analysis and comparison of the photodecomposition reaction
TRANSCRIPT
Kennesaw State UniversityDigitalCommons@Kennesaw State University
Master of Science in Chemical Sciences Chemistry & Biochemistry
Spring 5-1-2015
Analysis and Comparison of thePhotodecomposition Reaction of Methanol onCeO2 Impregnated TiO2 Surfaces Versus BareTiO2 SurfacesChristopher A. EstesKennesaw State University
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Recommended CitationEstes, Christopher A., "Analysis and Comparison of the Photodecomposition Reaction of Methanol on CeO2 Impregnated TiO2Surfaces Versus Bare TiO2 Surfaces" (2015). Master of Science in Chemical Sciences. Paper 3.
Analysis and comparison of the photodecomposition reaction of methanol on CeO2
impregnated TiO2 surfaces versus bare TiO2 surfaces.
by
Christopher Aaron Estes
Bachelor of Science Degree
Berry College, 2013
_____________________________________________________________________
Submitted in Partial Fulfillment of the Requirements
For the Degree of Master of Science in the
Department of Chemistry and Biochemistry
Kennesaw State University
2015
--------------------------------------------- ---------------------------------------------Committee Chair Graduate Program Coordinator
--------------------------------------------- ---------------------------------------------Committee Member Department Chair
--------------------------------------------- ---------------------------------------------Committee Member College Dean
ii
ACKNOWLEDGEMENTS
Kennesaw State University: Department of Chemistry and Biochemistry
Mark B. Mitchell, Ph.D.
John Bacsa, Ph.D. Emory University
Patricia Pollner
Laurianne Tremeau-Cayal
Fellow Graduate Students
iii
ABSTRACT
Titanium dioxide (TiO2) has been well-studied primarily due to its unique
photocatalytic activity. Different preparation techniques have been developed to increase
its photocatalytic activity by various means, including simply increasing the surface area,
changing the makeup of the framework, and modifying the surface with a dopant. This
study focuses on the effect cerium oxide (CeO2) has on the photocatalytic activity of TiO2
by depositing CeO2 on the TiO2 surface. Methanol is used as a probe reactant for these
experiments due to its wide use in heterogeneous catalysis research and the limited
products that can form. Three-dimensional CeO2 crystallites form on the 1 and 2
monolayer CeO2-TiO2 samples while crystals are not observed for the 0.5 or 0.25
monolayer samples. Diffuse Reflectance Infrared Fourier Transform Spectroscopy
(DRIFTS) was used to characterize the surface chemistry occurring on these samples
while a flow reactor was used to characterize the reaction products using a long path gas
cell in a Fourier transform infrared spectrometer (FTIR). Except for the 1 monolayer
sample, CeO2 seems to increase the reaction and decomposition of organic compounds,
especially chemisorbed methoxy groups, on the surface. The CeO2 modified samples
appear to adsorb molecular methanol less strongly than TiO2. Upon UV irradiation, the
CeO2 modified samples desorb methoxy as methanol more so than TiO2. All samples
consistently produce carbon dioxide, water, formaldehyde, dimethyl ether, and methanol
in the presence of ultraviolet light and oxygen.
iv
TABLE OF CONTENTS
ACKNOWLEDGEMENTS……………………...……………………………………….ii
ABSTRACT……………………………………………………………………………...iii
LIST OF TABLES/FIGURES…………………………………………….………….…...v
CHAPTER 1. INTRODUCTION…………………………………………………………1
CHAPTER 2. MATERIALS AND METHODS………………….…………….……......10
CHAPTER 3. CHARACTERIZATION RESULTS…………………………..………....16
CHAPTER 4. DRIFTS RESULTS…………………………………………………….....19
CHAPTER 5. FLOW REACTOR RESULTS…………………………………………....35
CHAPTER 6. DISCUSSION…………………………………………………………….45
CHAPTER 7. CONCLUSION………………………………………………...……........49
REFERENCES…………………………………………………………………………..51
v
LIST OF TABLES/FIGURES
Figure 1: Diagram of flow reactor system………………………………………………...15
Figure 2: X-ray diffraction patterns of the CeO2 modified surfaces……………………..18
Figure 3: DRIFTS thermal spectra of TiO2 without oxygen……………………………....19
Figure 4: DRIFTS thermal spectra of TiO2 with oxygen….…………………….………..22
Figure 5: DRIFTS thermal spectra of 2ML CeO2-TiO2 with oxygen …………………….23
Figure 6: DRIFTS UV spectra of TiO2 without oxygen …………………………………25
Figure 7: DRIFTS UV spectra of TiO2 with oxygen ………….…………………………..26
Figure 8: DRIFTS UV spectra of 1ML CeO2-TiO2 without oxygen………………………27
Figure 9: DRIFTS UV spectra of all powders (2700-3100 cm-1)…………………………28
Figure 10: DRIFTS UV spectra of all powders (3400-4000 cm-1)……………………….33
Figure 11: DRIFTS UV spectra of all powders (1200-1800 cm-1)………………….…….34
Figure 12: Gas phase spectra of photodesorption products with oxygen…………………37
Figure 13: Flow Reactor Data: TiO2 with methanol and oxygen…………………………38
Figure 14: Flow Reactor Data: 2ML CeO2-TiO2 with methanol and oxygen ……………39
Figure 15: Flow Reactor Data: TiO2 with methanol and oxygen exposed to UV…………40
vi
Figure 16: Flow Reactor Data: 0.25 ML CeO2-TiO2 with methanol and oxygen exposed to
UV………………………………………………………………………………………..41
Figure 17: Flow Reactor Data: 0.5 ML CeO2-TiO2 with methanol and oxygen exposed to
UV………………………………………………………………………………………..42
Figure 18: Flow Reactor Data: 1 ML CeO2-TiO2 with methanol and oxygen exposed to
UV………………………………………………………………………………………..43
Figure 19: Flow Reactor Data: 2 ML CeO2-TiO2 with methanol and oxygen exposed to
UV………………………………………………………………………………………..44
Table 1: Weight % Ce on CeO2 modified TiO2 samples………………………………….17
Table 2: Characteristic frequencies of each surface adsorbed species……………………29
Table 3: % loss of integrated peak area of characteristic IR C – H stretching region.........30
1
CHAPTER 1. INTRODUCTION
Heterogeneous catalysts serve many important roles in industrial and commercial
settings.1-3 The activity of many catalysts is induced by thermal energy, but catalysts exist
that can use other energy sources for catalysis. One type of catalyst that responds to other
energy sources are photocatalysts. As the name suggests, photocatalysts rely on light to
induce catalysis.4-6 When photocatalysis occurs, the energy from light is sufficient in
photocatalytic systems to promote an electron from the valence band to the conduction
band of the surface, which generates electron-hole pairs.7 The catalysis occurs when there
is a charge transfer between the electron-hole and the adsorbant.7 The particular spectral
region responsible for inducing this activity is dependent on the band gap of the particular
material.7 Titanium (IV) oxide (TiO2) is a photocatalyst with a relatively large band gap
(3.0 and 3.2 eV) with an active spectral region in the ultra-violet (UV).8,9 UV light
composes a mere 5% of solar light; nevertheless, TiO2 has been studied for applications
in photocatalysis for more than four decades.4
One of the first research efforts to analyze TiO2 photocatalysis was performed by
Fujishima and Honda where the observed the photocatalytic splitting of water on a
platinum electrode.21 The reactions below are reported to be the mechanistic steps of the
water splitting reaction described by Fujishima and Honda.
TiO2 + 2hv → 2e- + 2h+ (1)
2
When TiO2 is irradiated with ultraviolet light, electrons are promoted from the valence
band to the conduction band which creates a positive hole on the surface.
H2O + 2h+ → 0.5O2 + 2H+ (2)
Water will react with these positive holes to generate oxygen gas and protons, which will
then continue to react with the conduction band electrons to form hydrogen gas.
2H+ + 2e- → H2 (3)
Another study was performed on the photooxidation of cyanide on several TiO2
materials.10 The authors found that without UV irradiation, no oxidation occurred.
However, 30 minutes of 450 W UV irradiation was sufficient to oxidize 75% of the
cyanide on Anatase. They also performed an experiment where the cyanide dosed TiO2
was exposed to sunlight for two days. Two days of solar exposure provided enough
energy to oxidize 99% of the cyanide. This shows that, although slow, sunlight is
sufficient to induce TiO2 photocatalytic activity.
How TiO2 photocatalysis functions has been well studied.11-18 TiO2 has three
lattice structures with different optical properties: brookite, rutile, and anatase. Brookite
is studied less because it lacks photocatalytic activity. Anatase exhibits the most
photocatalytic activity with rutile closely following.7,9 Studies have shown that, although
the anatase phase demonstrates high photocatalytic activity as a single phase, the P25
mixed phase material by Degussa shows greater photocatalytic efficiency than any single
phase material.7,9,19 These Degussa samples are composed of approximately 80% anatase
and 20% rutile. Images have been obtained via TEM that suggest the two phases
agglomerate separately.9 P25 titanium (IV) oxide from Aldrich is the material used for
this investigation.
3
One use for photocatalytic materials is the degradation of potentially dangerous
bacteria or toxic chemicals.11,20-30 One study investigate the photodeactivation of E-coli
by TiO2.24 In the absence of radical scavenging groups, it was found that TiO2 is effective
at deactivating E-coli in water. Another investigation by Konstantinou and Albanis shows
how well TiO2 degrades various pesticides and monitors the intermediates formed.25
Although TiO2 is effective at breaking down these compounds, a number of different
products are formed. The photocatalytic reaction resulted in different byproducts
depending on the pesticide being degraded. Due to the complexity of the
photodecomposition mechanism, it is difficult to predict how TiO2 will react with
different organic compounds.
Alberici, et al. studied the reaction of aliphatic alcohols with P-25 and suggested
that reactions are negligible in the absence of ultraviolet light and the absence of oxygen
gas.20 These experiments also showed that photooxidation didn’t occur unless the
compound was preadsorbed to the surface before irradiation. This is believed to occur
because the photoejection and reinsertion event of electrons passing from the valence
band to the conduction band of TiO2 is so rapid that, unless the adsorbate is preadsorbed,
electrons don’t remain in the conduction band long enough to react.
Methanol is a common organic molecule used to investigate the surface chemistry
of heterogeneous catalysts. 2,31-40 As a well-studied probe molecule, the surface binding
and decomposition reactions on TiO2 have been investigated by various groups. In 1977,
Carrizosa and others published an article describing the thermal desorption/reaction of
methanol on TiO2.41 Reaction products desorbed below 400°C were determined to be
CH3OH, CH3OCH3 and H2O, whereas above 400°C, desorption products were
4
determined to be CO, CO2, H2, CH4 and C2H6. In 1988, Taylor, et al. analyzed the surface
characteristics of adsorbed methanol on TiO2 and the products desorbed through the use
of infrared spectroscopy and temperature programmed desorption.39 They found that
adsorbing methanol at different temperatures resulted in different surface adsorbed
species; physisorbed methanol and chemisorbed methoxy are observed to form at room
temperature whereas chemisorbed methoxy groups dominate during adsorption at high
temperature. Taylor suggests this determines a kinetic barrier for surface adsorbed
species that can be overcome by increasing the adsorption temperature to over 200°C. It
was found that different products are desorbed from the surface depending on the
coverage of methanol on the surface as described in the following reactions.
Where (ad) indicates adsorption to the surface. Assuming that chemical adsorption of
methoxy occurs due to a reaction with molecular methanol and a surface hydroxyl group
to create water and a surface adsorbed methoxy group, the surface coverage can be
explained in relation to previously existing surface hydroxyl groups. At lower coverages
of methanol, it is possible to have two surface hydroxyl groups nearby such as in
equation (4). In this case, a reaction would occur in which the two hydroxyl groups will
OH(ad) + OH(ad) → H2O(g) + O(ad) (4)
CH3O(ad) + OH(ad) → CH3OH(g) + O(ad) (5)
CH3O(ad) + O(ad) → CH2O(g) + OH(ad) + 2e– (6)
CH3O(ad) + CH3O(ad) → CH3OCH3(g) + O(ad) (7)
5
leave an oxide site on the surface and form water as a desorption product. Methoxy can
also react with a neighboring hydroxyl group to reform molecular methanol as described
in equation (5). Methoxy can also react with a neighboring oxide site to form desorbed
formaldehyde and a surface hydroxyl as described in equation (6). It was also determined
that two neighboring methoxy groups can react to form dimethyl ether as described in
equation (7). Another study proposed the following mechanism for the formation of
methyl formate on rutile TiO2.42
In equation (8), an adsorbed formaldehyde molecule reacts with a bridging oxygen from
the surface to form an adsorbed CHO species and a surface hydroxyl group. In equation
(9), the CHO species and a chemisorbed methoxy group react to form adsorbed methyl
formate, which can readily leave the surface. With such interesting surface chemistry
made possible by TiO2, other research groups continued to investigate TiO2.
In 1999, Chuang, et al., studied the photocatalytic reactions of methanol on TiO2
in the presence and absence of oxygen gas.34 It was found that without oxygen,
formaldehyde and methanol are desorbed upon ultraviolet irradiation. However,
formaldehyde is not mentioned as a desorption product for the reaction with oxygen gas
present. Instead, products reported are adsorbed formate, adsorbed water, carbon dioxide
and hydrogen gas. The production of carbon dioxide and hydrogen gas is minimal. A
mechanism is proposed for how methoxy reacts on a TiO2 surface in the presence and
CH3O(ad) + CH2O(ad) + O(ad)→ CH3O(ad) + CHO(ad) + OH(ad) (8)
CH3O(ad) + CHO(ad) → HCOOCH3(ad) (9)
6
absence of oxygen.34 In the absence of oxygen, the methoxy behaves as the electron-hole
capturing group which will react with the surface to produce formaldehyde. In the
presence of oxygen methoxy still behaves as the initial electron-hole capturing group.
However, oxygen will react with the radical species in such a way as to produce a
tetraoxide species that will decompose to form water, oxygen, and a chemisorbed
formate. A similar mechanism is proposed by Fujishima on the photoreaction of ethoxy
on TiO2 in the presence of oxygen.5,6 The only difference is that Fujishima suggests that a
surface hydroxyl group acts as the initial electron-hole capturing group.
Data indicating the presence of formate on the surface are supported by the
findings of Busca, et al.43 Experiments were conducted to examine formaldehyde
adsorption on different metal oxide surfaces and to find characteristic formaldehyde and
formate ion peaks for the different materials. For TiO2, those peaks appeared at 2885 cm-
1, 2880 cm-1, 1575 cm-1, 1560 cm-1, and several others between 1400 and 1300 cm-1. The
previous study shows some of the ways TiO2 photocatalysis occurs. However,
investigations have been performed to enhance the catalytic activity of TiO2
Some studies have investigated modification of TiO2 to determine any changes in
catalytic activity. Modifying the activity of a metal oxide can occur in various ways. One
example is incorporating a metal or nometal into the framework of the lattice. One such
study was conducted by Al-Mazroai and team in 2007.31 Nitrogen and iodine
incorporated into the lattice can shift the absorption spectrum of TiO2 into the visible
region, significantly increasing the ability of TiO2 to photodegrade compounds using
visible light.44
7
Other methods include impregnating the surface with metals. Other dopants that have
been investigated including noble metals such as gold and silver. Noble metal deposition
has been shown to uniquely enhance methanol reformation on TiO2 photocatalysts.31
Silver and gold deposition has also been shown to greatly enhance the visible light
absorption of TiO2, although visible light photocatalysis increased minimally.44 These
noble metals were actually recorded to inhibit photocatalysis of aromatic, weakly-
adsorbing compounds such as phenol in the UV region, whereas strong-adsorbing
chemicals such as oxalic acid were shown to degrade much more quickly in the presence
of these dopants.44 Palladium, platinum, and silver were used as dopant metals on TiO2
with goals of forming hydrogen gas and reforming methanol. Out of the considered
metals, palladium and platinum demonstrated a higher affinity for hydrogen gas
production than any of the other metals. Some metals were completely unreactive, but
iridium, silver, rhodium and rubidium all showed reactivity as well, albeit less activity
than palladium and platinum. The proposed explanation is based on the ease of
reducibility of noble metals. Although modification with noble metals has been shown to
increase reactivity, this study also showed how weight loading affected the reactivity of
the material. The highest reactivity was determined at 0.5 wt% Pd, whereas weight
loadings over 4 wt% Pd failed to produce any hydrogen gas. For methanol reformation to
occur, it was determined that a noble metal had to be present.
The area of research examined in the investigation described in this thesis is
surface modification with a secondary metal oxide. Mixed metal oxides have interesting
catalytic properties due to the heterogeneous combination of catalytic activities. One
study investigated the effects of modifying TiO2 with copper oxide.45 Cyanide was used
8
as a probe molecule to monitor how TiO2 would photooxidize cyanide when modified
with different quantities of copper (II) oxide. Photooxidation was shown to increase
slightly with copper oxide modification. However, once the quantity of copper (II) oxide
on TiO2 was increased to more than 0.1 at.%, photocatalytic enhancement was inhibited
due to copper (I) cyanide complexes competing for surface hydroxyl groups on the TiO2.
Our efforts focused on observing how cerium (IV) oxide can change the photocatalytic
activity of TiO2.
Cerium (IV) oxide (CeO2) is an important industrial and commercial
catalyst.18,27,48 CeO2 is most known for its use in the three-way catalyst, where its
oxidative properties are used to oxidize harmful molecules from automotive exhaust.45 It
also serves an important role in the dehydrogenation of ethylbenzene to styrene, which is
a precursor for polystyrene.48 Studies have also shown CeO2 is useful for removing soot
from diesel exhaust and helpful in waste-water management.48 Ceria catalysts have also
been shown to have improved activity upon modification with various noble metals. One
study is the water-gas shift reaction on copper and nickel modified CeO2. It was found
that a probe molecule (carbon monoxide) adsorbed to the doped metal clusters on the
surface while CeO2 provided oxygen to produce CO2.14 These results are supported by
the findings of Bunluesin, et al., where the water-gas shift reaction was analyzed on CeO2
surfaces modified with platinum, palladium, and rhodium. It was proposed that carbon
monoxide would adsorb to the dopant metal while CeO2 was responsible for reducing the
surface in the same fashion as reported by Li, et al.49,50
Modification TiO2 with CeO2 can be accomplished in two different ways. CeO2
can be incorporated into the TiO2 lattice with the intention of modifying the active
9
spectral region.51-54 The other way is to deposit CeO2 on TiO2 in a heterogeneous
fashion.55 Rather than modifying the surface in order to directly change the catalytic
behavior of TiO2, this method provides the environment for each metal oxides individual
catalytic activity to work together. The current study focuses on the latter to observe any
changes when the catalytic activity of TiO2 and CeO2 are combined.
10
CHAPTER 2. MATERIALS AND METHODS
Section A. Sample preparation and characterization
P25 titanium (IV) oxide, 99.5% from Aldrich, was used both as a control and as
an ingredient in the preparation of the CeO2 modified samples. Mixed oxide samples
composed of cerium oxide and titanium (IV) oxide were prepared with 1.3, 2.47, 4.84,
and 9.10 wt% Ce via a wet impregnation method known as precipitation deposition.
These weight loadings correspond to monolayer quantities where one monolayer is
equivalent to CeO2 replacing all surface hydroxyl groups on the P25 substrate56 thus, the
CeO2 modified TiO2 samples are composed of 0.25, 0.5, 1 and 2 monolayers of CeO2. To
prepare each modified material, TiO2 was calcined at 400 °C for 4 hours. For the 1ML
CeO2 modified TiO2, 0.318g of cerium (III) nitrate hexahydrate precursor, corresponding
to the monolayer weight loading desired, was dissolved in 200 mL deionized water in a
1000 mL beaker. Two grams of the calcined TiO2 was added to the 1000 mL beaker.
With continuous stirring, 0.1M NH4OH was added dropwise to the suspension until a pH
of approximately 10 was achieved. The solution was left to stir for an hour. The solid was
filtered from the suspension and dried in an oven at 80 °C for 24 hours. The samples
were removed, broken up in a mortar and pestle, and calcined at 400°C for 4 hours.
Cerium analysis of the CeO2 modified TiO2 materials was performed using
Inductively Coupled Plasma Optical Emission Spectroscopy (ICP-OES) by NSL
Analytical Services, Inc., Cleveland, OH. Powder XRD experiments on the CeO2
modified solids were carried out at Emory University. The data were collected on a
11
Bruker D8 Advance Diffractometer (Bruker AXS, Karlsruhe, Germany) in Bragg–
Brentano geometry using a primary 40 mm Göbel mirror producing a parallel beam. The
beam was collimated with a 0.6 mm diameter exit port. Cobalt Kα radiation (λ=1.7889Å)
was used and the instrument was operated at 35 kV and 40 mA. Data were collected in
reflection geometry with a Vantec PSD detector equipped with radial Soller slits and
configured with a 10° detector window. Data were collected between 10° and 95° 2-theta
using a 0.006° step size and variable scan times (over a 36 hour total time period). The
sample was mounted on a silicon zero background holder on the X,Y,Z stage of the
diffractometer. A laser/video-microscope alignment system was used to position the
samples. The data were processed and analyzed using EVA (Bruker-AXS, Karlsruhe,
Germany) and with REX (Mauro Bortolotti et al., 2009). The database used for
comparison with the experimental data was the International Centre for Diffraction Data
PDF-2, Release 2005.
Section B. DRIFTS experiment
Diffuse Reflectance Infrared Fourier Transform Spectra (DRIFTS) of the
modified and unmodified TiO2 solids with adsorbed methanol before and after UV
irradiation were measured using a Harrick Scientific Praying Mantis Diffuse Reflectance
Accessory in a Thermo-Nicolet 6700 FT-IR collected at 4 cm-1 resolution using a DTGS
detector.
For the thermal decomposition experiments in the absence of gaseous oxygen, a
sample powder was placed in the DRIFTS cell and a dome was placed over the sample
with two potassium bromide windows for the infrared beam and a quartz window for UV
irradiation. The sample cell was heated to 100°C with N2 flowing for 1 hour, then
12
allowed to cool and evacuated to approximately 1x10-3 Torr. A 1000 ppm gas mixture of
methanol in nitrogen was prepared in a glass vacuum manifold using 99.9% liquid
methanol from Fisher Scientific that had been cleaned via the freeze-pump-thaw method
and ultra high purity nitrogen, purchased from NexAir. The methanol adsorption was
carried out by allowing this gas mixture to flow over the 25°C sample for 3 hours at a
rate of 1 standard cubic centimeter per minute (sccm). After this surface pretreatment, the
sample cell was heated stepwise to 50, 100, 200, 300, and 400°C, and infrared spectra
were collected at each temperature.
For the thermal decomposition experiments carried out in the presence of gaseous
oxygen, a sample was placed in the DRIFTS cell and pretreated in the manner above.
This sample was heated stepwise as described with the previous sample, but with a 10.05
mole% O2/He gas mixture, purchased from Airgas, Inc., flowing over the sample at a rate
of 1 sccm during the stepwise heating phase.
For the UV irradiation experiments in the absence of gaseous oxygen, a sample
was placed in the DRIFTS cell and pretreated in the manner above. This sample was
irradiated with UV light at 1.5 W for 10 minutes. The UV source was a Dymax
BlueWave 200 UV lamp with a 0.5 m fiberoptic light pipe attached. The light pipe was
placed directly in front of the quartz window of the DRIFTS dome and the light directed
towards the sample.
For the UV irradiation experiments carried out in the presence of gaseous oxygen,
a sample was placed in the DRIFTS cell and pretreated in the manner above. The sample
was irradiated with UV light at 1.5 W for 10 minutes with the addition of the 10.05
13
mole% O2 gas mixture flowing over the sample at a rate of 1.0 sccm or the duration of
UV irradiation.
For spectral characterization of the sample, a background infrared spectrum was
obtained before methanol adsorption. A sample spectrum was obtained after methanol
adsorption. A spectrum was obtained at each temperature step for the thermal
experiments and after 10 minutes of UV irradiation for the photo experiments.
Section C. Flow Reactor Experiments
A flow reactor system was assembled for the purpose of characterizing the
reaction products that desorb from the solid surface. The system consisted of a gas
mixing manifold that carried the reactant gas mixture to the catalyst bed. After flowing
through the catalyst, the product gas stream flowed to a long-path (2.4 m) infrared gas
cell (Ultra-Mini Long Path Cell from Infrared Analysis, Inc.) in a Thermo-Nicolet Avatar
360 FT-IR. The gases were supplied by prepared gas mixtures: a 1000 ppm
methanol/nitrogen mixture purchased from PRAXAIR, and a 1000 ppm oxygen/nitrogen
mixture purchased from NexAir. The two gases were connected to separate flow meters
which allowed the regulation of the flow of each gas into the mixing manifold. The
mixing manifold was connected to a stainless steel reactor tube with a wire mesh stage
for the powdered sample. A mount was constructed on top of the tube to allow for the
fiber optic light guide from the UV light source.
The mass of TiO2 used for the flow reactor experiments was 200 mg and the mass
of the CeO2 modified powders was 800 mg to achieve similar sample volume in the
reactor. A sample powder was loaded on the wire mesh sample stage and flushed with the
14
1000 ppm oxygen/nitrogen mixture. The flow of the oxygen/nitrogen mixture was
reduced to 14 sccm. A macro was written to obtain infrared spectra at 4 cm-1 resolution
every 2 minutes until at most 100 spectra were obtained. The first five sample spectra
were obtained while the oxygen mixture flows at 14 sccm to establish a baseline. The
oxygen mixture was then turned off and the 1000 ppm methanol/nitrogen mixture was
allowed to flow through the sample at a rate of 14 sccm. This experiment was repeated on
a new sample while the sample was irradiated with UV light from the source mentioned
above at 50 mW intensity. A new sample was loaded in the cell and the baseline
spectrum are obtained in the same manner. After the fifth spectrum was obtained, the
oxygen flow was reduced to 7 sccm and the methanol flow was set to 7 sccm. This
experiment was repeated on a new sample while the sample was irradiated with 50 mW
UV light.
15
Figure 1: Diagram of flow reactor system. The gases exit their respective tanks and flow
is controlled using flow meters. Upon exiting the flow meters, the gases then enter a T-
connector and mix before entering the reactor. The gases enter the reactor through a T-
connector perpendicular to the reactor. The gas flows through the sample and continues
on to the gas cell in the FTIR. The ultraviolet light source enters through the top of this
second T-connector in order to shine directly onto the surface.
16
CHAPTER 3. CHARACTERIZATION RESULTS
As mentioned in Chapter 2, the theoretical weight percent (wt %) of cerium on
each of the CeO2 modified samples were calculated to be 1.3, 2.47, 4.84, and 9.10 wt %
Ce. The actual weight percent of cerium on the samples were determined by NSL
Analytical by ICP-OES to be 1.26, 2.45, 4.81, and 8.84 wt % Ce corresponding to the
0.25, 0.5, 1, and 2 ML samples respectively as observed in Table 1. Figure 2 displays the
XRD patterns for the four modified samples. The diffraction patterns indicated that the 1
and 2 ML samples showed evidence of the presence of cerium oxide crystals whereas the
0.5 and 0.25 ML samples showed no evidence of crystallite formation. Crystallization on
the 1ML sample could suggest that CeO2 does not deposit uniformly on the surface,
which would expose TiO2 sites. Also, the number of OH sites on the TiO2 powder could
be different from the literature reference used for calculations. Fewer OH sites could lead
to 3-dimensional sites forming.
17
Table 1: Weight loading of cerium on CeO2 modified TiO2
samples
Theoretical CeO2
Coverage(monolayers)
TheoreticalWt% Ce
ActualWt% Ce
.25 ML 1.3 1.26
.5 ML 2.47 2.45
1 ML 4.84 4.81
2 ML 9.10 8.84
18
Figure 2: X-ray diffraction patterns of the CeO2 modified surfaces. The spectra arestacked and spaced in an arbitrary fashion to distinguish one pattern from another. Frombottom to top, the 0.25, 0.5, 1, and 2 monolayer surfaces are presented. The vertical linesmark characteristic CeO2 crystal diffraction peaks. These peaks are easily visible on the2 monolayer surface and are just beginning to appear on the 1 monolayer surface.
2-Theta
Intensity
2ML
1ML
0.5ML
0.25M
L
19
CHAPTER 4. DRIFTS RESULTS
Diffuse Reflectance Infrared Fourier Transform Spectroscopy (DRIFTS) was used
for these experiments due to the surface sensitivity it exhibits compared to other infrared
analysis techniques.57 These experiments were used to characterize the surface chemistry
occurring on the various metal oxides examined. As discussed in Chapter 1, heating TiO2
powders to 400 °C is insufficient to completely decompose adsorbed methoxy groups on
the surface. In the current work, a methanol gas mixture was allowed to flow over the
sample in the DRIFTS cell until the surface was saturated with methanol and methoxy.
The flow duration was determined by obtaining infrared spectra every 30 minutes until
no change in the infrared spectrum was observed, which we determined to be 3 hours. An
examination of the thermal decomposition behavior was carried out for the bare TiO2
surface and for the 2 ML CeO2 modified surface. From Figure 3, it appears that in the
absence of oxygen, heating the bare TiO2 surface to 400 °C is insufficient to effectively
decompose methoxy groups on the surface. When oxygen is introduced to the system
methoxy is more efficiently decomposed upon heating as shown in Figure 4. In Figure 5,
the 2 ML CeO2 surface is capable of thermally converting methoxy to other products in
the presence of oxygen. However, effective methoxy conversion is not observed at
temperatures less than 400 °C.
The next experiments examined the decomposition reactions that take place with
UV irradiation, with and without oxygen present. To determine the effect that oxygen has
20
on the products from photocatalysis, photodecomposition experiments were performed
first on each of the modified samples in the absence of oxygen, and then while exposed to
oxygen. Figure 6 shows the photocatalytic degradation of methanol on the bare TiO2
surface in the absence of
21
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1440
2826
2852
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2846
2922
2951
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-0 .25
-0 .20
-0 .15
-0 .10
-0 .05
0 .00
0 .05
0 .10
0 .15
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0 .25
0 .30
Ab
sorb
ance
1 500 2000 2500 3000 3500 4000
Wavenumbers (cm-1)
Figure 3: DRIFTS thermal spectra of TiO2 without oxygen. Before heating, the surfaceis completely saturated with methanol.
Before heating
After heating
22
1360 1541
2824
2863 2894
2924
2955
1358
1438
2822
2848
2922
2947
-0 .20
-0 .15
-0 .10
-0 .05
0 .00
0 .05
0 .10
0 .15
0 .20
0 .25
0 .30
0 .35
0 .40
Ab
sorb
ance
1 500 2000 2500 3000 3500 4000
Wavenumbers (cm-1)
Before heating
Figure 4: Spectra from DRIFTS thermal reactions of methanol on TiO2 in the presenceof oxygen before and after heating. Before heating, the surface is completely saturatedwith methanol.
After heating
23
Figure 5: DRIFTS thermal spectra of 2ML CeO2-TiO2 with oxygen. Before heating,the surface is completely saturated with methanol.
-0 .4
-0 .3
-0 .2
-0 .1
-0 .0
0 .1
0 .2
0 .3
0 .4
0 .5
0 .6
Ab
sorb
ance
1 500 2000 2500 3000 3500Wavenumbers (cm-1)
f
25 °C
50 °C
100 °C
200 °C
300 °C
400 °C
24
oxygen, showing that in the absence of oxygen, UV irradiation is sufficient to decompose
surface methoxy groups, but there are notable differences.
After irradiation, approximately 30% of the integrated area of the C – H stretching
region (2700-3100 cm-1) remains. In Figure 7, TiO2 is much more effective desorbing
products from the surface when exposed to gaseous oxygen. In Figure 8, however, it
appears that the 1 monolayer surface actually hinders reaction products from leaving the
surface in the absence of oxygen. After UV irradiation, approximately 50% of the
absorption intensity in this region remains. In Figure 7, methanol is observed to interact
with TiO2 in much the same manner when exposed to UV light and oxygen as described
by Chuang, et al., with both adsorbed molecular methanol and methoxy being observed.34
Figures 9, 10, and 11 show the results of the photodegradation experiments in the
presence of oxygen with the CeO2 modified surfaces as well as the bare TiO2 surface. In
Figure 9a, note the absorption frequencies of the adsorbed molecular methanol and
methoxy groups. The symmetric (2825 cm-1) and antisymmetric (2925 cm-1) C – H
stretching frequencies, corresponding to chemisorbed methoxy, and the symmetric (2855
cm-1) and antisymmetric (2955 cm-1 ) C – H stretching peaks, corresponding to
physisorbed methanol, are assigned based on their similarities to the C – H stretching
frequencies previously reported by Chuang, et al.34 The characteristic peak intensity of
molecularly adsorbed methanol is less on samples with higher CeO2 weight loadings than
on those with lower weight loadings. Loss of IR intensity corresponding to physisorbed
methanol is not countered with an increase in intensity of methoxy characteristic peaks,
meaning there is less methanol adsorbed on the surface. After UV irradiation, a
25
2828
2926
1360
1438
2822
2848
2922
2951
-0 .30
-0 .25
-0 .20
-0 .15
-0 .10
-0 .05
0 .00
0 .05
0 .10
0 .15
0 .20
0 .25
Ab
sorb
ance
1 500 2000 2500 3000 3500 4000
Wavenumbers (cm-1)
Before UV irradiation
Figure 6: Spectra from DRIFTS photocatalytic reactions of methanol on TiO2 in theabsence of oxygen before and after UV irradiation. Before irradiation, the surface iscompletely saturated with methanol.
After UV irradiation
26
1360
1565
2832
2867
2930
2959
1360
1440
2826
2846
2926
2953
-0 .14
-0 .12
-0 .10
-0 .08
-0 .06
-0 .04
-0 .02
-0 .00
0 .02
0 .04
0 .06
0 .08
0 .10
0 .12
0 .14
0 .16
0 .18
0 .20
0 .22
0 .24
0 .26
0 .28
Ab
sorb
ance
1 500 2000 2500 3000 3500
Wavenumbers (cm-1)
Before UV irradiation
Figure 7: Spectra from DRIFTS photocatalytic reactions of methanol on TiO2 in thepresence of oxygen before and after UV irradiation. Before irradiation, the surface iscompletely saturated with methanol.
After UV irradiation
27
2828
2926
1356
1436 2826
2846
2924
2953
-0 .6
-0 .5
-0 .4
-0 .3
-0 .2
-0 .1
0 .0
0 .1
0 .2
0 .3
0 .4
Ab
sorb
ance
1 500 2000 2500 3000 3500 4000Wavenumbers (cm-1)
Figure 8: Spectra from DRIFTS photocatalytic reactions of methanol on 1ML CeO2-TiO2 in the absence of oxygen before and after UV irradiation. Before irradiation, thesurface is completely saturated with methanol.
Before UV irradiation
After UV irradiation
28
-0 .05
0 .00
0 .05
0 .10
0 .15
0 .20
0 .25
0 .30
0 .35
0 .40
Ab
sorb
ance
2 750 2800 2850 2900 2950 3000 3050 3100
Wavenumbers (cm-1)
-0 .02
-0 .00
0 .02
0 .04
0 .06
0 .08
0 .10
0 .12
0 .14
0 .16
0 .18
Ab
sorb
ance
2 750 2800 2850 2900 2950 3000 3050 3100
Wavenumbers (cm-1)
2ML CeO2-TiO2
Figure 9: DRIFTS UV spectra of all powders (2700-3100 cm-1). DRIFTSphotocatalytic experiments in the presence of oxygen before UV irradiation (a) andafter UV irradiation (b). Before irradiation, the surface is completely saturated withmethanol.
a
b
1ML CeO2-TiO2
0.5ML CeO2-TiO2
0.25ML CeO2-TiO2
Bare-TiO2
2ML CeO2-TiO2
1ML CeO2-TiO2
0.5ML CeO2-TiO2
0.25ML CeO2-TiO2
Bare-TiO2
29
Tab
le 2
: C
hara
cter
isti
c fr
eque
ncie
s of
eac
h su
rfac
e ad
sorb
ed s
peci
es Che
mis
orbe
d C
HO
2
Cur
rent
Stud
y
2886
2870 - - -
1372
,13
60
1566 -
G. B
usca
,et
al.43
2885
,28
80 - - -
1372
,13
46
1575
,15
60 -
Phy
siso
rbed
CH
2O
Cur
rent
Stud
y
- - - - - - -
M. E
l-R
oz, e
tal
.58
2896
2973
1508
1750 - - -
Che
mis
orbe
d C
H3O
Cur
rent
Stud
y
2825
2924
1475
-14
50 - - - -
Tay
lor,
et
al.39
El-
Roz
, et
al.58
2830
2930
1475
-14
50 - - - -
Phy
siso
rbed
CH
3OH C
urre
ntSt
udy
2849
2954
1475
-14
50 - - - -
Tay
lor,
et
al.39
El-
Roz
,et
al.58
2852
2952
1475
-14
50 - - - -
v s(C
H)
v as(
CH
)
Del
as(C
H)
v(C
O)
v s(C
OO
)
v as(
CO
O)
w(C
H2)
30
Table 3: Percent loss in the integrated peak area over the characteristic infraredC – H stretching region.
Material With oxygen ~% Loss (2800-3000 cm-1)
TiO2 No 75%
1 Monolayer CeO2/TiO2 No 45%
TiO2 Yes 75%
0.25 Monolayer CeO2/TiO2 Yes 75%0.5 MonolayerCeO2/TiO2 Yes 80%1 Monolayer CeO2/TiO2 Yes 65%2 Monolayer CeO2/TiO2 Yes 80%
31
significant loss of intensity is observed at ~2830, 2850, 2930, and 2950 cm-1 indicating
that the coverage of both chemisorbed methoxy and physisorbed methanol has
diminished. As can be seen in Table 1, the TiO2 surface experienced a loss of ~70% of
the integrated peak area in the C – H stretching region. The percent loss values are also
recorded for the CeO2 modified surfaces in Table 1.
Before UV irradiation, the C – H absorptions observed in the region between
2800 and 3000 cm-1, corresponding to methanol and methoxy groups, coincide with a
loss of intensity in the region between 3400 and 4000 cm-1 shown in Figure 10a, due to O
– H stretching frequencies for surface hydroxyl groups. This indicates a loss of surface
bound hydroxyl groups upon methanol adsorption. These findings are consistent with
those of Chuang et al. and Taylor et al. for the adsorption of methanol on TiO2.34,39
Methanol reacts with surface hydroxyl groups to form surface bound methoxy and
gaseous water. Before UV irradiation, the TiO2 surface has a peak at 3630 cm-1 that is not
observed on the CeO2 modified samples, indicating CeO2 removes sites typically
occupied by these missing hydroxyl groups. Figure 10b illustrates this region after UV
exposure. Hydroxyl groups return to the surface, but the stretching mode of the O – H
bonds have higher frequencies. Higher frequencies indicate the O – H bonds of the
hydroxyl groups are now stronger than before, which indicates the hydroxyl groups are
more basic.
Although our TiO2 surfaces interact with methanol in much the same manner as
TiO2 surfaces reported in other studies, peaks are observed at ~1440 cm-1 and ~1360 cm-1
before UV irradiation which also correspond to chemisorbed methoxy and physisorbed
methanol. After irradiation, these peak at 1360 cm-1 becomes larger and a peak at 1566
32
cm-1 appears whereas the peak at 1440 cm-1 disappears. Figure 11a and 11b provides the
region from 1000-2000 cm-1 for each of the samples before and after UV irradiation. The
peak at ~1360 cm-1 appears smaller on the CeO2 modified samples.
33
-0 .10
-0 .05
0 .00
0 .05
0 .10
0 .15
0 .20
0 .25
0 .30
0 .35
Ab
sorb
ance
3 500 3600 3700 3800 3900 4000
Wavenumbers (cm-1)
-0 .08
-0 .06
-0 .04
-0 .02
0 .00
0 .02
0 .04
0 .06
0 .08
0 .10
0 .12
0 .14
0 .16
0 .18
0 .20
0 .22
0 .24
0 .26
0 .28
0 .30
0 .32
Ab
sorb
ance
3 500 3600 3700 3800 3900 4000
Wavenumbers (cm-1)
Figure 10: DRIFTS UV spectra of all powders (3400-4000 cm-1). DRIFTSphotocatalytic experiments in the presence of oxygen before UV irradiation (a) andafter UV irradiation (b). Analysis of surface hydroxyl behavior. Before irradiation, thesurface is completely saturated with methanol.
a
b
2ML CeO2-TiO2
1ML CeO2-TiO2
0.5ML CeO2-TiO2
0.25ML CeO2-TiO2
Bare-TiO2
2ML CeO2-TiO2
1ML CeO2-TiO2
0.5ML CeO2-TiO2
0.25ML CeO2-TiO2
Bare-TiO2
34
-0 .04
-0 .02
0 .00
0 .02
0 .04
0 .06
0 .08
0 .10
0 .12
0 .14
0 .16
0 .18
Ab
sorb
ance
1 300 1400 1500 1600 1700 1800
Wavenumbers (cm-1)
-0 .06
-0 .04
-0 .02
0 .00
0 .02
0 .04
0 .06
0 .08
0 .10
0 .12
0 .14
0 .16
0 .18
0 .20
0 .22
0 .24
0 .26
0 .28
0 .30
0 .32
Ab
sorb
ance
3 500 3600 3700 3800 3900 4000
Wavenumbers (cm-1)
Figure 11: DRIFTS UV spectra of all powders (1200-1800 cm-1). DRIFTSphotocatalytic experiments in the presence of oxygen before UV irradiation (a) andafter UV irradiation (b). Before irradiation, the surface is completely saturated withmethanol.
a
b
2ML CeO2-TiO2
1ML CeO2-TiO2
0.5ML CeO2-TiO2
0.25ML CeO2-TiO2
Bare-TiO2
2ML CeO2-TiO2
1ML CeO2-TiO2
0.5ML CeO2-TiO2
0.25ML CeO2-TiO2
Bare-TiO2
35
CHAPTER 5. FLOW REACTOR RESULTS
The flow reactor was an essential tool for the characterization of the reaction
products from the photocatalytic experiments. Due to the enhanced photo-catalytic
activity due to oxygen exposure found in Chapter 4, the flow reactor experiments focus
on those including oxygen.
Figure 12 shows one of the gas-phase infrared spectra from the flow reactor
experiment that includes contributions from all of the gaseous reaction products. From
the spectrum shown, the characteristic frequencies of each product can be observed. The
region from 2200-2400 cm-1 is characteristic of carbon dioxide, 1600-1900 cm-1 is
characteristic of formaldehyde and methyl formate, 1100-1250 cm-1 is characteristic of
methyl formate and dimethyl ether, and 900-1000 cm-1 is characteristic of methanol.
Water vapor is also a reaction product, but the water vapor frequencies overlap with those
of formaldehyde. To generate the plots in Figures 13-19, water subtraction was
performed on every tenth spectra throughout the reaction process. The integrated
absorbance was determined for each of these characteristic regions in the water
subtracted spectra to produce the data plots in Figures 13-19. The difference between a
spectrum with water and one without water can be seen in Figure 12.
In Figure 13, TiO2 is exposed to a combined methanol and oxygen flow of 14
sccm, in the absence of UV light As time passes, different reaction/desorption products
appear to form, but it appears that formaldehyde and methyl formate are not produced in
36
the absence of UV light. Similar results are described in Figure 14 for the 2 ML CeO2
modified sample.
Similar result are observed for the CeO2 modified surfaces. The amount of CeO2
modified powders used for these experiments is a factor of four larger than the amount of
TiO2 used for the same experiments due to the density of these powders being
approximately a factor of four larger than the bare TiO2 powders. Therefore, the
timescale for the CeO2 modified surfaces is approximately four times that of the TiO2
surfaces. In the absence of ultraviolet light, the desorbed products progress sequentially
as described for the TiO2 surface.
Much different results were observed when the materials were irradiated with
ultraviolet light. In Figure 15, the TiO2 surface exhibits consistent product formation. In
the first hour of methanol flow, carbon dioxide is produced in a large quantity, but
slightly decreases in production until asymptotically approaching a limit. Formaldehyde
is formed, but appears to slowly approach a maximum limit. At about the time methanol
is detected in the gas cell, methyl formate appears as well.
The CeO2 modified samples exhibit similar results with what appear to be slightly
different proportions of products desorbed. It appears the lowest weight loading, 0.25
monolayers, is the most similar to the bare TiO2 surface in that it seems more favorable to
formaldehyde production when exposed to UV light as seen in Figure 15. It appears that
the higher the CeO2 loading, the more susceptible these materials are to allowing
methanol passage uninhibited rather than converting to formaldehyde. It also appears
bare-TiO2 produces CO2 more consistently.
37
methanol
methyl formate
formaldehyde
carbon dioxide
-0 .45
-0 .40
-0 .35
-0 .30
-0 .25
-0 .20
-0 .15
-0 .10
-0 .05
0 .00
0 .05
0 .10
0 .15
0 .20
0 .25
0 .30
0 .35
Ab
sorb
ance
5 00 1000 1500 2000 2500 3000 3500
Wavenumbers (cm-1)
Figure 12: Gas phase spectra of photodesorption products with oxygen. The topspectrum includes water vapor whereas the bottom spectrum has water vaporsubtracted. In order to observe formaldehyde in the gas phase, water must besubtracted from the spectra.
38
-2
0
2
4
6
8
10
12
14
16
18
0 10 20 30 40 50 60 70 80 90
Inte
grat
ed A
rea
Time (min)
CH3OH
CH2O
500 1000 1500 2000 2500 3000 3500Wavenumbers (cm-1)
CO2
Figure 13: Flow Reactor Data: TiO2 with methanol and oxygen. Time lapse of themethanol reactions on the TiO2 powder. The powder is introduced to methanol andoxygen gas, each at a rate of 7 sccm for a cumulative flow of 14 sccm. Two graphsillustrate product formation. The top graph displays product formation over time as afunction of integrated peak area. The bottom graph displays a 3D representation of thespectra gathered over the course of the reaction.
Wavenumbers (cm-1)
39
-2
0
2
4
6
8
10
12
14
16
18
20
0 20 40 60 80 100 120 140 160
Inte
grat
ed A
rea
Time (min)
CO2
CH3OH
CH2O
500 1000 1500 2000 2500 3000 3500
Wavenumbers (cm-1)
Figure 14: Flow Reactor Data: 2ML CeO2-TiO2 with methanol and oxygen. Timelapse of the methanol reactions on the 2 ML CeO2 modified TiO2 powder. Thepowder is introduced to methanol and oxygen gas, each at a rate of 7 sccm for acumulative flow of 14 sccm. Two graphs illustrate product formation. The top graphdisplays product formation over time as a function of integrated peak area. Thebottom graph displays a 3D representation of the spectra gathered over the course ofthe reaction.
Wavenumbers (cm-1)
40
-5
0
5
10
15
20
25
30
0 50 100 150 200 250
Inte
grat
ed A
rea
Time (min)
CO2
CH2O
CH3OH
CHOOCH3
Figure 15: Flow Reactor Data: TiO2 with methanol and oxygen exposed to UV. Timelapse of the methanol reactions on TiO2 powder. While irradiated with 50 mW UVlight, the powder is introduced to methanol and oxygen gas, each at a rate of 7 sccmfor a cumulative flow of 14 sccm. Two graphs illustrate product formation. The topgraph displays product formation over time as a function of integrated peak area. Thebottom graph displays a 3D representation of the spectra gathered over the course ofthe reaction.
wavenumbers (cm-1)
41
0
5
10
15
20
25
30
35
0 50 100 150 200 250
Inte
grat
ed A
rea
Time (min)
CHOOCH3
CH2OCH3OH
CO2
Figure 16: Flow Reactor Data: 0.25 ML CeO2-TiO2 with methanol and oxygenexposed to UV. Time lapse of the methanol reactions on the 0.25 ML CeO2 modifiedTiO2 powder. While irradiated with 50 mW UV light, the powder is introduced tomethanol and oxygen gas, each at a rate of 7 sccm for a cumulative flow of 14 sccm.Two graphs illustrate product formation. The top graph displays product formationover time as a function of integrated peak area. The bottom graph displays a 3Drepresentation of the spectra gathered over the course of the reaction.
wavenumbers (cm-1)
42
-5
0
5
10
15
20
25
30
0 50 100 150 200 250
Inte
grat
ed A
rea
Time (min)
CO2
CH2O
CH3OH
CHOOCH3
Figure 17: Flow Reactor Data: 0.5 ML CeO2-TiO2 with methanol and oxygen exposedto UV. Time lapse of the methanol reactions on the 0.5 ML CeO2 modified TiO2
powder. While irradiated with 50 mW UV light, the powder is introduced to methanoland oxygen gas, each at a rate of 7 sccm for a cumulative flow of 14 sccm. Twographs illustrate product formation. The top graph displays product formation overtime as a function of integrated peak area. The bottom graph displays a 3Drepresentation of the spectra gathered over the course of the reaction.
wavenumbers (cm-1)
43
-5
0
5
10
15
20
25
30
35
0 50 100 150 200 250
Inte
grat
ed A
rea
Time (min)
CO2
CH2O
CH3OH
CHOOCH3
Figure 18: Flow Reactor Data: 1 ML CeO2-TiO2 with methanol and oxygen exposedto UV. Time lapse of the methanol reactions on the 1 ML CeO2 modified TiO2
powder. While irradiated with 50 mW UV light, the powder is introduced to methanoland oxygen gas, each at a rate of 7 sccm for a cumulative flow of 14 sccm. Twographs illustrate product formation. The top graph displays product formation overtime as a function of integrated peak area. The bottom graph displays a 3Drepresentation of the spectra gathered over the course of the reaction.
wavenumbers (cm-1)
44
-5
0
5
10
15
20
0 50 100 150 200 250
Inte
grat
ed A
rea
Time (min)
CO2
CH2O
CH3OH
CHOOCH3
f
Figure 19: Flow Reactor Data: 2 ML CeO2-TiO2 with methanol and oxygen exposedto UV. Time lapse of the methanol reactions on the 2 ML CeO2 modified TiO2
powder. While irradiated with 50 mW UV light, the powder is introduced to methanoland oxygen gas, each at a rate of 7 sccm for a cumulative flow of 14 sccm. Twographs illustrate product formation. The top graph displays product formation overtime as a function of integrated peak area. The bottom graph displays a 3Drepresentation of the spectra gathered over the course of the reaction.
wavenumbers (cm-1)
45
CHAPTER 6. DISCUSSION
The characterization results provide information about the loading of cerium on
the materials and the nature of the cerium on the surface as a way of comparing and
contrasting the impact of the cerium loading. Knowing exactly how much CeO2
deposited on the surface, and the XRD patterns, we can speculate about the nature of the
effect of the cerium. The XRD patterns reveals that CeO2 crystallites form on the 1 and 2
monolayer surfaces. The x-ray results do not support the formation of crystallites for the
0.25 and 0.5 monolayer surfaces. For the 0.25 and 0.5 monolayer surfaces, this data
indicates that CeO2 deposits in an amorphous fashion where the CeO2 appears to be well
dispersed on the surface. For the 1 monolayer surface, a small, broad peak characteristic
of CeO2 crystallites was observed, indicating that 3-dimmensional crystals form
minimally on this surface. This indicates that CeO2 is almost completely dispersed over
the surface, which could nearly occupy all of the hydroxyl sites on the TiO2 surface. The
2 monolayer surface shows significant crystallization, indicating that as CeO2 is
deposited in quantities greater than 1 monolayer, CeO2 domains form. However, even at
2 monolayers, and with the formation of these domains, some TiO2 sites could remain
exposed for reaction with methanol.
The DRIFTS experiments performed provided information regarding the
chemistry occurring on the surfaces. In Figure 9a, it appears that samples with higher
weight loadings of CeO2 have a lower affinity for adsorption of molecular methanol than
46
the lower weight loading samples. There is no increase in methoxy formation observed.
An increase in methoxy formation would indicate that the CeO2 modified surfaces
convert more methanol to methoxy. However, it simply appears that methanol does not
physisorb on the modified surface. This may be due to CeO2 occupying surface sites on
the TiO2 surface that are the preferred sites for methanol physisorption. It is observed that
sites occupied by methoxy are not altered because the initial methoxy peak intensity is
very similar for all samples, regardless of cerium loading.
In Figure 10a, interesting changes in the hydroxyl absorptions are observed. There
is a significant loss of intensity in the hydroxyl region with the introduction of methanol.
A negative peak is observed at 3650 cm-1 on the bare TiO2 surface that is not observed on
the CeO2 modified surfaces when methanol is introduced. This frequency is
representative of acidic hydroxyl groups that appear to be present on the TiO2 surface
before methanol exposure, but not present on the CeO2 modified samples under the same
conditions. This could indicate that CeO2 occupies surface sites that are normally
occupied by acidic surface hydroxyls. These acidic surface hydroxyl groups could also be
the preferred sites for methanol physisorption. Upon UV irradiation, Figure 10b shows
that surface hydroxyl groups seem to return, but return at a higher frequency, which
indicates these new hydroxyl groups are more basic in nature and that the new surface
structure is different from the structure before adsorption and decomposition.
Before UV irradiation, Figure 11a gives two peaks at 1440cm-1 and 1360cm-1
which correspond to chemisorbed methoxy and physisorbed methanol respectively. The
peak at 1360 cm-1 is smaller on samples with CeO2 until it disappears on the 2 ML CeO2-
TiO2 sample. This indicates once more that methanol does not physisorb onto the CeO2
47
modified samples. After UV irradiation, Figure 11b indicates the formation of surface
adsorbed formate, which is consistent with prior studies as discussed in Chapter 1.
The flow reactor system was essential for the characterization of desorption
products. In Figure 12, methanol and oxygen are allowed to flow over TiO2 without UV
light, carbon dioxide is desorbed from the surface after just 5 minutes of methanol flow.
However, carbon dioxide production stops after approximately 30 minutes. When
methanol reacts with the surface to form chemisorbed methoxy, water vapor is formed.
Water begins desorbing from the surface after 10 minutes of methanol flow, which is the
product of methanol chemically adsorbing to the surface as methoxy.33 After 15 minutes
of methanol flow, methanol begins to appear in the infrared spectra, indicating there are
no other sites available on the TiO2 powder for methanol to adsorb. The reaction is
allowed to continue until the intensity of the characteristic methanol peak reaches a
constant maximum, which indicates gaseous methanol is not interacting with the surface.
Similar results are observed for the 2 ML CeO2 modified surface as described in Figure
13. The same products are observed to desorb from the CeO2 modified surface in the
same fashion. The only difference is the onset time when each product is detected. This is
due to using more solid for the CeO2 modified samples than for the TiO2 samples. The
density of the modified materials is much greater than the bare-TiO2, therefore we
increased the amount of material used to achieve approximately equal volumes of powder
per experiment. Equal volumes of material were needed so that there would be limited
difference in reactions due to methanol passing through the sample uninhibited.
These results are different when UV light is introduced. In Figure 14, TiO2 is
exposed to methanol and oxygen while being irradiated with UV light. There is an initial
48
large spike in carbon dioxide production which slows in production when water and
formaldehyde are observed. Methanol and methyl formate are observed to form shortly
thereafter. The initial spike in carbon dioxide production is most likely due to a reaction
with chemisorbed methoxy and water vapor. Previous studies suggest that methoxy
adsorbed to TiO2 reacts with water in the following method when irradiated with UV
light.33
When methanol becomes detectable, it indicates that the surface concentration of
methoxy is high. Similar results are observed for the CeO2 modified samples as well. In
Figures 15-29, the same desorption products are observed and they desorb in the same
way. The difference appears that the higher the weight loading of CeO2, the higher the
affinity to allow methanol to pass through the surface unhindered.
CH3OH(g) + OH(ad) → CH3O(ad) + H2O(g) (10)
CH3O(ad) + H2O(g) → CH2O2(ad) + H2(g) (11)
CH2O2(ad) → CO2(g) + H2(g) (12)
49
CHAPTER 7. CONCLUSION
The purpose of these experiments were to determine if the photocatalytic activity
of TiO2 could be increased by impregnating the surface with CeO2. Four samples were
prepared consisting of different amounts of CeO2 deposited on TiO2. Characterization
results revealed that at high weight loadings, CeO2 forms crystallite domains on the
surface whereas at lower loadings, CeO2 deposits in a more amorphous fashion. It
appears that CeO2 encourages desorption of methoxy from the surface during UV
irradiation as long as TiO2 sites are still exposed. For instance, the 1 monolayer (1 ML)
sample presented the lowest desorption activity, even compared to TiO2. The 1 ML
sample also shows evidence in the X-ray diffraction data to suggest the CeO2 deposits
evenly over the CeO2 surface with minimal crystallization. This indicates that it is likely
that most TiO2 hydroxyl sites are not availablev.
Although the CeO2 modified samples show a higher affinity for methoxy
desorption, it appears that TiO2 converts more methoxy to formaldehyde than the CeO2
modified samples. Rather, it appears that CeO2 modified samples will produce
formaldehyde, but will also react with methoxy to desorb molecular methanol. Other
products observed are carbon dioxide, water, and dimethyl ether.
Future work should consider using DRIFTS with a quadrupole mass spectrometer
in order to further characterize the reaction products. It has been reported in prior studies
50
that hydrogen gas is produced in small quantities by the photocatalytic decomposition of
methanol on TiO2. Because hydrogen does not have a dipole moment, it is not detectable
using infrared spectroscopy. Also, standard infrared absorbance spectra should be
obtained in the flow reactor system for the various products formed from these reactions.
This would provide necessary data to construct a Beer’s law plot to quantify the
desorption products.
Interesting results could be discerned from analyzing the modified samples
through scanning electron microscopy and elemental diffraction x-ray spectroscopy
(SEM-EDX) with which elemental maps from the images could be obtained. This would
confirm or refute our suspicions of how well dispersed the CeO2 deposits on the surface.
51
REFERENCES
1. Thompson, T. L.; Yates, J. T. Surface science studies of the photoactivation of TiO2new photochemical processes. Chemical Reviews 2006, 106, 4428-4453.
2. Wachs, I. E.; Madix, R. J. The oxidation of methanol on a silver (110) catalyst. SurfaceScience 1978, 76, 531-558.
3. Wachs, I. E.; Madix, R. J. The selective oxidation of CH3OH to H2CO on acopper(110) catalyst. Journal of Catalysis 1978, 53, 208-227.
4. Fujishima, A.; Honda, K. Electrochemical photolysis of water at a semiconductorelectrode. Nature 1972, 238, 37-38.
5. Fujishima, A.; Rao, T. N.; Tryk, D. A. Titanium dioxide photocatalysis. Journal ofPhotochemistry and Photobiology C: Photochemistry Reviews 2000, 1, 1-21.
6. Fujishima, A.; Zhang, X. Titanium dioxide photocatalysis: present situation and futureapproaches. Comptes Rendus Chimie 2006, 9, 750-760.
7. Linsebigler, A. L.; Lu, G.; Yates Jr, J. T. Photocatalysis on TiO2 surfaces: principles,mechanisms, and selected results. Chemical Reviews 1995, 95, 735-758.
8. Berger, T.; Sterrer, M.; Diwald, O.; Knözinger, E.; Panayotov, D.; Thompson, T. L.;Yates, J. T. Light-induced charge separation in anatase TiO2 particles. The Journalof Physical Chemistry B 2005, 109, 6061-6068.
9. Ohno, T.; Sarukawa, K.; Tokieda, K.; Matsumura, M. Morphology of a TiO2Photocatalyst (Degussa, P-25) Consisting of Anatase and Rutile Crystalline Phases.Journal of Catalysis 2001, 203, 82-86.
10. Frank, S. N.; Bard, A. J. Heterogeneous photocatalytic oxidation of cyanide ion inaqueous solutions at titanium dioxide powder. Journal of the American ChemicalSociety 1977, 99, 303-304.
11. Bickley, R.; Jayanty, R. Photo-adsorption and photo-catalysis on titanium dioxidesurfaces. Photo-adsorption of oxygen and the photocatalyzed oxidation ofisopropanol. Faraday Discussions Chemical Society 1974, 58, 194-204.
12. Bickley, R. I.; Gonzalez-Carreno, T.; Lees, J. S.; Palmisano, L.; Tilley, R. J. D. Astructural investigation of titanium dioxide photocatalysts. Journal of Solid StateChemistry 1991, 92, 178-190.
13. Hoffmann, M. R.; Martin, S. T.; Choi, W.; Bahnemann, D. W. Environmentalapplications of semiconductor photocatalysis. Chemical Reviews 1995, 95, 69-96.
14. Kumar, S. G.; Devi, L. G. Review on modified TiO2 photocatalysis under UV/visiblelight: selected results and related mechanisms on interfacial charge The Journal ofPhysical Chemistry A 2011, 115, 13211-13241
15. Matthews, R. W. Kinetics of photocatalytic oxidation of organic solutes over titaniumdioxide. Journal of Catalysis 1988, 111, 264-272.
16. Shen, S.; Wang, X.; Chen, T.; Feng, Z.; Li, C. Transfer of Photoinduced Electrons inAnatase–Rutile TiO2 Determined by Time-Resolved Mid-Infrared Spectroscopy.The Journal of Physical Chemistry C 2014, 118, 12661-12668.
52
17. Thompson, T. L.; Yates Jr, J. T. TiO2-based photocatalysis: Surface defects, oxygenand charge transfer. Topics in Catalysis 2005, 35, 197-210.
18. Zhang, Z.; Yates Jr, J. T. Direct Observation of Surface-Mediated Electron− HolePair Recombination in TiO2 (110). The Journal of Physical Chemistry C 2010, 114,3098-3101.
19. Deiana, C.; Fois, E.; Coluccia, S.; Martra, G. Surface structure of TiO2 P25nanoparticles: infrared study of hydroxy groups on coordinative defect sites. TheJournal of Physical Chemistry C 2010, 114, 21531-21538.
20. Alberici, R. M.; Jardim, W. F. Photocatalytic destruction of VOCs in the gas-phaseusing titanium dioxide. Applied Catalysis B: Environmental 1997, 14, 55-68.
21. Augugliaro, V.; Palmisano, L.; Sclafani, A.; Minero, C.; Pelizzetti, E. Photocatalyticdegradation of phenol in aqueous titanium dioxide dispersions. Toxicological &Environmental Chemistry 1988, 16, 89-109.
22. Carey, J. H.; Lawrence, J.; Tosine, H. M. Photodechlorination of PCB's in thepresence of titanium dioxide in aqueous suspensions. Bulletin of the EnvironmentContamination Toxicology 1976, 16, 697-701.
23. Huang, Z.; Maness, P.; Blake, D. M.; Wolfrum, E. J.; Smolinski, S. L.; Jacoby, W. A.Bactericidal mode of titanium dioxide photocatalysis. Journal of Photochemistry andPhotobiology A. 2000, 130, 163-170.
24. Ireland, J. C.; Klostermann, P.; Rice, E. W.; Clark, R. M. Inactivation of Escherichiacoli by titanium dioxide photocatalytic oxidation. Applied and EnvironmentalMicrobiology 1993, 59, 1668-1670.
25. Konstantinou, I. K.; Albanis, T. A. Photocatalytic transformation of pesticides inaqueous titanium dioxide suspensions using artificial and solar light: intermediatesand degradation pathways. Applied Catalysis B: Environmental 2003, 42, 319-335.
26. Maira, A. J.; Coronado, J. M.; Augugliaro, V.; Yeung, K. L.; Conesa, J. C.; Soria, J.Fourier Transform Infrared Study of the Performance of Nanostructured TiO2Particles for the Photocatalytic Oxidation of Gaseous Toluene. Journal of Catalysis2001, 202, 413-420.
27. Matthews, R. W. Purification of water with near—u.v. illuminated suspensions oftitanium dioxide. Water Research 1990, 24, 653-660.
28. Prairie, M. R.; Evans, L. R.; Stange, B. M.; Martinez, S. L. An investigation oftitanium dioxide photocatalysis for the treatment of water contaminated with metalsand organic chemicals. Environmental Science and Technology 1993, 27, 1776-1782.
29. Rusu, C. N.; Yates, J. T. Adsorption and decomposition of dimethylmethylphosphonate on TiO2. The Journal of Physical Chemistry B 2000, 104,12292-12298.
30. Schrauzer, G. N.; Guth, T. D. Photocatalytic reactions. 1. Photolysis of water andphotoreduction of nitrogen on titanium dioxide. The Journal of the AmericanChemical Society 1977, 99, 7189-7193.
31. Al-Mazroai, L. S.; Bowker, M.; Davies, P.; Dickinson, A.; Greaves, J.; James, D.;Millard, L. The photocatalytic reforming of methanol. Catalysis Today 2007, 122,46-50.
53
32. Aurian-Blajeni, B.; Halmann, M.; Manassen, J. Photoreduction of carbon dioxide andwater into formaldehyde and methanol on semiconductor materials. Solar Energy1980, 25, 165-170.
33. Chiarello, G. L.; Aguirre, M. H.; Selli, E. Hydrogen production by photocatalyticsteam reforming of methanol on noble metal-modified TiO2. Journal of Catalysis2010, 273, 182-190.
34. Chuang, C.; Chen, C.; Lin, J. Photochemistry of methanol and methoxy groupsadsorbed on powdered TiO2. The Journal of Physical Chemistry B 1999, 103, 2439-2444.
35. Jacox, M. E.; Milligan, D. E. Matrix isolation study of the vacuum-ultravioletphotolysis of methanol: The infrared spectrum of the CH2OH free radical. Journalof Molecular Spectroscopy 1973, 47, 148-162.
36. Park, M.; Hauge, R. H.; Kafafi, Z. H.; Margrave, J. L. Activation of O–H and C–Obonds of methanol by photoexcited iron atoms. Journal of the Chemical Society,Chemical Communications 1985, 1570-1571.
37. Shen, M.; Henderson, M. A. Identification of the active species in photochemical holescavenging reactions of methanol on TiO2. The Journal of Physical ChemistryLetters 2011, 2, 2707-2710.
38. Siokou, A.; Nix, R. M. Interaction of Methanol with Well-Defined Ceria Surfaces:Reflection/Absorption Infrared Spectroscopy, X-ray Photoelectron Spectroscopy,and Temperature-Programmed Desorption Study. Journal of Physical Chemistry B1999, 103, 6984-6997
39. Taylor, E.; Griffin, G. Product selectivity during methanol decomposition on titaniapowders. Journal of Physical Chemistry 1988, 92, 477-481.
40. Zhang, J.; Nosaka, Y. Photocatalytic oxidation mechanism of methanol and the otherreactants in irradiated TiO2 aqueous suspension investigated by OH radicaldetection. Applied Catalysis B-Environmental 2015, 166, 32-36.
41. Carrizosa, I.; Munuera, G.; Castañar, S. Study of the interaction of aliphatic alcoholswith TiO2: III. Formation of alkyl-titanium species during methanol decomposition.Journal of Catalysis 1977, 49, 265-277.
42. Phillips, K. R.; Jensen, S. C.; Baron, M.; Li, S.; Friend, C. M. Sequential photo-oxidation of methanol to methyl formate on TiO2 (110). Journal of the AmericanChemical Society 2013, 135, 574-577.
43. Busca, G.; Lamotte, J.; Lavalley, J. C.; Lorenzelli, V. FT-IR study of the adsorptionand transformation of formaldehyde on oxide surfaces. Journal of the AmericanChemical Society 1987, 109, 5197-5202.
44. Veréb, G.; Ambrus, Z.; Pap, Z.; Kmetykó, Á.; Dombi, A.; Danciu, V.; Cheesman, A.;Mogyorósi, K. Comparative study on UV and visible light sensitive bare and dopedtitanium dioxide photocatalysts for the decomposition of environmental pollutants inwater. Applied Catalysis A: General 2012, 417–418, 26-36.
45. Chiang, K.; Amal, R.; Tran, T. Photocatalytic degradation of cyanide using titaniumdioxide modified with copper oxide. Advances in Environmental Research 2002, 6,471-485.
46. Diwell, A. F.; Rajaram, R. R.; Shaw, H. A.; Truex, T. J. The Role of Ceria in Three-Way Catalysts. Studies in Surface Science and Catalysis 1991, 71, 139-152.
54
47. Idriss, H.; Diagne, C.; Hindermann, J. P.; Kiennemann, A.; Barteau, M. A. Reactionsof Acetaldehyde on CeO2 and CeO2-Supported Catalysts. Journal of Catalysis1995, 155, 219-237.
48. Trovarelli, A.; de Leitenburg, C.; Boaro, M.; Dolcetti, G. The utilization of ceria inindustrial catalysis. Catalysis Today 1999, 50, 353-367.
49. Bunluesin, T.; Gorte, R. J.; Graham, G. W. Studies of the water-gas-shift reaction onceria-supported Pt, Pd, and Rh: Implications for oxygen-storage properties. AppliedCatalysis B: Environmental 1998, 15, 107-114.
50. Li, Y.; Fu, Q.; Flytzani-Stephanopoulos, M. Low-temperature water-gas shift reactionover Cu- and Ni-loaded cerium oxide catalysts. Applied Catalysis B: Environmental2000, 27, 179-191.
51. Cai, T.; Liao, Y.; Peng, Z.; Long, Y.; Wei, Z.; Deng, Q. Photocatalytic performanceof TiO2 catalysts modified by H3PW12O40, ZrO2 and CeO2. Journal ofEnvironmental Sciences 2009, 21, 997-1004.
52. Liu, B.; Zhao, X.; Zhang, N.; Zhao, Q.; He, X.; Feng, J. Photocatalytic mechanism ofTiO2–CeO2 films prepared by magnetron sputtering under UV and visible light.Surface Science 2005, 595, 203-211.
53. Watanabe, S.; Ma, X.; Song, C. Characterization of structural and surface propertiesof nanocrystalline TiO2− CeO2 mixed oxides by XRD, XPS, TPR, and TPD. TheJournal of Physical Chemistry C 2009, 113, 14249-14257.
54. Yang, H.; Zhang, K.; Shi, R.; Tang, A. Sol–gel synthesis and photocatalytic activityof CeO2/TiO2 nanocomposites. Journal of the American Ceramic Society 2007, 90,1370-1374.
55. Shin, M.; Cha, J.; Shin, B.; Chun, H. H.; Lee, H. Activities of CeO2/TiO2 catalyst forSCR of NO with NH3 at low temperature according to operating conditions.Electronic Materials Letters 2013, 9, 71-76.
56. Mueller, R.; Kammler, H. K.; Wegner, K.; Pratsinis, S. E. OH surface density of SiO2and TiO2 by thermogravimetric analysis. Langmuir 2003, 19, 160-165.
57. Fuller, M. P.; Griffiths, P. R. Diffuse reflectance measurements by infrared Fouriertransform spectrometry. Analytical Chemistry 1978, 50, 1906-1910.
58. El-Roz, M.; Bazin, P.; Thibault-Starzyk, F. An operando-IR study of photocatalyticreaction of methanol on new* BEA supported TiO 2 catalyst. Catalysis Today 2013,205, 111-119.