the surface intermediate formation by protonation of alkane. The
protonation either to a C-H bond or to a C-C bond
forms pentacoordinated carbonium ion, and this carbonium ion liberates one hydrogen or one alkane molecule while leaving the carbenium ion as intermediate
in the isomerization [7].
The major concern of this research is t o elucidate
the effect of sulfate ion loaded on the properties of
Pt/SO;--ZrOz, in terms of the acidity and activity
on the isomerization of n-butane to iso-butane.
2. Experimental
2.1. Preparation of catalyst
The sulfate ion-treated SO;--ZrOz was prepared
by impregnation of zirconium hydroxide Zr (OH)4
with H2SO4 aqueous solution followed by filtration
and drying a t 110 "C. The concentration of the H2S04
aqueous solution was varied: 0.5 N, 1.0 N, 2.0 N, and
4.0 N. The SO;--ZrOz was obtained by calcination of
the SO;--Zr(OH)4 at 600 "C under static air for 3 h
~~91.
The Pt/SO;--ZrOz was prepared by impregnation of the SO;--ZrOn with H2ClsPt.BH20 solution
followed by drying and calcination at 600 "C in air
for 3 h and obtained 0.5wt% Pt in Pt/SO;--Zr02.
The X-ray powder diffraction pattern of the Sample was recorded on a JEOL X-ray diffractometer
JDX-3500 with a Cu-Ka 40 kV, 45 mA radiation
source. The diffraction spectrum range was 20-90'
and the scanning speed and step were 1 s and 0.050°,
respectively.
The ammonia T P D test was done on JEOL Multitask TPD-MS. The sample was pre-treated with He
flow at 400 "C for 2 h, it was then cooled with
He flow until 100 "C. The sample was outgassed at
100 "C until it reached
torr; the sample was
then exposed t o dehydrated ammonia at 100 "C for
30 min (30 torr) followed by purging with He flow
to remove excess ammonia. T P D was run at room
temperature up to 900 "C with a heating rate of
10 "C/min.
IR measurement was done using a Perkin-Elmer
Spectrum One FT-IR Spectrometer equipped with a
MCT detector. A self-supported wafer placed in an
in-situ IR cell was pre-treated at 400 "C for 3 h and
outgassed a t 400 "C in hydrogen flow for 3 h. The
measurements were done at room temperature. The
process was repeated with the samples with different
SO;- loadings.
Thermogravimetric analysis (TGA) was conducted by the T G analysis system (Perkin Elmer
Pyris Diamond TG/DTA). The sample was pretreated under N2 flow at 300 "C for 1 h. The sample
temperature was then cooled to 50 "C before the Sample was heated up to a final calcination temperature
of 950 "C at a heating rate of 5 "C/min under Nz flow.
2.3. Catalyst testing
Isomerization of n-butane was done in an online
microreactor system at 200 "C for 15 min in the presence of hydrogen (100 ml/min, atmospheric pressure)
and in saturated n-butane. 0.4 g of catalyst sieved
to 35-80 mesh was charged for each catalytic testing. The reaction products were analyzed by online
gas chromatography, using VZ-7 packed column with
hydrogen, nitrogen, and helium as the carrier gases.
3. Results and discussion
2.2. Characterization of the catalyst
3.1. Physical structure
The surface area and the pore distribution of the
catalyst were determined using a COULTER SA3100
apparatus. The sample was treated and outgassed at
300°C for 3 h before being subjected to nitrogen (N2)
adsorption.
After calcinations of pure zirconia a t 600 "C and
above, zirconia transforms into a monoclinic phase
(20=28.3' and 31.4O) from a tetragonal phase
(20=30.Z0) of zirconia [8]. However, the addition of
Special Column of the INRET 2006/Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
metal oxide or anion such as tungsten oxide, molybdenum oxide or sulfate ion stabilizes the tetragonal
phase of zirconia and suppresses the formation of the
monoclinic phase of zirconia. Figure 1shows the XFtD
patterns of the prepared samples with different concentration of sulfate ion loading. The results show
that the tetragonal phase is a dominant structure of
zirconia and the addition of sulfate ion up to 1.0 N
does not considerably change the crystallinity of the
catalysts. However, further addition of the sulfate ion
will collapse the structure of zirconia. In the addition
of 4.0 N, a new peak associated with the monoclinic
phase of zirconia was observed at 31.4'.
20
30
40
50
60
70
80
90
28/(0 )
Figure 1. XRD pattern of Pt/SOz--ZrOz with 0.5 N,
1.0 N, 2.0 N and 4.0 N sulfate ion loading
Table 1shows the specific surface area and the total pore volume for the Pt/SO:--ZrOz samples with
different sulfate ion loading. The sample with 1.0 N
sulfate ion catalyst obtained the largest surface area
and pore volume. This catalyst has a surface area of
118 m2/g and a pore volume of 0.127 ml/g. The addition of sulfate ion up to 4.0 N decreases the specific
surface area and the total pore volume to a great extent lower than that of 1.0 N sulfate ion concomitant
with the collapsing of the structure of the catalysts.
Table 1. Surface area and pore volume of the
prepared Pt/SOz--ZrOa samples
Sulfate ion loading
Surface area
Pore volume
0.5
116
0.116
1.0
118
0.127
2.0
38
10
0.049
4.0
0.016
249
Yori et al. suggested that all sulfate ions are at
the surface of ZrO2 [lo]. They proposed that the
sulfate species form at low coverage values (corresponding to samples with sulfate loading lower than
1.0 N sulfate ion), which correspond to isolated surface SO:- groups located in the crystallographic defective configuration (side terminations).
Sulfate species are formed up to the coverage
values (corresponding to samples with 1.0 N sulfate
ion), which correspond roughly to the completion of a
monolayer. These species are located on the patches
of low index crystal planes (top terminations of the
scale-like particles); these produce Lewis and some
Broensted acid sites.
When the sulfate concentration becomes higher
than an average monolayer, polynucleic surface sulfates appear, probably of the pyrosulfates (S20;-)
type, which are also mainly located on the regular patches of low-index crystal planes (top terminations), and originate from Broensted sites. This circumstance leads to a reduction of the surface area and
the pore volume for these samples.
Zalewski et al. reported that the most active catalyst system for isomerization was obtained when the
sulfate loading levels approached monolayer coverage
[ll]. This corresponds to a surface coverage of one
sulphur atom for every two zirconium atoms. The
Broensted acidity also maximizes at this level. The
concentration of Broensted acid sites increases as the
sulfate concentration approaches monolayer coverage.
The abrupt decrease in the specific surface area
for the higher sulfate species contents observed in Table 1 also correlates with the alteration of the crystal structure and the sulfate migration into the bulk
phase of the solid. At high sulfate ion levels, the special stabilization of the tetragonal form starts to diminish. Thus, at a sulfate loading level higher than
1.0 N, the sulfate ion monoclinic form begins to appear
PI
3.2. S t r u c t u r e of the sulfate ion
Figure 2 shows a peak at about 1395 cm-', which
is assigned to the asymmetric S=O stretching mode
of the sulfate groups bound by bridging oxygen atoms
to the surface. The S=O acts as an active acid site
where this species generating stronger acid sites [12].
With increasing sulfate ion content, the band at
1395 cm-' vanishes and shifts to a lower frequency,
which is because of the change in the type of the
sulfate species. Isolated structure of sulfate species
250
Sugeng Triwahyono et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
((ZrO)sS=O) may change to the bidentate and/or
polynucleic structure of sulfate species ((Zr0)zSOz)
~31.
1450
I400
1350
I300
I250
Wavenumber (cm-' )
Figure 2. IR spectra of the sulfur species stretching
region on Pt/SO:--ZrOz with 0.6 N, 1.0 N,
2.0 N and 4.0 N sulfate ion loading
3.3. Acid properties
Figure 3 shows ammonia T P D plots for Pt/SOt-ZrOz with different amounts of sulfate ion loaded.
At low temperature, the TPD plots consisted of two
peaks of ammonia adsorptions; the first peak at about
200 "C can be attributed to ammonia adsorbed over
the weak acid site, while the second peak at about
400 "C can be attributed to the adsorption of ammonia over the medium acid site. Pt/SOi--ZrOz with
0.5 N sulfate ion obtained the highest intensity for
weak acid site followed by 1.0 N, 2.0 N, and 4.0 N
sulfate ions. The increase in the amount of sulfate
ion decreases the intensity of the weak acid site and
develops a new peak at about 400 "C for 1.0 N and
new peaks at about 700-750 "C for 2.0 N and 4.0 N
sulfate ions.
f
.'
Zr-OH
+2.0 N
-f
-1.0N
D
--
tri-bridged OH S-OH S-OH
8
I
,
,
I
I
,
,
I
I
I
,
,
,
4J
N
-4.0
At high temperature, the peak appearing a t about
700-750 "C is because of the adsorption of ammonia on strong acid site. In this region, the catalyst
with 4.0 N sulfate ion obtained the highest amount
adsorbed followed by the catalyst with 2.0 N sulfate
ion. No-ammonia adsorption for catalysts with 0.5 N
and 1.0 N sulfate ions was observed at high temperature. This phenomenon can be explained that the
excessive amount of sulfate ion converted the weak
and medium acid sites to strong acid sites.
The phenomena of the decrease of weak and
medium acid sites and the generation of strong acid
site were also strongly evidenced by the results of IR
study. The spectral regions of interest on sulfated zirconia are the hydroxyl stretching regions (3000-3800
cm-l). Figure 4 shows the appearance of terminal
0 - H on zirconia (Zr-OH) and S-OH in the catalyst.
The addition of sulfate ion generates the hydroxyl
groups a t 3757 cm-' (Zr-OH) and 3652 cm-' (tribridged OH); further addition of sulfate ion causes
eroding of the Zr-OH and the tri-bridged OH groups
and develops new peaks assigned to the OH group
bonded to the sulfate species (S-OH) a t 3585 and 3548
cm-' . This phenomenon occurs because the surface
was covered by sulfate ion [12].
The physical adsorption of water on the catalyst
can be observed at about 1600 cm-'. Although it
is not shown in the present article, the intensity of
this peak decreases by the addition of sulfate ion accompanied by the decrease of the intensity of the
terminal and tri-bridged OH. This result suggests that
the hydroxyl groups are Broensted acid sites.
51
<
+ 0.5 N
I
,
I
,
,
I
I
I
,
I
I
,
,
,
I
,
l , , , , l , , , , ~ , , , , l , ' , , , l ' , , , , l , , , , l ~ , , ,
3800
100
200
300
400
500
600
700
800
900
Temperature (OC )
Figure 3. Ammonia T P D plot of Pt/SO:--ZrOa with
0.5 N, 1.0 N, 2.0 N and 4.0 N sulfate ion
loading
3750
3700
3650
3600
3550
3500
3450
3400
Wavenumber (6'
)
Figure 4. IR spectra in the Zr-OH and S-OH
stretching region on Pt/SO:--ZrOz
with
0.5 N, 1.0 N, 2.0 N and 4.0 N sulfate ion
loading
251
Special Column of the INRET 2006/Journal of Natural Gas Chernistr,y Vol. 15 No. 4 2006
Although it is not shown in the present article, the
TGA results explain the phenomena of weight loss,
which results from heating the Pt/SO:--ZrOz sample
in nitrogen flow. The curves for all samples exhibit
two distinct weight loss regions. The first weight loss
of about 0.15% occurs during heating up t o 550 "C
and this corresponds to water loss. The second weight
loss at high temperature is attributed to the decomposition of the sulfate species. The percentage weight
loss was about 28% for sample content 4.0 N sulfate
ion. The percentage weight loss decreases with the decrease of sulfate ion loading in the sample. After being
heated up to 950 "C, the percentage weight losses for
samples with 0.2 N, 0.1 N, and 0.5 N sulfate ion were
16%, 6%, and 5%, respectively.
The decomposition temperature shifted towards
low temperature with the increase of the sulfate ion.
The decomposition temperatures for 0.5 N, 1.0 N,
2.0 N, and 4.0 N sulfate ion are 650 "C, 538 "C, 524 "C,
and 517 "C, respectively. These results suggest that
the sulfate ion is more strongly held on the surface
at lower than at higher sulfate ion content because of
the formation of the monolayer structure of the sulfate
species.
polynucleic structure of sulfate species. The bridging bidentate structure of SO2 on the zirconia support ((Zr0)aSOa species) causes the strong inductive
effect on the Lewis acid site of Zr4+, which produces
an electronic deficiency markedly increasing the acidity of the zirconia cations. In the presence of hydrogen
atom, the Lewis acid sites are converted t o Bronsted
acid sites, which are required in the isomerization of
n-butane to iso-butane.
-
0.20 -
.
-(€
5
0.1s
1
& 0.10
-
0
2
D
%-
C
.0.05
1
c
0
"
I
"
,
~
'
~
"
'
'
"
'
"
~
"
'
'
It has been identified that sulfation of zirconia
causes the change in the surface area and the crystalline structure of zirconia and the additional sulfate
ion promotes the acidity and activity. Furthermore,
it is clear that the sulfate content of the catalysts depends on the quantity of sulphuric acid used for impregnation. It is seen that the sulphur content of the
catalysts increases with an increase in the quantity of
sulphuric acid solution used in the impregnation.
3.4. Isornerization of n-butane
Figure 5 shows the total conversion of n-butane
into iso-butane per specific surface area of the catalyst in the function of sulfate ion loading. The conversion of n-butane increases slightly with the increase
of sulfate ion up to 1.0 N, and further increase of sulfate ion increases the conversion of n.-butane linearly.
These results reveal that the acidity of catalysts roleplay the activity of catalysts towards the isomerization of n-butane directly. The excessive amount of
sulfate ion generates large number of strong acid sites,
which is favourable for isomerization. It is confirmed
by IR study, the shift of peaks in the range 1370-1400
cm-' to a lower frequency for 2.0 N and 4.0 N sulfate
ion loading catalysts may correspond to the transformation of isolated structure to the bidentate and/or
4. Conclusions
This research shows that the properties of sulfated
zirconia catalysts are greatly affected by the amount
of sulfate ion loading. The properties, particularly the
number and the strength of acid sites play an important role in the isomerization. For a limited range of
sulfate loading investigated in this study of sulfated
zirconia, it appears that the conversion of n-butane
per specific surface area of catalysts increases with
the sulfate ion loaded.
252
Sugeng Diwahyono e t al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
References
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151(2): 292
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[3] Hara S, Miyayama M. Solid State Zonics, 2004, 168(12): 111
[4] Ivanov A V, Vasina T V, Masloboishchikova 0 V,
Khelkovskaya-Sergeeva E G, Kustov L M, Houzvicka
J I. Catal Today, 2002, 73(1-2): 95
[5] Song X, Sayari A. Catal Rev Sci Eng, 1996, 38(3):
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[6] Yori J C, Parera J M. Appl Catal A: General, 1996,
147(1): 145
[7] Olah G A, DeMember J R, Shen J. J A m Chem Soc,
1973, 95: 4952
[8] Triwahyono S. In: Division of Molecular Chemistry
Graduate School of Engineering Hokkaido University.
2002
[9] Triwahyono S, Hattori H, YamadaT. In: Proceeding
of 13thSaudi-Japanese Catalyst Symposium Dahran.
Saudi Arabia, 2003
[lo] Yori J C, Parera J M. Appl Catal A: General, 1995,
129(2): L151
[ll] Zalewski D J, Alerasool S, Doolin P K. Catal Today,
1999, 53(3): 419
[12] Stevens J r R W, Chuang S S C, Davis B H. Appl Catal
A: General, 2003, 252(1): 57
[13] Tran M-T, Gnep N S, Szabo G, Guisnet M. Appl Catal
A, 1998, 171(2): 207
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Article
C0,-Free Hydrogen and Carbon Nanofibers Produced from Direct
Decomposition of Methane on Nickel-Based Catalysts
Siang-Piao Chai,
Sharif Hussein Sharif Zein,
Abdul Rahman Mohamed*
School of Chemical Engineering, Uniuersiti Sains Malaysia, Engineering Campus, Seri Ampangan,
14300 Nibong Tebal, SPS Penang, Malaysia
[Manuscript received October 18, ZOOS]
Abstract: Direct decomposition of methane was carried out using a fixed-bed reactor at 700 "C for
the production of C0,-free hydrogen and carbon nanofibers. The catalytic performance of NiO-M/SiOz
catalysts (where M=AgO, COO,CuO, FeO, MnO, and MOO)in methane decomposition was investigated.
The experimental results indicate that among the tested catalysts, NiO/SiOz promoted with CuO give
the highest hydrogen yield. In addition, the examination of the most suitable catalyst support, including
Alz03, CeOz, LazO3, SiOz, and TiOz, shows that the decomposition of methane over NiO-CuO favors
SiOz support. Furthermore, the optimum ratio of NiO to CuO on SiOz support for methane decomposition
was determined. The experimental results show that the optimum weight ratio of NiO to CuO fell at 8:2
(w/w) since the highest yield of hydrogen was obtained over this catalyst.
Key words: methane decomposition; hydrogen; carbon nanofibers; supported catalyst
1. Introduction
The utilization of natural gas, one of the world's
abundant resources, to produce valuable chemicals is
one of the desirable goals in the current natural gas
processing industry. In this regard, the production of
hydrogen from natural gas has attracted the interest
of industrialists and researchers. The conventional options of hydrogen production from natural gas, such
as steam reforming, partial oxidation, and autotherma1 reforming [l],involve CO, production at some
point in the technological chain of the process. Thus,
the gas needs further treatment to separate hydrogen
from the other gases such as CO and COz. Hence, it
is of importance to develop a simpler and less energy
intensive method t o produce hydrogen, which is free
of CO, formation as this can reduce the capital cost
in comparison to the conventional method. One way
to reach this objective is to use direct decomposition
of methane over catalytic materials for the production
of hydrogen [2-41.
* Corresponding author. Email: chrahmanOeng.usm.my
Direct catalytic decomposition of methane, the
main constituent of natural gas, offers the possibility
of producing two valuable chemical commodities:
pure hydrogen and carbon nanofibers. It is a technologically simple one-step process without energy and
material intensive gas separation stages and shows
the potential to be a C0,-free hydrogen production process. The hydrogen produced from methane
decomposition does not contain CO, impurities.
Undoubtedly, the produced hydrogen can be used
directly without any removal of CO,.
Carbon
nanofibers, a by-product, which are produced from
methane decomposition, are predominant mesoporous
materials with high surface area [4]. Such properties make these promising materials for applications
in the areas of catalysis, adsorption, and nanocomposites [5-71.
Recently, direct catalytic decomposition of
methane at lower temperature has received attention
as an alternative route to the production of hydrogen.
Several groups have studied methane decomposition
254
Siang-Piao Chai et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
at moderate temperature, such as, at 550 "C since
the catalysts possess high stability at this temperature [8-lo]. However, the hydrogen yield from
methane decomposition is low because of the thermodynamic limitation at temperatures below 700 "C [ll].
Decomposition of methane at higher temperatures has
recently attracted considerable attention because the
conversion of methane is higher [12]. However, catalysts lose their activity and stability at higher temperature (>700 "C) because of their quick deactivation.
It has been proved that the deactivation of catalyst
for methane decomposition is mainly caused by the
formation of encapsulating coke on the active sites of
the catalyst. Therefore, it is necessary to develop a
catalyst, which can resist its deactivation at higher
temperatures.
In the previous studies, it was reported that the
NiO/Ti02 catalyst is efficient for methane decomposition at high temperature with the calculated activation energy being one of the lowest ever reported in
the literature for this process [12]. In addition, the examination of the catalyst promoters for the NiO/TiOz
catalyst indicated that MnO, is the preferable promoter for the mentioned catalyst [13]. It was also
reported that the metal in oxide form provided better
conversion for methane during the methane decomposition reaction [13]. In continuation of this investigation, in this article, the different supported metal
oxide catalysts for methane decomposition are further
examined. This includes the investigations on the catalytic performance of NiO/SiO2 catalysts promoted
T
................................................
Gas mixer
h
Furnace
with other metal oxide species, the effect of catalyst
supports for NiO-CuO complex, the optimum weight
ratio of NiO t o CuO for methane conversion, and the
morphology of the carbon deposited on the used catalysts.
2. Materials and methods
Methane decomposition was studied using
Ni(N03)2*6HzO, Co(N03)~.6H20, F e ( N 0 3 ) 2 . H ~ 0
C U ( N O ~ ) ~ . ~Mn(N03)~.4H20,
H~O,
and (NH4)6Mo7024.4H20 as metal sources for the preparation of
the supported bimetallic catalysts. A1203, MgO,
CeOz, TiO2, La2O3, and SiOz (Cab-osil) were used
as catalytic supports in this study. All catalysts were
prepared using a conventional impregnation method
as reported elsewhere. The desired amounts of the
transition metal oxides were dissolved in deionized
water. The resulting paste was dried at 105 "C for
12 h and was calcined in air a t 600 "C for 5 h. The
catalysts were then sieved t o a size of 425-600 pm.
The production of C0,-free hydrogen and carbon
nanofibers via catalytic decomposition of methane
were carried out in a fixed-bed reactor a t 700 "C.The
detailed experimental setup and procedure have been
reported earlier [14]. The activity tests for the catalysts in the decomposition of methane were conducted
at atmospheric pressure in a stainless-steel fixed-bed
reactor (length and diameter of the reactor were 600
and 20 mm, respectively).
I1
11
f
......................... ....................... ............
0 %low
-
....... -Vent
Catalyst bed
Mass flow
controller
r:
I'
Reactor
B
Vent
Figure 1. Schematic diagram of the experimental apparatus
,
m
Data recorder
Special Column of the INRET 2006/Journal of Natural Gas Che1nistr.y Vol. 15 No. 4 2006
255
Packard Series 6890, USA). The schematic diagram
High purity methane (99.999% supplied by
of the experimental apparatus is shown in Figure
Malaysian Oxygen Sdn. Bhd.) was mixed with ar1. The carbon nanofibers deposited on the catalysts
gon (99.999% purity, supplied by Sitt Tatt Industrial
were analyzed using a transmission electron microGases Sdn. Bhd.) before entering the reactor and
scope (TEM) (Philips, CM12). During preparation
0.2 g of catalyst was put in the middle part of the
for the TEM experiments, a few samples of the spent
reactor for each run. The flow of methane was regcatalyst were dispersed in acetone (99.8% purity), and
ulated using a mass flow controller (MKS) and the
then a drop was deposited on a coated copper grid.
argon flow was regulated using mass flow controllers
The conversion of methane, and the hydrogen yield
(Brooks, model 58503). The product gases were anare defined as follows:
alyzed using an on-line gas chromatograph (HewlettMole of methane reacted
Methane conversion =
x 100%
Mole of methane input
Mole of hydrogen produced
Hydrogen yield =
Mole of metal oxide in a fresh catalyst
3.1. Effect of catalyst promoters
Table 1 shows the catalytic performance of the
NiO/Si02 catalyst promoted with different metal oxide species for methane decomposition at 700 "C. The
promoted catalysts are denoted as NiO-M/SiOz in
this article (where, M=AgO, COO,CuO, FeO, MnO,,
and MOO). The loadings of NiO-M were adjusted t o
10 wt.% with respect to the weight of the catalyst and
the weight ratio of NiO to M was set as 9:l. For all
catalysts, the methane decomposition proceeded selectively to form hydrogen as the only gaseous product. Methane conversion over the NiO/SiO2 catalyst
was high during the initial reaction. The conversion
decreased remarkably with time on stream, possessing
conversion of below 1%after 30 min of time on stream.
The addition of COO and MOO slightly improved the
catalytic lifetime of NiO/SiO:, for the methane decomposition. In contrast, the addition of Ago, FeO, and
MnO, shortened the catalytic lifetime of NiO/SiOa.
It was noticed that NiO/SiO2 promoted with CuO
showed significant increase in the catalytic activity
and the lifetime in methane decomposition. In the
Table 1. Methane conversions up to 90 rnin reaction
over 9NiO-1M catalysts supported on SiO2 at 700 "C
Catalyst
Methane conversion (%)
5 rnin
30 rnin
60 rnin
90 rnin
47.6
1.7
48.8
51.5
0.8
0.9
3.5
43.0
0.7
0.9
1.6
26.2
0.5
0.7
1.4
1.5
44.1
3.2
47.8
1.o
0.7
0.6
0.7
2.1
0.5
1.6
0.4
1.4
presence of CuO as the promoter for NiO/SiO2, the
stability and activity of the catalyst were enhanced
markedly.
Figure 2 shows the hydrogen yields over the tested
catalysts for 90 min of reaction. The yields were estimated from the obtained methane conversion, assuming that the methane conversion to hydrogen and
carbon proceeded stoichiometrically (CH4+C+2Hz).
The hydrogen yield for NiO/SiO2 was evaluated
as 140 rnolHz/mo1Nio2. The hydrogen yields for
9NiO-lCoO,9NiO-lFeO, and 9NiO-lMoO/Si02 were
157 molH,/m01(NiO-CoO) , 134 mOlHz /mol(NiO-FeO),
and 161 molHz/mol(NiO-MoO), respectively. However, NiO/SiOz promoted with A g o and MnO,
were less active in methane decomposition, giving hydrogen yields of 19 molHz/mol(NiO-AgO) and
14 molHz/mol~NiO-MnO,),respectively. According to
500
-
'5
-5
2
3
2
400
300
v
4
.-x
C
s
'r;
200
j
100
~~
NiO/SiOz
9NiO-lAgO/Si02
9NiO-lCoO/Si02
9NiO-lCuO/Si02
9NiO-lFeO/SiOz
9Ni0-1MnOZ/Si02
9NiO-lMoO/SiO?
0
(a)
(b)
(4
(4
(el
(f)
(9)
Figure 2. Hydrogen yields in the methane decomposition over NiOjSiOz and 9NiO-lM/Si02
catalysts at 700 'C
(a) NiO/SiOz; (b) 9NiO-lAgO/Si02; ( c ) 9NiO-lCoO/Si02;
(d) 9NiO-lCuO/Si02; (e) YNiO-lFeO/SiOz;
(f) 9Ni0-1MnOZ/SiO2; (9) 9NiO-lMoO/Si02
256
Takenaka et al. [lo], the addition of Cu into Ni/SiO2
lowered the hydrogen yield as compared to Ni/SiO2
itself for methane decomposition at 550 "C. In the
present study, the results shown in Table 1 and Figure 2 indicate that the addition of CuO into NiO/SiO2
improves the catalytic lifetime and the hydrogen yield
by more than three times that of NiO/SiO2, i.e.
494 rno1H2/rno1~Ni0-Cu0),
in the methane decomposition at a reaction temperature of 700 "C. It is likely
that the addition of CuO stabilized the NiO/SiOz by
forming CuO-NiO complex, which resulted in a long
catalytic lifetime and a high hydrogen yield in the
methane decomposition.
The hydrogen yields after 90 min of reaction decreased
in the order of SiOz>CeO2>TiO2>A1203>La203
supports with the obtained hydrogen yields of 494,
234, 77, 71, and 43 molHz/mol~NiO-CuO),
respectively.
500
5
h
-2
400
z
-
300
2
v
= 200
.K
u1
g
gc,
3.2. Effect of catalyst supports for bimetallic
100
catalysts
0
Table 2 shows the catalytic performance of 9Ni01CuO catalysts loaded on different supports. Many
oxides, such as A1203, CeO2, La203, Si02, and
TiO2, which are widely used as catalytic supports in
methane decomposition [10,12,16,17],were selected as
the supports in this study. For all the catalysts, the
loadings of NiO-CuO were adjusted to 10 wt.% with
respect to the weight of the catalyst. The initial activities of 9NiO-lCuO/La203 and 9NiO-lCuO/Ti02 for
the methane decomposition were low, having methane
conversions of 16.1% and 14.7%, respectively. After 30 min of reaction, 9Ni0-1Cu0 supported on
A1203, and La203 possessed methane conversions below 1%,revealing that both catalysts are not stable
in methane decomposition at this temperature. The
results shown in Table 2 indicate that the NiO-CuO
supported on Si02 is the most suitable for methane
decomposition, which showed higher catalytic activity
and stability in the methane decomposition.
(a)
(b)
(C)
(4
(e)
Figure 3. Hydrogen yields in the methane decomposition over 9Ni0-1Cu0 supported on
different oxides at 700 "C
(a) 9NiO-lCuO/A1203; (b) 9NiO-lCuO/Ce02;
(c) 9NiO-lCuO/La203; (d) 9NiO-lCuO/Si02;
(e) 9NiO-lCuO/Ti02
3.3. Effects of the weight ratio of active metals
to promoters
Figure 4 shows the changes of methane conversion
with time on stream over NiO-CuO/Si02 catalysts
with different weight ratios of NiO to CuO.
t9Ni0- lCuO/SiOz
8NiO-2CuO/Si02
7NiO-3CuO/SiOL
x 6Ni0-4Cu0/Si02
-Y- 5Ni0-5CuO/SiO2
-o- 4Ni0-6CuO/Si02
-m-
Table 2. Methane conversions up to 90 rnin reaction
over 9Ni0-1Cu0 on different supports at 700 "C
Catalyst
9NiO-lCuO/AI203
9NiO-lCuO/Ce02
9NiO-lCuO/La203
9NiO-lCuO/Si02
9NiO-lCuO/Ti02
Methane conversion (%)
5 min
26.1
42.3
16.1
51.5
14.7
30 min
60 min
0.6
23.2
0.3
43.0
5.2
0.3
0.6
26.2
2.0
90 min
0.6
1.5
0.6
The amounts of hydrogen produced over the
tested catalysts after 90 min of reaction are summarized in Figure 3. The hydrogen yield was the highest
for the NiO-CuO/Si02 catalyst among all the s u p
ported 9Ni0-1Cu0 catalysts examined in this study.
0
30
60
90
120
150
180
Time on stream (min)
Figure 4. Kinetic curves of methane conversions as
a function of time on stream over 9Ni0lCuO/SiOZ catalysts with different weight
ratios of NiO to CuO at 700 "C
The hydrogen yields in the methane decomposition were estimated from the kinetic curves of
methane conversion and are shown in Figure 5. It was
noticed that the addition of CuO into the NiO/SiO2
catalyst with the weight ratio of Ni0:CuO ranging
Special Column o f the I N R E T 2006/Journal o f Natural Gas chemistr.y Vol. 15 No. 4 2006
from 9:l to 5:5 remarkably prolonged the catalytic
lifetime. However, the addition of excess amounts
of CuO (weight ratio NiO:Cu0<4:6) decreased the
catalytic lifetime for the methane decomposition. As
shown in Figure 5, the hydrogen yield increased significantly with an increase in CuO loading into the
NiO/SiO2 catalyst and the yield attained the maximum at NiO:Cu0=8:2. The maximum hydrogen
yield was 536, which was about 4 times greater than
that of the NiO/SiOz catalyst. The hydrogen yields
obtained for the catalysts with the weight ratio of
Ni0:CuO as 9:1, 7:3, 6:4, 5:5, and 4:6 are 494, 452,
272, 208, and 128 molHz/rnol(NiO-CuO),respectively.
This shows that the increase of copper content improved the catalyst stability, but the continual increase of copper content resulted in a great decrease
of the catalyst stability.
600
01
lOi0
I
1
I
I
I
I
911
812
713
614
515
416
257
tization than those on the 8Ni0-2CuO/Si02 catalyst. Owing to the high activity and stability of
the 8Ni0-2CuO/Si02 catalyst in methane decomposition, the grown carbon nanofibers are denser, forming an interwoven coverage under TEM observation.
It is important to note that catalyst particles were
present on the tips of carbon nanofiber grown on the
NiO/SiO2 catalyst, whereas no catalyst particles were
observed on the tips of carbon nanotubes formed by
the 8Ni0-2CuO/Si02 catalyst. This shows that the
growth model of carbon nanofibers on these two catalysts was different. It is deduced that the carbon
nanofiber grown by the NiO/SiO2 catalyst followed
the tips-growth model where the catalyst particles located on the tips of the carbon nanofibers were active for methane decomposition. On the other hand,
the based-growth model was used for growing carbon
nanofibers on the 8Ni0-2CuO/Si02 catalyst, where
the catalyst particles attached on the Si02 support
were active for methane decomposition.
J
317
Ratio ofNiO / CuO
Figure 5. Hydrogen yields in the methane decomposition over NiO-CuO/Si02 catalysts with
different weight ratios of NiO to CuO at
700 "C
3.4. Characterization of used catalysts
Figure 6 shows the TEM images of carbon deposited by the methane decomposition over NiO/SiOz
and 8Ni0-2CuO/Si02 catalysts. It was found that
carbon nanofibers were formed on the surfaces of
both catalysts. However, there are differences in
the morphology of the grown carbon nanofibers. It
was noticed that the average diameter of the carbon
nanofibers grown on the 8Ni0-2CuO/Si02 catalyst
were clearly larger than those on the NiO/SiOz catalyst. F'urthermore, carbon nanofibers on the 8Ni02CuO/SiO2 catalyst seem more fragile when compared to those formed on the NiO/SiO2 catalyst.
This reveals that the carbon nanofibers grown on the
NiO/Si02 catalyst have a higher degree of graphi-
Figure 6. TEM image of carbon nanofibers produced
on (a) NiO/SiOz, and (b) SNi0-2CuO/Si02,
in methane decomposition at 700 "C
4. Conclusions
A comparison of the catalytic properties of the
NiO/SiO2 catalyst promoted with Ago, COO, CuO,
258
FeO, MnO,, and MOO showed that, the addition of
CuO increased the hydrogen yield and prolonged the
lifetime of bimetallic NiO-CuO catalysts in methane
decomposition a t 700 "C. The examination of the catalyst supports indicated that SiOz is the best catalyst
support tested for NiO-CuO. The highest hydrogen
yield, being 536 molHz/mol(NiO-CuO),was obtained
using this catalyst with the weight ratio of Ni0:CuO
as 8:2 on the SiOz support. With this weight ratio,
the NiO-CuO alloys formed on SiOz possessed high
stability in methane decomposition even though the
react,ion was conducted at high temperature. It is important to note that only hydrogen was detected by
the GC as gaseous product. Analysis of the deposited
carbons by TEM indicated that carbon nanofibers
were grown on the surfaces of NiO/SiOz and NiOCuO/SiOz catalysts. These carbon structures have
great potential applications as catalyst supports and
nanocomposites.
Acknowledgements
The authors gratefully acknowledge the financial support provided by the Academy of Sciences Malaysia under scientific Advancement Grant Allocation (SAGA)
(Project: A/C No. 6053001).
References
[l] Islam M N, Dixon M. AIChE J , 2003: 547
[2] Zhang T, Amiridis M D. Appl Catal A : Gen, 1998,
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[3] Ermakova M A, Ermakov D Yu, Kuvshinov G G. Appl
Catal A: Gen, 2000, 201: 61
[4] Avdeeva L B, Reshetenko T V, Ismagilov Z R,
Likholobov V A. Appl Catal A: Gen, 2002, 228: 53
[5] Reshetenko T V, Avdeeva L B, Ismagilov Z R,
Chuvilin A L. Carbon, 2004, 42: 143
[6] Shaikhutdinov, Sh K, Avdeeva L B, Novgorodov B N,
Zaikovskii V I, Kochubey D I. Catal Lett, 1997, 47:
35
[7] Madronero A, Ariza E, Verdu M, Brand1 W , Barba
C. J Mater Sci, 1996, 31: 6189
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167: 161
[9] Ermakova M A, Ermakov D Y, Chuvilin A K, Kuvshinov G G . J Catal, 2001, 201: 183
[lo] Takenaka S, Kobayashi S, Ogihara H, Otsuka K. J
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[ll] Parmon V N, Kuvshnov G G, Sadykov V A, Sobyanin
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[12] Zein S H S, Mohamed A R, Sai P S T. Ind Eng Chem
Res, 2004, 43: 4864
[13] Zein S H S, Mohamed A R. Energy Fuels, 2004, 18:
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[14] Chai S P, Zein S H S, Mohamed A R. Chem Phys Lett,
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[16] Reshetenko T V, Avdeeva L B, Ismagilov 2 R,
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Available online at w.sciencedirect.com
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Journal ofNatural Gas Chenuslry
SCIENCE PRESS
www.elsevier.codocateljngc
Article
Ethylene Conversion to Higher Hydrocarbon over Copper
Loaded BZSM-5 in the Presence of Oxygen
Ramli Mat1*,
Nor Aishah Saidina Aminl,
Zainab Ramli2,
W. Azelee W. Abu Bakar2
1. Dept. of Chemacal Engzneenng, Faculty of Chemzcal &Natural Resources Engzneerzng, Unaversata Teknologz Malaysaa, Skudaa,
81 310 Johor Bahru, Malaysza; 2. Dept. of Chemzstry, Unaversatz Teknologz Malaysza, Skudaz, 81 31 0 Johor Bahru, Malaysaa
Abstract: The successful production of higher hydrocarbons from methane depends on the stability
or the oxidation rate of the intermediate products. The performances of the BZSM-5 and the modified
BZSM-5 catalysts were tested for ethylene conversion into higher hydrocarbons. The catalytic experiments
were carried out in a fixed-bed micro reactor at atmospheric pressure. The catalysts were characterized
using XRD, NH3-TPD, and IR for their structure and acidity. The result suggests that BZSM-5 is a weak
acid. The introduction of copper into BZSM-5 improved the acidity of BZSM-5. The conversion of ethylene
toward higher hydrocarbons is dependent on the acidity of the catalyst. Only weaker acid site is required
to convert ethylene to higher hydrocarbons. The loading of Cu on BZSM-5 improved the selectivity for
higher hydrocarbons especially at low percentage. The reactivity of ethylene is dependent on the amount
of acidity as well as the presence of metal on the catalyst surface. Cul%BZSM-5 is capable of converting
ethylene to higher hydrocarbons. The balances between the metal and acid sites influence the performance
of ethylene conversion and higher hydrocarbon selectivity. Higher loading of Cu leads to the formation of
co,.
Key words: ethylene conversion; BZSM-5 zeolite; acidity; higher hydrocarbon
1. Introduction
The conversion of methane to higher hydrocarbons has been studied in detail; the initial formation of ethylene and its subsequent conversion to long
chain hydrocarbons is considered a possible mechanism. The conversion of methane to ethylene followed by the processing of ethylene over ZSM-5 is
an efficient and flexible route for the production of
synthetic hydrocarbons from either natural gas. Depending on the process conditions, the products can
be produced ranging from gasoline to distillate fuels [1,2]. The transformation of light alkenes over zeolites catalysts is of importance in the various petrochemical processes such as met hane-olefin-gasolinedistillate (MOGD) and methane to gasoline (MTG).
The mechanism of the reaction between methane
and oxygen to produce higher hydrocarbons over zeolite is postulated to start from the formation of methyl
radical from CH4 [3,4]. The methyl radicals combine to form ethane, which dehydrogenate to ethylene. Oligomerization and aromatization of ethylene
will produce higher hydrocarbons such as aromatics or
liquid fuels. However, ethylene may easily oxidize to
CO,. Therefore, the successful production of higher
hydrocarbons depends on the stability or the oxidation rate of the intermediate products. The catalysts
selected must have the ability to control the oxidation and must be able to oligomerize the intermediate
products.
Acidity is one of the most important characteristics of zeolites, which makes these extremely important materials in catalytic applications. The acidity of
zeolites is known t o depend on several factors: struc-
* Corresponding author. Tel: 607-5535567; Fax: 607-5581463; E-mail: ram1iQfkkksa.utm.my
260
Ramli Mat et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
ture, preparation method, chemical composition, impurities, Si/Al ratio, additives, and poisons. Considerable attention has been given to the relationship
between the acidity of zeolites and their catalytic activity. Successful production of C2+ from methane is
dependent on the zeolite acidity [5,6]. The synthesized BZSM-5 and the modified BZSM-5 present very
different acidity properties and consequently different
catalytic properties as reported in refs [7-91. Thus, it
is of considerable interest to compare the catalytic behaviors of the BZSM-5 and the copper loaded BZSM-5
zeolites for the transformation of ethylene. Ethylene,
which is the intermediate product of methane activation, undergoes subsequent oligonierization and aromatization on Bronsted acid sites of zeolites to form
higher hydrocarbons. Thus, the higher reactivity of
0 2 with C2H4 as compared to CH4 has been a major
problem while achieving higher selectivity for higher
hydrocarbons. In this context, the possible inhibition
of the oxidation of C2+ is of special interest. With the
aim of identifying the importance of ethylene as a n intermediate species for the oxidative dehydrogenation
of methane to higher hydrocarbons, ethylene has been
used as a probe performing several catalytic tests.
Ethylene was used as it is assumed to play the intermediate role in methane conversion to higher hydrocarbons. This study provides information on C2+
stability over these catalysts. These data are relevant
t o the design of the catalyst for direct methane conversion to higher hydrocarbons, since it is crucial that
the catalyst identified does not catalyze the decomposition of C2+ to CO,. Obviously, it is important
that C2+ must be stable over any potential catalyst,
under appropriate conditions.
2. Experimental
2.1. Catalyst preparation
BZSM-5 was first synthesized by Taramasso
et ad., [lo] using organic compounds of silicon and
tetrapropylammonium cation as the template. Recently, some publications in the open as well as in
the patent literature have dealt with boron-containing
pentasil type zeolite. The BZSM-5 was prepared following the procedure described by Plank et d. [12].
Generally, a gel was prepared by mixing: 69 g of
sodium silicate (Merck: 25.5%-28.5% Si02, 7.5%8.5%Na20), 0.53 g of boric acid (Merck), 7.53 g of
Tetra propyl-ammonium bromide (TPA-Br) as the
template (Fluka) and 160 g of distilled water to form
a reaction mixture with the following molar composition B203: 20Na20:70Si02:7TPABr:22OOH20.The
pH of the reaction mixture was maintained around
10-12 by the addition of sulfuric acid. The gel was
stirred a t room temperature for 3 h in a one-liter
stainless steel autoclave (PARR reactor). The gel wits
then heated in the oven at 160 "C for 5 days without
stirring. The TPA ion, present in the solution, was the
key reagent favoring the formation of the tetrahedral
units, which in turn played a key role in determining
the characteristic structure of the BZSM-5. The template was removed through calcination to yield the
template-free product.
Once the zeolite was formed (crystalline white
solid), it settled at the bottom of the autoclaves,
leaving a clear supernatant liquid. The crystalline
white solid product was filtered, washed thoroughly
with deionized water, and dried at 120 "C for 12 h.
The resultant material was calcined at 550 "C for five
hours t o remove the organic material and to obtain
the sodium form of the BZSM-5, Na-BZSM-5. The
Na-form obtained was converted into the NH4-form
by ion exchange using 1M solution of ammonium nitrate, NH4N03. For every l g of NaBZSM-5, it was
treated with 25 ml of 1M NH4N03 solution, and was
stirred under reflux for three hours at 80 "C. The procedure was repeated thrice, ending with refluxing the
solution for 12 h. Finally, the catalyst was dried and
calcined at 550 "C for five hours.
Recently, it has been found that the introduction
of copper can remarkably increase the activity of catalyst in methane conversion [6]; therefore, the effect
of copper needs to be investigated. The incorporation of copper into the calcined BZSM-5 was carried
out through the impregnation method. The BZSM-5
obtained was impregnated with copper nitrate solution to provide lwt% , 4wt% , and 9wt% of copper and was labeled as Cul%BZSM-5, Cu4%BZSM-5,
and Cu9%BZSM-5, respectively. The solid was then
calcined in the furnace at 550 "C for five hours.
2.2. Catalyst characterization
Three different types of characterizations were
performed in this study, namely, (i) X-ray diffraction
(XRD), (ii) infrared spectroscopy (IR), and (iii) acidity measurement. The XRD patterns were acquired
on a Siemens D5000 goniometer using CuK, radiation in the range of 20 from 2O to 60" a t a scanning
speed of 3 O p e r minute. The IR spectra were examined with a Shimadzu 3000 FT-IR spectrometer using
261
Special Column of the I N R E T 2006/Journal of Natural Gas Chernistr,y Vol. 15 No. 4 2006
the KBr wafer technique. The samples (0.25 mg of zeolite powder) were mixed with 300 mg of KBr powder
and were finely ground. These mixtures were placed
on a die and were pressed to make a transparent thin
pellet. The IR spectra in the range of 2400-400 cm-'
were recorded at room temperature.
For the acidity measurement, it is necessary t o
determine the amount, strength, and type of the acid
sites. In this study, three types of techniques, namely,
temperature programmed desorptiori (TPD) of ammonia, IR spectroscopy for the hydroxyl region, and
IR spectroscopy of adsorbed pyridine, were used to
evaluate the acidic properties of the catalyst. The
amount and strength were determined using temperature programmed desorption (TPD) of ammonia,
while the type of acid site information was obtained
using IR spectroscopy for the hydroxyl region and IR.
spectroscopy of adsorbed pyridine.
2.3. C a t a l y s t t e s t i n g
Conversion of ethylene
Selectivity for
=
c2-c4
=
Selectivity for CO,
=
The performances of the BZSM-5 and the
modified BZSM-5 catalyst were tested for ethylene
conversion into higher hydrocarbons. The catalytic
experiments were carried out in a fixed-bed micro reactor a t atmospheric pressure.
Gases of ethylene, compressed air and nitrogen
were supplied from individual gas cylinders. The
reactor was preheated at a reaction temperature of
800 "C under nitrogen flow for two hours to activate
the catalyst. Ethylene (purity 99.9%) and compressed
air were then fed into the reactor with 9% volume
of oxygen in the feed. The total feed flow rate was
200 ml/min. The catalyst weight used in this study
was 1 g. The reaction products were analyzed using
an on-line gas chromatograph. The GC analysis was
carried out using the thermal conductivity detector
(TCD) equipped with Porapak packed column.
The conversion of ethylene and the selectivity for
the higher hydrocarbons were determined according
to the following equations:
moles of C2H4 reacted x 100%
moles of CaH4 in feed
moles of hydrocarbon gas produced other than C2H4 x 100%
moles of C2H4 reacted
moles of CO and COz gas produced x 100%
moles of C2H4 reacted
3.1. C a t a l y s t characterization
Figure 1 shows the diffractogram of the
BZSM-5 impregnated with different loadings of copper ranging from lwt% to 9wt%. No significant
difference was found between the diffractogranis of
the parent zeolite (BZSM-5) and the catalysts after
impregnation with copper. All samples showed similar pattern and were highly crystalline. However, one
peak at 28=38.6', which is characteristic of CuO, was
detected for the higher loading of Cu (above 4%).
These results were consistent with the findings by
Torre-Abreu et al. [12], where the CuO peak was observed when the HZSM-5 was loaded with 5.5wt% Cu.
They reported that in the HZSM-5 with low copper
loading, the copper was mainly present in the form of
isolated Cu2+ ions. On the other hand, in the catalyst with high copper loading, CuO, isolated Cu2+
ions and also Cuf ions were detected using H2-TPR
and ESR. It was also verified that the concentration
of the CuO species increases when the catalyst copper
loading increases, which probably results in the formation of CuO aggregates. Nunes et al. [13] reported
similar findings in their study on the effect of copper
loading on the acidity of the Cu/HZSM-5 catalyst.
.$
+.
-
P)
J'
10
5 Cu4%BZSM-5
20
30
40
50
28/(" )
Figure 1. Effect of the XRD pattern on the different
copper loadings on BZSM-5. CuO peak was
observed at higher Cu loading
262
Ramli Mat et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
The NH3-TPD results are shown in Table 1. The
addition of lwt% copper into BZSM-5 resulted in the
increase of acidity; however, the acidity decreased by
about 30% in Cu4%BZSM-5 and Cu9%BZSM-5 catalysts as compared to the BZSM-5. This may be due to
the charge imbalance imposed by copper as proposed
by Tanabe [14]. A similar observation was reported
by Ismail et al., [15] and Anggoro [16] in their study
of Cu loaded into HZSM-5. When Cu was loaded into
the HZSM-5, it produced additional acid sites. However, when Cu was further loaded (i.e. 4% and 9%),
the total amount of acid seemed to diminish from the
catalyst surfaces as tabulated in Table 1. There is a
possibility that some acid sites were lost during further loading of metal because of the blockage of the
zeolite pore and the formation of the Cu cluster as
seen using XRD [13,17].
clearly shown in Table 2; this result suggests that
BZSM-5 has more terminal silanol groups or silanal
at the defect site. The silanol groups are recognized
as neutral or very weak acid. The silanol group existing on the external zeolites is considered a weak acid,
however, it showed a catalytic role on Beckmann rearrangement and also some reaction with olefin as
reported in the refs [18,19].
36,lO
Table 1. The results of the NHa-TPD experiments
Sample
BZSM-5
Cul%BZSM-5
Total amount of chemisorbed (mmol/g)
Cu4%BZSM-5
CuS%BZSM-5
1.619
1.426
2.302
For a complete characterization of the zeolite acidity, it is necessary to determine the nature and the
concentration of both Bronsted and Lewis acid sites.
This can be carried out by examining the hydroxyl
group region (3500-3750 cm-l) and the pyridine adsorption of IR spectroscopy. The infrared spectra
of the samples recorded at 303 K in the range of
OH stretching vibrations are shown in Figure 2. For
the HZSM-5 sample, the bands are observed at 3740,
3660, and 3610 cm-'. The band at 3740 cm-' is
attributed to the stretching of the terminal silanol
(SiOH) groups located at the boundaries of the zeolite crystal. The silanol groups are considered a
weak acid. The band at 3660 cm-l is assigned to
OH groups associated with extraframework Al, while
the band at 3610 cm-I is attributed to the stretching
vibration of the bridging hydroxyl groups associated
with framework aluminium (Si-OH-Al). This band is
normally more acidic; this hydroxyl group has acidic
properties since the band is not observed when pyridine is adsorbed on the zeolite.
In the case of the BZSM-5 sample, the stretching
vibration band at 3610 cm-l is due to the bridging
hydroxyl groups having lower intensity. Furthermore,
the band area for silanol groups was larger for BZSM-5
when compared to that of HZSM-5. The result is
-------
x-
7.867
I
4000
3800
,
I
,
,
,
#
I
,
,
,
3600
,
I
3400
I
,
,
,
3200
Wavenumber (cm-' )
Figure 2. IR spectra in the hydroxyl region of
the samples after dehydration at 400 "C,
lope MPa for 2 h
Table 2. Position of some characteristic OH vibration
bands and the integrated band area
Sample
Wavenumber (cm-l )
Al-non
Al-framework
Silanol
framework
HZSM-5
3610
(Integrated area)
BZSM-5
(Integrated area)
(7.21)
3611
Cul%BZSM-5
(Integrated area)
3661
(0.55)
3741
(0.09)
3658
3741
3610
(0.65)
3661
(0.21)
3741
(5.89)
(0.64)
(0.19)
(1.95)
For copper loaded on BZSM-5, the intensity of
the hydroxyl band at 3610 cm-' was slightly higher
but was still lower than that of HZSM-5. Thus, it can
be suggested that the acidity of BZSM-5 improved
with the introduction of copper into BZSM-5. This
result supports the earlier data from the NH3-TPD
analysis, which indicates that adding Cu species 0 1 1
BZSM-5 improved the acidity of BZSM-5.
The types of acid that were present in the zeolite
sample were further characterized using the pyridine
adsorption method. Pyridinium ion signals (pyridine
on Bronsted acid sites) appeared at wave numbers of
1638 and 1546 cm-' and pyridine on Lewis a,cid sites
Special Column of the INRET 2006,lJournal of Natural Gas Chemistry Vol. 15 No. 4 2006
appeared at 1450 cm-' [20]. The concentration of
the Bronsted acid sites and the Lewis acid sites were
calculated from the integrated area of the bands at
1540 and 1450 cm-' according to the formula proposed by Hughes and White [all. The results are
tabulated in Table 3. The acidity increased when
BZSM-5 was loaded with copper. More Lewis acid
sites were generated when more copper was loaded
on BZSM-5. The ratio of the Bronsted to Lewis acid
site considerably reduced a t higher loading of copper.
Two factors can explain this result: First, at higher
copper loading, the Bronsted acid sites were partly
covered by CuO. Second, part of the Cu species were
exchanged with the Bronsted acid sites that were capable of transforming into Lewis acid sites. Similar
results were reported by Wang et al. [22] in their
studies of different Mo loading on HZSM-5. They
claimed based on the IR and NH3-TPD studies that
the number of Bronsted acid centers decreased and
the Lewis acid centers increased after more Mo loading into the HZSM-5 zeolite.
Table 3. Concentration of the Bronsted and Lewis
acidity of the samples
Acidity (pmol/g)
Sample
BZSM-5
Bronsted(B)
(at 1545 cm-')
Lewis(L)
(at 1450 cm-')
Ratio
B/L
5
0
-
Cul%BZSM-5
105
8
13
Cu4%BZSM-5
6
40
0.15
The results of the various catalytic performances
for the reaction of ethylene and oxygen at 800 "C and
atmospheric pressure are given in Figure 3 . This
figure shows the ethylene conversion and the carbon selectivity towards CO,, c2&4 (exclude C2H4),
and Cs+(Cs hydrocarbon and above) for the catalyst
tested. BZSM-5 was slightly less effective for ethylene
conversion as compared to Cul%BZSM-5. The acidity of BZSM-5 is lower than Cul%BZSM-5. Thus, the
slightly lower activity of BZSM-5 may be due to the
lower acidity as explained by Sohn and Park [2]. Furthermore, the copper loaded in zeolite may act as a
bifunctional catalyst and is widely used in hydrocarbon conversion. On this catalyst, the transformation
of hydrocarbons involves both of the hydrogenation
and the dehydrogenation step on metal sites and the
rearrangement and/or the cracking step on acid sites.
Indeed, it has been shown very clearly that the balance between metal sites and acid sites remarkably
263
influenced the performance of the bifunctional catalysts in hydrocarbon conversion [23,24].
100
.-
80
-28
60
g
0
.->
r
DConversion
!ZZZlCO, selectivity
C,-C, selectivity
C5. selectivity
Q
E
.-2
e
40
>
u0
20
0
BZSMJ
Cul%BZSMJ Cu4YoBZSM-5 CU~YOBZSMJ
Figure 3. Catalytic performance of the ethylene reaction with 9 vol% of oxygen over BZSM-5
and modified BZSM-5 with copper catalyst
at reaction temperature of 800 "C and gas
hourly space velocity (GHSV) of 8000 h-'
under atmospheric pressure
The results show that copper loading a t low concentration promotes ethylene conversion; however, a t
higher copper loading, the ethylene conversion decreases. The difference could be interpreted in terms
of copper deposited over the BZSM-5 zeolite. At
1 wt% of copper loading, copper oxide may not be
totally deposited over the surface of the BZSM-5 zeolite. Many active sites of BZSM-5 zeolite remain vacant and are responsible for ethylene activation. At
high copper loading, the ethylene conversion decreases
because of the deposition of the metal species over the
acid sites together with the blocking of the channels.
BZSM-5 is seen to have reasonable activity but
slightly lower selectivity towards higher hydrocarbons
since mainly carbon oxides are formed. The selectivity towards C Z - C ~for all the catalyst was between
45% up to 66%, where BZSM-5 shows the highest
C2-C4 selectivity. On the other hand, Cul%BZSM-5
shows the lowest C Z - C ~selectivity. Furthermore,
Cul%BZSM-5 shows the lowest selectivity towards
carbon oxides when compared to higher loading of
copper on BZSM-5. These results were consistent
with the result obtained by Min and Mizuno [25] in
their study on the effects of copper additives on the selective oxidation of light alkanes. Min and Mizuno [25]
demonstrated that the addition of copper enhanced
the catalytic performance for oxidation light alkanes
under oxygen-poor condition to CO,.
Based on the above results, the only way to relate
the activity with the acidity of zeolite is by consid-
264
RarnJi Mat et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
ering the amount of acid of the catalyst. Figure 4
shows the effect of acidity on the Cg+ selectivity for
ethylene conversion. Cul%BZSM-5, which has more
acid sites, was observed to provide a higher conversion of ethylene and higher hydrocarbon selectivity
than the other catalysts. Following this observation,
the amount of acid can be used to explain the activity of the catalyst. This finding is in-line with
the result reported by Guisnet et al. [26], where the
oligomerization process was dependent on the zeolite
acidity. Nor Aishah and Anggoro [4] reported that
the oligomerization of ethylene was dependent on the
zeolite acidity. As the number of the acid sites increased, the oligomerization reaction increased and
thus more Cg+ were produced. However, the amount
of acid is not the only cause of the catalyst being
active, rather the presence of copper ions may also
affect its activity as in case of Cu4%BZSM-5. It is
seen from Figure 4 that the acidity of Cu4%BZSM-5
is lower than that of BZSM-5; however, its selectivity
towards Cg+ is higher. Only weak acid site is required t o activate ethylene or olefin to gasoline range
hydrocarbon [27]. Kitigawa et al. [28] studied the
requirement of acid site in the C-H bond activation
on weakly acidic borosilicates with ZSM-5 structures.
He reported t,hat Zn exchanged borosilicates did not
catalyze propane dehydrocyclodimerization but these
catalysts could convert propene t o aromatics at a significant rate.
50
I
Acidity
1'1
1.619
BZSM-5
4. Conclusions
The conversion of ethylene toward higher hydrocarbon is dependent on the acidity of the catalyst.
Only weaker acid site is required to convert ethylene
to higher hydrocarbons. Loading of Cu on BZSM-5
improves the selectivity for higher hydrocarbons, especially at lower percentage. The reactivity of ethylene is dependent on the amount of acidity as well
as the presence of metal on the catalyst surface.
Cul%BZSM-5 is capable of converting ethylene to
higher hydrocarbons. The balances between metal
and acid sites influence the performance of ethylene
conversion and higher hydrocarbon selectivity. Higher
loading of Cu leads to the formation of CO,.
Acknowledgements
The authors wish to thank Malaysian Ministry of Science, Technology and Environment for funding this work
through IRPA project No. 02-02-06-0031.
0
C,, selectivity
,861
Biscardi and Iglesia [31] reported that the introduction of Ga, Zn, or Pt species into zeolites increased
the rate and selectivity for aromatization reactions
and inhibited cracking side reactions that led to the
loss of carbons to undesirable products. They explained that the oligomerization and cracking of light
alkenes occurred readily on acid sites. Light alkenes
can be converted to a mixture of higher molecular
weight alkenes via a sequence of acid-catalyzed shapeselective oligomerization reaction over zeolites. Furthermore, aromatization of the product formed required a concerted reaction between the acid and
metal cation sites.
I .426
CUI YoBZSM-5 Cu4YoBZSM-5 Cu9%BZSM-5
Figure 4. Effect of the amount of acid (mmol/g) on
the C5+ selectivity for ethylene conversion
Introducing metal into the zeolite resulted in an
increase of the conversion and selectivity. This finding
was in agreement with the finding by Nishi et al. [29],
Liu et al. [30], and Biscardi and Iglesia [31]. They
claimed that introducing the metal component t o the
zeolite catalyst will improve the product selectivity.
References
[l] Heveling J, Nicolaides C P, Scurrell M S. Appl Catal
A , 1998, 173(1): 1
[2] Sohn J R, Park W C. Appl Catal A , 2002, 230(1-2):
11
[3] Han S, Maternak D J, Palerrno R E, Pearson J A ,
Walsh D E. J Catal, 1994, 148(1): 134
[4] Nor Aishah Saidina Arnin, Didi Dwi Anggoro.
J Natur Gas Chem, 2002, 11: 79
[5] Ernst S, Weitkamp J. In: Irnarisio G, Frim M, Bemtgen J M eds. Hydrocarbons Source of Energy. London: Graham and Trotman, 1989. 461
[6] Nor Aishah Saidina Amin, Didi Dwi Anggoro.
J Natur Gas Chem, 2003, 12: 123
[7] Liu H, Ernst H, Freude D, Scheffler F, Schwieger W.
Macroporous Mesoporous Mater, 2002, 54(3): 319
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[8] Beyer H K, Borbely G. In: Murakami Y, Iijima A,
Ward J W eds. New Developments in Zeolite Science
and Technology. Elsevier: Amsterdam, 1986. 867
[9] Ratnasamy P, Hedge S G, Chandwadkar A J. J Catal,
1986, 102: 467
[lo] Taramasso M, Perego G, Notari B. In: Rees L V eds.
Proc. 5th Intern. Conf. Zeolites, Heyden & Son, London, 1980. 40
[ll] Plank C J, Rosinski E J, Schwartz A B. UK Patent
1402981, 1974
[12] Torre-Abreu C, Ribeiro M F, Henriques.C, Delahay
G. Appl Catal B, 1997, 12(2-3): 249
[13] Nunes M H 0, d a Silva V T, Schmal M. AppZ Catal
A , 2005, 294(2): 148
[14] Tanabe K. Solid Acids and Bases: Their Catalytic
Properties. New York: Academic Press, 1970
[15] Saaid I M, Mohamed A R, Bhatia S. J Mol CataZ A ,
2002, 189(2): 241
[16] Anggoro D D. [Ph. D Thesis], Modified Oxidative
Coupling Process of Methane to Liquid Fuels Over
Metal Loaded ZSM-5 Catalyst. Skudai: Universiti
Teknologi Malaysia. 2003
[17] Choudhary V R, Mantri K, Sivadinarayana C. Mzcroporous Mesoporous Materials, 2000, 37(1-2): 1
[18] Roseler J , Heitmann G, Holderich W F. Appl Catal,
1996, 144(1-2): 319
265
[19] Kondo J N, Yoda E, Ishikawa H, Wakabayashi F,
Domen K. J Catal, 2000, 191(2): 275
[20] Tynjala P, Pakkanen T T . J Mol Catal, 1996, llO(2):
153
[21] Hughes T R, White H M. J Phys Chem, 1967, 71:
2192
[22] Wang J, Karig M, Zhang Z, Wang X. J Natur Gas
Chem, 2002, 11: 43
[23] Wang J, Li Q Z, Yao J D. AppZ Catal A , 1999, 184(2):
181
[24] Girgis M J, Tsao Y P. Ind Eng Chem Res, 1996, 35:
386
[25] Min J S, Mizuno N. Catal Today, 2001, 71(1-2): 89
[26] Guisnet M, Gnep N S, Vasques H, Ribeiro F R. In:
Jacobs P A, Jaeger N I, Kabelkova L, Wichterlova B
eds. Zeolite Chemistry and Catalysis. Amsterdam:
Elsevier, 1991. 321
[27] Wang D Z, Lu X D, Dou X Y, Li W B. Appl Catal,
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[28] Kitigawa H, Sendoda Y, Ono Y. J Catal, 1986, 101:
12
[29] Nishi K, Komai S, Inagaki K, Satsuma A, Hattori T.
Appl Catal A , 2002, 223(1-2): 187
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207
ScienceDirect
]<lUrMl of Natural Gas Chemistry
SCIENCE PRESS
www.elsevim.com/localeljngc
Article
Production of High Purity Multi-Walled Carbon Nanotubes
from Catalytic Decomposition of Methane
Kong Bee Hong,
Mahayuddin,
Aidawati Azlin Binti Ismail,
Abdul Rahman Mohamed,
Mohamed Ezzaham Bin Mohd
Sharif Hussein Sharif Zeiri*
School of Chemical Engineering, Engineering Campus, Chiversiti Sains Malaysia, 14300 Nabong Tebal,
Seberang Perai Selatan, Pulau Pinang, Malaysia
Abstract: Acid-based purification process of multi-walled carbon nanotubes (MWNTs) produced via
catalytic decomposition of methane with NiO/TiOz as a catalyst is described. By combining the oxidation in air and the acid refluxes, the impurities, such as amorphous carbon, carbon nanoparticles, arid
the NiO/TiOz catalyst, are eliminated. Scanning electron microscopy (SEM) and transmission electron
microscopy (TEM) images confirm the removal of the impurities. The percentage of the carbon nanotubes
purity was analyzed using thermal gravimetric analysis (TGA). Using this process, 99.9 wt% purity of
MWNTs was obtained.
Key words: multi-walled carbon nanotubes; purification; acid refluxes; oxidation; methane; decomposition
1. Introduction
Since their discovery by Iijima in 1991 [l],carbon nanotubes have been extensively researched and
have resulted in various potential applications [2-41,
thus opening a new chapter in nanoscale materials
science. However, a major issue that remains unresolved is its purification. Most synthesis methods of
the carbon nanotubes are based on the use of the catalyst and the as-synthesized carbon nanotubes are then
contaminated with metal catalyst and other carbonaceous materials such as amorphous carbon and carbon
nanoparticles [5]. These impurities are closely entangled with the carbon nanotubes and hence influence
the carbon nanotubes structural and electronic properties and thereby limit their applications [6]. Therefore, it is necessary to purify the as-synthesized carbon nanotubes t o enable their application in many
areas.
Several purification processes have been reported.
For example, Wiltshire et al. [7]used magnet t o separate ferromagnetic catalyst particles from an aqueous
surfactant solution of carbon nanotubes. The residual quantity of the Fe catalyst was 3 wt%. Moon
et al. [8] used a two step process of thermal annealing in air and acid treatment to purify single-walled
carbon nanotubes. This process provided carbon nanotubes with metal catalysts less than 1%. Strong
et al. [9] used a combination of oxidation followed
by acid washing and provided residue mass as low
as 0.73 wt%. A microwave-assisted digestion system
was used to dissolve the metal catalyst in organic acid
followed by filtration [10,11]. This method provided
99.9 wt% purity of the carbon nanotubes.
Although various purification methods have been
reported by researchers, which have shown high purity, no effective common method has yet been found
for the removal of impurities for all types of assynthesized carbon nanotubes. Therefore, the purification method depends on the specific type of cat-
* Corresponding author. Tel: 6045996442; Fax: 604-5941013; E-mail: [email protected]
Special Column of the INRET 2006/Journal of Natural Gas Chemistry VoJ. 15 No. 4 2006
alyst used in the synthesis of carbon nanotubes, the
reaction time, and the temperature [12].
Recently, our group had succeeded in obtaining a
higher yield in the synthesis of MWNTs from methane
decomposition using NiO/TiOz as the catalyst [13]
with activation energy, 60 kJ/mol, being the lowest
reported in the literature for this reaction [14]. To
enable their application in many areas, it was necessary to purify the as-synthesized MWNTs. In this
article, an acid-based purification process of the assynthesized MWNTs produced via catalytic decomposition of methane with NiO/TiOz as the catalyst
has been reported.
2. Experimental
2.1. Samples
Multi-walled carbon nanotubes (MWNTs) were
synthesized via the catalytic decomposition of
methane with NiO/TiOz as the catalyst. A complete
description of the synthesis of the catalyst and the carbon nanotubes are explained in detail elsewhere [13].
2.2. Purification
The acid-based purification process of multiwalled carbon nanotubes (MWNTs) produced via
catalytic decomposition of methane with NiO/TiOz
as the catalyst has been described.
The acid
refluxes/the chemical oxidation process and the acid
refluxes/the oxidation in air process have been compared. In the first step, 0.5 g of MWNTs was refluxed
in 100 ml of concentric acid (10 M) above boiling point
for 6 h. The effectiveness of nitric acid and sulfuric acid on the impurities were also compared in this
step under similar conditions. Then, the acid treated
MWNTs were either oxidized in air or chemically. Oxidation in air was done in a furnace at 350 "C for
2 h. Chemical oxidation was done using KMn04 and
HzSO4 at 80 "C for 1 h. The treated MWNTs were
then separated from the chemical solutions using microfiltration. The MWNTs obtained after the oxidation process were then dispersed in an aqueous solution of benzalkonium chloride. The mixture was then
sonicated for 2 h and the suspension was then separated from the solution using microfiltration. The
solid caught on the filter was then soaked in ethanol
to washout the surfactant. A final washing was done
with de-ionised water and then dried in an oven of
temperature 120 "C for 8 h.
267
2.3. Characterization
The morphology of the MWNTs before and after the purification process were examined using the
scanning electron microscope (SEM) system (A Leo
Supra 50 VP Fuel Emission) and the transmission
electron microscope (TEM) system (Philips Model
CMl2). The percentages of the impurities of the
MWNTs before and after the purification process
were analyzed using thermal gravimetric analysis
(Perkin Elmer TGA7 Thermogravimetric Analyzer).
Thermogravimetric analysis (TGA) is used to detect the percentage of MWNTs, metal catalysts, and
other impurities according to the combustion temperature difference between these materials. Figure 1
shows the TGA and the differentiated thermogravimetric analysis (DTG) curves of MWNTs before and
after purification. In Figure l(a),l(b), l(c), and l ( d ) ,
the solid lines and the dotted lines correspond to the
TGA curves and the DTG curves, respectively. Figure
l ( a ) shows the TGA of the as-synthesized MWNTs
and indicates that the weight starts to reduce near
510 "C. The MWNTs were completely burned a t
700 "C. The remaining materials were metal catalysts,
which were approximately 29% of the entire weight.
There was only one stepwise weight-loss, which indicates that the MWNTs did not contain amorphous
carbon. In the DTG curve, no peak was found in a
temperature below 500 OC, which again indicates that
the MWNTs did not contain amorphous carbon. The
peak at 620 "C in the DTG curve indicates the oxidation temperature of the MWNTs. Figure l(b) shows
the TGA results of MWNTs, which were purified using the nitric acid refluxes followed by chemical oxidation. Based on the TGA curve, the combustion temperature range between 0 "C and 100 "C is assumed to
be water vapor. There was a small peak in the DTG
curve a t temperature 200 "C, which indicates the presence of 4 wt% amorphous carbon in the MWNTs. The
MWNTs started burning at 450 "C and completed at
650 "C. In this temperature range, the weight percent of the sample dropped from 95 wt% to 75 wt%.
This shows that the sample contains only approximately 20 wt% MWNTs. This is considerably lower
than the as-synthesized MWNTs (Figure 1 (a)). This
maybe because the chemicals used for purification remained in the sample. The initial burning temperature of MWNTs (450 "C) is lower than that of the
268
Kong Bee Hong et a]./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
as-synthesized MWNTs (500 "C). This is because of
the metal catalysts that still remained in the MWNTs
and enhanced the combustion rate of the MWNTs and
thus reduced the combustion temperature [15]. Figure l(c) shows the TGA graph of MWNTs that were
purified using nitric acid refluxes followed by oxidation in air. There was no weight loss between 0 "C and
400 "C, which indicates that these MWNTs are free
of amorphous carbon. The MWNTs started burning
at approximately 500 "C and completed at 700 "C.
Thus, the purified MWNTs have purity of 84 wt%.
The metal catalysts that still exist were of 16 wt%.
Therefore, in this purification process, oxidation in
air is more suitable than chemical oxidation.
To remove the end caps of the multi-walled carbon
nanotubes and to expose the metal oxides for further
acid dissolving, oxidation in air was introduced prior
to acid refluxes. Figure l ( d ) shows the TGA graph of
the MWNTs after purification using oxidation in air
followed by nitric acid refluxes and then re-oxidation
in air. There was no mass loss between the temperature ranges of 300 "C and 400 "C, which indicates
that the purified MWNTs are free of amorphous carbon. The MWNTs started burning at 500 "C and
stopped at 835 "C. The residue at 835 "C amounted
to 8 wt% of the original mass and was attributed to
the NiO/TiOz catalyst. The total mass loss of this
sample was 92 wt%.
0
0
0
I00
-0.5
-20
-
95
.-C
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.-C
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h
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-30
80
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75
70
u
-50
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0
'Temperature (T)
200
400
600
Temperalure ("2)
0
200
400
600
Temperature ("C)
0
-10
80
0
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C
1-
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d
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4
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0
- 60
-60
0
200
400
600
Temperature ('C)
800
-71)
0
200
400
600
Temperature ('C)
800
Figure 1. TGA graphs of: (a) as-synthesized MWNTs, (b) MWNTs after purification using nitric acid
refluxes/chemical oxidation, (c) MWNTs after purification using nitric acid refluxes/oxidation in
air, (d) MWNTs after purification using oxidation in air followed by nitric acid refluxes and then
re-oxidation in air, (e) MWNTs after purification using oxidation in air followed by sulfuric acid
refluxes and then re-oxidation in air
gE
4
s
6
.-
Special Column of the I N R E T 2006/Journal of Natural Gas i3emistr.y Vol. 15 No. 4 2006
The effectiveness of sulfuric acid was also studied under similar conditions where MWNTs were purified using oxidation in air followed by sulfuric acid
refluxes and then re-oxidation in air. This is demonstrated in Figure l(e). The first total mass loss of
this sample was 2 wt%, which occured before 100 'C,
and which was probably due to water vapor. The
mass loss of MWNTs started at 500 "C. The residue
at 850°C amounted to 0.01 wt% of the NiO/TiOz catalyst. The purified MWNTs have purity of 99.9 wt%
of the total dry original mass. Thus, sulfuric acid has
higher catalyst (NiO/TiOz) dissolving efficiency than
nitric acid.
269
Figure 2 (a) and (b) show the TEM and SEM
images of the as-synthesized MWNTs, respectively.
The metal particles were evidently embedded in the
tip and between the MWNTs. The bright spots in
the SEM image shown in Figure 2 (b) indicate the
metal particles. Figure 3 (a) shows the TEM images
of the purified MWNTs. It clearly shows that all
tubes were opened and the metals embedded inside
the tubes were removed. Figure 3 (b) shows that the
SEM images of the purified MWNTs are free of bright
spots, which indicates that the purified MWNTs are
free of metal catalysts. Hence, these results show that
the MWNTs have high purity.
~
Figure 2. The images of the as-synthesized MWNTs: (a) TEM, (b) SEM
Figure 3. Purified MWNTs using oxidation in air followed by sulfuric acid refluxes and re-oxidation in air:
(a) TEM image, (b) SEM image.
4. Conclusions
Acid refluxes/oxidation in air provides higher purification efficiency of the as-synthesized MWNTs
than acid refluxes/chemical oxidation. Oxidation
in air prior to acid treatment can open the tips of
MWNTs and expose the metal particles inside the
tube for further acid solvating. Oxidation in air after
acid treatment helps to remove the amorphous carbon created after the acid treatment. In this study,
sulfuric acid provides a better result than nitric acid
to purify MWNTs produced via the catalytic decomposition of methane with NiO/TiOz as the catalyst.
Using this acid, purity of MWNTs as high as 99.9
wt% was obtained.
The authors acknowledge the financial support provided by Short Term Grant USM (Proiect:
A/C No:
~"
6035146) and Academy of Sciences Malaysia under Scientific Advancement Grant Allocation (SAGA) (Project:
A/C No. 6053001).
270
Kong Bee Hong et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
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ScienceDirect
Journal of Natural Gas Chcmmv
SCIENCE PRESS
www.elsevier.com/locate/jngc
Article
Microreactor for the Catalytic Partial Oxidation of Methane
Widodo Wahyu Puwanto*,
Yuswan Muharam
Sustainable Energy Research Group, Gas and Petrochemical Engineering Department, Faculty of Engineering,
University of Indonesia, Kampus Universitas Indonesia, Depok 1642, Indonesia
Abstract: Fixed-bed reactors for the partial oxidation of methane to produce synthetic gas still pose hotspot problems. An alternative reactor, which is known as the shell-and-tube-typed microreactor, has been
developed to resolve these problems. The microreactor consists of a 1 cm outside-diameter, 0.8 cm insidediameter and 11 cm length tube, and a 1.8 cm inside-diameter shell. The tube is made of dense alumina
and the shell is made of quartz. Two different methods dip and spray coating were performed to line
the tube side with the LaNi,O, catalyst. Combustion and reforming reactions take place simultaneously
in this reactor. Methane is oxidized in the tube side to produce flue gases (CO:! and HzO) which flow
counter-currently and react with the remaining methane in the shell side to yield synthesis g&. The
methane conversion using the higher-loading catalyst spray-coated tube reaches 97% at 700 "C, whereas
that using the lower-loading catalyst dip-coated tube reaches only 7.78% because of poor adhesion between
the catalyst film and the alumina support. The turnover frequencies (TOFs) using the catalyst spray-and
dip-coated tubes are 5 . 7 5 lop5
~
and 2 . 2 4 ~
lo-' mol/gCat. s, respectively. The catalyst spray-coated at
900 "C provides better performance than that at 1250 "C because sintering reduces the surface-area. The
hydrogen to carbon monoxide ratio produced by the spray-coated catalyst is greater than the stoichiometric
ratio, which is caused by carbon deposition through methane cracking or the Boudouard reaction.
Key words: microreactor; catalytic partial oxidation; methane; coating method
1. Introduction
Synthesis gas, which consists primarily of carbon
monoxide and hydrogen, is produced through steam
reforming, carbon dioxide reforming, and partial oxidation of methane. The H2/CO ratio produced by the
highly-endothermic steam and C02 reforming is not
suitable for use as the feedstock of methanol synthesis and the Fischer-Tropsch reaction [l].The partial
oxidation of methane is an interesting process to prevent the drawback posed by both reforming reactions
[2]. The other advantages of the partial oxidation of
methane are that the reaction is slightly exothermic
and the resident time is quick.
Fixed-bed reactors for the partial oxidation of
methane show a drawback, ie., hot spots are mostly
formed in the entry zone of the reactor. This is related
to the reaction mechanisms of the partial oxidation of
methane, the indirect mechanism depicted in Equations 1-3, or the direct one shown in Equation 4:
+
+
CH4 2 0 2 + C02 2H20
(AH298K = -801 kJ/mol)
+
+
+
+
+
+
CH4 CO2 H 2CO 2H2
(AH298~= 247 kJ/mol)
CH4 H2O H CO 3H2
(AH298~= 207 kJ/mol)
(1)
(2)
(3)
CH4 1 / 2 0 2 ++ CO 2H2
(4)
(AH298K = -36 kJ/mol)
For the indirect mechanism, the highlyexothermic total oxidation of methane (Reaction 1)
* Coressponding autor. Tel: 62-21-7863516; Fax: 62-21-7863515; E-mail: [email protected]
272
Widodo Wahyu Puwanto et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
occurs a t the entry of the reactor, followed by the
endothermic COz and steam reformings of unconverted methane (Reactions 2 and 3) occurring at the
remaining parts of the bed [3].
In better designs, a reactor for the partial oxidation of methane delivers heat released by the exothermic reaction in a manner such that the heat is simultaneously utilized by steam and C02 reformings.
One of these designs is the shell-and-tube microreactor [4]. Total oxidation takes place in the tube side,
while steam and COz reforming occuring in the shell
side of the reactor utilizes the heat released by the
oxidation.
To support this purpose, this study was undertaken with the aim of developing the shell-and-tube
microreactor for the catalytic partial oxidation of
methane over nickel-based catalyst, especially on the
impact of the coating methods and the operating conditions on the reaction performances.
2. Experimental
2.1. Reactor system
The shell and tube microreactor (Figure 1) was
made of dense a-alumina for the tubes, quartz for
the shell, and stainless steel for the unions. The catalyst length in the inside and outside of the tubes
was 11 cm; the outside and the inside diameter of the
tubes were 1 cm and 0.8 cm, respectively; and the
inside diameter of the shell was 1.8 cm.
The reactant flowrate was 300 ml/min with a
Products
Reactants
Catalyst
I
-$:;T;
t
I
ULI
I]
Union tee
Reducer union
Shell
CH4/02 ratio of 2.3. The gas leaving the reactor
passed through a water trap to condense the water
contained in the gases. A soap bubble was mounted
to meter the gas flow rates. The gas composition
was analyzed using a TDC-typed gas chromatograph
(Shimadzu GC-8A) with active carbon column.
2.2. Catalyst preparation and coating
The LaNi,O, catalyst was prepared using the solgel technique [5,6]. The precursor La(N03)3. 6Hz0
was mixed with Ni(N03)Z. 6H20 with a Ni/La molar
ratio of 10. The mixture was then solved in water.
Maleic acid was added gradually into the nitric solution produced until a maleic acid/total nitric molar
ratio of 1 was achieved. The solution was then stirred
a t 70 "C.
Coating of the catalyst on the tube was done
through dip-coating and spray-pyrolisis methods. The
alumina tube was cleaned by immersing it in a 50%
acetone solution and by treating it in an ultrasonic
bath containing 30% acetone solution for 20 min. Alumina was then dried in vacuum furnace at 140 "C for
1 hour. In the dip-coating method, the tube was
immersed in a catalyst sol. The catalyst-deposited
tube was heated to evaporate the solvent. The catalyst was then calcinated a t 900 "C for 3 h. In the
spray-pyrolisis method, the catalyst sol was calcinated
at 900 "C, then crushed into small size powder, and
solved in isopropyl alcohol (0.2 g/ml). The catalyst
solution was sprayed on the tube surface which was
already heated. The catalyst deposit was sintered at
sintering temperatures (900 and 1250 "C) for 4 h [7].
The catalyst samples sintered at 900 "C and 1250 "C
are called the spray-900 and the spray-1250, respectively.
The resulting catalyst powder were characterized
with BET (Autosorb 6 , Quantachrome) for surface
area, and EDX Scan (Oxford 6599) for catalyst composition.
Tube
3.1. Effect of coating method
Figure 1. Scheme and photograph of the microreactor
The test conducted in the blank reactor at 700 "C
shows that no gas products were observed. The catalyst prepared using the dip-coating method converted
7.7% of methane. The low conversion of methane is
caused by very low loading of the catalyst in the tube
(0.06 g in the outside and 0.02 g in the inside of the
Special Column of the INRET 2006/JournaI of Natural Gas Chemistry Vol. 15 No. 4 2006
tube), whereas the spray-900 and the spray-1250 catalysts provided high methane conversions (more than
97%) because of the higher loading of the catalyst in
the outside (0.4 g) and the inside (0.2 g) of the tube.
273
of the spray-900 catalyst (3.03 m2/g) is larger than
that of the spray-1250 catalyst (2.1 m2/g).
3.2. Effect of temperature
The effect of the reaction temperatures on the
conversion is shown in Figure 4.
100,
1
-
80
g
.-
spray-900
-
:
60
-
v)
"
n
,
F
6
0.00 0.00
Blank reactor
Dip-coating
Spray-900
20
Spray-I 250
Figure 2. Methane conversions at 700 "C for each
type of catalyst
It can be explained that the dip-coating method
provided imperfect catalyst film deposition on the
alumina surface because of unstable heating a t 70 "C
or fast evaporation, and the resulting sol was not homogeneous. The aged sol will form gel at drying,
which will produce a thick film and flakes or weak
adhesion. Meanwhile, the spray-pyrolisis method
provided strong adhesion between the support and
the catalyst layers and consequently produced higher
loading.
Figure 3 shows the turn over frequency (TOF) of
the three different catalysts at 700 "C representing
the reaction performance per weight of the catalyst.
The T O F with the spray-pyrolisis catalysts
and
provides a better performance ( 5 . 7 5 lW5
~
5.73~
mol/gcat. s) than the dip-coating catalyst
(2.24~
mol/gcat. s). The spray-900 catalyst has
higher activity when compared to the spray-1250 catalyst, which may be because the BET surface area
6 ,
40-
U
I
Figure 3. The turn over frequency (TOF) of the
different catalysts at 700 "C
-
0Spray-I250
80
-
60
-
9
c
.-
2
$
401
0
20 nL
500 'C
700 "C
600 'C
Figure 4. Effect of temperature on the reactant conversion
It is seen that the decrease in temperature will decrease the conversion of methane. At higher temperature (700 "C),the methane conversion is similar while
using the spray-900 catalyst and the spray-1250 catalyst (about 97%). The reaction is thermal controlled
at higher temperature, whereas at the same time,
the difference in methane conversion significantly occurs at lower temperatures (600 "C and 500 "C), in
which the methane conversion of the spray-900 catalysts is larger than that of the spray-1250 catalyst,
which indicates that the reaction is catalytic controlled. For the lowest temperature of 500 "C, most
of the nickel catalyst is required t o consume oxygen
to form NiO, so that the oxygen conversion is considerably higher than the methane conversion. The
selectivity of CO, C02, and H2 for the spray-pyrolisis
catalysts are shown in Figure 5. The H2 selectivity
increases, whereas the C02 selectivity decreases with
increasing temperatures. It can be explained that the
Gibbs free energy of methane combustion is more negative than that of steam and C02 reforming within
the temperature range of 500-700 "C. At lower tem-
2 74
Widodo Wahyu Puwanto et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
peratures, the complete oxidation of methane is dominant than reforming. The CO selectivity is relatively
constant, although the reforming reactions proceed
more spontaneously at higher temperatures. This is
0
2
4
6
8
I0
12
(keV)
Figure 7. EDX spectrum of the spray-900 catalyst
4. Conclusions
500 "C
600 *C
700 *C
Figure 5. Effect of temperature on product selectivity
Figure 6 exhibits that the Hz/CO ratios are always greater than 2, which are possibly caused by the
existence of methane cracking and by the Boudouard
reactions. The occurrence of carbon deposition was
proved by the appearance of COz when the used catalyst was flushed by oxygen and EDX scan as shown
in Figure 7.
4,
The following conclusions can be drawn based on
the above discussion:
The catalyst loading prepared using the spraypyrolisis method is higher than that of the dip-coating
method leading to better performance of partial oxidation of methane (the methane conversion is 97%,
s).
and the TOF is 5.75~10-~mol/g,,~.
Owing to the higher surface area (3.03 m2/g), the
performance of the catalyst sintered at 900 "C is better than that at 1250 "C,
At lower temperatures (500-600 "C),the methane
conversion using the spray-900 catalyst is higher than
that using the spray-1250 catalyst. At a higher temperature (700 " C ) ,the two catalysts convert almost
the same amount of methane. The occurrence of carbon deposit causes the hydrogen to carbon monoxide
ratios produced to be greater than the stoichiometric
ratio.
Acknowledgements
We especially thank M. Subhan for technical assistance in carrying out the most of experiments.
I
References
500 '
C
600 *C
700 *C
Figure 6. Effect of temperature on the hydrogen to
carbon monoxide ratio
[I] Basile A, Paturzo L. Catal Today, 2001, 67(1-3): 55
[2] Jin W, Li S, Huang P, Xu N, Shi J, Lin Y S. J Membr
Sci, 2000, 166(1): 13
[3] Swaan H M, Rouanet R, Widyananda P, Mirodatos
C. Stud Surf Sci Catal, 1997, 107: 447
[4] Piga A, Verykios X E. Catal Today, 2000, 6O(l-2): 63
[5] Norman A K, Morris M A. J Mater Process Technol,
1999, 92-93: 91
[S] Golden S J. US Patent 6 372 686. 2000
[7] Ritchie J T, Richardson J T, Luss D. AIChE J , 2001,
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ScienceDirect
Journal of Natural Gas Chemistry
SCIENCE PRESS
www.elsevier.cddjmnocateljngc
Review
Efficient Fixation of Carbon Dioxide by Electrolysis
- Facile Synthesis of Useful Carboxylic Acids Masao Tokuda*
Division of Molecular Chemistry, Graduate School of Engineering, Hokkaido University, Sapporo 060-8628, Japan
[Manuscript received October 18, 2006]
Abstract: Electrochemical fixation of atmospheric pressure of carbon dioxide to organic compounds is a
useful and attractive method for synthesizing of various carboxylic acids. Electrochemical fixation of carbon
dioxide, electrochemical carboxylation, organic halides, organic triflates, alkenes, aromatic compounds, and
carbonyl compounds can readily occur in the presence of an atmospheric pressure of carbon dioxide to
form the corresponding carboxylic acids with high yields, when a sacrificial anode such as magnesium or
aluminum is used in the electrolysis. The electrochemical carboxylation of vinyl bromides was successfully
applied for the synthesis of the precursor of nonsteroidal anti-inflammatory agents such as ibuprofen
and naproxen. On the other hand, supercritical carbon dioxide (SCCOZ)has significant potential as an
environmentally benign solvent in organic synthesis and it could be used both as a solvent and as a reagent
in these electrochemical carboxylations by using a small amount of cosolvent.
K e y words: carbon dioxide; fixation; supercritical carbon dioxide; electrolysis; sacrificial anode; carboxylic acid
1. Introduction
Efficient fixation of an atmospheric pressure of
carbon dioxide to appropriate organic substrates is
a very useful and attractive method for synthesizing
of various carboxylic acids. Electrochemical methods for the efficient fixation of carbon dioxide have
been studied as the electrochemical reactions usually
proceed under mild conditions. Among them, the
electro-chemical reductive method using a platinum
cathode and a sacrificial anode such as magnesium
or aluminum metal in a one-compartment cell was
found to be the most convenient and effective method
[l-21. The electrochemical fixation of carbon dioxide
to various types of organic compounds was studied by
the authors and a variety of useful carboxylic acids
were prepared in high yields. These electrochemical
fixations, electrochemical carboxylations, were suc-
cessfully applied for the efficient synthesis of the precursor of nonsteroidal anti-inflammatory agents such
as ibuprofen and naproxen.
On the other hand, supercritical carbon dioxide (scCO2) has significant potential as an environmentally benign solvent for the replacement of hazardous organic media in organic synthesis, as it is inexpensive, nontoxic, and can be readily recovered and
reused after the reaction. We found that supercritical
carbon dioxide could be used both as a solvent and as
a reagent in these electrochemical fixations.
In this article, our results on the efficient electrochemical fixation of an atmospheric pressure of carbon dioxide and its application for the synthesis of
various useful carboxylic acids including the precursor of anti-inflammatory agents are summarized. The
use of supercritical carbon dioxide both as a solvent
and as a reagent in electrochemical fixations is also
described in this article.
*Corresponding author. Present address: HanakawiGkita 4-2-1-4, Ishikari 061-3214, Japan
Tel: +81-133-74-2811; Fax: +81-133-74-2811; Email: tokudaQorg-mc.eng.hokudai.ac.jp
276
Masao Tokuda et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
2. Electrochemical fixation of carbon dioxide
The electrochemical fixation of carbon dioxide to
organic compounds is considered to proceed via two
pathways. The pathway A includes the generation of
anionic species by the two-electron reduction of the organic substrates (Scheme 1, equations (1)-(4)). The
carbanions undergo a nucleophilic attack on carbon
dioxide to yield the corresponding carboxylic acids.
This case occurs when the reduction potentials of the
organic substrates are more positive than that of carbon dioxide. The electrochemical fixation of carbon
dioxide occurs through pathway B when the reduction potentials of the organic substrates are more negative than that of carbon dioxide (Scheme 1, equations (5)-(8)). In this method, the anion radical of
carbon dioxide generated by the one-electron reduction reacts with organic substrates, typically such as
alkenes, to yield the corresponding carboxylic acids.
[Pathway A]
RX-%[RX]’--
R.+X-
(1)
-
R. +e_ RR-+ C02
(2)
(3)
R-CO;
R-COY +Hi_ R-CO2H
[Pathway B] C02
(4)
-
[CO,]’-
[co21’-+ RA
(5)
(6)
R k C O ’
R&
a;
4_
R/\/CO;
R&co;
(7)
(8)
S R - C 0 2 H
Scheme 1. Electrochemical fixation of carbon dioxide
to organic substrates
These electrochemical fixations of an atmospheric
pressure of carbon dioxide to organic substrates can
effectively occur when the electrolysis is carried out in
a one-compartment cell using a platinum cathode and
a sacrificial anode such as magnesium or aluminum
[l-21. Such fixation, the electrochemical carboxylation, results in the formation of various carboxylic
acids with good yields.
3. Electrochemical carboxylation of allylic and
propargylic halides
Electrolysis of allylic bromides (1)in the presence of an atmospheric pressure of carbon dioxide
with a platinum cathode and a magnesium anode regioselectivelyform the corresponding P,y-unsaturated
carboxylic acids 2 with 34%-71% isolated yields
(Scheme 2) [3]. The electrolysis was carried out in
an N , N-Dimethyl formamide (DMF) solution containing 0.1M Et4NC104 at constant current using a
one-compartment cell. This electrochemical carboxylation results in the formation of acids 2 and this reaction is of considerable significance as the carboxylation of allylic organometallic compounds derived from
allylic halides 1 usually yields P,y-unsaturated carboxylic acids 3 [4].
Electrochemical
carboxylation
of
(Y,Qdisubstituted propargyl bromides (4) yield allenic
acids 6 as the predominant product (Scheme 2) [5].
Carboxylation of propargylic organometallic compounds results in the formation of a mixture of acids
5 and 6 with very low yields.
n
Mg+
R*
I
P f
R
0.1 M E14NCI04-DMF
1
R
+ c02
--+R
X
4
*
3 Flmol, 25 mAicm2
[Y=34%-71%]
+
”&
R
3 (10%)
2 (90%)
n
Mg+
+-p
R
Pt-
b
0.1 M E14NCI04-DMF
5 Flmol, 25 rnAlcm2
[Y=36%-44%]
R‘ L C 0 2 H
,=.+
+
CO2H
5 (0-10%)
H02C
R
R
6 (1 00%-90%)
Scheme 2. Electrochemical carboxylation of allylic and propargylic halides
The electrochemical carboxylations of both the allylic and the propargylic halides would proceed via
the same pathway A as shown in Scheme 1, as both
these halides are more readily reduced than carbon
dioxide. In these cases, the corresponding carban-
ions that are derived from the allylic or propargylic
halides are produced as the intermediate species and
the regioselectivityyielding 2 or 6 as the predominant
product in the electrochemical carboxylations of 1 or
4 can be elucidated by the addition of more stable
277
Special Column of the INRET 2006,’Journal of Natural Gas Chemistr,y Vol. 15 No. 4 2006
carbanions t o carbon dioxide.
The hitherto unknown 3-methylenepent-4-enoic
acid ( 8 ) was synthesized by the electrochemical carboxylation of allylic tribromide 7 (Scheme 3) [6]. Tribromide 7 was readily prepared by the ll 4-addition of
bromide to isoprene followed by allylic bromination[7]
and it can work as a synthetic equivalent of isoprenyl
carbanion (A). Two-electron reduction of 7 results
in the formation of isoprenyl bromide (9), whereas
two-electron reduction of 9 generates the carbanion
A (Scheme 3).
-
Pl-
Br
I
I
t2e
MA‘
idative addition of Ni(0) to vinyl bromide would yield
vinyl nickel complex 12. Two-electron reduction of
the complex 12 yields the corresponding vinyl carbanion, which would be trapped by an atmospheric pressure of carbon dioxide to yield the corresponding a,@unsaturated carboxylic acids 11. Cyclic voltammetry
of 1-bromocycloheptene in the presence of NiBra(bpy)
showed the existence of a novel reduction peak at ca
-1.5 V, which has a higher positive potential than
that of the original vinyl bromide [lo].
n
PI- Mg’
-A
0.1 M Ei,NI-l)MF
5 Flmol. 10 mA!cm2
(MA: AlloyofMn.NiaiidCu)
CO2H
8 (57%)
t
R’
0.1 M 13uqNBl:4-DMF
3 Wrnol. 10 mA/cm2
Br
10
Electrochemical carboxylation of vinylic
bromides
Reduction potentials of aryl-substituted vinyl
bromides are more positive than that of carbon dioxide and therefore, their electrochemical carboxylation would proceed via the pathway A in Scheme 1.
Electrolysis of aryl-substituted vinyl bromides 10
in the presence of an atmospheric pressure of carbon dioxide with a platinum cathode and a magnesium anode gave the corresponding a , P-unsaturated
carboxylic acids 11 with isolated yields of 63%92% (Scheme 4)[8,9]. When R1, R2 or R3 are
alkyl or hydrogen atoms, the reduction potentials
of the vinyl bromides become highly negative and
the electrochemical carboxylations of them form a$unsaturated carboxylic acids 11 with low yields.
However, the addition of Ni(I1) catalyst in this electrochemical reduction significantly enhance the yield
of the desired a$-unsaturated carboxylic acids 11
and the isolated yields are 58%-82% (Scheme 4)[10].
Probable reaction pathways for the electrochemical carboxylation of aliphatic vinyl bromides in the
presence of Ni(I1) catalyst are shown in Scheme 5.
Reduction potentials of 1-bromocycloheptene and
NiBrz(bpy) are<-2.6 V and -1.25 V vs Ag/AgCl, respectively. Two-electron reduction of more easily reducible Ni(I1) catalyst yields Ni(0) species and the ox-
“HR3
R?
C02H
II
a ) R ’ , R’, R3 = Ph. alkyl or H
h) R’. R‘, R 3= alkyl or H
c ) R ’ , R‘, R3 = alkyl or H; 20rnol%NiBr2.hpy
Y=63%-92%
Y=14%-43%
Y=58%-82%
Scheme 4. Electrochemical carboxylation of vinylic
bromides
NiW)
Scheme 3. Electrochemical carboxylation of isoprenyl
anion equivalent 7
4.
*
Br
R’
10
12
11
Scheme 5. Proposed reaction pathways of the electrochemical carboxylation of aliphatic vinyl
bromide using Ni(I1) catalyst
Stereochemistry of the electrochemical carboxylation using Ni(I1) catalyst was also examined.
Electrochemical carboxylation of ( E ) - and (2)-@bromostyrene in the presence of 20 mol% of NiBr:!
(bpy) proceeded with the retention of the stereochemistry to yield the corresponding (19)and (2)cinnamic acids with high-stereoselectivities [ll].
High efficiency in the electrochemical fixation of
carbon dioxide using a sacrificial anode can be rationalized by the reaction pathways shown in Scheme 6.
At the cathode, a two-electron reduction of organic
halides occurs to yield the corresponding carbanions
(R-), which are trapped by carbon dioxide to yield
the corresponding carboxylates (RCOO-) (Scheme 6,
equations (1) and (2)). At the anode, on the other
hand, the dissolution of magnesium metal takes place
and this results in the formation of magnesium ion
(Mg2+) (equation (3)). The magnesium ion readily
captures carboxylates to yield stable magnesium car-
278
Masao Tokuda et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
instead of vinyl halides, was also examined and triflyl
function (-OTf=-OS02CH3) acts as a good leaving group in organic reactions. During the electrochemical carboxylation of vinyl triflates 13, two
types of reactions occurred and two entirely different
products, namely, a,@-unsaturatedcarboxylic acids
(14) or @-keto carboxylic acids (15) are obtained
(Scheme 7) [12-141. Electrolysis of phenyl-substituted
vinyl triflates in the presence of an atmospheric pressure of carbon dioxide formed acids 14 with good
yields, whereas the similar electrochemical carboxylation of alkyl-substituted vinyl triflates yielded carboxylic acids 15 as the predominant product. These
divergent electrochemical carboxylations are resulted
from the chemoselective cleavage of 0-S or C-0
bond of the vinyl triflates. In the case of phenylsubstituted vinyl triflates, a preferential reduction of
vinyl triflates occurs to yield vinyl carbanions, which
are trapped by carbon dioxide to yield 14. On the
other hand, in the case of alkyl-substituted vinyl
triflates, carbon dioxide is more readily reduced to
form its anion radical, which would attack the sulfur
atom of the vinyl triflates followed by 0 - S bond cleavage and this results in the formation of corresponding
enolate anions. These enolates would be captured by
carbon dioxide to exclusively yield 15.
boxylates ((RC00)2Mg or RCOOMgX) (equation (4)
or (5)). If magnesium ions were not present in the solution, most of the carboxylates would either undergo
decarboxylation to regenerate the original carbanions
or it would undergo decomposition at the anode to
yield radicals followed by Kolbe coupling reaction.
These decarboxylation and decomposition of carboxylate anions would result in the formation of carboxylic
acids in low yield.
-
at cathode (Pt)
R-X t 2e
R-+CO2
at anode (Mg)
Mg
overall
2RC00-
+ Mg2+
andor
RCOO-+ Mg2+ f X-
R-+X-
(1)
RCOO-
(2)
Mg2++2e
(3)
(RC00)2Mg
(4)
RCOOMgX
(5)
-
Scheme 6. Probable reaction pathways of an efficient
electrochemical fixation of carbon dioxide
using a sacrificial anode
5.
Electrochemical carboxylation of vinyl
triflates
Electrochemical carboxylation of vinyl triflates,
PtR 1R2J 7 R l
15
*03IFM/ ~BurNBF4-DMF
O I .10 m ~ i c r n '
[R'.,'R R'
= alkyl.
I -1Mg'.CO,
R'
My+. CO,
R3
pt-
*,
R2+Rl
HI
0.1 M BurNBFr-DMF
OTf
13
3 Fimol. 10 InA/cm2
[R'.RZ,R'
=ary. alkyl, HI
R2+R'
CO2H
14
Scheme 7. Divergent electrochemical carboxylation of vinyl triflates (13)
Carboxylation pathway of 13 to give 15 can completely be changed to another one giving 14 by the
use of Ni(I1) catalyst in the electrolysis. Such electrochemical carboxylation of lactone enol triflates (16)in
the presence of 20 mol% of NiBr2 (bpy) gave the corresponding cyclic a-alkoxyl-a,@-unsaturatedcarboxylic
acids 17, captodative cycloalkenes, with 63%-79%
yields (Scheme 8) [15].
w
16
0.1 M Bu4NBFa-DMF
3 Flmol, 10 &cm2
[20 mol% NiBrzbpyl
Fixation of two molecules of carbon dioxides can
take place when the electrolysis of phenyl-substituted
alkenes was carried out in the presence of an atmospheric pressure of carbon dioxide with a platinum
cathode and a magenesium anode. Various phenylsuccinic acids (19)were obtained by the electrochemical
dicarboxylation of alkenes 18 and the isolated yields
were 66%-91% (Scheme 9) [IS]. Dicarboxylation probably occurs via the pathway A in the case of stilbene
(18;R1=R3=H, R2=Ph) and via the pathway B in
17 (63%-79%)
Scheme 8. Electrochemical carboxylation of lactone
enol triflates (16)
6. Electrochemical carboxylation of alkenes
3 Flmol, 25 mAlcm'
18
..
I9 (66%-91 %)
Scheme 9, Electrochemical dicarboxylation of phenylsubstituted alkenes (18)
279
Special Column of the INRET 2006/Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
the case of styrene as shown in Scheme 1 (18;
R1=R2 =R3=H).
Electrochemical fixation of carbon dioxide t o
alkenes having more complex structure readily occurs to give the corresponding carboxylic acids with
d
. R'
good yields. Electrocheniical carboxylation of ringfused alkylidenecyclopropanes 20 in the presence of
an atmospheric pressure of carbon dioxide afforded either mono-(21) or dicarboxylic acid (22) in 44%-74%
yields (Scheme 10) [17].
R'R'CH
+ coz
20
C02H
d
or
21
22
R ' = H. alkyl
R' = C02R', COMe
(n = I. 2 )
Scheme 10. Electrochemical carboxylation of bicyclo[n.l.0]alkylidene derivatives (20)
7. Use of supercritical carbon dioxide in the
electrochemical carboxylation
Supercritical carbon dioxide (scCO2) can readily be attained under relatively moderate conditions
(Tc=31 "C, Pc=7.5 MPa). It has significant potential as an environmentally benign solvent to replace
the hazardous organic solvents, as it is nontoxic, inexpensive, and can be recovered and reused after the
reaction. The authors of this study developed a novel
method for electrochemical carboxylation in scCO2
by the use of small amount of acetonitrile (5-10 ml
to 155 ml of scCO2) as a cosolvent [18]. No current
flows in pure supercritical carbon dioxide due to its
poor conductivity.
Fixation of carbon dioxide to aromatic compounds successfully occurred in scCO2. Electrolysis
of aryl halides (Ar-X) in scCO2 containing a small
amount of acetonitrile with a platinum cathode and
a magnesium anode resulted in the formation of aryl
carboxylic acids (Ar-C02H) such as 23, 24, and 2 5
with high yields (Scheme 11) [19,20]. Electrochemical dicarboxylation of phenanthrene and anthracene
also proceeded efficiently in scCO2 and formed acids
26 and 2 7 with high yields (Scheme 12) [20]. Similar
electrochemical carboxylation of anthracene in acetonitrile solution in the presence of an atmospheric
pressure of carbon dioxide formed 27 and the yield
was only 32%. Most of the electrochemical carboxylations in scC02 gave higher yields of carboxylated
products than those using a n atmospheric pressure
of carbon dioxide. This is probably because of the
high diffusion rates of RCOO- and Mg2+ species in
scCO2 to form stable (RC00)zMg or RCOOMgX,
compared with the slow diffusion rate in DMF solution containing an atmospheric pressure of carbon
dioxide (Scheme 6 , equations (4) and (5)).
Electrochemical carboxylation of arylmethyl
halides, 1-aryl-1-bromoethenes, and aryl methyl ketones in scCO2 also proceeded efficiently to form the
corresponding carboxylic acids with high yields [20].
n
PI- Mg'
Ar-X + scC02
t
Ar-C02H
Bu4NBF4-CIiICN
3 F/mol. 25 mA/cn?
140 "C. 80 kgicm2]
X
23
24
25
Scheme 11. Electrochemical carboxylation of aryl
halides in supercritical carbon dioxide
PI-
Mg'
Bu4NBF4-CH3CN
3 F/mol. 25 mAIcm'
140 T,80 kglcm']
26 (93%)
+ SCCO?
3 I Imol. 25 mA/cm*
140 %, 80 kg/cm*]
CO2H
27 (93%)
Scheme 12. Electrochemical dicarboxylation of arenes
in supercritical carbon dioxide
280
Masao Tokuda et a]./ Journal of Natural Gas Chemistry Val. 15 No. 4 2006
8. Application to the synthesis of the precursor of anti-inflammatory agents
ample, the electrochemical carboxylation of vinyl bromide 28 formed the desired a,p-unsaturated carboxylic acid 29 with 93% yield [8],which can be readily transformed into (S)-ibuprofen by enantioselective
hydrogenation [21,22] (Scheme 13). The precursors of
naproxen (31),ketoprofen (32),and flurbiprofen (33)
were produced in good yields by similar electrochemical carboxylation (Scheme 13).
Efficient electrochemical carboxylation of vinyl
bromide using an atmospheric pressure of carbon
dioxide was successfully applied for the synthesis of
the precursor of anti-inflammatory agents. For ex-
28
29 (93%)
Naproxen precursor (31) (80%)
(S)-(
+)-lhuprqfin
30 (97%. 9 7 % ~ )
Ketoptwjen precursor (32) (74%)
Flurbiptwjen precursor (33) (75%)
Scheme 13. Synthesis of several precursors of anti-inflammatory agents by electrochemical carboxylation of vinyl
bromides
The electrochemical carboxylation in scCO2
was successfully applied for the synthesis of antiinflammatory agents. Electrolysis of benzylic chloride 34 in scCO2 formed racemic naproxen (35)
with 74% yield (Scheme 14) [19,20]. Electrochemical carboxylation of aryl methyl ketone 36 in s c C 0 ~
formed a-hydroxycarboxylic acid 37 with 78% yield,
which can afford (S)-naproxen (38)by dehydration
and enantioselective hydrogenation. Synthesis of (5')loxoprofen (41)is of considerable synthetic as well
34
[40 'C. KO kgicm2]
racemic Naproxen (35) (74%)
n
M y ' , scC02
t
Me0
36
n
PI-
M~'.~cCOL
-
O
&
:H
2'
PI-
Bu,NBF,-CH,CN
5 Fimol. 20 mAicm'
140 %.KO kg/cm21
as electrochemical significance as the usual chemical reaction cannot control the reactivity of the two
carbonyl groups in 39. In the electrochemical reactions; however, aryl methyl ketone is more readily reduced than cyclopentanone and the electrochemical carboxylation of 39 formed the desired ahydroxypropanoic acid (40)with 91% yield (Scheme
14). The product 40 can be transformed into ( S ) loxoprofen (41)[20].
Meo
\
/
Me0
37 (78%)
m
c
\
o
2
(S>Nuproxen (38)
-
Qco2H
CO2H
Bu4NBF,-CH3CN
5 Fimul. 20 mA/cm2
0
39
H
/
OH
[40 Z , KO kgicm?]
40 (91%)
(S)-Loxopro/en (41 )
Scheme 14. Synthesis of anti-inflammatory agents by electrochemical carboxylation in supercritical carbon dioxide
Special Column of the INRET 2006/3ournal of Natural Gas Che1nistr.y Vol. 15 No. 4 2006
Similar electrochemical carboxylation of vinylic
bromides in scC02 also gave the precursor of antiinflammatory agents 29, 31, and 33 with almost the
same yields as those using the atmospheric carbon
dioxide [20].
The precursors of various anti-inflammatory
agents such as 29, 31, 32, and 33 as well as those of
cicloprofen, indoprofen, suprofen, and loxoprofen were
prepared by the cross-coupling reaction of the corresponding aryl iodides with organozinc bromides obtained by the reaction of ethyl 2-bromoacrylate with
the electrogenerated highly reactive zinc [23-251.
9. Summary
The present electroclieniical method for an
efficient fixation of an atmospheric pressure of carbon dioxide to a variety of organic compounds has
several advantages: use of a simple one-compartment
cell by the use of a platinum cathode and a magnesium anode, simple electrolysis at a constant-current ,
high yields in the synthesis of useful carboxylic acids
and easy application to the synthesis of the precursor of nonsteroidal anti-inflammatory agents. This
procedure might be used for the industrial production of high value-added substances such as pharmaceuticals. Use of supercritical carbon dioxide for the
electrochemical fixation of carbon dioxide would be
useful in the future since an environmental problem
will become more important for us.
Acknowledgments
The author wishes to express his thanks to Professors
H. Suginome, H. Senboku, and N. Kurono and many excellent students including Dr. Aishah A. Jalil for their
helpful discussions and their hard work during this study.
This work was supported by Grants-in-Aid for Scientific
Research from The Ministry of Education, Science, Sports
and Culture, Japan.
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Springer-Verlag, 1998. 239
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2681
Yuan Kou*,
Wei Xiong,
Guohong Tao,
Hui Liu,
Tao Wang
PKU Green Chemistry Center, Beijing National Laboratory for Molecular Sciences, College of Chemistry and
Molecular Engineering, Peking University, Beij’ing 100871, China
[Manuscript received October 13, 2006; revised October 30, 20061
Abstract: A reversible storage-release process switched by a temperature difference of 10 “C around
room temperature can be realized. This fast, recyclable, energy efficient, low cost and green system within
a wide range of temperature and pressure is reported here for the first time. The system is believed to
open up a new route for the storage and homogeneous utilization of methane.
Key words: ionic liquid; methane; absorption; capture; polarity; solubility
1. Introduction
Energy efficient, cost economic and green methods for capturing methane into liquid and solid phase
have been great challenges to researchers[l-4]. The
simplest approach for the storage of CH4, like that
for COZ [5,6], is its absorption by a liquid. The solubility of CH4 in some of the volatile organic chemicals
(VOCs) is good, but in green solvents, such as nonvolatile ionic liquids (ILs), is generally poor [ 1,2]. Storage of CH4 using well-designed crystalline frameworks
has attracted considerable academic attentions[3,4],
but to what extent those advanced materials can be
used in a practical way is still an open question.
In general, ILs are believed to be highly polar
and therefore expected t o be more capable for replacing traditional polar organic solvents in various applications[7,8], for example, in catalytic homogeneous
arene hydrogenation[9]. Development of ILs that can
take the place of weakly or non-polar organic solvents
is still in its infancy. ILs are highly tailorable materials, i.e., their polarities may be tuned by altering the structure of either the cation or the anion, or
both. It is interesting to note that our previous work
has revealed that typical non-polar CSOmolecules can
easily be dissolved in [C4mim]Tf2N1a commercially
available IL, with a very good solubility[lO], implying
that low polarity ILs may exhibit great potentiality in
capturing and storing methane in a very simple, hence
energy and cost efficient way. Due to the fact that
ILs are salts of low melting points, a smart approach
for dissolving methane in an IL is that the dissolution
is accompanied with a reversible phase transition between the liquid and the solid upon varying the temperature, making the fixation of CH4 as easy as its
dissolution.
Here we report on the invention of a low-polar IL
system for the absorption of CH4 and demonstrate
the formation of a CH4-IL complex under a certain
pressure around room temperature. The complex is
a stable solid under ambient conditions and can release CH4 by mild heating. Such a reversible system
should facilitate the storage, transportation, homogeneous utilization and catalytic conversion of methane.
2. Experimental
2.1. Synthesis Of [N8888]Tf2N ionic liquid
[NssssIBr (10.94 g, 0.02 mol) was dissolved in
deionized water (30.0 ml), then LiNTf2 (5.74 g,
* Correspondence author. Tel: 86-10-62757792; Fax: 86-10-62751708; E-mail: yuankouQpku.edu.cn
This work was financially supported by the National Science Foundation of China (Project No.20533010).
Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
0.02 mol) was added with vigorous stirring. After
the mixture was stirred a t room temperature for
1 h, the upper aqueous layer was decanted. The
residual ionic liquid (IL) layer was then dissolved in
dichloromethane (50.0 ml), and washed with deionized water for three times. The dichloromethane solution was removed by rotary evaporation to leave a
colourless transparent liquid (yield 13.00 g, 87%).'H
NMR (CDC13, G/ppm): 3.14 ( t , 8H, 5=8.4, -NCH2),
1.60 ( m , 8H, -CHz-), 1.28-1.35 (m, 40H, -(CHz)s-),
0.89 ( t , 3H, J=6.6, -CH3).13C NMR (CDC13, G/ppm):
58.76, 31.55, 28.90, 26.10, 22.53, 21.82, 14.01.
2.2. Solubility tests
The IL [N8888]Tf2N was put in a n autoclave and
then the system was heated and monitored to a desired temperature. After the temperature was stable,
the autoclave was degassed under vacuum and washed
with methane for three times to make sure that no air
was in the system. The amount of methane dissolved
in the [Nssss]Tf2N during the degas procedure was
negligible because the contracting area between gas
and liquid is small. Then methane was charged into
the system under a desired pressure. The stirring was
started and the gas pressure decreased quickly. The
procedure was carried out at least for one hour until the pressure became stable. After the system had
reached its equilibrium, we attained the final temperature and pressure. The mole fraction of methane
was calculated according the Virial equation and the
Virial constant in the literature.
2.3. Formation of methane-ionic liquid complex
After the IL had absorbed methane under a certain pressure and temperature in the liquid state, the
methane-IL solution was cooled by ice or other refrigerating materials. When the temperature had decreased to near 0 T, the IL was frozen to the solid
state, and the resulting white solid was the methaneIL complex.
2.4. Release of methane from the complex
After the formation of the complex solid, the
methane which remained in the autoclave was released
until the pressure went down to 1 atm. A solid material was obtained after opening the autoclave. When
the as-prepared methane-IL complex solid was heated,
283
numerous bubbles would be given out from it immediately, and the methane fixed in the IL was released.
MS and FT-IR were used to make sure that the
released gas is methane. After reducing the pressure
of the autoclave to 1 atm, 10 atm N2 was introduced
into the autoclave for four times to dilute the methane
remaining in the gas phase. Then, by heating to the
melting point of the methane-IL complex solid, the
gas collected was detected by MS and FT-IR, respectively. Blank experiments (No IL in autoclave) were
also carried out using the same procedures.
For molecular liquids, single parameter empirical
polarity scales methods such as observing shifts in the
absorption maxima of solvatochromic or fluorescent
dyes have been developed. These methods have also
been used to analyze the polarity of the ILs[ll-151,
but definitive results have not been obtained so far,
since the values all fall to within a narrow range and
the order varies somewhat with the dye employed[l6171. The IR method has the advantage of simplicity
and has wide applications in the characterization of
solvent effects [18]. We have recently shown how the
polarity of the ILs can be correlated with the shift of
the IR absorption bands using acetone or Fe(C0)S as
spectroscopic probes, and demonstrated that the vibrational modes associated with the carbonyl group
are affected by the surrounding solvent molecules[l9].
The C=O stretching frequencies of acetone dissolved in different ILs (Their structures shown in Figure 1) are shown in Table 1. In each case, a red shift
from the value of pure acetone (1715 c111-l) is observed. Though the data are indicative of lowest polarities for quaternary ammonium-based ionic liquids
containing the Tf2N- anion, the red shifts observed
are similar to and therefore cannot be differentiated
from each other. Iron pentacarbonyl, a more sensitive probe molecule, is therefore employed to compare
the polarity of the ILs listed in Table 1. In an acetone solution, Fe(CO)5 gives peaks at 2022 cm-' and
1996 cm-', corresponding to CO stretching modes.
On adding the acetone solution of Fe(CO)5 to the ILs,
the above two peaks are found to be accompanied with
the appearance of two characteristic shoulders at the
higher and lower wave number sides. Based on the
previous results [19], the shift of the lower wave number shoulder, peak 4, suggests the polarity of the ILs,
284
Yuan Kou et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
"X2221TfZN
"x4441Tf2N
Figure 1. The structure o f ionic liquids
i.e., a higher value indicates a lower polarity. For
imidazolium-based ionic liquids using C4 as the alkyl
substituent in cations, it can be seen from the Table 1 that, when changing the anions from PF, and
BF, to TfiN-, peak 4 shifts from 1987 cm-' and
1987 cm-' to 1994 cm-', giving the polarity order
[C4mim]PF6~[C4mim]BF4>[C4mim]TfiN.
Using Cg
as the alkyl group, similar trends can be observed.
ILs containing TfzN- anions seem to have a lower
polarity, whilst the increase in length of the alkyl
suhstituents in the cations leads to the lowering of
the polarity of the ILs. This conclusion is further
proved by quaternary ammonium-based ionic liquids.
For quaternary ammonium-based ionic liquids, when
TfZN- was used as anions, it can be seen from the
Table 1 that the increase in the total number of carbon atoms in the cations leads to a peak 4 shifting from 1993 cm-' to 1998 cm-', giving the polarity order [N4222]Tf2N> [N8~22]Tf2N-[N8444]Tf2N
> [Ns666]Tf2N>[ N g g g g ] T f 2 N . Since the salts combine
these quaternary ammonium cations with BF, or
PF, anions are not in a liquid state around room
temperature, we did not discuss them in the Table 1.
It is important to note that the [N888g]Tf2Nis for the
first time demonstrated to be an IL with the lowest
polarity.
Table 1. C=O stretching frequencies o f acetone and Fe(CO)5 and property o f Cao i n different ILs
VC=O
of Fe(CO)5 (cm-l)b
Solvent
vc=o (cm-l)a
Acetone
[CqmimIPFs
[CsrnimlPFs
[C4mim]BF4
[Cemim]BF4
[C,mim]TfzN
[CsmimITfzN
1715
1710
2041
2019
2001
1987
1711
1710
1711
1712
1714
2037
2038
2019
2002
1991
2018
2019
2001
2002
2004
2004
[N4zzz]TfzN
1713
1714
2039
2036
2021
[Nszz2]Tf2N
[Ns444ITfiN
[Nsass]TfzN
[Nssss]TfzN
2004
2004
1714
1714
1714
2036
2036
2035
2020
2021
2021
a
Peak 1
Peak 2
Peak 3
2022
2035
2037
2035
Peak4
-
-
3 . 4 ~
<0.1
pale yellow
6.1 x
<O. 1
1987
1992
1994
1995
2.9x10-'
8.0~
<0.1
pale yellow
pale yellow
6.9~
8 . 9 lop2
~
<0.1
<0.1
<0.1
1993
1994
4.1~10-'
6 . 1 lo-'
~
<0.1
<0.1
1994
1996
1998
3 . 5 lo-'
~
1.2
1.8
<0.1
0.10
0.18
1996
2021
2020
2020
The values are C=O stretching frequencies of acetone,
Based on their low polarity characteristics, we attempted to use CGOas the probe molecule to demonstrate the unique properties of the [N*gsg]Tf2N. By
virtue of consisting purely carbon atoms, fullerenes
are archetypal non-polar materials. A poor solubility of the c 6 0 in most organic solvents has been
2004
2004
2003
c 6 0 property in c6O-ILS
Solubility
Color
(mg/ml)
(mol)%
~
pale yellow
pale yellow
pale yellow
pale yellow
pale yellow
yellow
pale purple
purple
in acetone solution.
one of the main impediments for its applications.
Combination of the unique properties of fullerenes
with those of the ILs should be of great interest.
But there have been few such reports in the literature, perhaps due to the fact that ILs are believed t o be very polar materials and the solubility
285
of fullerenes in polar organic solvents is kriowii to
be very low [20]. We prepared a series of C60-IL
solutions according t o a reported method [lo], and
measured the solubilities by UV-visible absorption
spectra. Table 1 gives the solubility order of c60,
[C4mim]BF4< [C4mim]PF6< [N422z]Tf2N< "82221Tf2N [C8mim]PF6< [Cqmim]Tf~N<[C~mini]BF4
< [Csmim]TfZN<[N8444]TfZN<[N6666]TfZN<"88881Tf2N. The lowest-polarity IL, [N8888]Tf2N,could dissolve much more c60 than the higher polarity ones.
The solubility in [N8888]TfiNcan be up to 1.8mg/ml,
which is higher than those of the most common organic solvents, and about 20 times higher than the
best results reported[lO].
Generally, CH4 is very difficult to be dissolved
in HzO. It can be seen from Table 2 that only less
than 0.1 mol% CH4 can be dissolved in H 2 0 under
4 MPa at 30 "C for 5 h. Though nearly 5 mol% CH4
can be dissolved in a nonpolar solvent, CCl4, under
the same conditions, the solubility of CH4 in the conventional IL is still very low, less than that in CCl4
(Table 2). Exposure of the [N8888] TfzN to gaseous
CH4 under 0.5-4 MPa in a temperature range of 3090 "C for 3-5 h caused the formation of a solution of
CH4 in the IL. Figure 2 shows that the solubility of
CH4 increases consistently with the increasc of pressure. Under higher pressures, the solubility of CH4
decreases slightly, but it decreases consistently with
the increase of the temperature. The highest solubility obtained under 4 MPa CH4 pressure a t 30 "C
for 5 h was 27 mol%. Note that the conditions used
here represent the simplest operation, for example,
N
the pressures of 0.5-4 MPa are in the range of a safe
and cost-effective pressure limit [3],and the temperature of 30 "C is close to room temperature. It is
also worth noting that the highest solubility obtained,
which was 27 mol%, was almost 5 times of that in a
CC14 solution or in a COz-enhanced IL system[4] under comparable conditions.
Table 2. Solubilities of methane in different solvents
Entry
Solvent
Solubility (mol%)
1
H2O
cc14
<0.1
[Czmim]BF4
[C*mim]BF4
0.47
3.9
[C4mim]PFa
[CqmimITfzN
INxxnxlTfzN
3.6
5.2
27
2
3
4
5
6
7
5
Reaction conditions: time 5 h, T 30 "C, P 4 MPa.
0.30
--
0.25
u
8
.
-a.l
8
-
-5.-x
-
F *
0.20
0.15
0.10
ZJ
rA
0
0.05
0
20
30
40
50
60
70
Temperature ('C)
80
90
100
Figure 2. Isobaric solubility curve of methane in
[Nssss]TfzN
(1) 0.5 MPa, (2) 1.5 MPa, ( 3 ) 3 MPa, (4) 4 MPa
Figure 3. The releasing methane from [NssssITfzN.
(a) White CH4-IL complex, (b) Bubbles of methane being released, (c) Ionic liquid after releming methane (very small bubble still
remained in the IL)
286
Yuan Kou et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
After cooling the C&-[N8888] Tf2N solution down
to room temperature, a white solid stable under ambient conditions was obtained, and as perfect as expected (Figure 3a). This white solid was readily converted back into liquid by heating to 30 "C. Following
its melting, CH4 bubbles were clearly observed (Figure 3b). After releasing CH4, the IL turned back to
the pure state (Figure 3c). The IL can repeatedly
dissolve and release CH4 for more than hundred times
without any decrease in activity. TGA analysis on the
solid being kept in a refrigerator for two days yielded a
weight loss of 0.56 wt%, with a raising of the temperature from 30 "C to 40 "C, corresponding to 26 mol%,
almost equal to the highest solubility obtained. We
conclude that a reversible storing-releasing process
switched by a temperature difference of 10 "C around
room temperature has been realized.
4. Conclusions
In conclusion, we have designed a low-polar functionalized IL for the capture and storage of methane.
A fast, recyclable, energy efficient, low cost and green
system within a wide range of temperatures and pressures is reported here for the first time. The system is
believed to have opened up a new route for the storage
and homogeneous utilization of methane.
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Article
Catalytic Combustion of Methane over Col-,Mg,O/A1203/FeCrAl
Monolithic Catalysts
Liping Zhao,
Shengfu Ji*,
Fengxiang Yin,
Zexiang Lu,
Hui Liu,
Chengyue Li*
State Key Laboratory of Chemical Resource Engineering, Beijing University of Chemical Technology, Beijing 100029, China
[Manuscript received April 19, 2006; revised May 22, 20061
Abstract:
A series of Col-,Mg,O/A1203/FeCrAl
catalysts (z=O-l) were prepared. The structures
of the catalysts were characterized using XRD, SEM, and TPR analyses. The catalytic activity of the
catalysts for methane combustion was evaluated in a continuous flow microreactor. The results indicated
that the active washcoats adhered well on the FeCrAl foils. The phases in the catalysts were Col-,Mg,O
solid solutions, a-A1203, and y-A1203. The surface particle size of the catalysts varied with variations
in the molar ratios of Co to Mg. The Co component of the Col-,Mg,O/Alz03/FeCrAl catalysts played
an important role in the catalytic activity for methane combustion. In the Co1-.Mg,0/A1~03/FeCrAl
series catalyst (z=0.2-0.8), the catalytic activity in terms of z was in the order of 0.5>0.2>0.8 under the
experimental conditions. The presence of Mg in these catalysts could promote the thermal stability to
a large extent. There were strong interactions between the Col-,Mg,O oxides and the Al203/FeCrAl
supports.
Key words: catalytic combustion; methane; metallic monolithic catalyst; XRD; SEM; T P R
1. Introduction
Catalytic combustion has been proposed as an
effective method for the oxidation of fuel/air lean mixtures with low emission of NOz, CO, and unburnt
hydrocarbons, due t o the lower combustion temperature in comparison with the conventional thermal
combustion method [l-31. According to the active
components, the catalysts of methane combustion can
be classified into the noble metal catalysts (Pd, Pt,
Rh, and Au) and the various metal-oxide catalysts
(single metal-oxides, perovskites, solid solution, and
hexaaluminates) [4,5]. Although noble metal-based
catalysts show very high specific activity, their utilization in combustors is limited by their high cost, high
volatility of pure metals and their oxides, and the tendency toward sintering at moderate temperature [l51. In the development of more suitable catalysts for
methane combustion, considerable attention is paid to
the solid solution catalysts, which are the new promising hydrocarbon combustion materials, owing to their
high thermal stability and relatively low cost. As a
good active catalyst for methane combustion, Co304
has been extensively investigated [6-91, and it was
generally formed as solid solutions, with MgO, ZrO2,
TiOz, and A1203 stabilizing the cobalt ions to avoid
or retard the sintering process. MgO specifically can
form a solid solution with COOover a wide concentration range because the Mg2+ and Co2+ ions have similar sizes. In these solid solutions, MgO displays a high
melting point and thermal stability and maintains a
relatively high surface area under extreme reaction
conditions [9]. However, the conventional fixed-bed
reactor randomly packed with the pellets of catalysts
has a high pressure drop and poor heat transfer that
can induce hot spots during the exothermic reaction.
Hence, in the high gas hourly space velocity (GHSV)
* Corresponding authors. Tel: +86-10-64412054; Fax: +86-10-64419619; E-mail: jisfQmail.buct,edu.cn, [email protected]
288
Liping Zhao et d./Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
and in highly exothermal catalytic gas-solid reactions
such as the catalytic combustion of methane, the catalysts are prone to sintering, which results in the drop
in activity.
Recently, monolithic catalysts and reactors, especially those using the metal as the catalyst supports,
have received considerable attention [10,11]. The catalysts and reactors can provide some suitable order
flow channels in many forms based on the reaction
type; they can even be made into monolithic catalytic
reactors with honeycomb structures. Compared with
conventional fixed-bed reactors loaded with catalyst
pellets, those using metallic monolithic catalysts have
lower pressure drops, a smaller size of the reactor,
a higher thermal conductivity and mechanical shock
resistance, and a lower temperature gradient [12-151.
Hence, the metallic monolithic catalysts have promising applications for the reactions with high gas hourly
space velocity and heat effect, such as the methane
catalytic combustion reaction.
Usually, the monolithic catalyst is composed of
a high surface area inorganic oxide carrier, i.e., yA1203, upon which the active components are dispersed [13]. There are some reports on the methods
for depositing the A1203 washcoats onto the FeCrAl
alloy foil [14-161. However, at present there is little
investigation on the methods for supporting the active
phase onto the metal supports [17]. In the previous
work, a series of the Cel-,Cu,O~-,/A1203/FeCrAl
monolithc catalysts (z=O-l) [18]were prepared, and
the effect of the mass transfer and heat transfer as
well as the hydrodynamics on the metallic monolithic
reactor performances of the catalytic combustion was
investigated by modeling and simulation with the
CFD software [19,20]. In this study, a series of the
Col-,Mg,O (2=0.2-1) solid solution catalysts and
the Co304-based metallic monolithic catalysts were
prepared, and their structure and the catalytic activity for methane combustion were investigated; the
effects of the monolithic support (Al203/FeCrAl) and
the Mg component on the activity and stability of the
catalysts were discussed.
2. E x p e r i m e n t a l
2.1. Preparation of the catalysts
The Col-,Mg,O solid solutions were prepared using Co(N03)2.6H20 and Mg(N03)2.6H20 as precursors according to the urea combustion methods [21].
Urea was added to the Co and Mg salts in the desired
ratio, and the mixture was mulled at room temperature for 1 h. The mixture was then calcined at 650.850 "C for 10-30 min t o obtain the Col-,Mg,O solid
solutions.
The monolithic catalysts were prepared using the
FeCrAl alloy foils (OC404, Sandvik Steel, Sweden)
as supports. The FeCrAl alloy foil flats were rolled
into cylinders of different diameters and a length of
100 mm. These were cleaned in ethanol, and in acidic
and basic solutions, respectively, to remove the oil, the
primary oxides, and the other superficial impurities.
The cylindrical alloy foils were then thoroughly rinsed
in deionized water and were finally calcined at 950 "C
for 15 h in air to form the oxidized metallic supports.
Subsequently, the heat-treated metallic supports were
immersed in a boehmite primer sol, with a withdrawal
velocity of 3 cm/min to ensure uniformity and then
dried at room temperature in air and were finally calcined at 500 "C for 4 h. The boehmite primer sol was
prepared according to the literature [16], which was
used as the first washcoat layer to improve the adhesion between the washcoat layers and the heat-treated
metallic supports. Similarly, the y-A1203 slurry used
as a second washcoat layer to increase the surface area
was washcoated onto the samples. The y-A1203 slurry
was prepared by wet milling according to the literature [14,15]. The monolithic support A1203/FeCrAl
was then formed after which the mixture slurry of the
Col-,Mg,O solid solutions and the y-A1203 was deposited onto the monolithic supports. The mixture
slurry was prepared by wet milling. The weight ratio
of the solid solution t o y-Alz03 was l:lO, and nitric
acid was used as the stabilizing agent. Finally, the
Col-,Mg,O/A1203/FeCrAl catalysts were prepared
with the weight of the solid solution layers and the
A1203 washcoat layers as ca. 8 wt% and ca. 10 wt%,
respectively.
2.2. C h a r a c t e r i z a t i o n of the catalysts
X-ray diffraction (XRD) was performed to examine the surface phase composition of the coatings using a Rigaku D/Max 2500 VB2+/PC diffractometer
with Cu K , radiation, operated at tube voltage 40 kV
and current 200 mA. The diffraction spectra were
scanned between 10-80' (28) at the rate of 10°/min.
Scanning electron microscopy (SEM) was used to observe the morphologies of the catalysts on a Cambridge Instruments Streoscan 250MK3 scanning electron microscope. The temperature programmed reduction (TPR) was carried out using a Thermo Elec-
289
tron Corporation TPDRO 1100 Series Catalytic Surface Analyzer equipped with a thermal conductivity
detector (TCD). The samples were preheated with
10 vol% O2/He at the rate of 20 "C/min from room
temperature to 500 "C and were then cooled to room
temperature in N2 flow, and thereafter the temperatures were reduced stepping up from 40 to 1000 "C at
20 "C/min in the gas stream of 5 vol% H2/N2 a t
20 ml/min. The water produced by the sample reduction was condensed in a cold trap before reaching
the detectors. Only H2 was detected in the outlet
gas confirming the effectiveness of the cold trap. The
coating adherence was measured by an ultrasonic vibration test using KQ-400 DB. The washcoated samples were treated in an ultrasonic bath for 30 min a t
160 W, which showed that the weight loss of both
the A1203 layers and the activity layers was less than
3 wt%.
aluminum [14,15]. From Figure 1(3),it is found that
diffraction peaks of 7-A1203 are observed at 37.0°,
45.9O, and 67.0°, respectively, in addition t o the characteristic peaks of FeCr and cu-Al203. This indicates
that, the monolithic support (A1203/FeCrAl) is obtained by coating the 7-A1203 slurry on the surface
of the metallic supports. The peaks of -pA1203 are
wider, suggesting that the particles of y-Al2O3 are
dispersed finely on the surface of the metallic supports.
*
FeCr
T
Co@,
v
a-A120;
A
v
MgO
y-Al,O,
rn
*.
Co,,Mg,,O
v
*.
I
2.3. Catalytic activity test
The catalytic combustion of methane was carried
out using a conventional quartz reactor (i.d, 8 mm;
length, 300 mm) in a fixed-bed continuous-flow system at atmospheric pressure. The monolithic catalysts were fixed in the center of the reactor. The
reaction gas, composed of 2 vol% CH4 and air t o
loo%, was fed into the reactor with a gas hourly
space velocity (GHSV) of 10700, 21400, and 42800
ml/g(A1203+cat).
h. The outlet products were analyzed
using gas chromatography (Beijing East and West
Electronics Institute, GC-4000A) after the methane
combustion was stabilized for 30 min a t the required
temperature. Carbon dioxide and water were the
only reaction products detected during the entire experiment. Following the completion of the catalyticactivity tests, the catalysts were used continuously for
100 h to test their stability for the methane combustion.
3. Results
3.1. XRD of the samples
The XRD patterns of the catalysts are shown in
Figure 1. It can be seen that there are characteristic
peaks of FeCr (JCPDS 34-0396, Figure l(1)). After the thermal treatment at 950 "C for 15 h, the
peaks of a-A1203 (JCPDS 88-0826, Figure l ( 2 ) ) appear, indicating that a-A1203 forms on the surface of
FeCrAl because of the segregation and oxidation of
i
10
20
iii
30
40
i i
50
60
ii
70
80
281(Q)
Figure 1. XRD patterns of the preoxidized metallic
support and the fresh monolithic catalysts
(1) the FeCrAl foils, (2) the FeCrAl foil pre oxidized at
950 "C for 15 h, (3) A1203/FeCrAI, (4) bulk MgO,
(5) MgO/A1203/FeCrAI, (6) Coo.zMgo.~O/A1~03/FeCrAI,
(7) bulk Coo.5Mgo.50, ( 8 ) C O O . ~ M ~ O . ~ O / A ~ Z O ~ / F ~ C ~ A
(9) Coo.8Mgo.zO/A1203/FeCrAI, (10) C0304/A1~03/FeCrA1,
(11) bulk Co304 catalysts
The XRD patterns of the Col-,Mg,O/Al203
/FeCrAl samples (2=0.2-1) and CogOq/A1203
/FeCrAl can be observed in Figure 1(5), 1(6) and Figure l(8)-l(l0). Herein, the XRD patterns of MgO
(Figure 1(4)), Co0.5Mgo.50 (Figure 1(7)), and Co304
(Figure l(11))powders are provided for better understanding of the phase structures of these samples.
According to literature [22], the major peaks at 37.0°,
43.0", and 62.3" over the bulk MgO (Figure l ( 4 ) ) are
identified as MgO. The major peaks a t 36.5O, 42.5O,
61.7", 73.9", and 77.7" over the Coo.sMgo.50 solid solution (Figure l ( 7 ) ) are assigned to the Coo.sMgo.50
solid solution [7]. However, it is difficult to detect the peaks of MgO in Figure l(5) and those of
the Col-,Mg,O solid solutions in Figure 1(6), l ( 8 )
and 1(9), suggesting that the MgO and Col-,Mg,O
290
Liping Zhao et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
on the A1203/FeCrAl monolithic supports are highly
dispersed. From the XRD patterns of Co304 (Figure 1(11)), it can be found that the peaks at 19.2',
31°, 36.8', and 59' are identified as Co304. At
the same time, the peak of Co304 can be observed
in the Co304/A1203/FeCrAl catalyst (Figure 1( lo)),
whereas no diffraction peaks of Co304 are seen in
the Col -,Mg, O/Alz 0 3/FeCrAl monolithic catalysts
(x=0.2-0.8) (Figure 1(6), l(8) and l(9)). It is likely
that the concentration of Co304 is very less and the
phase is highly dispersed.
3.2. Surface morphology of the samples
Figure 2 shows the surface morphologies of the
heat-treated FeCrAl support, the Co304/A1203/Fe
CrA1, and the Col-,Mg,O/AlzO?,
/FeCrAl (2=0.21) monolithic catalysts. It is seen that after the FeCrAl metal support was calcined at 950 "C for 15 h
in air, several porous whiskers were formed on the
oxidized surface (Figure 2(a)). The analysis results
of XRD show that the porous whisker materials are
the a-A1203 layers, which are formed due to the aluminum segregated from the inner layer and oxidized
into the many regularly arranged a-alumina crystal
clusters on the surface. A certain specific surfacearea of the oxidized layer is still maintained, which
can possibly be attributed to the adherence of the
successive washcoat on the heat-treated FeCr Al.
Figure 2. SEM images of the preoxidixed metallic support and the fresh monolithic catalysts
( a ) FeCrAl at 950 "C for 15 h, (b) MgO/A1203/FeCrAl, ( c ) Co0,zMgo,80/A1203/FeCrAl, (d) Coo.5Mgo.50/A1203/FeCrAl,
( e ) Coo.~Mgo.~O/A1203/FeCrAl,
( f ) Co304/A1203/FeCrAl
291
As for the Col-,Mg,O/A1203/FeCrAl
catalysts, the surface morphologies differ significantly
with the value of 2 .
For the morphology of
MgO/A1203/FeCrAl (Figure 2(b)), it is evident that
the surface of the MgO/A1203/FeCrAl sample is quite
nonhomogeneous with high porosity, and the particle heap is loose, and the clusters of the particles
are large. The surface morphologies of the catalysts change significantly when the Co content increases. Figure 2(c)-2(e) show the morphologies of
the Col-,Mg,O/Al203/FeCrAl (x=O.2-0.8) sample.
When x=0.8, namely Coo,2Mgo.~O/Alz
0 3 /FeCr A1
(Figure 2(c)), the dispersion is considerably more homogenous, the particles on the surface are small and
are finely dispersed, the clusters of the particles disappear, and the particle heap is compact and no
cracks appear. When the Co content in the catalysts increases to 2=0.5 (Figure 2(d)), the particle
size remains almost unchanged and no cracks appear. However, for the catalyst with 2=0.2, namely
Coo.~Mgo,2O/A1~03/FeCrAl
(Figure 2(c)), there is a
crack on the surface of the catalyst and the particle
sizes are similar to each other, about 0.2-1 pm. The
morphology of the C0304/A1203/FeCrAl sample is
shown in Figure 2(f), from which it is evident that
the cluster of the particles is large and the particle
heap is loose.
3.3. Catalytic activity
The catalytic activities of the samples are shown
in Figures 3, 4, and 5. The temperatures of 10%
methane conversion ( T I o,)50% methane conversion
( T ~ o and
) , 90% methane conversion (Tgo) are summarized in Table 1. It is found that the catalytic
activity of the catalysts depends strongly on the
value of x in the Col-,Mg,O solid-solution oxides
and on the reaction conditions. When the GHSV
is 10700 ml/g(Alz03+cat).h(Figure 3, Table l), the
MgO/A1203/FeCrAl catalyst shows the lowest activity over the methane combustion. The T ~and
o Tgo are
552 "C and 741 'C, respectively, higher than those of
all other catalysts. The Co30d/A1203/FeCrAl sample has a higher activity, its TIOis 510 "C, and T ~ o
is 682 "C. For the Col-,Mg,O/A1203/FeCrAl monolithic catalysts, minor differences in the combustion
activity of methane are only present when x=0.2, 0.5,
and 0.8 as shown in Table 1, and the order of the
catalytic activity in terms of x in the catalysts was
0.5 >0.2>0.8 under the experimental conditions.
I00
R
+Co,O,/Al,O,/FeCrAl
--t
Co, ,Mg, ,O/AI,O,/FeCrAI
450
400
500
550
600
Temperature ( 'C)
650
700
Figure 4. Methane
conversion
over
the
Coi-,MgzO/A1203/FeCrAl
catalysts at
GHSV = 21400 rnl/g(Al,Ostcat).h
100,
Co,O,/Al,O,IFeCrAI
i
-Co,,Mg,,O/AI,O,/FeCrAI
-o- Co,,M&,,O/AI,O,/FeCrAI
+Co,,Mg,,O/AI,O,/FeCrAI
-m-
Figure 3. Methane
conversion
over
the
Col-,Mg,O/A1203/FeCrAl
catalysts at
GHSV = 10700 ml/g(ALa03+cat).h
B
-m-
-o-
400
Temperature (-C)
750
Co,O,/Al,O,/FeCrAl
Co,,*Mg,,,O/AI,O,/FeCrAI
Co,,,Mg,,,O/AI,O,/FeCrAI
450
500
550
600
650
700
750
Temperature ("C
)
Figure 5. Methane
conversion
over
the
Coi-.MgzO/Al~03/FeCrAl catalysts at
GHSV = 42800 ml/g(Ala03+cat).h
292
Table 1. The Tlo, T s o , and T g o of the Col-,Mg,O/AlaOs/FeCrAl
catalysts under the various GHSV
GHSV/ml/(g.h)
MgO/A120g/FeCrAl
552
Coo.~Mgo.~O/AI~Og/FeCrAl 508
Coo.5 Mgo.5O/A1203 /FeCrAl
503
Co~.~Mgo.~O/A120g/FeCrAl 516
Cog 0 4 /A1203 /FeCrAl
510
691
74 1
612
711
-
625
738
-
610
683
556
680
736
577
702
739
589
673
554
655
732
5 74
675
734
598
677
574
668
736
603
692
740
607
682
558
668
736
585
690
739
From Figures 3 to 5, it is found that with
increasing GHSV, the activity of all the catalysts
decrease to different extents. For example, with
the GHSV increasing from 10700 ml/g(A1203+cat).h
to 21400 rnl/g(AlzOsfcat)~h(Table l ) , the T50 of
the MgO/A1203/FeCrAl catalyst increases from
691 "C to 711 "C. For catalysts with x=0.2, 0.5,
and 0.8, their T50 increases from 598 "C, 589 "C,
and 610 "C to 668 "C, 655 "C, and 680 "C, respectively. The TSOof C0304/A1203/FeCrAl increases
from 607 "C to 668 "C. When the GHSV increases
further to 42800 ml/g(A1203+cat).h(Table l ) ,the T50
of the MgO/A1203/FeCrAl catalyst increases from
711 "C to 738 "C. For catalysts with x=0.2, 0.5, and
0.8, their T ~ increases
o
from 668 "C, 655 "C, and
680 "C to 692 "C, 675 "C, and 702 "C, respectively. The
T ~ofoCo304/Al203/FeCrAl increases from 668 "C to
690 "C.
Based on the above catalytic activity results of the
Col _,Mg,O/Al203/FeCrAl catalysts (x=0.2-1) and
the Co304/A1203/FeCrAl catalysts, it can be concluded that under the experimental conditions, the
order of the catalytic activity in terms of the value of
x was 0.5>0.2>0.8>1. However, the order is different
from that of the catalysts in the literatures [7,9],suggesting that the catalytic activity of the samples is
affected not only by the value of x but also by other
factors, such as the A1203 washcoats and the FeCrAl
support.
lysts showed better stability in the reaction and
stabilized to an almost constant value. When the
reaction time was 0.5 h, the methane conversion
on the Co~-,Mg,O/A1203/FeCrAl catalysts with
x=0.2, 0.5, and 0.8 was 99.7%, 98.6%, and 97.6%,
respectively. During the 100-hour activity test, the
methane conversion of these catalysts remained almost unchanged, and when the reaction time was
100 h, the methane conversion for the catalysts
with x=0.2, 0.5, and 0.8 was 97.9%, 99.4%, and
97.3%, respectively. The methane conversion on the
Co304/A1203/FeCrAl catalysts decreased from 99 to
95% after the reaction was run for 5 h, then decreased
to 91% for the next 35 h, and finally, the value remained almost at 91% for the next 60 h.
8 95 c+
z
:
I
Co,O,/AI,O,/FeCrAl
t
--
c Coo,Mg,,OIA1,O1/FeCrAl
Co,,M&,O/AI,O,/FeCrAI
t
_0
20
40
60
Reaction time (h)
80
I00
Figure 6. Methane conversion vs time on stream at
740 "C, GHSV=10700 m l / g ( ~ l , ~ , + , , t ) . h
3.4. The stability of the catalysts
The results of the 100-hour stability test of the
catalysts at 740 "C under 10700 ml/g(A1203+cat).h
gas hourly space velocity are shown in Figure 6.
It was found that at first, the methane conversion
over the MgO/A1203/FeCrAl catalyst sharply decreased from 92% to 86.9% after the reaction was
run for 5 h, and then the methane conversion slightly
decreased from 86.9% to 81% for the next 95 h.
The Col -,Mg,O/A1203 /FeCrAl (x=O.2-0.8) cata-
According to the literatures [9,23], Co304 undergoes severe sintering above 527 "C so that its
thermal stability is rather limited. However, when
compared with the result of the stability test of
the Co304/A1203/FeCrAl catalyst, it is inferred
that the thermal stability of Co304 is improved
by the A1203 washcoats and the FeCrAl support. From the results shown in Figure 6, it, is
found that the Col-,Mg,O/A1203/FeCrAl catalysts
(x=O.2-0.8) have better thermal stability in the pres-
293
ence of MgO.
reduction peaks observed at around 899 "C over
the Col-,Mg,O/A1203/FeCrAl
catalysts (2=0.20.8) (Figure 7(2)-7(4)) in this study can be attributed
to the reduction of the Col-,Mg,O solid solution.
3.5. H2-TPR analysis
The H2-TPR profiles of the Co304/A1203/FeCrAI
and Col-,Mg,O/A1203/FeCrAl
(2=0.2-1) catalysts are presented in Figure 7. For comparison,
the H2-TPR of pure Col-,Mg,O solid solution and
Col-,Mg,O/Al203
were measured, and the T P R
profiles are shown in Figures 8 and 9, respectively.
In Figure 7(5), the peak at 488 "C can be attributed to the large crystalline Co304, as inferred
from the profile of H2-TPR of bulk Co304 (Figure
8(5)). It has been reported that for Co304, there is
one main reduction peak, or two reduction peaks of
H2 below 500 "C, which can be assigned to the twostep reduction of Co3+ to Co2+ and then to the Co
metal [22,24-261. Hence, one main reduction peak
observed a t 488 "C for Co304/A1203/FeCrAl in the
Figure 7(5) can be assigned t o the overlap of the twostep reduction of Co3+ to Co2+ and then to the Co
metal. The reduction peaks a t 683 "C observed in all
the samples can be considered to be the reduction of
surface A13+ on 7-Al203 [27,28]
488
:
553
1
0
200
1
1
,
1
400
1
I
,
l
600
.
899
,
I
I
800
,
,
I
I000
Temperature ( % )
Figure 7. TPR patterns of the Col-,Mg,O/AlzOa
/FeCrAl catalysts
(1) MgO/A1203/FeCrAl, (2) C00.z Mgo.8O/Alz 03/FeCrAl,
(3) Coo.sMgo.50/A1203/FeCrAl, (4) Coo..~Mg0,20/A1203
/FeCrAl, ( 5 ) Co304/Al203/FeCrAl
According to the literatures [6,22], the peak observed at 553 "C in the Figure 7(2)-7(4) can be
attributed t o the reduction of MgCo204. Ruckenstein and coworkers [29,30] have concluded that
a complete reduction of the CoO-MgO solid solution phase requires a temperature higher than
lOOO"C, and Furusawa et al. [22] concluded that
it was difficult to completely reduce the CoO-MgO
solid-solution phase until 900 "C. Hence, the small
1
482
43 I
(1)
I
.
.
,
A.
I , , ,
400
709
.
I
I
,
,
'
,
,
I
,
,
,
,
800
I000
Figure 8. TPR patterns of the Col-,Mg,O
lution
(1) MgO, (2) Coo.2Mgo.eO, (3) Coo.sMgo.50,
(4)Coo.sMgo.20, ( 5 ) Co304
solid so-
0
200
,
600
Temperature ( % )
Figure 8 shows the profiles of H2-TPR of the
C03O4 and Col-,Mg,O samples. For the bulk MgO
(Figure 8(1)), the intensities of the two reduction
peaks at 431 "C and 709 "C were considerably lower
than the others, and the H2 consumption was negligible. The main peak in Figure 8(5) can be assigned
to Co304. With z increasing, the stronger interaction between cobalt and magnesium led t o two separate peaks (between 300 "C and 600 "C, Figure 8(2)8(4)) during the T P R of the Col-,Mg,O samples.
Wang et al. [30] reported that Co304 and MgCo204
have the same spinel structure, but the latter has a
higher lattice energy than the former, because of the
partial substitution of Co by Mg. Consequently, the
latter requires a higher reduction temperature than
the former. Hence, the two peaks (between 300 and
600 "C, Figure 8(2)-8(4)) can be attributed to Co304
and MgCozO4, respectively [6,30,31]. However, no
diffraction peaks of Co304 and MgCozO4 were found
in the XRD result, suggesting that their content was
very less.
Figure 9 shows the T P R patterns of
the Col -,Mg,O /A1203 samples.
In the
Col-,Mg,O/Al203 samples, the reduction peaks at
434 "C and 502 "C (Figure 9(5)) can be attributed to
the bulk Co304. According to the literatures [32,33],
the reduction peaks at 706 "C with a shoulder observed on C0304/A1203 (Figure 9(5)) at 663 "C can
be described as disordered, X-ray amorphous sur-
294
face overlayers of the Co oxide. The T P R profiles
significantly change with the increase of the Mg content. The temperatures of the peaks a t 434 "C and
502 "C increase with the increase in x. The intensities of the peaks increase with x increasing upto
0.5 and then decrease with x increasing further, and
the peak at 502 "C disappears when x is 0.8. The
peak at 663 "C shifts to a lower temperature and the
intensity increases when x increases upto 0.2; it then
decreases with further increase in x. The temperature of the peak at 706 "C becomes higher, and the
peak becomes sharper with increase in x. The reduction peak at 874 "C, which can be attributed to a
cobalt-aluminum oxide compound, similar to cobalt
aluminate (CoA1204) [24,32,33],remains almost unchanged with the increase in the Mg content.
706
663 A
4. Discussion
The Col-,Mg,0/A1203/FeCrAl
catalysts have
been shown to be active toward high-temperature
methane catalytic combustion (Figure 3 to Figure 5).
The MgO/A1203/FeCrAl catalyst is less active in
comparison with other catalysts that are used for
methane combustion under the conditions studied; it
would appear that the Co component plays a n important role in the metal monolithic catalysts. According to the literature [7], after the cobalt entered
the lattice of MgO forming a uniform solid solution,
the highly dispersed Co2+ forms the active site for
methane combustion, or the presence of Co2+ in the
MgO lattice has a synergic effect on the combustion of methane, thus decreasing the methane light-
off temperatures in the reaction [7]. The catalytic
activities of the catalysts with x=0.2 and 0.8 were
lower than those of the catalyst with x=0.5, suggesting that the catalytic activity decreased when the Co
content in the catalysts was either very high or very
low. For the cobalt-magnesium solid-solution catalysts, the catalyst activity depends on the number of
active sites. A higher dispersion of the catalyst leads
to more active sites, and a t the low loadings, only
a small proportion of the Co that enters the solid
solution is present on the exposed surfaces. However, as the Co content of the sample increases, a
larger amount is exposed on the surface. These Co2+
sites are accessible in comparison with those held in
the bulk, and hence, the activity increases with the
surface Co2+ content [7,9]. With the considerable
increase in cobalt content, the agglomeration of Co
(e.g. Co304) takes place, and the catalyst particles
become larger, which reduces the available Co surface area and limits the observed catalytic activity [7].
Thus, there is an optimal Co content, and the catalyst
with the optimal Co content shows the best activity.
According to the results of the T P R measurement
in Figure 8, three reduction peaks are observed for
C03O4, MgCo204, and Col-,Mg,O; with increase in
x, the intensity of the C03O4 peak decreases. For
example, for the Coo.sMgo.50 solid solution (Figure 8(4)), the area of the Co304 (481 "C) peak is
smaller than that of the Co0.8Mgo.20 solid solution, whereas for Coo.2Mgo.80, it is negligible. Thus,
for the Col-,Mg,O/A1203/FeCrAl (x=O.2-0.8) catalysts, the Co serves as the active surface site in
the metallic monolithic catalysts, and the synergetic
effect of Co/Mg is possibly responsible for improving
the activity of the catalyst. At the same time, the activity order of the Col-,Mg,O/A1203/FeCrAl series
catalysts is different from that of the catalysts in the
literatures [7,9],suggesting that the catalytic activity
of the samples is not only affected by x but also by
the A1203/FeCrAl support.
When GHSV was 10700 ml/g(A120s+cat).h
(Figure 3), the catalytic activity of the
MgO/A1203/FeCrAl catalyst was far lower than
that of the other catalysts, and at the same time,
the catalytic activity of the CogO*/A1203/FeCrAl
catalyst was close t o Coo,2Mgo,8/A1203/FeCrAl,
which was lower than the solid-solution monolithic catalysts with x=0.2, 0.5. When the GHSV
was increased (Figures 4 and 5), the activities
of the Co304/A1203/FeCrAl catalyst and the
COI-, Mg,O/A12 0 3 /FeCrAl (x=O.2-0.8) catalysts
decreased more than that of the MgO/Al203/FeCrAl
catalyst.
Hence, the catalytic activity of the
MgO/A1203/FeCrAl catalyst was close to that of
the other catalysts, suggesting that when the contact
time of the reactant and the catalysts decreases, the
activity of the catalysts decrease.
COO and MgO are known to form solid solutions
in which the Co2+ ions are located at octahedral positions, substituting the Mg2+ sites in the MgO lattice. This solid solution maintains the structure of
MgO, a NaC1-type crystal structure [9]. This structure is beneficial to the stability of the Co2+ ions as
magnesium oxide has a high melting point and thermal stability. Hence, this structure is able to maintain a relatively high surface area under extreme reaction conditions and improve the stability of the catalysts [7-91. The results of the stability test of the
catalysts show that the Col _,Mg,O/Al203/FeCrAl
catalysts (x=0.2-0.8) have better thermal stability
than the Co30d/A1203/FeCrAl catalyst. It indicates
that the Co2+disperses well in the MgO lattice in
the catalysts, which significantly improves the thermal stability of the catalysts. The results of the
TPR in Figure 7 also conform this. The H2 consumption of the Col-,Mg,O
(x=0.2-0.8) solid solution (Figure 7(2)-7(4)) is considerably lesser than
that of Co3O4 (Figure 7(5)), and the temperature
of the peak of the solid solution (899 "C) is much
higher than that of Co304 (488 "C). This indicates that the Col-,Mg,O/A1203/FeCrAl
catalysts
(x=0.2-0.8) are more difficult to be reduced than
C0304/A1203/FeCrAl because of the effect of MgO.
By comparing the TPR profiles of the
Col-,Mg,O/AlzO3/FeCrAl
(2=0.2-0.8) catalysts
(Figure 7(2)-7(4)) with those of the Col-,Mg,O
(x=O.2-0.8) samples (Figure 8(2)-8(4)), it is found
that the reduction peaks of thc Col-,Mg,O solid
solution shifted to the lower temperature and the intensity of the peak decreased sharply. At the same
time, the reduction peak of Co304 is negligible, indicating that there was strong interaction between the
Col-,Mg,O solid solution and the A1203/FeCrAl
monolithic support after the Col-,Mg,O solid solution was loaded onto the monolithic supports.
This interaction changed the redox properties of
Col-,Mg,O/A1203/FeCrAl
and influenced the catalytic activity and the stability of the catalysts.
According to the above result, the formation of
Col-,Mg,O solid solutions, the interaction of the
Col-,Mg,O oxide and the A1203/FeCrAl support,
and the synergetic effect of Co/Mg are equally re-
295
sponsible for improving the performance of the catalyst.
5 . Conclusions
A cobalt-magnesium solid-solution-based monolithic catalysts on FeCrAl alloy foils supports with
A1203 washcoats have been prepared.
Based
on the results, it can be concluded that: (i)
For the Co~-,Mg,O/A1203/FeCrAl (x=0.2-1) and
C0304/A1203/FeCrAl catalysts, the phase structures are ai-A1203, yAl2O3, and Co304. The surface particle-shape and size are related to x in the
Col-,Mg,O solid solution-type oxides. The cobaltmagnesium solid solution is well dispersed on the surface of the metallic supports. (ii) The formation of
the Col-,Mg,O solid solutions, the interaction of
the Col-,Mg,O oxide and the Al203/FeCrA1 support, and the synergetic effect of Co/Mg synthetically
affect the activity of the catalysts. Under the experimental conditions, the order of the catalytic activity
in terms of x was 0.5>0.2>0.8>1.With the increase
in GHSV, the catalytic activities of the different catalysts decreased t o different extents. (iii) In the 100hour activity tests, the Col-,Mg,O/A1203/FeCrAl
moiiolithic catalysts (x=0.2-0.8) show better stabilization than the Co304/A1203/FeCrAl monolithic
catalysts because of the inclusion of MgO. (iv) The results of the T P R analysis indicate that strong interactions between the Col-,Mg,O solid solution and the
A1203/FeCrAl monolithic supports can significantly
change the redox properties.
Acknowledgements
Financial funds from the Chinese Natural Science
Foundation (Project No.: 20376005) and the Specialized
Research Fund for the Doctoral Program of Higher Education (Project No.: 20030010002) are gratefully acknowledged.
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ScienceDirect
Juurnalof Natural Cas Cheniistrg
SCIENCE PRESS
www.elseviermdocate/joge
Article
Promotional Effect of Bismuth as Dopant in Bi-Doped
Vanadyl Pyrophosphate Catalysts for Selective
Oxidation of n-Butane to Maleic Anhydride
Y. H . Taufiq-Yap'*, Y. Kamiya2,
K.P. Tan'
1. Department of Chemistry, Universiti Putra Malaysia, 43400 UPM Serdan,g, Selangor, Malaysia; 2. Graduate
School of Environmental Earth Science, Hokkaido University, Sapporo 060-0810, Japan.
[Manuscript received April 25, 2006; revised July 25, 20061
Abstract:
Bismuth-promoted (1% and 3%) vanadyl pyrophosphate catalysts were prepared by refluxing
Bi(N03)3.5Hz0 and VOP04.2HzO in isobutanol. The incorporation of Bi into the catalysts lattice increased the surface area and lowered the overall V oxidation state. Profiles of temperature programmed
reduction (TPR) in Hz show a significant shift of the maxima of major reduction peaks to lower temperatures for the Bi-promoted catalysts. A new peak was also observed at the low temperature region for
the catalyst with 3% of Bi dopant. The addition of Bi also increased the total amount of oxygen removed
from the catalysts. The reduction pattern and reactivity information provide fundamental insight into the
catalytic properties of the catalysts. Bi-promoted catalysts were found to be highly active (71% and 81%
conversion for 1% and 3% Bi promoted catalysts, respectively, at 703 K), as compared to the unpromoted
material (47% conversion). The higher activity of the Bi-promoted catalysts is due t o that these catalysts
possess highly active and labile lattice oxygen. The better catalytic performance can also be attributed to
the larger surface area.
Key words: bismuth; promoter; vanadyl pyrophosphate; n-butane oxidation
1. Introduction
Selective oxidation of n-butane to maleic anhydride over the vanadium phosphorus oxide (V-P-0)
catalyst is still the only industrial catalytic process
for partial oxidation of an alkane [1,2]. The catalytic
performance may be improved by adding specific d o p
ing agents t o the V-P-0 composition. It is considered that the promoters have a twofold structural role,
namely, to enable the formation of the required V-P-0
compounds and decrease the formation of deleterious
phases, and to enable the formation of solid solutions
that regulate the catalytic activity of the solid [3]. A
wide range of cations have been added as the modifier
and some beneficial effects have been claimed. Many
published data have also indicated their influence on
the yield, on the selectivity for maleic anhydride for-
mation and on the reaction rate over these catalysts
[3-51.
In this study, we explore the modification of
(VO)zPzO7 catalysts induced by bismuth doping on
the physico-chemical characteristics and the catalytic
performance for n-butane oxidation to maleic anhydride. The relationships between the reduction behaviour and reactivity of the catalysts will also be
described and discussed.
2. Experimental
2.1. Preparation of catalysts
2.1.1. Unpromoted vanadyl pyrophosphate
Vanadyl phosphate dihydrate, VOP04.2HzO was
prepared by refluxing VzO5 (12.0 g from Fisher) and
* Corresponding author. Tel: +603-8946 6809; Fax: +603-8946 6758; E-mail: [email protected]
298
Y . H . Taufiq-Yap et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
ortho-phosphoric acid (115.5 g, 85% from Fisher) with
water (24 ml HzO/g solid) for 8 h at 393 K. The resulting VOP04.2H20, which was in yellow colour, was
recovered by filtration, washed with water and dried
at 383 K for 16 h. This solid was confirmed by XRD.
VOP04.2HzO with 4.0 g was refluxed with isobutanol (80 ml from BDH) for 21 h at 393 K with continuous stirring. The light blue slurry was recovered
by filtration, washed and dried at 423 K overnight to
obtain the precursor, VOHP04.0.5H20 (denoted as
PVPD).
2.1.2. Bismuth promoted vanadyl pyrophosphate
For the preparation of the Bi-promoted precursor, bismuth nitrate salt with the required mass was
previously dissolved in isobutanol, prior to refluxing
VOP04.2HzO with isobutanol. The precursors obtained were denoted as PVPDBix where x=1 and 3.
Both unpromoted and Bi-promoted precursors
were then undergone calcination in a reaction flow
of n-butanelair mixture (0.75% n-butane in air) for
75 h at 673 K. The unpromoted catalyst was denoted
as VPD, whereas the Bi-doped catalysts were denoted
as VPDBiz, where x=1 and 3.
DRO 1110 apparatus with a thermal conductivity detector (TCD).
2.3. Catalytic oxidation of n-Butane
Oxidation of n-butane was carried out a t 703 K
in a flow reactor (Pyrex tube, 10 mm of inside diameter) with a mixture of n-butane (1.5 vol%), 0 2
(17 vol%), and He (balance) under atmospheric pressure. After the catalyst (0.2 g) was placed in the
reactor, the reactant gas was introduced at a rate of
10 cm3/min (W/F=540 g.h/mol(bUt,,,)). The temperature was raised t o 703 K a t a rate of 5 K/min,
and when the temperature had reached 703 K, the
gas at the outlet of the reactor was analyzed using
on-line gas chromatography. For n-butane and MA,
an FID GC (Shimadzu GC8A) with a Porapak QS column (1 m) was used. A high speed GC (Aera M200)
with Porapak Q and Molecular Sieves 5A columns was
utilized for the analysis of CO, C02, and 0 2 in the
gas-phase .
After the conversion and selectivity had reached a
stationary state (about 35 h), the W / F dependence of
conversion was determined (W is the catalyst weight
and F is the total flow rate) by changing the total
flow rate.
2.2. Catalyst characterisation
The total surface areas of the catalysts were measured by the BET method using nitrogen adsorption
at 77 K. This was done by a Micromeritics ASAP 2000
nitrogen adsorption/desorption analyzer.
The bulk chemical composition was determined
by using a sequential scanning inductively coupled
plasma-atomic emission spectrometer (ICP-AES) of
Perkin Elmer Emission Spectrometer model Plasma
1000.
The average oxidation numbers of vanadium in
the sample bulk were determined by redox titration
following the method of Niwa and Murakami [6].
The X-ray diffraction (XRD) analyses were carried out using a Shimadzu diffractometer model
XRD 6000, employing CuK, radiation to generate
diffraction patterns from the powder crystalline samples at ambient temperature.
SEM images were taken by a Jeol JSM-6400 electron microscope. The samples were coated with gold
using a Sputter Coater.
T P R (Temperature Programmed Reduction) experiments were done by using a ThermoFinnigan TP-
3.1. BET surface area measurement, chemical
analysis and redox titration
Addition of the Bi promoter had increased the
BET surface area of the VPDBil and VPDBi3 catalysts (Table 1) to 25.0 and 24.8 m2/g, respectively,
compared to 20.3 m2/g for unpromoted VPD. Doping
with Bi into the VPD catalyst has somehow altered
the development of the basal (100) (V0)2P207 face
which is the interesting feature of the high BET surface area of the catalysts. Chemical analysis using
Inductively Coupled Plasma (ICP) indicated that the
ranges of P / V ratio for unpromoted and Bi-promoted
catalysts were between 1.09 and 1.16 and the amount
of Bi was 1%and 3.3% for VPDBil and VPDBi3, respectively. Doping with Bi promoter has resulted in
a decrease of the V5+ contribution from 24% for unpromoted VPD to 8% and 6% for VPDl and VPD3,
respectively. The average oxidation state of vanadium reduced from 4.24 for VPD to 4.08 and 4.06
for VPDBil and VPDBi3, respectively (see Table 1).
299
Table 1. Specific BET surface area, chemical properties, average vanadium valence and percentages
of V4+ and V6f oxidation states present in undoped and Bi-doped VPD catalysts
Specific BET surface
area (m2/al
20.3
Catalyst
VPD
VPDBil
VPDBi3
Atomic ratio
P IV
BiIV
1.09
-
v4+
(%)
v5+ (%)
Average vanadium
valence
76
24
25.0
1.12
0.010
92
8
4.24
4.08
24.8
1.16
0.033
94
6
4.06
3.2. X-ray diffraction
The structures of the promoted precursors and
catalysts were studied by X-ray diffraction (XRD).
Both unpromoted and Bi-promoted precursors (Figure 1) are identical to VOHP04.0.5H20 with peaks
at 28=15.75O, 19.83O, 24.42O, 27.28O and 30.63”. The
XRD patterns for the catalysts are shown in Figure 2. All these materials have patterns similar to
a well crystallized (VO)zP207 with main peaks appearing at 28 =22.9O, 28.4’ and 29.3O which are correspond to the (020), (204) and (221) planes, respectively [7]. The addition of Bi into the catalyst leads
to the reflection a t 28=22.9O which indexed to (020)
plane being more intense, as compared to the unpromoted material.
I
10
20
I
,
,
,
I
I
,
,
.
30
l
.
,
,
,
l
l
l
l
50
40
,
60
20ip )
Figure 2. XRD patterns of VPD,
VPDBi3 catalysts
20
10
30
40
50
60
281(O )
Figure 1. XRD patterns of the undoped and Bi-doped
precursors
VPDBil
and-
Table 2 shows the linewidth of the reflections of
(020) and (204) planes. The parameter used to determine the crystal size is the half width of the (020)
peak. The linewidth increases with the decreasing size
of the crystallites. The decrease in the full width at
half maximum (FWHMs) of the (020) reflection indicates that the thickness of the particles in the (100)
direction decreases. The half width of the (204) peak
changes slightly as well, reflecting a constant crystalline order of the bc plane. The thickness of (204)
is only indicative of the mean “length” at the (204)
face, while the thickness of (200) is more representative of the actual thickness [8]. The particle thickness of (020) was increased from 69.09 A(VPD) to
136.46 A(VPDBi1) and 177.30 A(VPDBi3).
Table 2. XRD data of doped and Bi-doped VPD catalysts
Catalyst
Linewidtha (020)/(O )
Linewidt hb(204)/(O)
ThicknessC(020)/A
Thickness‘ (204)/A
VPD
VPDBil
VPDBi3
1.1600
0.5873
0.4520
0.4616
0.3541
0.3488
69.09
136.46
177.30
173.62
226.32
232.38
a
FWHM of (020) reflection
FWHM of (204) reflection
Plate thickness by means of Scherrer’s formula: T (A)=(0,89xX)/(FWHM x cos 0)
300
Y. H. Taufiq-Yap et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
3.3. Scanning electron microscopy
The morphologies of the unpromoted and Bipromoted catalysts by scanning electron microscopy
(SEM) are shown in Figure 3. The principle structure of the catalysts is the same, consisting of platelike crystals which are arranged into the characteristic rosette-shape clusters. Bi-promoted VPD catalysts (Figures 3b and 3c) show a higher compact
structure with more layered plate-like crystals which
are formed at the surface of clusters. The size of the
rosette-shape clusters observed is smaller than the unpromoted counterpart. This could explain the fact
that a higher surface area was obtained for the Bipromoted catalysts. The layered structure increased
the exposure of the basal (100) ( V O ) Z P Z Oface.
~
3.4. Temperature programmed reduction
Figure 4 shows the temperature programmed reduction (TPR) profiles in a HZ/Ar stream (5% H2 in
Argon, 0.1 MPa, 25 crn3.minp1) using a fresh sample of the catalyst and raising the temperature from
ambient to 1223 K at 5 Ksmin-' in that stream.
The unpromoted catalyst gave a characteristic pattern with three reduction regions, namely a , p and
y, with peak maxima occurring at 863, 1011 and
1143 K. The VPDBil catalyst also gave three reduction peaks. However, the first two peak maxima significantly shifted to 798 and 906 K. An increase of
Bi addition to 3% induced the formation of a new
peak (a') at 838 K despite the a and ,B peaks appearing at lower temperatures of 796 and 915 K, respectively. This effect may be due to the decrease of
the lattice energy induced by the presence of the Bi
dopant. Both the a and p peaks are assigned to the
removal of oxygen species from V5+ and V4+ phases,
respectively [9]. The total amount of 0 2 removed
was 2 . 5 ~ 1 0 ' Pa/g
~
for the unpromoted VPD catalyst. The introduction of Bi increased the amount to
2 . 8 ~ 1 0 ' Pa/g
~
for the VPDBil and 3 . 3 ~ 1 0 " Pa/g
for the VPDBi3 (Table 3). The lowering of the reduction peak temperature and the increment of the
oxygen atoms removed from the lattice by the reaction of Hz may suggest that the oxygen species in the
Bi-promoted catalysts are more reactive and will give
a higher conversion for n-butane oxidation compared
to the unpromoted catalyst.
a
~
I
I
I
.
l
l
,
l
,
l
,
,
,
,
l
,
,
,
,
I
,
.
.
,
I
,
,
I
I
Temperature (K)
Figure 3. SEM micrographs for (a) VPD, (b) VPDBil
and (c) VPDBi3
Figure 4. TPR Profiles
VPDBi3
of
VPD,
VPDBil
and
301
Table 3. Total number of oxygen atoms removed from the undoped and Bi-doped VPD catalysts
by reduction in H2/Ar
Oxygen atom removed from
the catalyst (mol/g)
Tmax/K
Peaks
from the catalyst (atom/g)
Oxygen atom removed
Coverage
(atom/crn2)
7.6 x 1 0 - ~
8.1 x 10''
1.3 x loz1
4.6 x lozo
4.0 x 1015
6.2 x 1015
2.3 x 1015
4.2 x 10W3
2.5 x loz1
1.3 x 10l6
1.4 x 10-3
1.8 x 1 0 - ~
8.7 x lozo
1.1 x 1021
3.5 x 1015
4.4 x 1015
VPD
(Y
P
Y
863
1.3 x 10W3
1011
1143
2.1 x 10-3
Total oxygen atom removed
VPDBil
P
798
906
Y
1151
ff
1.3 x 10-3
8.0 x lo2"
3.2 x 1015
4.6 x
lop3
2.8 x 102'
1.1 x 10'6
1.3 x
lop3
1.3 x lop3
8.1 x 10'"
5.7 x 1020
1.1 x 1021
8.0 x lozo
3.3
2.3
4.3
3.2
4.1 x 10p3
3.3 x 1021
1.3 x 1016
VPDBi3
796
838
ff
a'
P
9.5 x 10-4
1.8 x 1 0 - ~
915
1151
Y
x 1015
x 1015
x 1015
x 1015
Surface area: VPD=20.3 m2/g; VPDBil=25.0 m2/g; VPDBi3=24.8 m2/g;
Weight: VPD=0.0379 g; VPDBil=0.0375 g; VPDBi3=0.0282 g
3.5. Selective oxidation of n-butane
80
t
Figures 5 and 6 show the time courses of the oxidation of n-butane over VPD, VPDBil and VPDBi3.
Figure 7 presents the W / F dependence of the conversion for n-butane oxidation at 703 K , where W is the
weight of the catalyst (g) and F is the flow rate of
n-butane (mol/h). Among these catalysts, VPDBi3
with 81% conversion was found to be the most active, followed by VPDBil (71%) and VPD (47%).
The great difference in catalytic activity for the
Bi-promoted catalysts compared to the unpromoted
tVPD
t-VPDBil
--e VPDBi3
501
'1
'
0
'
'
I
10
'
'
'
'
'
'
'
'
'
I
20
30
Reaction time (h)
'
'
'
'
I
40
'
'
'
'
50
Figure 6. Time course changes of MA selectivity in
oxidation of n-butane over VPD, VPDBil
and VPDBi3 catalysts
100
tVPD
tVPDBi I
401,
I
0
J
'
I
10
'
'
I
5
1
20
30
Reaction time ( h )
'
I
j
5
I
40
"
"
I
50
Figure 5. Time course changes of conversion in oxidation of n-butane over VPD, VPDBil and
VPDBi3 catalysts
material is attributed to the reactivity of the oxygen
atoms in the lattice. As reported earlier, selective oxidation of n-butane over V-P-0 catalyst was shown
to proceed via a redox mechanism [lo]. Therefore
the reduction behaviour of this catalyst is a reliable
method as an indicator for the catalytic performance.
It should be noted that two major reduction peaks
of VPDBi3 and VPDBil (from T P R profiles in Figure 4) occurred significantly at the lower temperature
region compared t o VPD (Table 3). These two reduction peaks ( a and p) corresponding to the removal
of oxygen species linked t o V5+ and V4+ phases, re-
Y.H. Tauf iq-Yap et al./ Journal o f Natural Gas Chemistry Vol. 15 No. 4 2006
302
spectively, played a major role in the performance of
the catalysts, as shown in the increment of the conversion. A significant improvement in the oxygen's
reactivity by lowering the reduction activation energy
with the addition of Bi leads to the enhanced catalytic
performance for n- but ane oxidat ion. Further more,
the combination of these two types of oxygen species
also favoured the hydrocarbon's activation. Figure 8
gives the change in selectivity t o MA as a function
of the conversion of n-butane and demonstrates that
VPDBi3 is the most selective catalyst.
4. Conclusions
The doping of Bi into the lattice of a vanadyl pyrophosphate catalyst increased the total surface area
of VPDBil and VPDBi3 from 20.3 m2/g (unpromoted
VPD) t o 25.0 and 24.8 m2/g, respectively. This increment was effected by an increase in the rosette-type
platelets subtending the (100) face. The presence of
Bi also reduced the overall oxidation state of the vanadium. The addition of 1% and 3% Bi significantly
enhances the activities of the catalysts. The effect is
attributed to the highly active and labile lattice oxygen. It is also in part due to a structural effect, as the
surface area of the Bi-promoted catalysts is increased
by 25%.
I
loo
-s-
80 -
'g
60-
Acknowledgements
Financial assistance from Malaysian Ministry of Science, Technology and Innovation is gratefully acknowledged.
C
G
U
-
40
tVPD
tVPDBil
- 4
-A-
VPDBi3
References
Centi G. Catalysis Today, 1993, 16: 5
Wang D X, Kung M C, Kung H H. Catalysis Letters,
2000, 65: 9
80
1
I
h
55
1
20
-.- VPDBil
-VPDBi3
Hutchings G J. Applied Catalysis A: General, 1991,
72: 1
Brutovsk? M,Kladekov6 D, Kosturiak A. Chemical
Listy, 1995, 89: 682
Hutchings G J, Higgins R. Journal of Catalysis, 1996,
162: 153
Niwa M, Murakami Y. J Catal, 1982, 76: 9
Taufiq-Yap Y H, Waugh K C, Hussein M Z. Oriental
Journal of Chemistry, 1998, 14(1): 1
Kesteman E, Merzouki M, Taouk B, Bordes E, Contractor R. In: Poncelet G et a1 eds. Preparation of
Catalysts VI. Amsterdam: Elsevier Science B V, 1995.
707
40
60
Conversion (YO)
80
I00
Figure 8. Changes in selectivity t o maleic anhydride
as a function of conversion of n-butane
Pierini B T, Lombard0 E A. Material Chemistry and
Physics, 2005, 92: 197
Centi G,Trifiro F, Ebner J R, Franchetti V M. Chemical Review, 1988, 88: 251
ScienceDirect
Journal of Natural Gas Chcmirtry
SCIENCE PRESS
www.elsevier.mmilmte/jngc
Article
Molybdosphoric Acid Mixed with Titania Used as a Catalyst to
Synthesize Diphenyl Carbonate via Transesterif icat ion of
Dimethyl Carbonate and Phenol
Tong Chen'>2? Huajun Han' , Zhiping Du' , Jie Yao' , Gongying Wangl*?
Dachuan Shi3, Desheng Zhang3, Zhiming Chen3
1. Chengdu Institute of Organic Chemistry, Chinese Academy of Sciences, Chengdu 610064, Sichuan, China;
2. The College of Guiyang Traditional Chinese Medicine, Guiyang 550001, Guizhov, China;
3. PetroChina Company Limited, Jilin Branch, Jilin 13,2021, Jilin, China
[Manuscript received April 28, 2006; revised June 13, 2006)
Abstract:
The 12-molybdosphoric acid mixed with titania (MPA-Ti02) was found t o be a novel
and efficient catalyst for the synthesis of diphenyl carbonate (DPC) via transesterification of dimethyl
carbonate (DMC) and phenol. The X-ray diffraction (XRD) and infrared (IR) techniques were employed
to characterize the prepared catalysts. The effect of the weight ratio of the 12-molybdosphoric acid to
titania on the transesterification was investigated. A 13.1% yield of DPC and an 11.6% yield of methyl
phenyl carbonate (MPC) were obtained over MPA-Ti02 with the weight ratio of MPA to Ti02 as 5:l.
Key words: transesterification; 12-molybdosphate acid; titania; methyl phenyl carbonate; diphenyl
carbonate
1. Introduction
Polycarbonates (PCs) are important engineering thermoplastics, with good mechanical and optical properties as well as electrical resistance that
are useful for many applications [I]. PCs are commercially produced by the reaction of phosgene and
bisphenol-A. The use of the conventional process
results in few serious environmental problems such
as the use of the highly toxic phosgene, the formation of a stoichiometric amount of NaCl or HC1, and
the use of copious amounts of methylene chloride
as the solvent. Melt polymerization of bisphenol-A
and diphenyl carbonate (DPC) is the most practical
nonphosgene process for manufacturing PCs; thus,
DPC is the key material for the green production
of PCs. Several available nonphosgene processes for
the manufacture of DPC have been proposed, and
the major ones include carbonylation of phenol and
transesterification of dimethyl carbonate (DMC) and
phenol [2]. Oxidative carbonylation of phenol is a
prospective route for the synthesis of DPC, but the
use of noble-metal catalysts and the low yield of DPC
limit its industrialization. At present, the transesterification of DMC and phenol is thought to be the
most suitable method for the industrial production of
DPC. This route is a two-step process, which involves
the transesterification of DMC and phenol to methyl
phenyl carbonate (MPC) (Equation (1))and the further transesterification of MPC and phenol (Equation (2)) or disproportion of MPC to DPC (Equation
(3)). The reaction suffers from low yield and selectivity even a t an elevated temperature because of a
critical thermodynamic limitation in the formation of
MPC (3x lo-* a t 180 "C) and because of the low reaction rate [3]. Therefore, active catalysts are essential for the transesterification. Generally, the transesterification of DMC with phenol is carried out in the
* Corresponding author. Tel: 028-85215405; E-mail: gywangQcioc.ac.cn.
304
Tong Chen et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
liquid phase using homogeneous catalysts such as organic Sn, Ti, Al, and Fe compounds [4&7]. However,
homogeneous catalysts are toxic (for example, organic
Sn), unstable and are difficult t o be separated from
the final products. Therefore, considerable attention
has been paid to the development of active solid catalysts, which can simplify the purification process to a
large extent. Unfortunately, the reports on solid catalysts are very few. In previous literatures, the solid
catalysts are mainly Mo, Ti, Si, Pb, and rare-earth
metal oxides [8-121.
Heteropoly compounds (HPCs), as environmentally benign catalysts, are widely utilized, specifically
for materials with Keggin structure, in homogeneous and heterogeneous catalytic reactions because
of their unique physicochemical and catalytic properties. Catalysis using heteropoly acids (HPAs) and
the related compounds is a growing field with increasing significance, since HPA-based catalysts possess bifunction with strong acid and redox abilities and have
higher activity than the known traditional catalysts
[13-181.
This study explores the transesterification reaction of dimethyl carbonate with phenol [4,5].
Herein, 12-molybdosphoric acid (H3PMo12040,
abbreviated as MPA) mixed with titania (MPATiO2) as a novel solid catalyst for the synthesis of
DPC from dimethyl carbonate with phenol via transesterification has been reported.
DMC
+
QO
O
!Q
CH30H
DPC
*QOKOCH,
e O ! O Q
=
MPC
2. Experimental
2.1. Chemical reagents
Dimethyl carbonate (Huasheng Co. Ltd, China)
was fractionally distilled and stored over a molecular
sieve (4A). The phenol and the 12-molybdosphoric
acid were of analytical reagent (AR) grades and were
used without further purification.
The catalysts were prepared as follows: the Ti02
powders were added to the aqueous solutions of MPA
with stirring for 0.5 h. The mixture was dried at
100 "C and was then calcined for 3 h at 300 "C for activation. The catalyst was denoted as MPA-Ti02. The
Ti02 was prepared by hydrolysis of Tic14 followed by
dehydration of the Ti(OH)4 xerogel at 100 "C (named
as TiO2-100).
2.3. Characterization of the catalyst
Powder X-ray diffraction patterns of the samples
DPC
0
+ CH30COCH3
II
DMC
were recorded on a DX-1000 diffraction instrument
using Cu K , radiation a t a wavelength of 0.154 nm.
A scan rate of O.O3O/S with a step size of 0.03' was
used for data collection. Framework vibration spectra in the range of 2000-400 cm-' were recorded on
an FT-IR (NICOLET MX-1E) spectrometer in KBr
disks at room temperature.
2.4. Reaction procedure
The liquid-phase reaction was carried out in a
100 ml three-necked, round-bottomed flask equipped
with a nitrogen inlet, a drop-wise filler, and a fractionating column connected to a liquid dividing head.
The phenol and the catalyst were added under a nitrogen atmosphere with stirring and slow heating. After
the mixture was heated to 178 "C, DMC was added
in drops. The reaction temperature was maintained
between 150 "C and 180 "C with refluxing under atmospheric pressure. At the end of reaction, the mixture was cooled, and the catalyst was filtered. The
filtrates and the azeotrope of methanol with DMC
during the reaction were quantitatively analyzed us-
305
ing gas chromatography (Shimadzu GC14-B) with a n
FID detector and a capillary column (30 rn), and
the product mixtures were characterized by GC-MS
(HPGC/MC 6890/5973) for the confirmation of DPC
and MPC.
3.2. X-ray diffraction
The diffractograms of MPA-Ti02 in various proportions that were calcined at 300 "C are shown in
Figure 2 . For comparison, the Ti02-100 was calcined
a t 300 "C (named Ti02-300), and the diffractograms
of Ti&-300 are also given in Figure 2 .
3.1. Infrared spectra
The FT-IR spectra of the samples are shown in
Figure 1. All the MPA-Ti02 samples exhibited the
Keggin structure at 1063, 961-964.9,866.7-862.3, and
785.5-801 cm-', which can be assigned to the stretching modes vaS(P-Od), vas(Mo-Ot), va,(MO-ObMo), and v, (Mo-0,-Mo),
respectively, where ob
represents corner-sharing oxygen and Oc represents
edge-shar ing oxygen. Meanwhile, the characteristic
peaks of a-Mo03 could not be identified on the MPATi02 samples. It is indicated that the MPA-Ti02
samples calcined at 300 "C preserve the Keggin structure. Comparing the MPA, the stretching vibrations
of MPA-Ti02 samples produce a small shift owing to
the interactions of MPA and TiOz. The stretching
frequency of the characteristic IR bands of the samples are shown in Table 1.
1800
1600
1200
1400
800
1000
600
400
Waven titn ber ( c & )
Figure 1. IR spectra of the samples with various
weight ratios of MPA to TiO2: (1) 1:1, (2)
3:1, (3) 5:1, (4)1O:l
Table 1. Stretching frequency of the characteristic
IR bands of the samples
Stretching frequency (cm-l )
MPA-~O~
(weight ratio)
MPA
10:1
5:l
3: 1
1:l
P-0
M-0,
1063.9
Mo=Ot
962.4
1063.5
960.7
1063.0
961.3
Mo-ob
867.5
866.7
866.4
1062.6
963.4
865.7
795.2
964.9
862.3
801.0
1063.0
789.4
785.5
791.6
.-C
B
4-
10
20
30
40
50
60
28/(' )
Figure 2. XRD patterns of the samples with various
weight ratios MPA to TiO2: (1) 1:1, (2) 3:1,
(3) 5:1, (4)1O:l
Anatase Ti02 was detected on Ti02-300, and
amorphous Ti02 was observed on Ti02-100. The
MPA-Ti02 samples show only the characteristic fraction peaks of the Keggin structure of the heteropoly
compound, and no characteristic fraction peaks of the
anatase Ti02 or @-Moos could be identified. This
result corresponds t o that of IR. It is indicated that
the Keggin structure keeps intact in the MPA-Ti02
samples calcined at 300 "C, and the Ti02 is microcrystalline or highly dispersed; that is, the MPA prevents
the production of the Ti02 crystal phase even when
the sample was calcined at 300 "C. The characteristic
peaks of the Keggin structure weaken with the decrease in the amount of MPA in the catalyst, and the
partial diffraction peaks disappear, which implies the
mild, dispersed effect of Ti02 on MPA.
3.3. Catalytic behavior
The catalytic behavior of MPA-Ti02 with various
weight ratios of MPA to Ti02 are investigated and
compared with that of bulk MPA and Ti02 in the
transesterification, and the catalytic-activity results
are summarized in Table 2.
306
Tong Chen et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
From Table 2, it can be observed that the only
by-product is anisole with the exception of the two
main products MPC and DPC, and that the transesterification selectivity exceeds 90% over all catalysts. Furthermore, the data in Table 2 clearly
show that bulk PMA has low conversion of 12.2%
for phenol, and the pure Ti02 is poorly active with
only 0.8% conversion of phenol. The conversion of
phenol over MPA-Ti02 is higher than that of the
bulk MPA. This illustrates that Ti02 acts as a promoter in the catalyst. The yield of transesterification
(MPC+DPC) increases slightly with the increase in
the ratio of MPA to Ti02 up to 5:1, and then it
slightly decreases. When the optimal weight ratio of
MPA to Ti02 is 5:1, the yield of MPC and DPC of
24.7% and the transesterification selectivity of 99%
T a b l e 2. T h e c a t a l y t i c activities of a series of M P A - T i 0 2 f o r the transesterificationa
Yields (%)=
MPA:Ti02
Amount of
(weight ratio)
MPAb(mol%)
(%I
MPC
DPC
Bulk MPA
0.20
12.2
1:1
0.10
0.15
0.17
0.19
0.00
23.3
23.9
25.0
25.3
0.8
6.2
13.3
12.3
11.6
12.3
0.8
5.5
9.6
10.8
13.1
11.3
3: 1
5: 1
1O:l
Ti02 (anatase)
Phenol Conv.
-
Transesterif ication
AN
0.5
0.4
0.8
0.3
1.7
-
selectivity (%)
96
98
97
99
93
100
a. Reaction conditions: phenol (160 mmol), DMC (160 mmol), catalyst (1.2 g), reaction time (8 h), and reaction
temperature (150-180 "C);
b. The molar ratio of MPA t o the total reactants;
c. MPC=methyl phenyl carbonate; DPC=diphenyl carbonate; AN=anisole.
are attained. On combining this with the results of
XRD,this total yield continuously increases with the
increase in the amount of MPA, which may be ascribed to the impact of accumulation of MPA. After
the ratio becomes 5:1, the catalytic activity does not
increase, which indicates that the functions of Ti02
begin to weaken because of being enveloped in excess
MPA.
4. Conclusions
(1) MPA-Ti02 is a n environmentally benign and
efficient heterogeneous catalyst that is used for the
synthesis of diphenyl carbonate (DPC) via transesterification of dimethyl carbonate (DMC) and phenol. The yield (MPC and DPC) and selectivity of the
transesterification over MPA-Ti02, with the weight
ratio of MPA to Ti02 as 5:1, was 24.7% and 99%,
respectively.
(2) MPA is the main active component in the
MPA-Ti02 catalyst. Ti02 is a promoter, which is
dispersed by the MPA, and no characteristic fraction
peaks of anatase Ti02 are seen.
(3) The MPA-Ti02 that is calcined at 300 "C preserves the Keggin structure.
Acknowledgements
The work was supported by the National Technology
Research and Development Program of China (863 plan)
(2003AA321010) and China Petroleum & Natural Gas Co.
Ltd.
References
Ono Y. Appl Catal A , 1997, 155: 133
Mei F,Li G, Nie J, Xu H. J Mol Catal A , 2002, 184:
465
Shaikh G, Sivaram S. Ind Eng Chem Revs, 1992, 31:
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Du Z, Kang W, Chen T. J Mol Catal A , 2006, 246:
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Niu H Y , Yao J, Wang Y, Wang G Y. J Mol Catal A ,
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Mei F,Li G, Nie J, Xu H. J Mol Catal A , 2002, 184:
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Lee H, Bae J Y , Kwon 0-S. J Organomet Chem, 2004,
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Kim W B, Lee J S. J Catal, 1999, 185: 307
Mei F,Pei Z, Li G. Process Res Dev, 2004, 8: 372
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Zhou W, Zhao X, Wang Y , Zhang J. Appl Catal A,
2004, 260: 19
Kozhevnikov I V. Chem Rev, 1998, 98: 171
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Izumi Y,Hisano K, Hida T. Appl Catal A , 1999, 181:
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Kishore G D, Baskaran S. J Org Chem, 2005, 70: 4520
Li X-K, Zhao J, Zhang Z-B. J Catal, 2006, 237: 58
Abd M M M, Said A A. J Mol Catal A , 2005, 109
ScienceDirect
Journal of Natural GdS Chermstry
SCIENCE PRESS
wmv.elsevier.dmk/jngc
Article
Kinetic Rates of the Fischer Tropsch Synthesis on
a Co/NbaOs Catalyst
Victor R. Ah&,
Paulo L. C. Lage, Carlos D. D. de Souza,
Fabiana M. Mendes, Martin Schmal*
Federal Unaversaty of Rao de Janearo Programa d e Engenharaa Quimaca, COPPE/UFRJ C.P. 68502,
21941-972, Rao de Janezro, R J , Brazal
[Manuscript received May 15, 2006; revised June 27, 20061
Abstract:
The kinetics of the Fischer-Tropsch reaction over a Co/Nbz05 catalyst in a fixed bed
reactor was investigated experimentally. Experiments were carried out under isothermal and isobaric
conditions (T=543 K, P=2.1 MPa) and under different conditions of several Hz/CO feed molar ratio
(0.49-4.79), space velocities (0.2-3.8 h-'), mass of catalyst (0.3-1.5 g), and CO conversion (10%-29%).
Synthesis gas conversion was measured and data were reduced to estimate the kinetic parameters for
different Langmuir-Hinshelwood rate expressions. Differential and integral reactor models were used for
the nonlinear regression of kinetics parameters. One of the rate equations could well explain the data. The
hydrocarbon product distributions that were experimentally determined exhibited an unusual behavior,
and a possible explanation was discussed.
Key words: Fischer-Tropsch; natural gas; kinetics; selectivity
1. Introduction
The transformation of syngas (CO+Hz), through
the Gas-to-Liquids (GTL) technology, into a wide
spectrum of linear and branched hydrocarbons has
become an excellent alternative for the use of natural
gas t o obtain liquid transportation fuels. Besides, the
possibility of obtaining clean-burning fuels, satisfying
stricter environmental regulations, has introduced a
renewed interest in the study of the Fischer-Tropsch
(FT) synthesis.
A brief literature review on the FT synthesis
shows that the kinetics of the FT reaction has been
the main focus of many researchers, and several kinetic models from different mechanisms have been
proposed. However, only a few kinetic studies on
cobalt-based catalysts can be found, such as the studies of Sarup and Wojciechowski [l],Keyser et al. [2],
and Zennaro et al. [3]. Generally, kinetics for cobalt-
based catalysts present different expressions when
compared t o those for the iron-based catalysts, and
kinetic equations are based on the rate-determining
step, which involves a dual-site surface reaction, resulting in a squared expression in the denominator in
the rate equation.
Several studies have tried t o explain the deviations of FT hydrocarbon selectivity from the
Anderson-Schulz-Flory distribution model (ASF),
usually due t o the a-olefin readsorption with secondary chain propagation (Iglesia et al. [4], van
der Laan and Beenackers [5], Schulz and Claeys [6])
and the existence of two chain propagation mechanisms or sites (Madon and Taylor [7], Sarup and
Wojciechowski [l], Donnelly et al.
[8], Patzlaff
et al. [+lo]). Although a-olefin readsorption with
secondary chain propagation is regarded as the correct explanation of non-ASF product distributions for
ruthenium catalysts (Iglesia et al. [4],
Patzlaff et al.
* Corresponding author. Tel: +55 21 2562 8352; Fax: +55 21 2562 8300; E-mail: [email protected]
308
Victor R Ah6n et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
[9]),the existence of two different mechanisms or active sites for chain propagation seems to be the major
effect for iron and cobalt catalysts [9-lo]. Recently, a
selectivity model that combines both the effects has
been proposed [ 111.
The main objective of this study is to study the kinetics of the CO hydrogenation of the Co/NbnOs catalyst, which presents high selectivity toward C13-C17
paraffins and to explain the unusual behavior of hydrocarbon product distributions through the ASF distribution model, using experimental data and kinetic
rate expressions.
2. Experimental
The Nbz05 support was obtained by calcination
of niobic acid (CBMM -Companhia Brasileira de Metalurgia e Mineraqao) in air at 823 K at a heating rate
2 K/min for 3 h, as a result of which it is transformed
from an amorphous phase to crystalline TT or T form
of niobium pentoxide. The catalysts were prepared by
incipient wetness impregnation of the support with an
aqueous solution of cobalt nitrate, containing 5 wt%
Co. After impregnation, the samples were dried at
393 K for 16 h and calcined in air at 673 K for 2 h.
2.2. Reactor
The reactor was a vertical stainless steel tube with
approximately 1.5 cm ID, wall thickness of 2 mm, and
length of 15 cm. A stainless steel screen was placed
5 cm from the bottom of the reactor t o hold the catalyst sample. The gas flows down through the catalyst.
A thermocouple type J was adapted for monitoring
the reaction temperature, inserted into a 3.18 mm
diameter tube at 2 cm above the screen. The thermocouple tip reached the center of the reactor. A ceramic oven was used with a thermocouple connected
to a temperature programmer and controller.
The reaction conditions and procedures were similar to those described elsewhere (Mendes et al. [12],
Frydman et al. [13], Soares et al. [14]). The CO hydrogenation was performed in a fixed bed at 543 K,
containing a mass of 0.3-1.5 g of catalyst at 2.1 MPa.
The total feed flow rate was chosen to reach isoconversion, using a mixture of He, H2/CO feed ratios 0.49-4.79. Helium was used as internal standard
to calculate the total CO conversions. Before the
reaction, the samples were reduced under hydrogen
flow a t 773 K for 16 h. The reaction products were
analyzed by on-line gas chromatography (ShimadzuGC-l7A), equipped with a CP-PoraBOND Q, 50 m
(TCD) and a CP-SIL, 50 m (FID) columns. The products were analyzed using a flame ionization detector
(FID).
Typical experimental results for hydrocarbon distribution are as presented in Table 1 after running for
48 h with time on stream, at 2.0 MPa and temperature of 543 K, varying the space velocity.
Table 1. Experimental results for different reaction conditions
H2/CO (ratio)
X/%
WHSV (h-l)
CZ-c4 (%)
Diesel (%)
Ci9+ (%)
0.49
12.91
400
26.27
24.45
23.78
2.03
16.26
8000
3.08
2.39
24.95
2.83
24.01
4000
17.97
13.83
20.06
42.49
5.64
4.79
23.73
4000
46.54
33.56
11.98
1.99
5.93
2.4. Determination of the kinetic parameters
The experimental results for cobalt catalyst
showed that n-alkanes and 1-alkenes are the predominant products of the reaction. Therefore, the reaction
stoichiometry may be approximated as
%co
where
Product distribution
G a ~ o l .(%)
CH4 (%)
-
-
-I-( N c -tN H / ~ ) H-+z CGHK
% is the
+K H z O
(1)
average carbon chain length of the
.
7.67
17.83
48.88
20.71
hydrocarbon product and
is the average number
of hydrogen atoms per hydrocarbon molecule.
The reaction rate is defined, in this study, as the
number of moles of carbon monoxide converted per
time per mass of catalyst (-Rco). In the literature,
several expressions for the reaction rate of the FischerTropsch synthesis over cobalt-based catalyst were reported (Sarup and Wojciechowski [l],Yates and Satterfield [15],Kuipers et al. [16]). It was found that a
Langmuir-Hinshelwood equation, which involves a bi-
molecular surface reaction, could adequately describe
the Fischer-Tropsch reaction. These rate expressions
can be generalized as
where k is a kinetic rate constant, a and b are the reaction orders of the rate-determining step; Ki is the
adsorption constant for the i-th adsorption term, and
ci and di represent the dependency of surface coverage
on the reactant pressure of the i-th adsorption term,
P H and
~ PCO are the partial pressures of hydrogen
and carbon monoxide, respectively.
Based on the results obtained by Sarup and Wojciechowski [l] and Yates and Satterfield [15] and
due t o the differences in their functional dependence
on P H and
~ PCO,the following rate equations were
chosen for interpreting the experimental data in the
present analysis.
309
For a differential reactor, it is assumed that the
conversion is sufficiently less and does not affect the
reactant compositions along the reactor. Therefore,
the feed compositions can be used t o calculate the reaction rates along the whole reactor, giving (-Rco)O,
which is constant. Then, Equation (8) becomes
or
For the integral reactor, Equation (8) must be integrated. Because the compositions were measured
only a t the inlet and outlet reactor, the best approximation of the integral term in equation (8) is obtained
using the trapezoidal quadrature:
+ (-Rco)'
l
l
(4)
(11)
Then, Equation (8) becomes
(5)
For the determination of the unknown kinetics parameters for each rate expression, a nonlinear regression procedure was used. Due to the relatively high
conversion obtained in the several experiments, these
parameters were calculated considering the fixed bed
separately as differential and as an integral reactor.
The material balance of CO in the fixed bed tubular reactor is
- ~ F c o = (-RCo)PcatdV
(6)
where FCO is the CO molar flow rate (mol.min-l),
(-Rco) is the CO rate of consumption
(mol.g,at,-'.min-l),
peat is the catalyst density
(gcat,+cm-3),and V is the reactor volume (cm3).
Moreover,
Fco = Poqoxco/Q
(7)
where PO is the feed density (g.cmP3), qo is the volumetric feed rate (cm3.min-l), xco is the CO molar
fraction, and % is the average molar mass (g.mol-l).
After substitution of equation (7) into equation (6)
and integration along the reactor length, it becomes
--
(8)
Experimental results are presented in Table 2 for
different reaction conditions. First, all these 12 experimental data were used for kinetic studies. Table 3 shows the results of the nonlinear regression for
the rate expressions described by Equations (3)-(5),
considering the reactor as differential (Equation 10)
and integral (Equation 12). The significance of each
parameter was assessed by using the t-test (Froment
and Bischoff [17]), which is the ratio of the regressed
parameter to its standard deviation. The statistical
significance of the global regression was expressed by
means of the F-ratio, based on the ratio of the mean
calculated sum of squares and the mean regression
sum of squares. It can be seen that parameters k
and K and are not statistically adequate and must be
rejected.
310
Victor R Ahdn et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
Table 2. Selected experimental parameters for the reaction
Experiments
1
mcat/g
1.50
Fgo/(mol. min-')
2 . 4 7 lop4
~
Pi /MPa
P&/MPa
1.267
0.621
P&/MPa
0.212
(H2/CO)'
0.49
Xco (%)*
2
3
4
5
10.07
10.98
12.91
0.30
1.38~10-~
0.959
0.472
0.669
2.03
25.56
28.52
0.30
5.43~10-~
1.438
0.507
0.154
2.83
20.35
21.65
0.30
1.25~10-~
0.385
0.080
1.634
4.79
6
27.50
7
8
9
10
11
12
22.31
25.16
23.82
23.99
*CO conversion: Xco = (Fgo - Fco)/Fgo,FCO (mol.min-') volumetric flow rate at the outlet
Table 3. Regression results for the kinetics models using all data sets
Regression
results
ka
K
ck
Differential reactor
Equation (3)
2.70~10-~
0.62
5.57~
Equation (4)
1.84~10-~
0.36
2 . 5 6 10W2
~
Equation (5)
Equation (3)
8 . 4 0 lop3
~
0.50
1.9ox10-2
2.77~
0.61
5.57x 10-2
Integral reactor
Equation (4)
1.98~
0.37
2.75~10-~
Equation (5)
9 . 4 0 lop3
~
0.53
2.15~10-~
*K
1.12
0.40
0.82
tk
0.49
0.72
0.44
1.10
0.50
0.42
0.72
0.88
0.44
tK
F-ratio
0.55
5.31
0.90
6.78
0.61
5.42
0.55
5.39
0.60
5.42
R2
0.19
0.29
0.21
0.19
0.88
6.88
0.29
*
0.19
aEquation 3 (kmol/kgcat-h-bar), Equation 4 ( k m ~ l / k g c a t - h - b a r ~ . ~
Equation
),
5 (kmol/kgcat-h-bar2);
bCritical t value for 95% confidence that the parameters are significant is t0.975,10=2.228
The mass balance deviations for different ratios were calculated from the experimental data, and
they are listed in Table 4. Experimental results for
(Hz/C0)'=2.03 presented larger deviations in the
mass balance than in the others, which is thus improper for the kinetic interpretation.
Table 4. Mass balance deviation (%'ow) for different
reaction conditions after 70 h with time on stream
Experiments
(Hz/CO)O
Mass balance deviation
after 72 h on stream
2
6
8
12
0.49
2.03
2.83
4.79
-0.80
-21.27
0.15
0.42
Therefore, the kinetic parameters were determined again, but excluding the experiment for
(Hz/C0)'=2.03. These regression results are shown
in Table 5. In this case, the adjusted parameters obtained by nonlinear regression were statistically significant for rate expression given by Equation ( 5 ) .
The parameters of Equations (3) and (4) are still inad-
equate because their calculated t-values do not exceed
the corresponding tabulated value for confidence at
the 95% probability level. From the t-values, it is clear
that the experimental data are better described by the
rate expression given by Equation (5). Besides, values of the variance-covariance matrix are 2.03 x
and 2 . 0 3 lop2
~
indicating that parame2 . 1 4 lop4,
~
ters are not correlated.
The regression results presented in Tables 3 and 5
were calculated using both the differential reactor and
the integral reactor approximated solutions, Equations 10 and 12, respectively. It can be seen that
both reaction models lead to the same results within
the experimental statistical uncertainty. This implies
that the differential reactor approximation is valid for
the present data.
The hydrocarbon product distributions obtained
for different feed ratios were also calculated from gas
chromatographic (FID and TCD) analyses. They are
shown in Figure 1 for the different H2/CO ratios analyzed.
311
Table 5. Regression results for the kinetics models using a reduced data set
Regression
results
ka
K
uk
UK
t k -valueb
tK-valueb
F-ratio
R2
Differential reactor
Equation (4)
Equation (5)
7.85~10-~
3 . 7 0 lop3
~
0.38
0.52
4 . 1 2 ~ 1 0 ~ ~ 1.47~10-~
0.16
0.15
1.91
3.10
2.38
3.47
31.67
134.09
0.81
0.95
Equation (3)
2.70~
lo-'
0.62
5.58~
1.12
0.48
0.55
14.78
0.63
Equation (3)
1.19x10-2
0.77
1.30~10-~
0.69
0.92
1.12
14.51
0.60
Integral reactor
Equation (4)
7.471
~ 0-~
0.41
3 . 8 8 10-3
~
0.16
1.93
2.56
29.87
0.79
Equation (5)
3.83x lop3
0.59
1.43~10-~
0.15
2.68
3.93
130.15
0.95
aEquation 3 (kmol/kgcat-h-bar), Equation 4 ( k m o l / k g ~ a t - h - b a r ~ Equation
,~),
5 (kmol/kgcat-h-bar2);
bCritical t value for 95% confidence that the parameters are significant is t0.975,7=2.365
-0.49
0.1
u'
-
M
0.01
.\.
IE-3I'"""
0
5
'
'
~
10
'
*
'
"
'
15
"
'
1
"
"
20
1
25
Carbon number
Figure 1. Hydrocarbon product distributions for
(Hz/CO)O at 0.49, 2.03, 2.83, and 4.79
As expected, the total hydrocarbon molar compositions cannot be interpreted as an ASF distribution. Although the c 2 - C ~anomalies and the change
of the chain growth probability in the range from C3
to C8-C12 range may be explained by a-olefin readsorption with secondary chain propagation and the
existence of two mechanisms for chain propagation,
the increasing selectivity observed between C8 and
C12 for all experiments cannot be explained by these
mechanisms, except if a carbon number-dependent
product desorption rate constant is admitted.
Kuipers et al. [16] observed similar results with
Co/SiOz catalyst and suggested that for a non-ASF
behavior up to C13 is independent of the flow, whereas
Iglesia et al. [4] showed that the extension of a-olefin
reinsertion is dependent on the chain length and
influences the total product distribution, which is
also partly dependent on the contact time, due to
transport limitations. Although transport limitation occurs due t o porous distribution, it seems that
the umbilical cord mechanism suggested by Kuipers
et al. [16]would explain the behavior of the proposed
system. Therefore, reinsertion into the chain growth
occurs at a growth site, where the a-olefin reinsertion is physisorbed, can adsorb prior to desorption,
and therefore will be independent of the flow. As
observed, paraffin production prevails, which means
that, for high carbon number, all a-olefins are reinserted and do not desorb until they reach a critical
length and finally terminate as paraffins, which desorb to the vapor or liquid phase. On the other hand,
the drop in growth probability for NC around 14, observed in all experiments, can be attributed also to
product retention and condensation, which according
t o the literature, is not feasible due to the sigmoid
curve, which is more likely attributed to a hydrogenolysis process. In fact, the C15-C17 region can be hydrogenolyzed to lower carbons such as C11. Moreover,
product retention and condensation are not expected
to have occurred because the overall mass balance indicates that the reactor operates steady state, with
the exception of the experiments at Hz/CO=2.03.
As reported previously [12,13], the Co/Nbz05
catalyst presents different surface properties when reduced at 557 K and 773 K. In the later case, the
Nb205 is also reduced and affects the surface where,
besides NbO, patches, Co'f and Coo coexist. The
metallic cobalt prevails and is necessary for the initiation of the reaction; however, the presence of Co2+
species favors the reinsertion a-olefins due t o the electronic structure. The growth sites necessary for the
reinsertion of a-olefins and physisorption is sustained
by interfacial sites, probably a t NbO,/CoZf species
[12,13],which allows the co-adsorption of long-chain
molecules and probably favors the reaction to the formation of long chains at the surface. In addition, results have shown narrow selectivity between C13 and
(217, where the non-ASF behavior is observed, indicating that longer paraffins are easily desorbed, suggesting that bonding strengths decrease at the interfacial sites. It also disfavors the hydrogenation and
312
Victor R Ah6n et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
suggests again that it is independent of the flow rate
or condensation.
Particularly, the ratio Hz/CO=2.03 displayed
high selectivity for the C I ~ - C Ihydrocarbon
~
formation. All experiments were obtained at isoconversion around 10%-15% of CO conversion. The deviation from the ASF behavior and mass balance deviation, which is gradually reduced with time on stream,
is not clear. There are some possibilities concerning these experiments: condensation or retention of
products in the tubes and carbon deposition on the
catalyst, resulting in carbon mass balance deviation.
The first hypothesis can be disregarded because lines
were heated above condensation temperature of heavier products. It is noteworthy that a t 70 h the catalyst
was stable and deviation of mass balance was around
20%. In addition, during stabilization, samples after 48 h were taken, and coke was analyzed by TG,
as shown in Figure 2. It displays the loss of approximately 16% of carbon due to the coke deposition,
which allows us to explain the deviation observed in
the mass balance for the ratio Hz/CO=2.03.
-201
0
'
'
'
'
'
200
'
'
'
'
'
'
'
'
400
Temperature ("C)
'
'
600
'
'
'
'
I
800
Figure 2. Thermoanalysis after running for 48 h
with time on stream
4. Conclusions
Experimental rate data of gas conversion using
Fischer-Tropsch synthesis over a Fischer-Tropsch catalyst in a fixed bed reactor were obtained for different
conditions of Hz/CO feed ratio, space velocities, mass
of catalyst and CO conversion. A kinetics study
with three different two-parameter rate LangmuirHinshelwood expressions, was performed. The parameters were estimated by nonlinear regression considering both a differential and an integral reactor. One
of the equations showed the best fit for the analyzed
data, and its parameters were shown to be statistically significant.
The hydrocarbon product distributions of
Co/Nbz05 catalyst exhibit a n unusual non-ASF behavior, which suggests that NbO,/Cobalt interfaces
play an important role for chain growth of molecules.
Mass balance deviation occurred depending on
the feed ratio. The highest deviation in the ASF distribution equation was observed for a H2/CO ratio of
2.03, which apparently presented the best hydrocarbon distribution in the diesel fraction C13-C18. This
deviation can be attributed to the carbon deposition
on the catalyst.
Acknowledgements
This work was sponsored by PETROBRAS and
FINEP. The authors would also like to thank CNPq for
their financial support.
References
[I] Sarup B, Wojciechowski B W. Canadian Journal of
Chemical Engineering, 1988, 66: 831
[2] Keyser M J, Everson R C, Espinoza R L. Industrial
and Engineering Chemistry Research, 2000, 39: 48
[3] Zennaro R, Tagliabue M, Bartholomew C H. Catalysis
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[4] Iglesia E, Reyes S, Madon R J. Journal of Catalysis,
1991, 129: 238
[5] van der Laan G P, Beenackers A A C M. Catalysis
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[6] Schulz H, Claeys M. Applied Catalysis A : General,
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[7] Madon R J, Taylor W F. Journal of Catalysis, 1981,
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[8] Donnelly T J, Yates I C, Satterfield C N. Energy &
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[9] Patzlaff J, Liu Y, Graffmann C, Gaube J. Applied
Catalysis A : General, 1999, 186: 109
[lo] Patzlaff J , Liu Y, Graffmann C, Gaube J. Catalysis
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[ll]Ah6n V R, Costa Jr E F, Monteagudo J E P, Fontes C
E, Biscaia J r E C, Lage P L C. Engineering Science,
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[12] Mendes F T, Noronha F B, Schmal M. Studies in Surface Science and Catalysis, 2000, 130: 3717
[13] Frydman, A, Soares R R, Schmal M. Studies in Surface Science and Catalysis, 1993, 75: 2797
[14] Soares R R, Frydman A, Schmal M. Catalysis Today,
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[15] Yates I C, Satterfield C N. Energy & Fuels, 1991, 5:
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[16] Kuipers E W, Scheper C, Wilson J H, Vinkenburg I H,
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[17] Froment G F, Bischoff K B. Chemical reactor analysis
and design. 2nd ed, New York: Wiley, 1990
ScienceDirect
Journalof Ndlunl Gas Chemistry
SCIENCE PRESS
www.elsevier.codmljcate/jnge
Article
A Novel Carbon Nanotube-Supported NiP Amorphous
Alloy Catalyst and Its Hydrogenation Activity
Yan J u ,
Ferigyi Li*
Department of Chemistry, Nanchang University, Nanchang 330031, Jiangxi, China
[ Manuscript received May 16, 2006; revised July 13, 20061
Abstract: A carbon nanotube-supported NiP amorphous catalyst (NiP/CNT) was prepared by induced
reduction. Benzene hydrogenation was used as a probe reaction for the study of catalytic activity. The
effects of the support on the activity and thermal stability of the supported catalyst were discussed based
on various characterizations, including XRD, TEM, ICP, XPS, H2-TPD, and DTA. In comparison with the
NiP amorphous alloy, the benzene conversion on NiP/CNT catalyst was lower, but the specific activity of
NiP/CNT was higher, which is attributed to the dispersion produced by the support, an electron-donating
effect, and the hydrogen-storage ability of CNT. The NiP/CNT thermal stability was improved because
of the dispersion and electronic effects and the good heat-conduction ability of the CNT support.
Key words: carbon nanotube; catalyst support; catalytic property; Ni; P; hydrogenation; benzene
1. Introduction
Carbon nanotube (CNT), discovered by Iijima in
1991 [I],is a novel support material. Currently, several investigations have been focusing on the deposition of metal catalysts (Ni, P t , Ru) on CNT, and the
materials thus obtained have displayed good catalytic
behavior [2-41. Amorphous alloy catalysts, such as
NiB and NIP, that are produced by chemical reduction
have attracted considerable attention because of their
superior catalytic activity and unique selectivity [5,6].
However, the thermal stability of the amorphous alloys needs to be improved further in order t o find the
applications in industry. Depositing amorphous alloys
on a suitable support is one of the promising routes
that can be used to improve their thermal stability
[7-91. However, thus far, no study has reported on
the use of CNT to support an NiP amorphous alloy.
In this study, an NiP amorphous alloy is deposited on
CNT using an induced reduction method. For com-
parison, NiP amorphous alloy is also prepared. Both
supported and unsupported catalysts are investigated
using XRD, TEM, ICP, XPS, H2-TPD, and DTA.
Their catalytic activities in the reaction of benzene
hydrogenation were examined. The objective of this
study is t o determine the effect of CNT support on
an NiP amorphous alloy.
2. Experimental
The CNT (id.: 20-40 nm, 0.d.: 40-60 nm, specific
surface area: 106.5 m2/g) used in this study was produced by methane decomposition over coprecipitated
Ni-Cu-A1 catalyst with the addition of sodium carbonate [lo]. It was left overnight to dissolved in 4 mol/L
HC1 solution to remove metal catalyst particles, followed by refluxing at 373 K for 10 h in 6 mol/L HN03
solution to create surface complexes on the CNT sur-
* Corresponding author. Tel: +86-791-8305436; E-mail: [email protected]
Supported by the National Natural Science Foundation of China (No. 20263003) and Natural Science Foundation of Jiangxi
province (No. 0250009)
314
Yan Ju et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
face to improve the surface hydrophilicity, and then
calcined in H2 flow (purity of 99.999%) at 673 K for
2 h to remove adsorbed impurities.
NiP/CNT amorphous catalyst was prepared by
induced reduction. First, a very small amount of NiB
was supported on CNT by impregnation, following
chemical reduction with KBH4 aqueous solution [ll].
The formed NiB amorphous alloy on CNT was used
as the inducing agent to create crystalline nuclei of
metallic nickel on the surface of CNT, and then NiP
alloy can be grown on it. The loadings of Ni and
B on CNT determined by ICP were 0.42 wt% and
0.02 wt%, respectively.
After being washed thoroughly with distilled water, the received CNT was immersed to NiC12. 6 H 2 0
solution for 4 h. Then, H3PO2 (30 wt%) aqueous solution was added to the mixture under vigorous stirring at 303 K to reduce Ni2+. Amine (such as ntripropylamine etc) was used to adjust the initial pH
of the solution to 10. The reaction lasted about 2 h
until no significant bubbles were observed in the solution. The resulting NiP/CNT amorphous catalyst
was washed thoroughly with distilled water and then
dried at 363 K in a N2 atmosphere.
For comparison, the NiP amorphous catalyst was
also prepared by reducing Ni2+ in aqueous solution
with H3PO2. The reaction was induced by adding a
drop of 3 mol/L KBH4 aqueous solution with stirring at 303 K t o the mixture (pH=lO) containing
NiC12. 6H20, H3PO2, and amine. All preparation
parameters, such as the molar ratio of P/Ni, reaction temperature, the initial pH value of the solution, and reaction time, were similar to those used
for the preparation of the NiP/CNT amorphous catalyst. The washing of sample was also identical t o the
procedure described above.
binding energy values were calibrated using the value
of contaminant carbon (C1,=284.6 eV) as the reference. The XPS Peak Fitting Program for WIN95
was used for spectral deconvolution. Temperatureprogrammed desorption of hydrogen (H2-TPD) was
carried using a self-designed apparatus that included
gas chromatographic and data processing systems.
For each experiment, the impurities on the surface
of the sample were removed by nitrogen flow (purity
of 99.98%) at 473 K for 1 h and the sample was then
cooled t o room temperature in the same flow. After hydrogen (purity of 99.999%) was preadsorbed by
the catalysts t o saturation, the hydrogen flow was replaced with nitrogen, which was maintained at room
temperature for 2 h. Then H2-TPD was carried out at
temperatures ranging from room temperature up to
873 K with a heating rate of 5 K/min controlled by a
temperature-programmed method. The thermal stability was determined by differential thermal analysis
(DTA, PYRIS Diamond).
2.2. Catalyst characterization
A D/max-IIIA polycrystalline X-ray instrument
with CuK, radiation (X=0.15418 nm) was used for
X-ray diffraction (XRD) studies. TEM pictures were
acquired using a Hitachi H-600 transmission electron
microscope. The composition and the Ni loading
of the as-prepared catalysts were analyzed using inductively coupled plasma (ICP, AES OPTIMA 5300
DV). The specific surface area (SBET)
of the catalysts was studied by low-temperature adsorption of
nitrogen at 77 K using ST-2000 instrument. The surface electronic states were determined by X-ray photoelectron spectroscopy (XPS, Kratos XSAM800). All
3.1. C h a r a c t e r i z a t i o n
2.3. Activity test
Benzene hydrogenation was used t o evaluate catalytic activity. The reaction was carried out in a
microreactor (stainless steel U-shaped tube, 4 mm
i.d.) under atmospheric pressure. Hydrogen (purity
of 99.999%) was used as the reductive gas and the
carrier gas. The catalyst was reduced in situ under
hydrogen flow (30 ml/min) at 473 K for 1 h before the
reaction was carried out. A volume of 1 pl benzene
was injected into the reactor containing 0.1 g catalyst
each time. The product and reactant were analyzed
using an on-line gas chromatograph (102G) equipped
with TCD and data integration.
The XRD pattern of NiP alloy (Figure 1) shows
a broad peak only around 28=45". This is assigned
t o the amorphous structure of the NIP alloy. Several
crystalline peaks corresponding to C (002), C (004),
and C (100) appear in the XRD pattern of CNT.
There is no significant change in the XRD pattern of
NiP/CNT catalyst, but the peak appearing around
28=45O is broadened compared with that of CNT.
Thus, it can be concluded that the NiP on the support
exists in an amorphous alloy.
Journal of Natural Gas Chemistry VoJ. 15 No. 4 2006
C(100)
(2)
(1)
I
20
30
40
,
,
,
,
50
I
60
,
,
,
,
I
,
,
70
,
,
80
2e1r )
Figure 1. XRD patterns of the NiP alloy (l), CNT (2),
and NiP/CNT catalyst (3)
315
The TEM morphology (Figure 2(a)) clearly shows
that the NiP particles are spherical with a wide range
of size of 10-120 nm, and some have aggregated into
large particles to reduce the surface energy. When
supported on CNT (Figure 2(b)), the NiP particles
disperse homogeneously, and the diameter of the particle shows a slight decrease (average size around
40 nm). Obviously, this can be attributed to the
dispersion produced by CNT that effectively inhibits
the aggregation of small NiP particles. Furthermore,
most of the NiP particles are tightly adhered to the
surface of the CNT with a spherical shape (Figure
2(c)), indicating that there is a strong interaction between the NiP amorphous alloy and the CNT.
Figure 2. TEM morphologies of NiP alloy (a) and NiP/CNT catalyst (b) and ( c )
Some of the characters of NIP amorphous alloy
and NiP/CNT catalyst are summarized in Table 1.
No significant change in the bulk composition of the
NiP alloy is observed because both NiP amorphous
alloy and NiP/CNT catalyst are prepared under similar conditions using a similar procedure. But the
of the NiP alloy is very
specific surface area (SBET)
low because of the considerable high aggregation of
NiP particles.
Table 1. Composition and SBET of the NiP alloy
and the NiP/CNT catalyst
Ni-loading
Composition
(wt%)
(mole ratio)
NiP
90.80
Ni83.9Pi6.i
NiP/CNT
16.52
Ni84.7pi5.3
Sample
S B E T ( ~g-')
'.
11.54
101.12
The XPS spectra of the NiP amorphous alloy and
the NiP/CNT catalyst are shown in Figure 3. Both
elemental and oxidized Ni and P can be observed in
the XPS spectra. For the NiP amorphous alloy, the
peaks of Ni 2p312 electron binding energy at 858.0 eV
and 852.4 eV can be assigned to oxidized and metallic Ni, respectively, and the peaks of P 2p electron
binding energy at 132.2 eV and 131.0 eV can be assigned to oxidized and elemental P, respectively. For
the NiP/CNT amorphous catalyst, the peaks of Ni
2p312electron binding energy at 854.3-857.8 eV and
851.7 eV can be assigned t o oxidized and metallic Ni,
respectively, and the peaks of P 2p electron binding
energy at 132.7 eV and 130.6 eV can be assigned to
oxidized and elemental P, respectively. In comparison
with the NiP amorphous alloy, the binding energy
peak of metallic Ni in the NiP/CNT catalyst shifts
negatively by 0.7 eV, implying that in the NiP/CNT
catalyst there is electronic interaction between Ni and
316
CNT. Menon et al. [12] studied the interactions of
Ni atom with carbon nanotube walls using a tightbinding molecular dynamics method. They reported
that there is a strong interaction between the carbon
layer and the Ni atom and there is electron transfer
between C and Ni. According to C 1s spectra of the
860
855
850
Binding energy (eV)
845
860
NiP/CNT catalyst, the peak at 284.5 eV corresponding to the binding energy of sp2 hybridization (C=C)
shifts positively by 0.5 eV from CNT (284.1 eV), indicating that CNT may experience a partial loss of
electrons. Therefore, in the NiP/CNT catalyst, CNT
can donate electrons to Ni to form electron-rich nickel.
855
850
845
Binding energy (eV)
H2-TPD spectra of the NiP amorphous alloy,
CNT, and NiP/CNT catalyst are shown in Figure 4.
The Hz-TPD spectrum of CNT has one broad peak
between 584 K and 750 K. It is attributed to the
hydrogen-storage ability of CNT. Only one peak
around 502.6 K is found in the H2-TPD spectra of
the NiP amorphous alloy, indicating that only one
type of hydrogen adsorption sites exists over the NiP
amorphous alloy. Three peaks were found to appear
in the H2-TPD curve of the NiP/CNT catalyst: two
peaks at lower temperature and one peak a t higher
temperature, which were assigned to the desorption
of hydrogen adsorbed on Ni atom and stored in CNT,
respectively. This indicates that there is a strong in-
140
135
I30
I25
Binding energy (eV)
teraction between CNT and the NiP amorphous alloy. It is this strong interaction between CNT and
the NiP alloy that changes the type of hydrogen adsorption sites over the NiP/CNT catalyst. Figure 4
also shows that the desorption of hydrogen occurs at
lower temperature from the NiP/CNT catalyst than
that from the NiP amorphous alloy, indicating that
the strength of Ni-H adsorption bond is weakened in
the NiP/CNT catalyst. Generally, hydrogen is easy to
dissociate and can easily be adsorbed on the electrondeficient Ni [13],resulting in a relatively strong Ni-H
bond. Thus, the electron-rich Ni centers would reversibly lead to weak Ni-H bond. From the above XPS
analysis, it is obvious that the relatively weaker Ni-H
317
bond in the NiP/CNT catalyst can be attributed t o
the electron-donating effect of CNT on the NiP alloy.
1
I
683.8
neighboring Ni active sites and, in turn, enhance the
surface hydrogenation reactions. (3) The hydrogenstorage ability of CNT makes the NiP/CNT catalyst
adsorb more hydrogen than the NIP amorphous alloy,
which favors the improvement of the catalytic activity
of NiP/CNT catalyst.
70
I,,,A,,
1
668.9
426.1 485.5
300
400
500
,
600
700
800
,
? ) ,
900
Temperature (K)
Figure 4. Ha-TPD spectra of the NIP amorphous alloy, CNT, and NiP/CNT catalyst
(I) CNT, (2) Nip, (3) NiP/CNT
3.2. Catalytic hydrogenation activity
Cyclohexane is found to be the only product over
the NiP amorphous alloy and NiP/CNT catalyst in
benzene hydrogenation. Figure 5 shows the benzene
conversion and the specific activity of the two catalysts in benzene hydrogenation, respectively. Because
the NiP amorphous alloy and NiP/CNT catalyst are
the two catalysts with different Ni-loading, in order
to compare these two catalysts correctly, the concept
of specific activity is introduced. Here, specific activity represents the ratio of benzene conversion t o
Ni loading of catalysts. Although the benzene conversion of the NiP/CNT catalyst is lower than that of
the NiP amorphous alloy, its specific activity is higher
than that of the NiP amorphous alloy. According t o
the above characterization, this can be ascribed to
the following factors. (1) The dispersion produced
by the CNT support can prevent the aggregation of
NiP particles and therefore decrease their average diameters. Generally, both higher dispersion of active
components and smaller diameters are beneficial in
improving the reaction activity. (2) More electrons
transfer from CNT to Ni atom because of the electrondonating effect of support t o form electron-rich Ni.
On the one hand, electron-rich Ni relatively brings advantage to benzene hydrogenation because the active
site of hydrogenation reaction is elemental or metallic
Ni. On the other hand, electron-rich Ni centers would
weaken the Ni-H bond. The relatively weak Ni-H
bond can facilitate the transfer of the adsorbed hydrogen atoms to the benzene molecules adsorbed on the
0
1
393
I
I
I
I
403
413
423
433
Reaction temperature (K)
Figure 5. Benzene conversions (a) and specific activities (b) of the NiP amorphous alloy and
NiP/CNT catalyst in benzene hydrogenation
(1) NIP, (2) NiP/CNT
3.3. The thermal stability
The thermal stability of the samples is investigated by DTA (Figure 6). The NiP amorphous alloy
has one sharp exothermic peak at 589 K and a shoulder peak a t about 633 K , whereas the NiP/CNT catalyst has only one crystallization peak at 652 K. The
crystallization temperature of the NiP/CNT catalyst
is higher than that of the NiP amorphous alloy. This
indicates that supporting the NiP amorphous alloy on
CNT can improve its thermal stability. The promoting effect of CNT support on the thermal stability
of the NiP amorphous alloy can be explained by the
following three aspects. (1) The dispersion effect of
the CNT support on the NiP alloy can effectively inhibit the diffusion and aggregation process of the NiP
particles, thus increasing the temperature of crystallization. (2) The electronic interaction between the
NiP alloy and CNT support can prevent the contact
318
of the NIP alloy with oxygen and also prevent the
aggregation of NiP alloy particles at high temperature [14]. (3) In addition, CNT is a good conductor
of heat. In the presence of the CNT support, the
heat released during the crystallization of the NiP alloy can be quickly transferred from the surface to the
support matrix, which inhibited the increase of the
surface temperature and in turn improved the thermal
stability of the NiP amorphous alloy on the support
surface.
,
300
400
500
600
,
I
I
700
,
,
I
I
I
800
I
.
,
900
Temperature (K)
Figure 0. DTA curve of the NiP amorphous alloy (1)
and NiP/CNT catalyst (2)
4. Conclusions
Depositing the NiP amorphous alloy on CNT can
improve its catalytic activity, which is attributed to
the dispersion produced by the support, an electrondonating effect, and the hydrogen-storage ability of
CNT. The thermal stability of the NiP/CNT catalyst
is better than that of the NiP amorphous alloy. The
dispersion and electronic effect and the good heatconducting ability of CNT are the main factors responsible for the increase of the thermal stability for
the NiP/CNT catalyst.
References
[l] Iijima S. Nature, 1991, 354: 56
[2] Yin S F, Xu B Q, Zhu W X, Ng C F, Zhou X P, Au
C T. Catal Today, 2004, 93-95: 27
[3] Reshetenko T V, Avdeeva L B, Ismagilov Z R,
Chuvilin A L. Carbon, 2004, 42(1): 143
[4] Li C H, Yu Z X, Yao K F, Ji S F, Liang J. J Mole
Catal A : Chem, 2005, 226(1): 101
[5] Chen Y. Catal Today, 1998, 44(1-4): 3
[6] Deng J-F, Li H X, Wang W J. Catal Today, 1999,
51(1): 113
[7] Xie S H, Qiao M H, Li H X, Wang W J, Deng J-F.
Appl Catal A , 1999, 176(1): 129
[8] Wang W J , Qiao M H, Li H X, Deng J-F. Appl Catal
A , 1998, 166(2): L243
[9] Wang W J, Qiao M H, Yang J, Xie S H, Deng J F.
Appl Catal A , 1997, 163(1-2): 101
[lo] Ju Y, Li F Y, Wei R Z. J Serb Chem SOC,2005, 70:
277
[ll] Wang M W , Li F Y, Zhang R B. Catal Today, 2004,
93-95: 603
[12] Menon M, Andriotis A N, Roudakis G E. Chem Phys
Lett, 2000, 320(5,6): 425
[13] Yoshida S , Yamashita H, Funabiki T, Yonezawa T. J
Chem Soc Faraday Trans I, 1984, 80(6): 1435
[14] Li H X, Wang W J, Li H, Deng J-F. J Catal, 2000,
194(2): 211
ScienceDirect
Journalot Natural Gar Chemistry
SCIENCE PRESS
www.eIsevier.codocate/jllocsteljngc
Article
Study on the Nanosized Amorphous Ru-Fe-B/ZrOa Alloy Catalyst
for Benzene Selective Hydrogenation to Cyclohexene
Shouchang Liu*,
Zhongyi Liu,
Shuhui Zhao,
Yongmei Wu,
Zheng Wang,
Peng Yuan
Department of Chemistry, Zhengzhou University, Zhengzhou 450052, Henan, China
[Manuscript received May 29, 2006; revised July 31, 20061
Abstract:
A novel nanosized amorphous Ru-Fe-B/ZrOz alloy catalyst for benzene selective hydrogenation to cyclohexene was investigated. The superior properties of this catalyst were attributed to the
combination of the nanosize and the amorphous character as well as to its textural character. In addition,
the concentration of zinc ions, the content of ZrOa in the slurry, and the pretreatment of the catalyst were
found to be effective in improving the activity and the selectivity of the catalyst.
Key words: Ru-Fe-B/ZrOz amorphous catalyst; benzene selective hydrogenation; cyclohexene
1. Introduction
Selective hydrogenation of benzene to cyclohexene
has received considerable attention owing to its environmentally benign process, atomic economy, and
potentially wide industrial applications [l-71. However, it appears that such selective hydrogenation is
very difficult because cyclohexene is chemically active due to the presence of the double bond, on which
further hydrogenation to cyclohexane can occur easily. Therefore, considerable effort has been made t o
improve the selectivity and increase the yield of cyclohexene, and these efforts have met with significant
progress [8-131. Recently, amorphous alloy material
as a good-quality, novel type of catalyst has received
increasing attention because of its excellent activity
and superior selectivity in many hydrogenation reactions [14,15], and it has also been used in the selective hydrogenation of benzene to cyclohexene [16-181.
However, a large part of the work on the amorphous
alloy catalyst has been carried out only in laboratories; to date, there are no published reports regarding its application in industries. As the amorphous
alloy is thermodynamically metastable, the crystal-
lization deactivation process of the amorphous alloy
could occur spontaneously during the reaction, especially at high temperatures, which limits the industrial application of amorphous alloy catalysts. In this
study, a novel nanosized amorphous Ru-Fe-B/Zr02
alloy catalyst prepared by chemical reduction for benzene hydrogenation to cyclohexene was developed in
a pilot study, which exhibited higher activity and better selectivity. In our previous study, reaction conditions such as suitable temperature, appropriate pressure, optimal ratio of water to benzene in the reaction
system, and stirring rate that favored benzene selective hydrogenation over the amorphous alloy Ru-FeB/ZrOz catalyst were studied in detail [19]. It was
found that a chemical environment around the catalysts is crucial to improve the selectivity t o cyclohexene. In this study, the catalysts were subjected to
various characterizations using XRD, TEM, SAED,
and N2 physisorption. In particular, the effects of a
chemical environment on the performance of this RuFe-B/ZrO:! catalyst under conditions for a pilot study
were investigated. The aim of this study is to investigate the prospects of industrial application for the
catalysts.
* Corresponding author. Tel and Fax: +86-371-67763706; E-mail: Iiushouchang@zzu,edu.cn
320
Shouchang Liu et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
2. E x p e r i m e n t a l
2.1. C a t a l y s t p r e p a r a t i o n
The Ru-Fe-B/ZrOz catalyst sample was prepared
as follows. An appropriate amount of zirconium dioxide was added to 50 ml RuC13 and FeSO4 solution
(0.05 mol/L) with stirring for 30 min; 50 ml NaBH4
solution (0.5 niol/L) was then added drop by drop to
the above solution (mass ratio Ru/Zr02=10-20%).
The agitation was continued for 5 min. The black
precipitate was kept in liquor solution for a while and
then filtered and washed thoroughly with distilled water until neutrality; the nanosized amorphous Ru-FeB/ZrOz alloy catalyst was obtained.
2.2. C a t a l y t i c test
The selective hydrogenation of benzene was carried out in a 1-L autoclave. A total of 280 ml H20,
19.6 g ZnS04.7H20, and 4 g catalyst were introduced.
The autoclave was sealed and then filled with Hz more
than four times t o exclude air. Initially, the stirring
rate was adjusted to 600 r/min and the H2 pressure
was maintained at 3.0 MPa. When the temperature
increased up to 413 K, 140 ml of benzene was charged
into the reactor. After this, the stirring rate was adjusted to 1000 r/min and the pressure of H2 was elevated to 5.0 MPa and the reaction was considered
to have begun. The reaction process was monitored
by taking a small amount of reaction mixture a t intervals, followed by analysis in a gas chromatograph
equipped with FID. The quantification of benzene,
cyclohexene, and cyclohexane was carried out using
calibration curves.
Definition:
benzene converted by 1 g of catalyst per hour when
the conversion of benzene is at 40 mol%, and it is
a general industrial target for evaluating catalytic activity. S40 indicates cyclohexene selectivity when benzene conversion is 40 mol%. Benzene conversion and
the cyclohexene selectivity were plotted as a function
of time, respectively. S40 would be obtained from this
plot. 7 4 0 is given by:
where, V is the volume of benzene (ml); ~ B is
Z the
density of benzene (0.88 g/ml); XBZis the conversion
of benzene (40 mol%); t40 is reaction time (h) a t 40
mol% benzene conversion; Adcat.is the mass of active
component Ru in the catalyst (g).
2.3. Pilot s t u d y of the Ru-Fe-B/ZrOz catalyst
A total of 14 L of distilled water and 0.2 kg of
Ru-Fe-B/ZrOz catalyst were introduced into a 50-L
autoclave. Then the autoclave was sealed and air was
flushed out using nitrogen gas several times. Next,
hydrogen replaced the nitrogen and its pressure was
maintained at 3.0 MPa. The initial stirring rate was
300 r/min, and the temperature was raised at the rate
of 80 K/h. When the desired temperature, 403 K,
was reached, 7 L of benzene was introduced. Subsequently, the pressure was adjusted to 4.5 MPa, the
stirring rate was elevated to 600 r/min, and the reaction was considered to have begun. The reaction
temperature was controlled at 4 1 3 f 2 K. The analysis
of the product was carried out as mentioned above.
2.4. C h a r a c t e r i z a t i o n methods
XBZ =
mole of reacted benzene
x 100%
mole of initial benzene
SHE=
mole of cyclohexene formed
x 100%
mole of reacted benzene
YHE=
mole of cyclohexene formed
x 100%
mole of initial benzene
Where, XBEis benzene conversion, SHEis cyclohexene selectivity, YHE is cyclohexene yield.
The catalyst activity and selectivity is given by
740 and 5’40, respectively. 740 indicates the mass of
The phases of the Ru-Fe-B/ZrOz catalyst were determined by X-ray diffraction (XRD) using Cu K, radiation; the tube voltage was 40 kV,and the tube current was 40 mA. The surface morphology of the active
component on the support and the particle size were
determined with the aid of a high-resolution transmission electron microscope (HRTEM, JEM-201 l),using
an accelerating voltage of 100 kV. The amorphous
character of the as-prepared catalysts was verified by
selected area electron diffraction (SAED). The textural character of the as-prepared catalyst was determined by N2 physisorption a t 77 K on a Micronieritics
TriStar 3000 apparatus.
Journal of Natural Gas Chemistry Vol. 15 N o . 4 2006
Characterization of the Ru-Fe-B/ZrOZ
catalyst
3.1.
3.1.1. XRD and TEM of the Ru-Fe-B/ZrOz
Figure 1 shows the changes in the XRD patterns
of the Ru-Fe-B/ZrOZ catalyst during the heat pretreatment. As shown in Figure 1, when the sample
was treated at temperatures below 673 K, no significant change in the XRD patterns was observed.
Only the diffraction peaks of monoclinic zirconia were
observed. No distinct peaks corresponding to Ru
phase was seen in the patterns from 293 to 673 K.
Therefore, it can be safely assumed that the Ru-Fe-B
amorphous alloy is quiet stable below 673 K, which
was attributed to the high dispersion of the Ru-FeB amorphous alloy particles on the ZrO2 matrix and
the stabilizing effect of a small amount, of the additive of Fe. This is consistent with that reported
in the literature [16]. However, various crystalline
diffractional peaks corresponding to metallic Ru appeared on the XRD patterns of Ru-Fe-B/ZrOz sample
when the temperature was above 673 K , indicating the
crystallization of Ru-Fe-B amorphous alloy and the
formation of crystallized Ru. The intensity of these
Ru crystallized peaks increased gradually with the increase in treating temperature from 673 K to 873 K.
Therefore, it could be concluded that the crystallization process of the Ru-Fe-B/ZrOa amorphous catalyst
proceeded stepwise during which the crystallized Ru
phase formed simultaneously.
32 1
Figures 2 and 3 show the TEM and SAED images
of the Ru-Fe-B/ZrOz catalyst, respectively. The light
gray circular or elliptic flakes shown in Figure 2 were
naiiosized ZrOz crystallites, and the black particles
were the active components comprising RUB and Fe
amorphous alloys, with the particle diameter ranging
from 3 to 6 nm. As shown in Figure 3, the Ru-FeB/ZrOs sample showed a number of diffraction circles
and some small, white flecks on the diffraction circle
that were identified as Zr02 crystallites, indicating
the typical amorphous structure.
Figure 2. TEM image of the Ru-Fe-B/ZrOz catalyst
Figure 3. SAED pattern of the Ru-Fe-B/ZrOz catalyst
3.1.2. Texture character of the Ru-Fe-B/ZrOz
catalyst
.-0
m
C
c
10
20
30
40
50
60
70
281(O )
Figure 1. XRD patterns of the Ru-Fe-B/ZrOa catalyst at different temperatures
(1) 293 K , (2) 373 K , (3) 473 K, (4) 573 K , (5) 673 K ,
(6) 773 K , ( 7 ) 873 K
Figure 4 shows the curves of the N2 adsorptiondesorption isotherm of the Ru-Fe-B/ZrOz catalyst. It
can be concluded from the hysteresis curves that the
shape of the pores of this catalyst is similar to that of a
capillary tube, with both sides open and the pore size
distribution mainly ranging from 2 to 50 nm. The relative pressure ( p / p o ) at the separate region in the adsorption curves and desorption curves was more than
0.8, indicative of the catalyst with bigger pore diameter. Figure 5 shows the differential curve of the pore
size distribution of the Ru-Fe-B/ZrOZ catalyst. It wa5
322
observed that the pore size distribution of the catalyst
is between 2 nm and 100 nm and the most probable
pore diameter is about 28 nm. The results of the
measurement show that the BET surface area of RuFe-B/ZrOz catalyst is about 29 ni2.g-l , the specific
pore volume is about 0.20 cm3.g-', and the average
pore diameter ( 4 V I A by BET) is about 28 nm, which
is in good agreement with the data obtained from the
differential curve of pore size distribution.
100
I
d
B
f
s:
%
$
C
so
0
0.2
0.4
0.6
Relative pressure @ l p , )
1 .o
0.8
90
85 -
.
$
......
80
:
7s
-
70
-
i
mi
65
, , , , I , , ,
,
I
,
,
,
.
I
,
,
,
,
'
~
,
(
,
1
~
~
,
~
l
~
~
~
,
l
Figure 4. The NZadsorption and desorption isotherm
of the Ru-Fe-B/ZrOz catalyst
C
I
-
I
J
I
\
x
0
I
I
10
100
Dlnm
Figure 5. The differential curve of pore distribution
of the Ru-Fe-B/ZrOz catalyst
Activity and selectivity of the Ru-FeB/ZrOz catalyst
3.2.
Figure 6 shows the catalytic performance of the
Ru-Fe-B/ZrOz in a stirring autoclave. It was observed
that with the increase in benzene conversion, the cyclohexene yield was more than 50% (see Figure 6 (a)),
and the selectivity to cyclohexene over the catalyst is
beyond 85% in the initial stage (see Figure 6 (b)).
From Figure 6, it can be observed that t40 is 20.4 min
and that ,940 is 81.6% a t the benzene conversion of
40% by interpolation method. The mass of active
component Ru in the catalyst is 0.64 g. On the basis
of these data and according to equation (1) mentioned
above, 740 can be calculated as follows: 740=226 h-' ,
i.e. 226 g of benzene is converted by 1 g of catalyst
per hour at benzene conversion of 40%. In comparison
with the data reported thus far in the literature, tlit
Ru-Fe-B/ZrOz catalyst has exhibited higher activity
and better selectivity to cyclohexene. According to
the various characterizations described above, the superior catalytic properties of the Ru-Fe-B/ZrOZ catalysts could be attributed to the following. First, the
combination of the ultrafine size and the amorphous
character, which could offer more active centers for
benzene hydrogenation. Second, its textural character, which is helpful to the exterior diffusion of the intermediate cyclohexene, thereby avoiding its further
hydrogenation to cyclohexane. Third, the promoting
effect of the iron and boron species dispersed among
~
,
,
~
323
Journal of Natural Gas Chemistry VoJ. 1 5 No. 4 2006
the ruthenium particles is also an important factor because iron has low electronic affinity and boron is a n
electron-deficient element, which could promote the
water adsorption on the catalyst surface, thus greatly
enhancing the hydrophilicity of the catalyst and suppressing further hydrogenation of cyclohexene to cyclohexane [20].
3.3.
Operation conditions of the Ru-FeB/ZrOz catalyst
The reaction for benzene selective hydrogenation to cyclohexene over the Ru-Fe-B/ZrOz nanosized catalyst comprises four phases: vapor (hydrogen), oil, aqueous, and solid catalyst. Besides, suitable amounts of zinc sulfate and zirconium dioxide
are always added into the aqueous solution. Water
can form a stagnant layer around the catalyst surface, competing with cyclohexene on surface adsorption because of the low solubility of benzene in water,
thereby not favoring the hydrogenation of cyclohexene to cyclohexane. Catalyst hydrophilicity can be
improved by adsorbing zinc sulfate on the catalyst
surface, which is beneficial for enhancing the selectivity to cyclohexene. Zirconia as a dispersing agent
can prolong the life of the catalyst and improve the
selectivity to cyclohexene.
Our previous studies have shown that for the selective hydrogenation of benzene over the catalysts,
the suitable temperature was 408-413 K, the appropriate pressure was 4.5-5.0 MPa, the optimal volume
ratio of water to benzene in the reaction system was
2:1, and the stirring rate should be high enough t o
exclude the effect of diffusion. On the basis of the
above-mentioned detailed studies, we further investigated in particular the influences of the composition
of the reaction system, including the concentration
of zinc ions in the aqueous solution, the content of
ZrOz in the slurry, and the catalyst pretreatment, on
the properties of the catalyst in pilot units so as to
acquire some valuable information for industrial application for the catalysts. Then, the operation condition of the Ru-Fe-B/ZrOa catalyst is determined.
3.3.1. Influences of concentration of zinc ions
in the aqueous solution
The influence of the concentration of zinc ions in
the aqueous solution, in the absence of ZrOa, on the
performance of the catalysts is shown in Table 1.
Table 1shows that benzene conversion was higher,
whereas the selectivity to cyclohexene was lower in
the absence of zinc sulfate in the slurry, whose pH
value was 7.2. With the addition of zinc sulfate to
the slurry, the pH value declined, benzene conversions
decreased, and selectivities to cyclohexene showed a
dramat,ic increase. However, the concentration of zinc
up to 0.8 mol/L led to the highest benzene conversion and lower selectivity to cyclohexene, indicating
that the yield of the byproduct cyclohexane increased.
It was determined that the optimal concentration of
Zn2+ in the slurry is 0.50-0.60 mol/L, with the pH
values of 5.4-5.5. Under these conditions, both higher
selectivity and yield of cyclohexene can be obtained
at an acceptable benzene conversion rate.
Table 1. Selectivity and yield over Ru-Fe-B/ZrOz
catalyst for conversion of benzene to cyclohexene
at different concentrations of the zinc ion
0
7.2
50.9
35.2
17.6
0.10
6.1
25.0
0.30
5.8
26.1
71.9
76.7
18.0
20.0
0.50
5.5
37.6
74.0
0.60
5.4
39.9
70.6
27.9
28.2
0.80
5.2
44.6
62.3
27.8
Reaction conditions: 280 ml H20, 140 ml C6&, 4 g Ru-FeB/ZrOz catalyst (0.64 g Ru), 140
p H , = 5 MPa, 1000 r/min,
without pretreatment at reaction time of 5 min.
"c,
According to the above-mentioned results, the
pilot study was carried out in a 50-L stirring autoclave with the amount of [Zn2+] being 0.50 and
0.60 mol/L, respectively, with other conditions remaining unchanged. The results of the reactions are
shown in Figure 7.
From Figure 7 (c), it can be seen that when
the amount of [Zn2+]in the slurry was 0.5 mo1.L-'
(curve Y H E ( ~a) maximum
),
yield of cyclohexene of
35 mol% was obtained at the reaction time of 45
min, corresponding to benzene conversion of 66 mol%
(Figure 7 (a)); when the amount of [Zn2+] in the
slurry was 0.6 mo1.L-' (curve Y H E ( ~a) maximum
),
yield of cyclohexene of 39 mol% was obtained at the
reaction time of 19 min, corresponding to benzene
conversion of 64 mol% (Figure 7 (a)). Therefore, it
can be concluded that in the absence of ZrO2, bett,er results were obtained when the concentration of
zinc ions in the slurry was 0.6 mo1.L-l compared with
0.5 mo1.L-l.
324
80 -
.
s.
y:
60 40
-
20
-
0 0
0
0
10
20
30
40
50
60
I
,
10
l
,
l
,
l
30
20
40
,
l
.
50
.-
l
60
0
10
20
rlmin
rlmin
30
40
SO
60
rlmin
Figure 7. Benzene conversion (a), cyclohexene selectivity (b) and yield (c) over Ru-Fe-B/ZrOz catalyst at
different zinc ion concentrations
( I ) [Zn2+]=0.5mol.L-l, (2) [Zn2+]=0.6 mo1.L-l
3.3.2. Influences of the content of ZrOz i n the
slurry
Apart from the presence of zinc sulfate, nanosized
zirconia as a dispersing agent is added t o the slurry
100 -
90
80
-
80
60
-
40
-
in the presence of different concentrations of [Zn''].
Figure 8 shows the influence of content of zirconia in
the slurry on the performance of the Ru-Fe-B/ZrOz
catalyst in the presence of 0.5 mo1.L-l Zn2+.
70
s.
> 60
2
50
20 -
0
40
L
..
0
10
20
30
tlmin
40
50
60
0
20
40
60
X,, I mol%
80
100
0
10
20
30
40
50
60
70
tlmin
Figure 8. Effect of ZrO2 on the performance of Ru-Fe-B/ZrO2 catalyst at [Zn2+]=0.5 mo1.L-'
(a) Benzene coversion, (b) Cyclohexene selectivity, (c) Cyclohexene yield; (1) without ZrO2, (2) Cat./ZrOz=l:l, (3) Cat./ZrO2=1:2,
(4) Cat./ZrOz=1:2.5
From Figure 8 (a), it can be seen that with the
increase in the amount of ZrOz in the slurry, there is a
corresponding, simultaneous increase in benzene conversions because of the dispersing effect of zirconia on
the catalyst. From Figure 8 (b), it can be observed
that with the increase in the amount of ZrO2, the selectivity to cyclohexene varies in a complicated way
and too high benzene conversion results in a decrease
in selectivity at the mass ratio of catalyst to zirconia
of 1:2.5 (curve 4), and among them, the highest selectivity to cyclohexene was observed at the mass ratio
of catalyst to zirconia of 1:2 (curve 3). From Figure
8 (c), it can be seen that the highest yield of cyclo-
hexene can be achieved at the mass ratio of catalyst
to zirconia of 1:2. From Figure 8 (a)-(c), it can be observed that the optimal amount of ZrOz in the slurry
is equivalent to the mass ratio of catalyst to zirconia
of 1:2 in the presence of 0.5 mo1.L-' Zn2+, a t which
the highest yield of cyclohexene achieved is 46 mol%,
with 65% selectivity t o cyclohexene a t 70 mol% benzene conversion and with the corresponding reaction
time being approximately 22 min. In comparison with
the data shown in Table 1 in the absence of zirconia,
it is suggested that a suitable amount of zirconia in
the slurry can not only enhance activity of the catalyst but also significantly improve the selectivity and
325
yield of cyclohexene.
The hydrogenation reaction was carried out by
fixing the mass ratio of catalyst to zirconia at 1:2
and by altering the concentration of zinc sulfate from
0.5 t o 0.6 mo1.L-l. The results show that the maximum yield of cyclohexene achieved was 43 mol% with
62% selectivity to cyclohexene at 70 mol% benzene
conversion and the corresponding reaction time being
approximately 17 min.
In contrast to the two results discussed above, it,
can be observed that the performance of the catalyst
for benzene selective hydrogenation to cyclohexene is
closely related to the chemical environment around
the catalyst particles. When zirconia and zinc sulfa,te
are used in combination in the slurry, the variations in
benzene conversion and in the selectivity and yield of
cyclohexene are considerable different from the case
where only zinc sulfate is present in the slurry. It
should also be noted that the experiment was carried
out in a pilot study, and few similar studies have been
reported previously.
3.3.3. Influence of pretreatment to the Ru-FeB/ZrOz catalyst
The pretreatment of the catalyst was also carried
out in the pilot study. The operation is that under
reaction conditions the catalyst runs for a period of
time in tJheslurry in the absence of benzene. Figure 9
shows the comparison of performance of the runs with
and without pretreatment for 12 h for the nanosized
amorphous Ru-Fe-B/ZrO:! alloy catalysts.
50
0
10
20
30
40
50
60
I
0
10
20
30
40
50
60
tlmin
timin
Figure 9. Selectivity and yield for conversion of benzene to cyclohexene on Ru-Fe-B/ZrOn catalyst without
pretreatment (1) and with pretreatment for 12 h (2)
From Figure 9 (a), it can be seen clearly that
in the case of the pretreated catalyst, the benzene
conversion showed a dramatic decrease, whereas the
selectivities to cyclohexene showed a dramatic increase. From Figure 9 (b), it can be seen that for
the pretreated catalyst the yields of cyclohexene increased gradually until the reaction time was about
70 min, and the maximum yield of cyclohexene was
not achieved. This is considerably different from the
case of the catalyst with no pretreatment. In the
case of the pretreated catalyst, when the benzene conversion reached 40 mol%, a cyclohexene selectivity
of more than 80% and a cyclohexene yield of more
than 32 mol% were achieved in a reaction time of 40
min. Moreover, our pilot study also showed that the
nanosized amorphous Ru-Fe-B/ZrOz alloy catalysts
became more stable as a result of pretreatment. However, sedimentation and separation performance of
the pretreated catalysts could evidently be improved,
thereby effectively avoiding the loss of catalysts during the continuous process of product separation. It is
believed that the improvement of the catalytic properties of the nanosized amorphous Ru-Fe-B/ZrOz alloy
catalysts may well be responsible for changing from
hydrophobicity to hydrophilicity via the pretreatment
of catalyst.
On the basis of the results of the pilot studies, it
is proposed that the Ru-Fe-B/ZrOz catalysts would
have potential industrial application.
4. Conclusions
Using various characterizations and studies on the
performance of the novel nanosized amorphous Ru-FeB/ZrOz alloy catalyst prepared by chemical reduction
with NaBH4 in a pilot study, the main conclusions are
as follows.
The Ru-Fe-B/ZrOz catalyst belongs t o an amor-
326
Shouchang Liu et a1./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
phous alloy material of nanosize. The textural character shows that the shape of the pores of this catalyst
is similar to a capillary tube, with both sides open
and the pore size distribution mainly ranging from 2
to 50 nm, with the most probable pore diameter being
approximately 28 nm. The high activity and excellent
selectivity of this catalyst for benzene selective hydrogenation to cyclohexene is the result of its structural
as well as textural characters.
The hydrogenation reaction over the Ru-FeB/ZrOZ catalyst has to be carried out in a suitable
chemical environment that aids in the formation of
cyclohexene molecules.
The slurry containing zinc sulfate and zirconia is
necessary for enhancing the activity and selectivity
of the catalysts to cyclohexene. The results of pilot studies show that the appropriate concentration
of zinc sulfate in the aqueous solution is 0.50 mo1.L-’
and the suitable amount of zirconia in the slurry is
equivalent to the mass ratio of catalyst t o zirconia of
1:2. It also shows that the performance of the catalysts can be considerably enhanced by pretreatment.
These conclusions are considered to be very important
information for the potential industrial application of
the Ru-Fe-B/ZrOz catalyst.
References
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[16] Li H X, Wang W J , Li H, Deng J F. J Catal, 2000,
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[17] Li H X, Li H, Dai W L, Qiao M H. Appl Catal A ,
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Appl Catal A , 1999, 176(1): 129
[19] Liu S C, Luo G, Wang H R, Xie Y L, Yang B J, €Ian
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47
ScienceDirect
Journalof Natural Gas Clirniistrv
SCIENCE PRESS
www.elsevier.comnocate/jngc
Article
Syngas Production by Methane Reforming with Carbon
Dioxide on Noble Metal Catalysts
M. R e ~ a e i l > ~S., M. Alavil,
S. Sahebdelfar2 ,
Zi-Feng Yan3*
1. Chemzcal Engineering Department, Iran University of Science and Technoloqy, P. 0. Box 16315-67, Tehran, Iran;
2. Petrochemical Research & Technology Company ( N P C - R T ) , Tehran, Iran;
3. State Key Laboratory for Heavy Oil Processing, Key Laboratory of Catalysis, CNPC,
China University of Petroleum, Dongying 257061, Shan,dong, China
[Manuscript received June 6, 2006, revised August 14, 2006 ]
Abstract: A series of noble metal catalysts (Ru, Rh, Ir, Pt, and Pd) supported on alumina-stabilized
magnesia (Spinel) were used to produce syngas by methane reforming with carbon dioxide. The synthesized
catalysts were characterized using BET, T P R , TPO, TPH, and H2S chemisorption techniques. The activity
results showed high activity and stability for the Ru and Rh catalysts. The T P O and TPH analyses
indicated that the main reason for lower activity and stability of the Pd catalyst was the formation of the
less reactive deposited carbon and sintering of the catalyst.
Key words: noble metal; syngas; dry reforming; carbon dioxide; methane
1. Introduction
Carbon dioxide reforming of methane to synthesis gas, which converts two of the most abundant and
carbon-containing greenhouse gases (CH4 and C02)
into a useful chemical product, has received considerable attention in recent years [1-6]. This reaction
has the following important advantages: (i) the formation of a suitable H2/CO ratio for use in FischerTropsh synthesis [7], (ii) the reduction of CO2 and
methane emissions, as both gases cause heavy greenhouse effect [8],and (iii) better use in chemical energy
transmission [9]. Research on the nickel catalysts used
for this reaction has mainly focused on the intrinsic
activity of the metal phase, stability towards carbon
formation, the type of the support most suitable for
improving the efficiency of the catalyst, and the reaction mechanism. Although the extensively developed
nickel catalysts have shown very high activity from
the industrial point of view, they are completely deactivated within a few hours of reaction due to the
formation of stable and inactive carbon on the surface [10-12]. Recently, several studies on the dry reforming of methane focused on the noble metal catalysts, which exhibit better activity and very high stability due t o the less sensitivity to carbon deposition
[13,14]. Rostrup-Nie1sc.n [14] has compared catalysts
based on nickel, ruthenium, rhodium, palladium, iridium, and platinum and has reported that ruthenium
and rhodium showed high selectivity for carbon-free
operation.
In this article, a series of noble metal catalysts
supported on alumina-stabilized magnesia were used
in methane reforming with carbon dioxide for the production of synthesis gas and the activity and stability
of different catalysts were investigated.
2. Experimental
2.1. Materials
* Corresponding author. Tel: +86 546 8391527; Fax: +86 546 8391971;
E-mail: zfyancatehdpu.edu.cr1
328
M. Rezaei
ef; aI./
Ru(NO)(NO3)3 as precursors of Pd, P t , Rh, Ir, and
Ru, respectively.
The supported noble metal catalysts were prepared by the impregnating of pellets of aluminastabilized magnesia (the molar ratio of Mg/A1=7/1)
with solutions of metal precursors to obtain 1% of
metals. Before impregnation, the alumina stabilizedmagnesia was calcined a t 975 "C for 4 h. After impregnation, the pellets were dried at 80 "C and calcined at 450 "C for l h. Before the reaction, the
different samples were reduced with a flow of pure hydrogen gas (GHSV=2000 L/kg,,,.h) at a heating rate
of 10 "C/min from room temperature to 525 "C and
then maintained at 525 "C for 4 11.
2.3. Characterization
The surface areas (BET) were determined by nitrogen adsorption at -196 "C using an automated gas
adsorption analyzer (The Tristar 3000, Micromeritics). The pore size distribution was calculated from
the desorption branch of the isotherm using the Barrett, Joyner, and Halenda (B,JH) method. The XRD
patterns were recorded on an X-ray diffractometer
(PANalytical X'Pert-Pro) using a Cu-Ka monochroniatized radiation source and a Ni filter in the range
20=1Oo-8O0. The surface area of the metals was determined by chemisorption of hydrogen sulfide, as
described elsewhere [15] (conditions: the volume ratio of H2S/H2=15x10W6, 550 "C, 100 h). The surface area of the metal was calculated by assuming
a sulfur monolayer of 4 4 . 5 ~ 1 0 -gS.cmP2
~
for nickel,
which corresponds to 0.5 sulfur atom per nickel atom
(S/Ni=0.5) on the (100) surface [15]. This was
assumed to be close to the composition of the
rnonolayers on the noble metals. In other words,
the surface area of the metal was calculated using & = 4 4 0 ~ 1 0 - ~ ,equivalent t o 1 rn2.g-', So being the sulfur capacity of the catalyst (PgSlgmetal).
The mean diameter of a metal particle can thus be
estimated from Equation (1).
where dlncta1is given in nm and Xmeta1is the weight
percent of metal in the reduced state. Temperatureprogrammed reduction (TPR.) was carried out using
an automatic apparatus (ChemBET-3000 TPR/TPD,
Quantachronie) equipped with thermal conductivity
detector. The fresh catalyst (200 mg) was subjected to a heat treatment (10 "C/min) in a gas flow
(30 ml/min) containing a mixture of H2:Ar (1090).
Before the TPR experiment, the samples were heat
treated under a n inert atmosphere at 350 "C for 3 h.
Temperature-programmed oxidation (TPO) profiles
of the used catalysts were carried out using a siniilar apparatus by introducing a gas flow (30 ml/min)
containing a mixture of 02:He (5:95) ovcr 50 rrig
of the used catalysts and the temperature was increased up to 800 "C at a heating rate of 10 "C/niin.
Temperature-programmed hydrogenation (TPH) of
the used catalysts was carried out using the apparatus
that was used for TPO. The used catalyst (25 mg) was
subjected to heat treatment (10 "C/min up to 800 "C)
in a gas flow (30 ml/min) containing a mixture of
H2:Ar (10:90). Before the T P H experiment, the samples were heat treated under an inert atmosphere at
300 "C for 3 h.
2.4. Evaluation of catalytic performance
Activity measurements were carried out at atmospheric pressure in a fixed-bed continuous-flow reactor made up of a quartz tube of 7 mm i.d. The
reactor was charged with 200 mg of the prepared catalyst. A reactant gas feed consisting of a mixture
of CH4 and CO2 (CH4/C02=50/50 vol.%) was introduced into the reactor, and the activity tests were carried out at different temperatures ranging from 500 to
700 "C in steps of 50 "C that were maintained for 30
min a t each temperature. The loss in catalyst activity
at 700 "C was monitored for up to 5 h on stream. The
gas composition of the reactants and products was
analyzed using a gas chromatograph equipped with a
TCD and a Carbosphere column.
3.1. Structural properties of the catalysts
Figure 1 shows the pore size distributions of the
noble metal catalysts supported on alumina-stabilized
magnesia (Spinel). As can be seen, Rh, R.u, and
Pd catalysts showed a narrow pore size distribution,
whereas Pt and Ir catalysts showed a broad pore size
distribution. Table 1 shows that the Pt and Ir catalysts posses a bigger pore size than do Pd, Rh, and
Ru catalysts.
329
Journal of Natural Gas Chemistry Val. 15 No. 4 2006
0.14
-5
0.4
0.12
=E
0.10
.
-
--.0.08
v
'5
.
.
-5
55-
0.3
bo
&I
0.2
9
1
0.06
0.04
z
0.1
0.02
0
0
4
6
8
10
I2
14
2
4
12
6
8
10
Pore diameter (nm)
Pore diameter (nm)
14
Figure 1. Pore size distribution of the reduced catalysts
Table 1. Structural properties of the reduced and used catalysts
Catalyst
BET surface area
Pore volume
Pore diameter
Sulfur capacity
Metal area
Metal crystallite
(m2.g-')
(cm3 .g- )
(nm)
(wt%)
(m2/d
size (nm)
Reduced
Used
'
Reduced
Used
Reduced
Used
Reduced
Reduced
Reduced
0.238
0.222
0.258
0.333
10.66
4.96
10.02
13.90
295
450
0.67
1.02
4.24
1.53
Pd/Spinel
85.56
179.10
57.68
64.53
Rh/Spinel
159.00
145.95
0.276
0.249
6.95
6.30
1650
3.75
1.28
Ru/Spinel
120.84
56.22
0.253
0.240
7.53
11.66
700
1.59
3.96
Ir/Spinel
95.33
48.79
0.262
0.251
8.95
15.10
1.86
36.60
-
0.142
-
14.33
-
625
-
1.40
Spinel
-
-
Pt/Spinel
The BET measurements of the catalysts showed
a higher specific surface area for the Ru, Rh, and Ir
catalysts compared with the Pt catalyst (Table 1).
The pore volume for all the catalysts was higher than
200 ml/kg. It is noteworthy the specific surface area
of the catalysts after reduction was much higher than
that of the support. The specific surface area of the
This could be atsupport was about 36.6 rn2.g-'.
tributed to the transformation of the pore size dis-
18000
0.008
tAlumina-stabilized
0.007
0.006
.
9
0.003
1
s
0.002
0.00 I
0-
tribution after the impregnation of the support with
metal salts and calcination and reduction of the impregnated catalysts. During the calcination process,
the decomposition of metal salts led to a change in
the pore structure and also in the pore distribution of
thc support. Figure 2a shows the pore size distribution of the fresh support before impregnation, which
is completely different from the pore size distribution
of the catalysts (Figure l a and lb).
1
0
v;.
magnesia (Spinel)
(a)
I6000
I
I4000
I2000
I0000
4000
2000
0
10
10
Pore diameter (nm)
Figure 2. Pore size distribution (a) and XRD pattern (b) of the Spinel
20
30
28/(0
40
)
50
60
70
330
M . Rezaei et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
For t,he catalysts, the pore size shifted to smaller
values and the pore volume to higher values, especially for Rh, R u , and Pd catalysts (Table 1);therefore, these catalysts showed a higher specific surface
area compared with the Ir, Pt, and the fresh support.
The met a1 crystallite sizes, determined from hydrogen sulfide chemisorption analysis, showed a smaller
size than the pore size of the Spinel (Table 1). This
indicates tliat active metals could be introduced into
thr pores of thc Spinel, which would change the pore
structure and also the pore size distribution of the
support. This led to a reduction in the values of the
pore size and an increase in the specific surface area
of the catalysts (Table 1).
n
inn
200
300
500
400
600
700
Temperature (T)
Figure 3. TPR profiles of the noble metal catalysts
R h ( l ) , Ru(2), Ir(3), Pt(4), and Pd(5)
3.2. Temperature-programmed reduction
Figure 3 shows the temperature-programmed reduction (TPR) profiles of the noble metal catalysts.
It, was observed that Rh, Ir, and Ru catalysts showed
the lowest reduction temperature, whereas Pd and P t
catalysts showed the highest interaction with the support, as the maximum temperature of the reduction
pcak shifted to higher temperatures. For the Ru and
Ir catalysts, one major peak was observed at about
160 and 275 “C, respectively, which indicates that the
major fractions of Ru arid Ir have a low reduction
temperature, whereas the other catalysts showed several peaks in their TPR profiles. This showed that
the active metal in these catalysts has to be present
in several species with different types of interaction
with the support and therefore brings about different
reducibility.
3.3. Catalytic performance
Figure 4 shows the CH4 arid COz conversions
over the noble metal catalysts supported on aluminastabilized magnesia (Spinel) at different reaction temperatures. The results obtained showed an increase
in CH4 and COz conversion with increasing reaction
temperature. It was observed that Ru and Rh showed
the highest activity for methane reforming with carbon dioxide. Under these reaction conditions, the following order of activity was observed for different catalysts:
Rh-Ru>Ir>Pt>Pd
For all the catalysts, the COz conversioii was
higher than the CH4 conversion (Figure 4b) due to
the reverse water gas shift reaction.
80
(a)
RuiSpiiiel
RhiSpinel
tIrlSpinel
-mPUSpinel
--c PdiSoinel
.
-e-
(b)
-&-
-
60 -
g
-e-Ru/Spinel
-ARhiSpinel
tIr/Spinel
+PtiSpinel
--c PdiSpinel
S
.-
VI
8
:8:
40-
0
V
20 -
so0
ssn
600
650
Temperature (T)
700
n
l
son
550
l
.
,
I
I
,
,
600
leinperature ( “C )
,
/
/
,
650
Figure 4. CH4 (a) and COz (b) conversion of the different noble metals, GHSV=1.5x104 ml/(h.g,,t)
I
,
,
I
70(
33 1
Journal of Natural Gas Chemistr.y Vol. 15 N o . 4 2006
Figure 5 shows the variation of stabilit,y (Figure
5a) and H2/CO molar ratio (Figure 5b) over different
noble metal catalysts with time on stream.
reaction occurs simultaneously with CO:! reforming of
CH4. The lowest H2/CO molar ratio was observed for
the Pt catalyst, which indicates that the reverse water
gas shift reaction is much favorable on this catalyst.
Table 1 shows that the Rh and Ru catalysts have the
highest sulfur capacity and highest active metal surface area. Pt and P d catalysts have the lowest sulfur
capacity and therefore the lowest metal surface area,
which led to the larger average metal diameter for
these metals. The pore size distributions of the used
catalysts are shown in Figure 6.
-
tRuiSpinel
RhiSpinel
--t IriSpinel
--c PUSpinel
+PdiSpiiiel
-A-
-m-
PdiSpinel
tI'tiSpincl
-5,
'
'-g
E
--c RhiSpinel
tRuiSpinel
IriSpinel
+
0.12
0. I 0
1
0.08
0
L
tRhiSpinel
2
0.06
(b)
Irispinel
tP t f p i n e l
--c RuiSpinel
+PdiSpinel
-A-
0.04
0.02
1 .00 -
n
...
7
8I"
0.95
4
6
8
10
12
14
Pore diameter (nm)
L
Figure 6. Pore size distribution of the used catalysts
0.90
-
o . s 0 " ' " " " ' " " ' " ' " ~ " ' ' ~ ' ' " ~
0
50
100
150
200
250
300
Time on stream (inin)
Figure 5. Stability (a) and Hz/CO molar ratio (b)
of the noble metal catalysts at 700 " C ,
GHSV=1.5x104 ml/(h.g,,t)
The results obtained showed a high stability for
Ru, Rh, and P t catalysts. These catalysts showed
a very high stability without any decrease in CH4
conversion with time on stream. Ir catalyst showed
a slight decrease in CH4 conversion with time on
stream. The lowest stability was observed for the P d
catalyst. This catalyst showed a high degree of deactivation, especially during 2 h of reaction.
The H2/CO molar ratio for all of the catalysts
was less than one because the reverse water gas shift
It was observed that the pore size distributions for
all the catalysts changed, whereas for the Rh catalyst,
the pore size distribution showed almost no change.
Table 1 shows that there was a slight decrease in the
specific surface area of the used Rh catalyst. For
the other catalysts, the surface area showed a marked
change. The loss of specific surface area could be attributed to the sintering of active metals because the
support has a very high thermal stability and it was
calcined a t a temperature that was higher than the
reaction temperature (975 "C).
3.4. Effect of GHSV
The effect of gas hour space velocity (GHSV) on
t~hecatalytic performance of different noble metals
was studied by maintaining the reaction temperature
and feed ratio in the system constant at 700 "C. Table 2 shows that increasing the GHSV leads to a decrease in CH4 and COa conversions and also in CO
and H2 yields.
332
M. Rezaei et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
T a b l e 2. Effect of GHSV on the activity and selectivity of noble m e t a l c a t a l y s t s
Catalyst
CH4 conversion (%)
GHSV (ml/h.g,,t)
Pt/Spinel
Ru/Spinel
Rh/Spinel
Ir/Spinel
Pd/Spinel
COz conversion (%) Hz yield (%) CO yield (%) Hz/CO (molar ratio)
7500
62.95
87.61
50.62
75.28
0.672
12000
62.78
72.68
57.83
67.73
0.854
15000
59.22
69.86
53.9
64.54
0.835
18000
56.36
67.34
50.87
61.85
0.822
21000
52.36
64.60
46.24
58.48
0.791
7500
73.08
75.38
71.93
74.23
0.969
12000
71.76
74.70
70.29
73.23
0.923
15000
70.54
73.98
68.82
72.26
0.952
18000
67.38
74.88
63.63
71.13
0.895
21000
64.50
74.46
59.52
69.48
0.857
7500
69.8
79.88
64.76
74.84
0.865
12000
70.68
76.42
67.81
73.55
0.922
15000
70.28
74.72
68.06
72.50
0.939
18000
67.70
74.48
64.31
71.09
0.905
21000
66.72
73.60
63.28
70.16
0.902
7500
67.80
81.85
60.77
74.83
0.812
12000
68.20
74.48
65.06
71.34
0.912
15000
65.66
73.22
61.88
69.44
0.891
18000
64.20
70.44
61.08
67.32
0.907
21000
59.34
67.84
55.09
63.59
0.866
7500
54.54
76.06
43.78
65.30
0.670
12000
41.96
55.26
35.31
48.61
0.726
15000
36.86
46.06
32.26
41.46
0.778
18000
32.56
39.74
28.97
36.15
0.801
21000
28.52
35.30
25.13
31.91
0.788
Reaction conditions: temperature=700 "C, CH4/C0~=1:1
C02/CH4. The results also showed a decrease in
COZ conversion with increase in the COz/CH4 ratio because COz was available in excess. The Hz/CO
molar ratio was also decreased with increase in the
C02/CH4 ratio, because of the reverse water gas shift
reaction, as the WGSR takes place at a higher order
at a higher ratio of C O ~ / C H ~ .
3.5. Effect of feed ratio
Table 3 shows the effect of feed ratio on the molar ratio of CH4, c o 2 , and H2/CO for different noble
metal catalysts at 700 "c. The results ShOwed that
with imrease in the C02/CH4 ratio, the CH4 Conversion also increased. The RU catalyst showed a
complete conversion of methane a t a ratio of 3 for
Table 3. Effect of feed r a t i o on the catalytic a c t i v i t y of noble m e t a l catalysts
CH4 conversion (%)
COz conversion (%)
Hz/CO (molar ratio)
Catalyst
at different COz/CH4 ratios
1
2
Pt/Spinel
59.22
80.73
Ru/Spinel
70.54
89.38
Rh/Spinel
70.28
90.19
Ir/Spinel
65.66
Pd/Spinel
36.86
3
1
2
3
1
2
3
90.28
69.86
53.67
43.37
0.835
0.718
0.638
73.98
56.36
42.79
0.952
0.768
0.751
95.44
74.72
56.32
42.75
0.938
0.778
0.706
82.26
94.44
69.44
54.25
45.04
0.891
0.724
0.645
43.34
50.32
46.06
32.22
26.04
0.778
0.608
0.567
~
~~~~
100
Reaction conditions: GHSV=l.Sx lo4 ml/(h.g,,t), temperature 700 "C
333
3.6. Temperature-programmed oxidation
Figure 7 shows the temperature-programmed oxidation (TPO) profiles of the used catalysts. It
was observed that Pt and Ir catalysts showed one
peak in the TPO profile a t different temperatures
and with different intensities. For the Pd catalyst,
two peaks were observed in the T P O profile. Ru
and Rh catalysts showed a different profile. They
showed two peaks in the T P O profile, which were
located at lower temperatures compared with the
other catalysts. The results obtained are indicative of the fact that different types of carbonaceous
species are formed on the catalysts. The peaks observed at temperature of about 340-390 "C were related to oxidation of amorphous carbon, which usually
requires lower temperatures for oxidation compared
with whisker carbon (peaks located at temperatures
higher than 600 "C).
I
340
110
.-b
m
tained in T P O analysis were in agreement with the
results of catalytic performance because the catalysts
with higher amount of C, (Ru and Rh) showed higher
activity and stability in the dry reforming reaction.
The coke deposition analysis showed a low percent of
coke deposition over the noble metal catalysts (0.3,
0.5, 0.2, 0.2, and 0.1 wt% for Pt, P d , Ir, Ru, and Rh,
respectively).
3.7. Temperature-programmed hydrogenation
The temperature-programmed hydrogenation
(TPH) profiles of the used catalysts are shown in
Figure 8. The results obtained clearly showed the
existence of two types of carbonaceous species on
the catalysts with different reactivities towards hydrogenation. A first peak is observed at different
temperatures between 345 and 385 "C in different
catalysts. This peak is related t o amorphous carbon located a t the interior of the active metal particles [20]. A second peak is observed at temperatures
above 600 "C, except that of the Pt catalyst, identified as whisker-like filamentous carbon. According
to previously published reports, this species is produced by the adsorbed carbon atoms derived from
methane decomposition and CO dissociation.
F
630
385
"
I
n
0
100
200
300
400
500
600
700
345
800
Temperature ( % )
Figure 7. TPO profiles of the used catalysts
(1) Pt/Spinel; (2) Pd/Spinel; (3)Ir/Spinel; (4)Ru/Spinel;
(5) Rh/Spinel
The T P O results indicated that the whisker carbon was formed in the Pd catalyst and could be responsible for the lower activity and stability of this
catalyst in dry reforming reaction. For the Ru and
Rh catalysts, another peak was observed at about
100 "C, which could be related to C , according t o
Bartholomew [16] and was described by Mirodatos
et al. [17,18] as a superficial carbide. Chen and
Ren [19] proposed that these species may be the reaction intermediates and their reactivity is correlated
with the catalytic activity. This is in agreement with
the assignment to this carbonous species, which is responsible of the CO formation [19]. The results ob-
0
100
200
300
400
500
600
700
800
Temperature ("C)
Figure 8. TPH profiles of the used catalysts
(1) Pt/Spinel; (2) Pd/Spinel; (3)Ir/Spinel; (4) Ru/Spinel;
(5) Rh/Spinel
The results obtained also showed that in Ru catalyst, the amorphous carbon was more reactive because it showed a lower temperature for hydrogenation than did the other catalysts. For the Rh catalyst,
as well as these peaks, one small peak was observed
at a temperature of approximately 85 "C, which could
33.1
hi. Rczaej ct al./ Journal of Natural Gas Ctiernistr.y Vol. 15 No. 4 2OOfi
1 ) ix,lat,t.tl
~
to superficial carbidic carbon that shows
high rcmtivity towards IiydrogcntAoii. The P d catalyst sliowtd oiily a very weak peak at temperature of
al,pr(ixiiiiatc!ly 650 “C, despite higher amount of deposited ca.rl.)on. This indicates that the acciiniulated
car1)oii on this cw.talyst,litis a low reactivity with hy(IrOg(Y1.
4. Coriclusions
‘I’lic iiol)lc riirt,al catkdysts support,ctl on alurninast,al)ilizctl iriwgiiesia (Spinel) were used in rnetharie reforiiiing wit,h carbon rlioxidc. The results obtained
sliowcd a high dcgrte of x t i v i t y arid stability for
the R u , 1111, and P t catalysts. The T P O and T P H
iiiial,yst)s s1iowt:tl different types of deposited carbon
011 tliffcmtiit catalysts. Tlie T P O aiialysis showed that,
t,he Iiiglicr activit,y itnd stability of t,he Rii arid Rh cat;tlyst>coiild h t at,tributed to the forriiatiori of highly
rcxtivc carbon (C,) in comparison to the other catit1yst.s. Tho T P O and TPH analyses showcd that ttic
iiiaiii r c ~ s o i for
i the lower stability of the Pd catalysts
was tJhe forrriatioii of less roactivc carbon in compariso11 t,o t,l1c 0t~llc.rcatalysts.
References
[ l ] I~ostriip-NiclscnJ R., Bak Hanseii J H. J Catal, 1993,
144(1): 58
[2] hfark M F, Maier W F. J Catal, 1996, 164(1):122
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1997, 152(1): 83
[4] Hegarty M E S , O’Coiiiier A M,Ross .I R 1-1. C ‘ d d
Today, 1998, 42(3): 225
[5]Bradford M C .I, Vannicc M A . .I C h t u l , 1999, 183( I ) :
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[6] Portugal U L, Marques C M P, Ara.iijo E C! C, hloralcs
E V, Giott,o M V , Biieiio J hl C. Appl Cntd A , 2000,
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[<I] Fish J D, I-lawii D C. J Sol En(
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[lo] Osaki T, Horiuchi T. Suzuki I<, Mori T. Caful LcJtl.
1995, 35(1 2): 39
[ l l ] Wang S B, Lu G Q. A p p l Catul l3, 1998, 19(S 4): 267
[12] Montoya <J A , Rorncro-Pasciial E, Ginioii C, Do1 A1igc4
P, Moiizon A. Catal Today, 2000, 63( 1): 71
[13] Ashcroft A T, Chcethan A K , Green M 1, €I! V(!~IIIJIII
’I D F. Scie,nce, 1991, 352: 225
[14] Rostrup-Nielscri J It. Stud S 1 7 f Sci Cattrl, 19!&i, 8 1 :
25
[ 151 Rostrup-Nielseii J R. In: Cat,alyi.ic
IIaldor Topsee A/Y, Nyriiollevej 55,
Denmark, 1984
[lti] Bartholomew C H. Ca.tul Ren Sci E7149, 1982, 24: 67
[17] Mirodat,os C, Praliaud H, Priiriei. M . .I
107: 275
[18] Swaaii H M, Kroll V C H, Martin C A, Mirodat.os C’,
Catal Today, 1994, 21: 571
[19] Clieri Y G, R.en *J. Catal Lett, 1994, 29(1-2): 3 9
[20] Pereira E B, Mart,iri G A . Appl Catal A , 1‘394, 115( I ) :
135
ScienceDirect
SCIENCE PRESS
Jourrial of Nwt,iiral Gas C'heiiiistry 15(2006)335-339
Articlc
Effects of the Different Supports on the Activity and Selectivity of
Iron-Cobalt Bimetallic Catalyst for Fischer-Tropsch Synthesis
Abstract: Silica. aluriiiria, arid activated carhoii siipport.etl iron-cobalt catalysts wc:rc prq)ared by iiicipim t , wetness inipregiiat.ion. These cata1yst.s have l)em charactcrizcd by BET, X-ray diffraction (XRD), and
tcrripeI.atiirc-progritrrtrned retluctioii (TPR.). Activity a d selcc.tivit,y of iroii-cobalt supported on different,
carriers for CO hytlrogeriat,ioii were stutlictl riricicr t,he coiiditioiis of 1.5 hlPa. 193 K , 630 1i-l aiid Hz/CO
ratio of 1.6. The rcwilts indicate t,liilt thc activit.y. Ct olc!fiii/(Ct olcfin+Cq paraffin) ratio, and ("5
olefin/(Cs olefiii+C; paraffin) decrc:asc iii t,lic! order of Fo-Co/SiO2, Ft,-Co/hC:l, Fe-Co/A1203 ancl FeCo/AC2. The activity of Fe-Co/SiOz rrached a niaxirriuni. The, results of TPR show that the Fc-Co/SiOn
catalyst is to some extent, diffcrciit. X R D patterns show that t , t r Fc!-Co/SiO2 cai,al,yst,differs significaiitly
froiii the others; it has two diffraction peaks. The active spiric.1 p h a s c is correlat(:d with t.he supports.
Key words: Fischer-Tropsch synthesis; birnct,allic cat,alyst; iron; cobalt; support,; silica; alumina; active
carboii; syngas
~
1. Introduction
It, is wvll kiiowii that the conversioii of it niixt~iirc
of CO and H2 (syngas) t o a range of hytlrocarhons
u i n t,lic Fisclier-Tropscli syiit1ic)sis (FTS) lias I)ecii
considcwd its an c~iviror~tr~entally
friendly prot
GasoIjii(~.dicwl fi-a.ct,ioris,chcmiical intwinrdiates, aiid
heavy wax ran he obtaitictl using Fisclier-Tropsch syntlicsis. Jii t,he coniirig fiitiire. it is cxpwtotl that
tlierc will lw iiicreasirig tlciiiiiiid for fuels tlcrivotl fro111
coal atitl natural gas a i d tlie Fischcr-Tropscli (F-T)
will I)<, at1 dtctriiat,ive roiitc t,o obtain liqiiitl fucls
and clicniicitls (ill part,iculiir litirar I-nlkmcs) [I21.
Thorc have been coiisic1er;hlt. vffort,s toviLrtl c l ~ \ ~ ~ l oping catalysts for tlic cotivcrsicni o f syiigas to liquids [s 121. Iron atid coGalt, arc tlie triiditio~iidcatitlyst,s iistd iii Fischer-Tropscli sytit,liesis. Ft.-I)ased catalysts l ~ u t,o
l tlie foriiintion of iiiore olofinic prod~
~
li
Corrcspoiitling a u t h o r . Tc.1: +86-021-61252192; Fax: +8(i-O21-li 125 1002: E - I I I ; ~wyiiig~cr(:ciist
:
.odii.cn
l.'uiidation item: Doctorid Foiuiclatioii(NO. 20050251OOfi)
336
Xiangdorig &fa et al./ Journal of Natural Gas Chclnistr,y Vol. 15 No. 4 2006
but t,lierc are no reports about the effects of different
supports on the iron-cobalt supported catalysts under similar reaction conditions and using the siniilar
rnctliod of preparation.
In this study, different supports were impregnated
with Fe-Co cosolution. The objective of this study is
to show in a simple manner the effects of support
on Fisdier-Tropsch synthesis activity and selectivity
of different supported iron-cobalt binictallic catalysts
under siniilar rextion conditions iirid using a similar
niet,liod for catalyst preparation.
D/Max2550 VB/PC with Cii K a radiation at, 40 kV
and 100 mA a t a scan rate of 0.02"/rnin (20).
Teinperature-prograiiinied reduct>ion (TPR) experiments were carried out using a Aut,oCheniI1292(1
Micromeritics instrument. About 0.16 g of catalyst,
precursor was loaded into the quartz rcactor. A niixture of 10%HZ/Ar was allowcd t,o flow over tlic freshly
prepared catalyst,. The tcrnpcraturc was incre;tstvl
from 312 K to 1000 K at a rate (If 5 K/niiii. Thtl
effluent gas was monitored by TCD. HzO etc:. foriiiccl
during T P R wcrc removed by a cold trilp.
2. Experimental
2.3. Catalyst test
2.1. Materials
Two different types of active carbon were supplied by Shanghai act,ivated carbon factory, y-alumina
was obtained from Shanghai Supeng Clicniical Material. Ltd, and silica gels were obtained from Qingdao
Haiyaiig Cheiriical Co. Ltd. Coconut shell activated
carbon is designated AC1, and active carbon from
coal is designated AC2. Active carbon was added
to 1 niol/L KOH solution at 353-363 K, stirred for
1 1.1, filtered, and washed with deionised water till
pH=7. It was stirred in 1 mol/L H N 0 3 solution
for 1 h a t 353-363 K, then filtered and washed till
pH=7 with deionised water. y-aluniina was stirred
in 10 g NH4NO:1 per 100 ml aqueous solution for 2 h
at, 353-363 K , washed till pH=7 with deionised water. They wcre crushed and sieved to obtain particles
that were between 20 and 40 mesh for use as catalyst support, which were calcined at 673 K in an oven
for 6 h. Iron (111) and cobalt (11) nitrate cosolution
WBH impregnated incipiently onto the support pellets.
ARer impregnation, the wet Fe-Co/ACl, Fe-Co/AC2,
Fe-Co/y-alumina, and Fe-Co/silica gel catalysts were
maintained at room temperature for more than 24 11 to
allow water to evaporate and then the catalysts were
dried at, atmospheric pressure at 353 K for 8 h. Thereafter, Fe-Co/y-alumina and Fc-Co/silica gel catalysts
were calcined at. 673 K in air for 6 h, and Fe-Co/AC
catalysts were calcined a t 673 K under nitrogen for
6 h.
Catalyst tests were carried out in a fixed h c d
(12x650 mm) rcactor. The catalyst ( 5 inl) wiis rvduced in a flow of hydrogen a t 0.5 MPa, 673 K , ;ind
1000 h-'for 6 h, then cooled to teniperatmcs 1)rlow
373 K before switching to H2/CO rriixtiire at 1.5MPa.
493 K , 630 h k l , anti H2/CO ratio of 1.6. The <:iit,itlyt,ic
results were recorded a t the steady state &tr s t a h lization of more than 12 11. The CO, Hz, ( 2 0 2 , aiitl
CH4 products were analyzed using TCD after sc'I)>lrii.tion of a carbosieve picked column. The gas hydrocarbons were separated by the capillary coliirnn (stationary liquid: A1203) and then detectod usiiig FID.
Liquid products arid wax were collected in a, cold t , i q
and a hot, trap, respectively.
The CO conversion (niol%), the distri1)utioii of
hydrocarbons ( w t x ) , and t,lie wat,er gas shift (WCS)
extent was calculated as follows:
CO conversion (niol%)=
moles of transformed CO
x 100
niolcs of initial CO
Distribution of hydrocarbons to C, (wt%)=
lKlaSS
(C,)
total mass of hydrocarbons
The WGS extent
=
x 100
mass (COZ)
mass ( ( 2 0 2 HzO)
+
2.2. Characterization techniques
The surface area, average pore volurne, and pore
size distribution of the samples were nieasured using
ASAP 2010 Micromeritics instrumelit.
XRD data were collected using Rigaku
The activity of the catalyst in CO hydrogenation is a function of support, nictal loading, and preparation.
The order of decr(v~sirig CO hydrogenation activity at 0.1 MPa and
498 K for the catalysts containing 3 wt%j c ~ ) l ) d t
337
Jouriial of Natural Gas Chemistr,y Vol. 15 No. 4 2006
is Co/TiO2>Co/Si02>Co/A120:3 >Co/C>Co/MgO.
This was observed by Rciiel ef nl. [20]. Saib et a l .
studied the effect of pore dianictcr of support on
silica-supported cobalt Fisclier-Tropscli cataly
1111der tlie following conditioiis: H2/CO=2, P=1.5 LlPii,
T=493 K , and the C5+ aiitl riietliaiie selectivity passed
through a rnaxiniuni and minimum a t tlic 10 nrri supported catalyst, respectively [21].
Tlic catalytic rcsults obtaiiied are suiiiiiiarized in
Table 1. It. was shown that the act,ivity of t,hc catalyst
was sigiiificaritly dependent on the natiirc of the sup-
port,. The order of dccrea.sing a,ctivit,y a t an approxiiiiatcly 9wt%Fe9wtl'%~Co
loading level for this cata.lyst
on difftwrit, supports is as follows: catalyst, 1 > catalyst, 2>catalyst 3>cat,alyst 4. The catalytic results
of t,liesc catalysts arc shown in Tahlc 1. Tlie textural
structure of t,hc various fresh catalysts is shown in Tablc 2, and it, is c1ca.r that tlic na.tiirc of tho support
striict,urc has u proiiouiiccd offect, oii the x t i v i t y of
t,kic catalyst. Tlie activity of cat,alyst 1 reached a rnaxiiiiurn aiid may result from the average pore diameter
of 9.6 iim being t,he optirnum valuc.
Table 1. Catalytic result of the four catalysts
Catalyst
1
2
3
4
24.6
CO conversion (niol%)
50.3
38.8
28.7
C:O2 selectivity (inol%)
0.6
1.5
0.9
1.2
WGS extent,
0.012
0.015
0.027
0.014
Distribution of hydrocarbon (wt?%)
C1
7.4
17.0
6.1
11.4
c,
0.3
0.2
4.2
0.3
c2
1.5
2.0
2.0
2.4
0.2
0. I
0.7
0.1
2.0
3.4
2.4
1.9
1.3
1.9
1.5
1.6
o/(O+P)
c,=
0.6
0.0
0.6
0.6
1.6
I .7
1.1
0.6
CZ
1.3
1.7
1.5
1.0
O/(O+P)"
0.6
0.5
0.5
0.4
c,
0.9
1.1
0.6
0.5
c5
0.7
1.0
0.7
0.8
o/(0+P)
c,
c3
O/(O+P)"
0.6
0.5
0.5
0.4
C6-t
83.1
70.0
80.4
79.8
"Olefin/(olefin+paraffin) ratio.
Reaction conditions: 1.5 MPa, 393 K. 630 h - ' , H2/CO rnt,io o f 1.6
Table 2.
Textural property of the different fresh catalysts
Catalyht
Property of catalyst
1
2
3
4
Surface area ( m 2 / g )
198 2
270.2
208.4
385.3
Average pore volume (ml/g)
0.25
0.1 G
0.56
0.20
Average pore diameter (nm)
9.6
24.0
19.4
25.7
With regard to the hydrocarbon distribution,
from Table 1 it is clear that, C,i olefin/(C,i olefin+C,
paraffin) ratio (from 0.6: to 0.4) decrease in the followiiig order: catalyst 1, cat,alyst, 2 , cat,alyst 3, catalyst, 4, arid C5+ hytlrocwrbon distribution of catalyst 1
reached 83.1%.
T P R can trace the reductioii of t,hc oxide phase
providc infonnat,ion about metal-support and
riiet,al-nict,al interaction. The nature of the interaction
Iietwceii oxitlc a i d slipport affects the reduction. The
T P R profiles of tht: four catalysts a.rc shown in Figiirc 1. The four catalysts have t h e resolved reduct,ioii peaks arid show some sirriilarities t,o the reported
TPR spectra of Fe-co alloys and Fc-CO supported 011
iiiitl
Xiangdong Ma et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
338
Ti02[13,16]. The third maximum temperature (T3m)
increased from 753 K to 767 K in the order of catalystl, catalyst 2, catalyst 3, and catalyst 4. This
indicates that the oxide-support interaction of catalyst 1 was weaker than that of the other catalysts, and
also shows that its metal oxide dispersion and the size
of the particles are more effective for Fischer-Tropsch
synthesis. In contrast, the first maximum temperature (TI,) and second maximum temperature (T2,)
of the catalyst 1 were higher. This is attributed t o its
lower pore diameter. The lower pore diameter makes
the diffusion rate of the water in the pores formed by
reduction more difficult and thus results in an inhibition of the reduction process [22].
I
753
The XRD spectra of the four catalysts after the
catalytic test are shown in Figure 3. The spectra
showed that all catalysts had formed Fe-Co alloy of
spinel phase. The Fe-Co alloy of spinel phase was
formed on reduction of the Fe-Co systems, and
this was observed previously in an in situ XRD
study by Duvenhage [13]. Figure 3 clearly shows
that the four catalysts of different supports had
different diffraction peaks. The catalyst 1 showed
two diffraction peaks; catalyst 2 also showed two
diffraction peaks of Fe-Co alloy of spinel phase but
its diffraction peaks were smaller and sharper than
those of catalyst 1. Catalyst 3 and catalyst 4 showed
one diffraction peak of Fe-Co alloy of spinel phase,
respectively. These results indicate that the active
phase formation seems to be correlated with the supports.
~
~~
Fe-Co alloy of spinel phase
300
400
500
600
700
800
900
1000
Temperature (K)
Figure 1. TPR profile of the fresh catalysts
(1) Catalyst 1, (2) Catalyst 2, (3) Catalyst 3, (4) Catalyst 4
The XRD spectra of the four fresh catalysts (Figure 2) showed that Fez03 and c0304 are present
and this is similar to the XRD result reported
previously [131.
I
10
,
,
,
20
,
1
30
,
,
,
,
1
,
40
,
/
,
I
50
1
/
1
1
1
I
I
60
~
I
70
28/(0 )
Figure 3. XRD spectra of the catalysts after test
4. Conclusions
Activity and hydrocarbon selectivity of ironcobalt supported on silica, alumina, and active carbon carriers for CO hydrogenation were studied at
1.5 MPa, 493 K , 630 h-l, and H2/CO rat,io of 1.6.
The results indicate that for these catalysts thc activities in the order of catalyst l>catalyst 2>catalyst 3>
catalyst 4.
Hydrocarbon distribution, olef in/ (olefin psraf f in)
ratio and C5+, especially Cq olefin/(C4 olefin+C4
paraffin) ratio and Cg olefin/(Cg olefin+Ca paraffin)
ratio vary with the support. For catalysts containing 9wt%Fe9wt%Co, the order of CO hydrogenation
+
10
20
30
40
50
60
70
80
281(" )
Figure 2. XRD spectra of the fresh catalysts
activity is also catalyst 1, catalyst 2 , catalyst, 3, and
catalyst 4.
T P R results show that the reduction temperature
of the catalysts also varies with the support, which is
due to an interaction between Fe-Co species and the
support. The third maximum temperature (T3m)of
Fe-Co/SiOz catalyst shifts toward lower temperature
and the first maximum temperature (Tlnl)and second maximum temperature (Tz,,) shift toward higher
temperature, which might be ascribed to the effect of
the supports.
The XRD peaks of catalysts after catalytic
test show that the Fe-Co/SiOz catalyst differs significantly from the others. Its diffraction peak of FeCo alloy of spinel phase is larger and sharper; therefore, its activity is higher than the other three catalysts and its C5+ reached a maximum in the four catalysts. The active spinel phase is closely correlated
with the supports.
Acknowledgements
We gratefully acknowledge the financial support
provided by the Doctoral Foundation and Shanghai
Yankuang Energy R&D Co. Ltd.
References
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ScienceDirect
Joiirriitl of Natural Gas Chemistry 15(2006)340 347
SCIENCE PRESS
Article
Methane to Liquid Hydrocarbons over Tungsten-ZSM-5
and Tungsten Loaded Cu/ZSM-5 Catalysts
Didi Dwi Anggoro'*,
Nor Aishali Saidina Arnin2
1. Department of Chernical Erky.irieering, University of Diponegoro Tcmbalang, Semarang 502.39, 1ndo.riesm;
2. F a c d t y of Chemical trnd Natural Resouices Engineering, lJn.zversiti Teknologi Malaysia,
Johor Bahru 81 130, Malaysiu
[Manuscript received April 17, 2006; revised July 17, 20061
Abstract:
Metal containiiig ZSM-5 can produce higher hydrocarbons in metallaneoxidation. Many
researchers have studied the applicability of HZSM-5 and modify ZSM-5 for ruethane coriversion to liquid
hydrocarbons, but their research results still lead to low conversion, low selectivity and low heat, resistancc.
The modified I-IZSM-5, by loading with tungsten (W), could eriharice its heat resistant performance, arid
the high reaction temperature (800 "C) did not lead to a loss of tlie W coniporicnt by sublirna,tion. The
loading of HZSM-5 with tungsten and copper (Cu) resulted in an iiicrement in the methane coriversion as
well as COz and Cg+ selectivities. In contrast, CO, Ca-3 and H z 0 selectivities wcre reduced. The process
of converting methane to liquid hydrocarbons (C5+) was deperideiit on the metal surface area and the
acidity of the zeolite. High methane conversion and Cg+ selectivity, arid low HzO selectivit,y arc obtained
over W /3.OCu/HZSM .
K e y words: methane; liquid hydrocarbons; HZSM-5; W-ZSM-5; W-Cu/ZSM-5
1. Introduction
Tii gcIiera1, there are two routes for converting
Irl(.thitlie to gasoline: indirectly or/and directly. The
iiitiircct route is a two-step process whereby natural
gas is first converted into synthesis gas (a mixture of
1.12 aiid CO), and then int,o hydrocarbons of the gasoline range. The direct, route is a one step process in
which the natural gas reacts wit,h oxygen (or another
oxidizing species) to give the desired product directly.
The use of the HZSM-5 zeolite as a support of the
metal oxide phase is very interesting due to three reasoils: it,s thermal stability, its high surface area that
eiiaMes high rnet,al oxide loading and tlie presence of
acid sites that, could lead to the formation of certain
act,ive rnr:tal oxide species [I].
Ernst and Weitkamp [2] reported in a paper on
the conversion of methane over zeolite-based ciitalysts
that the presence of strong acid sites in the z e o h
catalyst is detriment,al for the select,ivt: oxidation of
rnethane to higher hydrocarbons; otherwise oxidized
products of CO, (CO, CO2) will predoniinatc. Wheii
the acidity is reduced by exchanging the zeolite with
alkali m e h l cations: tlie se1ectivit)y to higher hydrocarbons is slightly enhanced. Hari et nl. [3] derrioustrated the sucxessful production of higher hydrocnrboils from rnethaiie oxidation using a ZSM-5 zeolitc,
catalyst containing metal oxides. T h e metal oxides
with sufficiently high dehydrogerition and low olc!fiii
oxidation activities reduce the acidity of the ZSM-5.
As a result, thc metal containing ZSM-5 can product:
higher hydrocarbons in methane oxidation.
De Lucas et a1. [1,4,5] discovered that the illtroductiori of Cn(I1) ions by an ion-exchangc method
* Corresponding author. Tel: 6224-7460058; Fax: 6224-7460055; E-mail: [email protected]
Supported by Ministry of Science, Technology and Environnicnt, Malaysia.
Jonrnal of Natural Gas Chrmistr:~Vol. 15 N o . 4 2006
could rcniarkably iiicrea,se tlie activity of Mo/HZSM-5
for riietliaric aromatizatioii and iiiiprove it,s stabilit,y
to sonie extent. Mo species is tlie most act,ive component for nietharie iiori-oxidative aromatizatioii so far.
but its activity and stability riccd to be irriprovctl.
Xiong fd al. [6,7] studied the incorporation of rnetals Zii, Mn, La, arid Zr into tlic WIHZShl-5 cat,alyst. Under react,iori condit,ion of 0.1 MPa, 1073 K ,
GHSV of feed gas CH4 arid 10% Argon at 960 lip'.
tlie conversion of methane reached 18% 23% in the
first two hours of reaction! and tlie <:orresponding selectivity t,o benzenc: naphthalene, cthylonc and coke
was 48% 56%, IS%),5% and 2291, respectively. Ding
et u1. [8] reported the non-oxidative rnetha,ne reactioii
over WIHZSM-5 to produce C2 C12 hydrocarbons.
Under the condition of 700 "C, flow ratc: CH4 and
Argon a t 12.5 cni3/min, the C2 C12 selectivity was
70%-80%. However. the rnetliaiic conversion was low.
just bctwceri 2% arid 3%. On tho 1)asisof the chemical
siniilarities between k1oO:j and WQ3, it seems reasoriable to cxpcct a. parallelism in their catalytic properties. Mo species is the most, act,ivc coiriponents for
methane non-oxidat,ive aroiriatizatiori so far, but, its
activity arid stJa,bilitjyneed t o be improved. Recently,
we have fourid that the introdiict,ioii of Cu(1I) ions
by an ion-exchange method can remarka.bly increase
the activity o f MoIHZSAl-5 for methane arornatiz;ttion ant1 can improve itjs stability to some extent.
Cu loaded ZShl-5 catalysts ,tin acidic ion exchange nicthod have been iderit,ifietl to be potential cat,alyst,sfor the conversion of nit:t,liane to liquid
fuels [Y]. However, infrared study of metal loaded
ZSM-5 cat,alysts indicated that tlie catalysts are not
resistant to high temperatures. Previous st,udies have
indicated that. metal loaded ZSM-5 did not exhibit, the
vibration band at 3610 ciri-l and 3660 c i ~ i - ~exccpt
.
for the ZSM-5 which showed a weilk vi1)r;~t~iori
band at
3666 c n - l . The result suggested that, fraiiiework arid
non-framework aluminurn were eit,lier extracted to the
acidic solution or changed to silanol defect, form when
calcincd at, 800 "C and made tlie catalysts inactive
[lo]. In our previous studies [11,12] it, was indicated
that the Cii loaded W/ZSnil-5 catdyst was thernially
stable a t the reaction temperature (700-800 "C).
In this study ZSM-5 was niodified wit,h tungsten and copper and the catalyst pcrformancc was
tested for the oxidation of 1liethit1i(>t,o liquid hydrocarbons. The catalysts were characterized by XRD,
T P R . TPD and N2 Adsorption iiieasurerrients. Wc:
would like t,o irivestigate the resistivity of tungsten
modified HZSM-5 t o high teniptiratures and its cat,-
341
itlytic act,ivity for the coiiversion of rnet,hane.
2. Experimental
ZShl-5 zeolite with a SiO2/A120:3 mole ratio of
30 was supplietl by Zeolyst Intcrriat,iorial Co, Ltd,
Nrt,lierlmds. The siirface area of tlie zeolite is
400 iii2/g. The W (3%) weight)-HZSM-5 catalyst
was prcparcd 11y impregnating a certain amount
of t.lie HZSM-5 xeolit,e carrier with an annrioriiiini
tiirigstat,c 1iydr;itc solution [6,7]. The ariinioriiuiri
t,uiig'stat,r hydrate solution was prepared by dissolving (NH4)2WO4 in deionized water and adding a small
aniount of H2SO4 to regulate the pH value of the solutioii to 2 3. Thtl sample (10 rril of solution pes' gram
zcolite) was dried in an ovcm a t 120 "C for 2 h and
then calcined at 500 "C for 4 11.
The W loaded Cu/HZSM-5 was prepared by first
impregnating a certain iLtiiou1it of the HZShl-5 zeolite
carrier with a calculated miourit of copper nitrate in
q i i c ~ ) u ssolutions, followcd by drying at 120 "C for
2 h and calcining at 400 "C for 4 h, and siibseqiieritly
iiiipregiiating with a calculwted a,iriouiit of II2S04
;tcitlified (NH4)2WO4 aqueous solution (pH=2 3 ) .
Fiiidly, tlic sample WAS dried at, 120 "C for 2 h and
calcined a t 500 "C in air for 5 11.
2.2. Characterization and testing of catalysts
X-ray diffractioii (XRD), H2-tcniperat,ure progrninnied reduchoii (H2-TPR), NHs-teniperaturc
prograniined desorptioii (NH3-TPD), N2 adsorption
and FT-IR were utilized for the characterization of
t,hc catalysts. XRD arid FT-IR were employed to detoriiiine the zeolite structure. NH3-TPD provided the
acidity of the catalyst samples. H2-TPR, data were
pertiiieiit t o the zeolit,e morphology.
Thc. performance of tlie cat,alysts was tested for
nict,lianr conversion to liquid hydrocarbons (LHC)
'via a single step reaction in a fixed-bed micro reactor. Methane with 99.9% purity was reacted a t
atmospheric prcwiire arid various temperataims arid
oxygen concentration. A n on-line gas chromatography wit,li R TCD arid a Porapak-N column was utilized
to malyze the gas. Thc liquid product was analyzed
iisiiig the G C FID arid a.ii HP-1 column.
3.1. Characterization of catalysts
342
Didi Dwi Anggoro et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
X-ray diffraction (XRD) and nitrogen adsorption (NA) were employed to determine the morphology of the catalysts. The XRD diffractograms of
HZSM-5, W/HZSM-5 and W loaded Cu/HZSM-5 catalysts with different Cu loadings calcined at 550 "C are
shown in Figure 1. The peaks at 20=41° indicated
tungsten oxides [13]whilst copper oxides are indicated
a t 2H=34O. The intensities of these peaks increased
with increasing copper loading.
cuo
spectively. These values are lower compared to the
values of other samples. All the metal and hi-metal
ZSM-5 zeolite catalysts have surface and micropore
areas smaller than the parent zeolite. The reduction
in surface area of the metal-loaded HZSM-5 indicates
a strong interaction between the surface of the zeolite
and the copper and tungsten species, which enables
a good dispersion of the metals on the surface [ 5 ] .
Among the samples, the BET surface area and the
micropore area of W/3.0Cu/HZSM-5 are the lowest,
being 236 m2/g and 213 m2/g, respectively.
Table 1. Crystallinity and surface area of the
catalysts
wo,
Catalyst
10
20
30
40
50
60
281(O )
Figure 1. XRD pattern of different HZSM-5 catalysts
(1) HZSM-5, (2) W/HZSM-5, ( 3 ) W/0.5Cu/HZSM-5,
(4) W/l.OCu/HZSM-5, (5) W/1.5Cu/HZSM-5,
(6) W/2.0Cu/HZSM-5, (7) W/3.0Cu/HZSM-5
The crystallinity values calculated from the XRD
diffractograms and the areas of the samples from
NA analysis are tabulated in Table 1. The crystallinities of W/l.OCu/HZSM-5, W/2.0Cu/HZSM-5
and W/3.0Cu/HZSM-5 are 89%, 88% and 69%, re-
Crystallinity
(%I
BET surface
HZSM-5
100
area (m2/g)
403
W/HZ
W/0.5Cu/HZ
W / 1.OCu/HZ
W/1.5Cu/HZ
W/2.0Cu/HZ
W/3.0Cu/HZ
100
94
89
101
88
69
280
266
286
267
285
236
Micropore
area (m2/g)
373
257
243
261
244
260
213
The results in Table 2 pertain to the total volume,
micropore volume, average pore diameter and acidity
of the catalysts. Tungsten and copper species easily
entered or partially blocked the channels of the ZSM-5
zeolite pores and thus, reduced the volume of the catalysts. The average pore diameters of the metal-loaded
HZSM-5 zeolites are larger than the parent zeolite, as
revealed by the results in Table 2. The average pore
diameter of the W/3.0Cu/HZSM-5 is the largest, and
as indicated in Table 2, the percentage of micropore
volume in the catalyst surface has been reduced to
50%.
Table 2. Total volume, pore distribution and acidity of the catalysts
Catalyst
HZSM-5
W/HZ
W/0.5Cu/HZ
W/l.OCu/HZ
W/1.,5Cu/HZ
W/2.QCu/HZ
W/3.0Cu/HZ
Total volume
Micropore volume
Ratio of micropore
Average pore diameter
Acidity
(cm3/d
0.245
0.187
0.176
0.175
0.179
0.191
0.176
(cm3Id
0.149
0.106
0.101
0.110
0.101
0.109
0.088
to total volume (%)
61
(A)
24.3
26.8
26.5
(mol /kd
0.87
0.81
0.91
24.4
26.9
1.01
1.01
26.6
29.8
0.98
1.19
The ammonia-TPD spectra of the catalyst providetl useful information about the intensity and the
concentration of the acid sites on the catalyst sur-
57
57
63
56
57
50
face, as tabulated in Table 2. The concentration of
the surface acid sites (acidity) of the metal-loaded
HZSM-5 is higher than that of the ZSM-5 zeolite.
343
This could probably be attributed to the average
pore diameters of the metal-loaded HZSM-5 which
are larger than the pore diameters of the HZSM-5
zeolite itself, as shown in Figure 2 . This is similar
with the results of Koval e t al. (1996). They reported that the acidity decreased with the decreasing
micropore volume of HZSM-5 [14]. As revealed by
the result in Figure 2 , the strongest acidity shown by
the W/3.0Cu/HZSM-5 zeolite coincides with the fact
that it has tlie largest pore diameter (29.8 A). However, the acidity of the W/l.OCu/HZSM-5 is higher
(1.01 mol/kg) than HZSM-5 (0.87 mol/kg), although
the difference in pore size between W / 1 .OCu/HZSM-5
and HZSM-5 is not large.
1.4
medium strength Briinsted acid sites. This is similar
with the results found by previous researchers [6-7],
where the incorporation of Mg2+ and Zn2+ into the
W/HZSM-5 host catalyst resulted in the elimination
of strong surface Bronsted acid sites and the generation of new medium-strength acid sites.
I
430 'C
250 'C
I
t
w13.ocu
-
OIJ 1.0
L
Y
g
0.8
.0
V
.- 0.6
4
0.4
0.2
n
"."n
24.3
24.4
26.5
26.6
Pore size
26.8
26.9
29.8
(A)
Figure 2. Effect of pore size on acidity of catalysts
200
300
400
SO0
TIT
Figure 3 depicts the ammonia-TPD spectra of
the HZSM-5, W/HZSM-5 and W loaded Cu/HZSM-5
with different amounts of Cu loading. For the
HZSM-5 and W/HZSM-5, the ammonia-TPD peaks
appeared at about 250 "C and about 430 "C, which
may be ascribed to the desorption of two kinds of ammonia species adsorbed on weak acid (mostly Lewis
acid) sites and strong acid (mostly Brijnsted acid)
sites, respectively [15].
The addition of 0.5% Cu to HZSM-5 led to a
reduction of the intensity of the high temperahre
(- 430 "C) peak and a small downshift of its position, as revealed in Figure 3 . When the amount of
Cu loading was increased to l.O%, the high temperature peak disappeared, indicating that most of the
surface Brijnsted acid sites had vanished. However,
as the copper loading was further increased to 1.5%
and 3.0%, the ammonia-TPD peaks appeared again
at
430 "C. One interesting fcature of Figure 3 is
the ammonia-TPD spectra of the W/2.0Cu/HZSM-5
that indicated a peak at 400 "C (Figure 3 (6)). This
small peak may be attributed to the emergence of
-
Figure 3. Ammonia-TPD
spectra
of
different
HZSM-5 catalysts
(1) HZSM-5, (2) W/HZSM-5, (3) W/O.SCu/HZSM-5,
(4) W/ 1.OCu/HZSM-5, (5) W/1.5Cu/HZSM-5,
(6) W/2.0Cu/HZSM-5, (7) W/3.0Cu/HZSM-5
Thc niigrat,ion of the W arid Cu species was
indirectly studied by infrared (IR) spectroscopy.
W/HZSM-5 and W/Cu/HZSM-5 samples showed
in the OH stretching vibration three IR bands at
~ 3 6 1 cm-'
0
, due to bridge Si-OH (Al) acidic groups,
at ~ 3 6 6 0cm-' due to non-framework A1 sites, or
octahedral, and at ~ 3 7 4 0cm-' which is attributed
to terminal Si-OH non-acidic groups [lo]. The vibrations for the (OH) region of the IR spectra of
tlie W/HZSM-5 and W/3.0Cu/HZSM-5 zeolite catalysts are shown in Figure 4, where all the fresh
samples have hands at ~ 3 6 1 0cm-l, ~ 3 6 6 0cm-',
and z 3740crK1. The spectra indicated that all the
samplcs have aluminum framework, silanol, and aluminum non-frarnework groups. In addition, the IR
spectra demonstrated that the intensity of the band
a t about 3610 cm-' of fresh W/HZSM-5 is stronger
344
Didi Dwi Anggoro ct al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
tliaii frcsli W/3.0Cu/HZSM-5. This is probably due
to the W or W arid Cu species that have migrated in
thc zeolite frarncwork and occupied the H+ position.
It rwulted in a large tlccrernent of the 3610 crri-l IR
band intensity, as shown in Figure 4 (3).
I
4000
I
3740 3660 3610
1800
3600
3400
tion, W6++c-+W5+ and W5++(>-+W1+ [17]. Tlic
reducibility of this type of catalyst dccr cased i t s
thc strength of the1 interaction between tho iiictnl
oxidc species and the surfacc of the support increased. The existence of a single reduction pwk
at 550 "C for W/IIZSM-5, W/0.5Cu/HZSM-5 i ~ ~ i d
W/1 .OCu/HZSM-5 sariiples iii?iy be duc to the singlc
electron reduction of the W6f species tlerivetl froin
the (W06)"- precursor with octahcdral coordination.
Wo++e-+W5+ [6,7]. This pvak (at 550 "(2) d i s a p
pearcd if thc ainoimt of CII loading on the HZSM-5
was more than 1.0 wt%. The observcd hydrogen-TPR
peak at 370 "C could bc due to the reduction of Cu"+
species and the intensity of the peak hccaiiie stroiipt>r
a5 the copper loading increascd.
3200
Wavenumbcr (crn-l )
Figure 4. IR spectra in the OH region of fresh and
used catalysts
h c 4 1 (1) W/IIZSM-5 and ( 3 ) W/3 OCu/HZSM-5, used (2)
W/HZSM-5 and (4)W/3 0Cii/HZSiVI-5 for reaction at 800 "C
It2 is clear that the roles of W on W/HZSM-5
and W i-~ndCu on W/Cu/HZSM-5 is to reduce
t1hc arnoiiiit of Brorist,ed acid sites. The amount,
of Broiisted acid sites ( N B A C can
) bc calculated by
hridgiiig tlic! OH groups (at 3609 crri-' band) according to [10.16]. The NBAC for the HZSM-5,
W / H ZSM-5 and W /3.OCn/HZSM-5 iire 2.9 / m o l / g ,
0.5 pnoI/g arid 0.3 p i o l / g , respcctivcly.
Further react,ion with inetharie arid oxygen at,
800 "C for five lioiirs resulted in the disappcari i I l ( * ( ! of the 1)illid ilt about, 3610 cr1i-l for both of
the W/HZSM-5 and W/Cu/HZSM-5 samples (Figures 4 (2) and 4 (4)). This is probddy due t o the
ext,ractiori of aluininiiin in the zeolitic framework into
the lion-framework or diic to the deposition of carhniicc!oiis residues. Thc deposition of coke led to
cil,t,;tlysttleactivatiori after five hours of reaction.
The TPR profiles of W/HZSM-5 and W loaded
Cii/HZSM-5 catalyst,s are dcpictcd in Figure 5 . As
observed, all the curves contain several peaks in the
tcrnporature range of 200-900 "C. The T P R patteriis of all catalysts exhibited two peaks, with the
iiiaxiiniiin at 700 "C and 780 "C. These peaks rnay
I)c ascribcd to t,he two subscquent, steps of singleeloctron reduction of the W6+ species derived frorii
the ( WOd)2- precursor with tetrahedral coordina-
200
300
400
500
600
800
700
900
TIT'
Figure 5. Hydrogen-TPR
spectra
HZSM-5 catalysts
of
different
(1) W/HZSM-5, (2) W/0.5Cu/HZSM-5, ( 3 ) W/l.0Cu/llZShl-T,.
(4)W/1.5Cu/HZSA/I-5, (5) W/2.0Cu/HZSM-5,
(6) W/3.0Cu/HZSR.I-5
In Table 3 thc quantitativr. rcbults of thc TPR is
summarized. The TPR software automatically calculated the metal surface area, dispersiori of metal,
and mean particle diameter h s i n g on W md Cu foi
the bimetallic W/Cu/I-IZSM-5 catalysts in tlic rbxpcrimental part. The tungsten content for all ('tltiilyhth
is constant. The copper content kind percent of copper dispersed for W/Cu/HZSM-5 catalysts iiicrcywd
with increasing copper concentratiori. Such srrlidl particles (0.03-0.05 nin or 3-5 A), particularly tIir twigstcn particles, shouId be localized t o inside the Leolitc
pores [IS],wliero the pore diameter of ZSM-5 is h x i t
5.6 A.
345
Table 3. Metal state in the reduced catalysts
klotal content
Catalyst
(iL"'Wd
cu
W/HZ
W
163
W/O.SCu/IIZ
Metal surface
Dispersion of
area (rn2/g)
r u e t d (%)
W
55
CU
0.04
100
4.00
0.07
236
3.11
315
4.18
472
4.77
CU
-
W
2.66
163
79
6.02
W/l.OCu/HZ
163
157
W/1.5Cu/HZ
163
W/2.0CU/HZ
163
W/3.0Cu/HZ
163
The percentage of metal dispersion in Table 3 reveals that Cu is 100% dispersed for the
W/3.0Cu/HZSM-5 sample. Figure 6 shows that
the BET surface area of all the samples decreased
with increasing dispersion of Cu except for the
W/2.0Cu/HZSM-5. This is probably due to some
tungsten or copper species on W/2.0Cu/HZSM-5 that,
have formed timgsten oxide or copper oxide on the zeolite surface.
cu
24
0.0287
1.559
68
22
0.0419
1.721
0.16
65
40
0.0444
0.930
0.48
76
79
0.0379
0.475
0.92
86
100
0.0332
0.373
~
40
70
I 00
Dispersion of CLI(%)
Figure 6. Effect of dispersion of Cu (%) on BET surface area of catalysts
The results in T a l k 3 indicated the dispersion of
Cu on the W/l.OCu/HZSM-5 is the smallest (about
22%) compared to the other samples, owing to the
largest mean particle diameter of Cu (1.721 rim) on
W/1 .OCu/HZSM-5. The small percentage of Cu being
dispersed on the W/1 .O/HZSM-5 catalyst led to only
a sniall amount of Cu ions being exchanged with H+.
As a consequence, t,he acidity of the W/1 .O/HZSM-5
is high (1.01 mol/kg), although its pore size is small
(24.4 A). It is clear that the role of Cu on t,he
W/Cu/HZSM-5 is not only to reduce the WE+ species
derived from the (W06)''- precursor with octahedral
coordination, but also to have an effect on the acidity
of the ZSM-5 zeolite, as revealed by the ammoniaT P D result.
~
-
The T P R profiles of W loaded 3.0Cii/HZSM-5
catalysts before and after reaction are depicted in Figure 7. The spectra denionstrated thc existence of a
single reduction peak of Cu oxidr at, 370 "C. However,
after reaction at 800 "C the peak disappeared, probably due to the reduction of CU'~+.This is probably
due to the partiid oxidation of the catalyst during the
reaction.
1
24
(rim)
W
0.0518
200
22
Mean particle
diameter
1
,
1
~
1
300
#
1
:
1
11
400
1
1
1
,
/
500
/
,
,
~
600
1
~
,
~
,
,
700
1
,
1
/
800
7'1 "C
Figure 7. Hydrogen-TPR
spectra
W/3.0Cu/HZSM-5 catalyst before
after reaction at 800 "C
of
and
Nevertheless, the peak of W"+ did not change
before and after the reaction. The TPR profile for
tungsten revealed that the addition of tungsten has
increased the thermal stability of the catalyst, as the
W component has not lost due t o sublimation after
reaction at 800 "C.
3.2. Performances of catalysts
The methane corivtrsiori increased due to the increasing copper content and copper surface area. This
result demonstrated that the methane conversion is
related to the copper surface area, as shown in Figurr 8. The methanr conversion of HZSM-5 is about
13% and increased to 21% at copper surface area of
1
1
Did; Dwi Anggoro et al./ Journal of Natural Gas Chemistry Vol. 15 No. 4 2006
346
0.92 m"/g for W/3.0Cu/HZSM-5 (Table 3). However, methane conversion over the W/l.OCu/HZSM-5
is lower than over the W/0.5Cu/HZSM-5, although
the Cu surface area of the former is larger. It is clear
that t,he dispersion of Cu on the active surface of the
catalyst affected the methane conversion. Since the
percentage of Cu dispersed on the catalyst surface is
the smallest for W/l.OCu/HZSM-5 (see Table 3) the
methane conversion is the lowest using this catalyst.
catalysts is very low (l%-2%), as depicted in Figure 9. This is probably due to that all catalysts
enhanced the oligomerization of C Z hydrocarbons tjo
C5+ liquid hydrocarbons and reduced the amount of
Cz-C3 hydrocarbons. The C5+ hydrocarbons selectivity increased with increasing acidity, as shown in
Figure 10. The selectivity to Cg+ hydrocarbons using
W/3.0Cu/HZSM-5 is the highest (34%), owing to it,s
highest acidity (1.19 mol/kg) and copper surface area
(0.92 m2/g) among other samples.
=
iXXX
Carbon monoxide DCarhon dioxide
C,, liquid hydrocarbons
C'.? hydrocarbons
Water
I00
80
h
g
m
.'2
.-I
-
60
0
s
40
20
0
0.04
0.07
0.16
0.48
n
0.92
cu surface area (rn'/g)
%.Effect of Cu BET surface area on methane
conversion
The products of the reaction between methane
and oxygen over HZSM-5, W/HZSM-5 and
W/Cu/HZSM-5 with different concentrations of copper are CZHZ,CZH4, CZH6, C3H6, CO, COz, HzO
and liquid hydrocarbons. Figure 9 summarizes the
product selectivities of all the catalysts. The results
in Figure 9 shows that over HZSM-5 and metal loaded
HZSM-5, the selectivity to carbon monoxide is higher
than that of carbon dioxide. This indicates that the
partial oxidation of methane has occurred with carbon
monoxide and hydrogen as the products. However,
hydrogen possibly reacted with carbon dioxide to
form water in the Reversed Water Gas Shift (RWGS)
reaction.
The water selectivity of HZSM-5 is higher than
the metal loaded HZSM-5, because the HZSM-5 catalyst has no metal content or nietal surface area, while
the metal loaded HZSM-5 catalysts have tungsten oxide and copper oxide contents. Over tungsten oxide
or copper oxide the reaction of hydrogen and carbon
monoxide produced hydrocarbon gases (CZ and C3),
which could oligornerize to liquid hydrocarbons (C5+)
in the presence of the HZSM-5 zeolite catalyst.
The selectivity for C2-C3 hydrocarbons for all
Figure 9. Distribution of product selectivities over
different catalysts
40
,
a
g
a
30-
-8
a
a
.-2> . .-c
0
u
2
20 -
8
8
8
8
C
.-
8
6
a
2
$
10
-
0
0
'
~
'
"
'
"
'
'
"
'
'
"
"
'
'
"
"
'
~
~
Figure 10. Effect of acidity on methane conversion and
liquid hydrocarbons C,+ selectivity
(H) methane conversion,
(0)
C,+ selectivity,
(0)
methane conversion over W/1.5Cu/HZSM-5,
( o ) C,+ selectivity over W/1.5Cn/HZSM-5
The result in Table 4 indicates that the coniposition of C5- 10 hydrocarbons increased as the nii~n-
ber of Bronsted acid sites ( N B A C )decreased. The
gasoline range (Cs-10) composition of ZSM-5 was
about 96% a t a NBAC value of 2.9 and increased t o
100% with NBACequal to 0.3 for the W/3.0Cu/ZSM-5
sample (Table 4). From the ammonia-TPD (Figure 4) spectra the strength of Bronsted acid sites
of W/3.0Cu/ZSM-5 is the strongest among others.
These results suggest that higher composition of gasoline formed from methane depends on the strength of
Branstfed acid sites.
Table 4. Number of Bronsted acid sites and
composition of c5-10 and C1,+ hydrocarbons
over HZSM-5, W/ZSM-5 and W/3.0Cu/ZSM-5
Catalyst
HZSM-5
W/HZSM-5
Wl3.OCulHZSM-5
No. of Bronsted acid Composition (%)
sites (NBAC)(pmol/g) Cs-Cio
2.9
96
C11+
4
0.5
99
1
0.3
100
0
4. Conclusions
The loading of HZSM-5 with tungsten and copper decreased the crystallinity, surface area, and also
total volume of the catalysts. However, the average
pore diameter and the acidity of the zeolites increased
as a result of the modification with the metals. Such
metal particles are smaller than the average pore size,
and the metal particles should be localized t o the inner side of the zeolite pores. T P R patterns indicated
that modified HZSM-5 by loading with tungsten enhanced its heat resistant performance, so the high reaction temperature (800 "C) did not lead to loss of the
W component by sublimation.
While loading HZSM-5 with tungsten and copper enhanced the methane conversion, C02 and Cg+
products, however, reduced the CO, C2-3, and HzO
selectivities. The process of converting methane t o
liquid hydrocarbons (Cs+) is dependent on the metal
surface area and the acidity of the zeolite. The
W/3.0Cu/HZSM-5 is the potential catalyst, because
over this catalyst high methane conversion and C5+
selectivity, and low H20 selectivity are obtained.
347
Acknowledgements
The authors gratefully acknowledge the financial support received in the form of a research grant from the Ministry of Science, Technology and Environment, Malaysia.
References
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