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

1360

1440

2826

2852

2926

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2820

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

0 .20

0 .25

0 .30

Ab

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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

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