Chemical Doping of Metal Oxide Nanomaterials And

CHEMICAL DOPING OF METAL OXIDE NANOMATERIALS AND

CHARACTERIZATION

OF THEIR PHYSICAL-CHEMICAL PROPERTIES

JUNWEI WANG

Submitted in partial fulfillment of the requirements

for the degree of Doctor of Philosophy

Dissertation Advisor: Dr. Clemens Burda

Department of Chemistry

Case Western Reserve University

May, 2012

CASE WESTERN RESERVE UNIVERSITY

SCHOOL OF GRADUATE STUDIES

We hereby approve the thesis/dissertation of

______Junwei Wang______

Candidate for the ______PhD______degree*.

(signed) ______Carlos E. Crespo Hernández______(chair of the committee)

______Alfred B. Anderson______

______John E. Stuehr______

______Jeffery Dyck______

______Clemens Burda______

______

(date)______3/28/2012______

*We also certify that written approval has been obtained for any proprietary material contained therein. TABLE OF CONTENTS

TABLE OF CONTENTS ...... 1

LIST OF TABLES ...... 4

LIST OF FIGURES ...... 5

ACKNOWLEDGEMENTS ...... 13

LIST OF ABBREVIATIONS ...... 14

ABSTRACT ...... 16

CHAPTER 1 INTRODUCTION ...... 18

1.1 MOTIVATION: ENERGY AND ENVIRONMENT CRISIS ...... 18

1.2 POTENTIAL SOLUTIONS: PHOTOCATALYSIS ...... 20

1.3 CURRENT RESEARCH ON METAL OXIDES AS PHOTOCATALYSTS ...... 24

1.4 ABOUT THIS DISSERTATION ...... 28

1.5 REFERENCES ...... 30

CHAPTER 2 HYDROTHERMAL NITROGEN-DOPED TITANIUM DIOXIDE WITH

SWITCHABLE SURFACE WETTABILITY ...... 31

2.1 INTRODUCTION ...... 31

2.2 EXPERIMENTAL PROCEDURES ...... 32

2.3 RESULTS AND DISCUSSION...... 34

2.4 CONCLUSION ...... 41

2.5 REFERENCES ...... 42

CHAPTER 3 MODIFYING TITANIUM DIOXIDE WITH TUNGSTEN...... 44

3.1 INTRODUCTION ...... 44

3.2 EXPERIMENTAL PROCEDURES ...... 45

3.3 RESULTS AND DISCUSSION...... 47

3.4 CONCLUSION ...... 52

3.5 REFERENCES ...... 53

CHAPTER 4 A COMPARATIVE STUDY OF COBALT ION DOPING OF TIO2 AND

TIO2-XNX AT THE NANOSCALE: EVIDENCE FOR ROOM TEMPERATURE

CONVERSION FROM ANATASE TO RUTILE PHASE ...... 54

4.1 INTRODUCTION ...... 54

4.2 EXPERIMENTAL PROCEDURES ...... 55

4.3 RESULTS ...... 56

4.4. DISCUSSION ...... 71

4.5 REFERENCES ...... 76

CHAPTER 5 FACILE CHARGE TRANSFER, PHASE TRANSFORMATION, AND

THE ANALOG OF NITROGEN FIXATION IN HEAVILY IRON ION DOPED

TITANIUM OXIDE AND OXYNITRIDE NANOCOLLOIDS ...... 79

5.1 INTRODUCTION ...... 79

5.2 EXPERIMENTAL PROCEDURES ...... 80

5.3 RESULTS ...... 83

5.4 DISCUSSION ...... 106

5.5 CONCLUSION ...... 110

5.6 REFERENCES ...... 111

CHAPTER 6 SYNTHESIS AND CHARACTERIZATION OF NANO-SCALED

GALLIUM, NITROGEN CO-DOPED TITANIUM DIOXIDE WITH HIGH

NITROGEN CONCENTRATION...... 114

6.1 INTRODUCTION ...... 114

6.2 EXPERIMENTAL PROCEDURES ...... 115

6.3 RESULTS ...... 117

6.4 CONCLUSION ...... 126

6.5 REFERENCES ...... 127

CHAPTER 7 INTERACTIVE IRON-SILICON OXIDATION/REDUCTION

PROCESSES ON FUMED SILICA ...... 129

7.1 INTRODUCTION ...... 129

7.2 EXPERIMENTAL PROCEDURES ...... 130

7.3 RESULTS ...... 131

7.4 CONCLUSION ...... 146

7.5 REFERENCES ...... 148

CHAPTER 8 SUMMARY AND OUTLOOK ...... 149

BIBLIOGRAPHY ...... 151

LIST OF TABLES

Table 2.1 Relative XRD peak intensities (at 2θ = 25.2º, 53.8º and 55.0º) for the prepared samples and standard anatase TiO2 (For comparison, the intensities of the peak at

2θ = 55.0º are set to be 1.0)...... 36

Table 2.2 Contact angle measurements for the N-doped TiO2 samples before and after visible-light illumination ...... 39

Table 4.1 Composition of the prepared samples ...... 57

Table 5.2 2s and 2p Core level XPS Binding Energies as Function of Fe2+ Doping Level

(eV) - for several mole fractions (peak intensities for 2p3/2 (2p3) and 2p1/2 (2p1) are indicated) ...... 88

Table 5.3 Calculated 2s and 2p Binding Energies in eV...... 89

Table 6.1 Particle size calculated based on Scherrer equation...... 120

Table 6.2 Percentage and Binding Energies of N1s in the samples based on XPS analysis.

...... 122

Table 7.1 XPS de-convoluted peaks for Fe Cab samples using Gaussian functions ..... 135

Table 7.2 XPS results of the Silicon 2p in the Fe-CAB-O-SIL samples ...... 136

Table 7.3 Predicted and Observed Fe(II) and Fe(III) amounts on the Cab-O-Sil Surface.

...... 140

Table 7.4 Detailed analyses of IR peaks...... 143

LIST OF FIGURES

Figure 1.1 Crude oil prices from January 2nd, 1986 to March 6th, 2012. Data from

http://www.eia.gov/dnav/pet/hist/LeafHandler.ashx?n=PET&s=RWTC&f=A

...... 19

Figure 1.2 Average gasoline prices in US from 1986 to 2011...... 20

Figure 1.3 Illustration of photocatalytic water splitting[9]...... 22

Figure 1.4 Solar light energy distribution by wavelength[10]...... 23

Figure 1.5 Suggested pathways for photocatalytic decomposition of organic dyes. [12] . 24

Figure 2.1 XRD patterns of N-doped TiO2 prepared by a hydrothermal method. Shown in

the figure are samples with growing times ranging from 6 to 24 hours. In the

bottom of the figure are the standard XRD peaks for anatase TiO2 (pink,

JCPDS file NO. 21-1272) and brookite TiO2 (wine, JCPDS file NO.03-0380).

...... 34

Figure 2.2 SEM images of N-doped TiO2 samples after the indicated reaction time of 6,

12, 18 and 24 hours...... 35

Figure 2.3 UV-vis diffuse reflection spectra of N-doped TiO2 prepared by hydrothermal

method. Shown in the figure are the curves for 4 different growing times,

namely, 6, 12, 18 and 24 hours...... 38

Figure 2.4 Photographs of the water droplets on an N-doped TiO2 film with various

surface wettabilities...... 38

Figure 2.5 Contact angle measurements of the 24-hour N-TiO2 sample in 3 hydrophobic-

hydrophilic cycles (in air-exposed system). Values shown are averages from

three different measurements. It is noteworthy to point out that the same

illumination on the surfaces that are kept in vacuum or argon environment

did not result in the wettability change at all, indicating the important role of

water in the wettability transition process...... 40

Figure 2.6 Visible-light induced photocatalysis of methylene blue by nitrogen-doped TiO2

with various surface wettabilities. The numeric values of these surface

contact angles are included in the figure legend. For comparison, the

photocatalytic performance of Degussa P25 (a commercial available TiO2

product) under the same condition is also shown...... 41

Figure 3.1 Bands alignment of TiO2 (anatase) and WO3 (monoclinic)...... 45

Figure 3.2 TEM images of various samples prepared by single sol-gel method. The

uncalcined TiO2 is shown in negative...... 47

Figure 3.3 XRD patterns for TiO2 and TiO2 with 1 and 2 layers WO3 coating. Letter “A”

indicates the position of antase TiO2...... 48

Figure 3.4 UV-Vis diffuse reflection spectra of the samples made by single sol-gel

method...... 48

Figure 3.5 TEM images of various samples prepared by double sol-gel method...... 49

Figure 3.6 UV-Vis diffuse reflection spectra of the samples made by double sol-gel

method...... 50

Figure 3.7 SEM image of WO3 loaded TiO2 nanotube in low (13.0K) and high (100K)

magnification. The section view of TiO2 nanotube was also given (right) for

comparison...... 50

Figure 3.8 XRD patterns for the TiO2 nanotube with and without WO3 loading. The dots

in the figure correspond to the characterization peak position of metal

titanium. All other peaks correspond to anatase TiO2...... 51

Figure 3.9 UV-Vis diffuse reflection spectra of TiO2 nanotubes with and without WO3

loading...... 52

Figure 4.1 Raman spectra for a TiO2 nanocolloid (<10 nm) prepared with various

2+ concentrations of Co using (A) CoCl2 and (B) Co(NO3)2 at room

temperature. The Raman signal was obtained using a 1 μm spot size and a

power of 25 mW or less.[19,22] The typical vibration modes of rutile TiO2

are indicted by solid vertical lines...... 58

2+ Figure 4.2 Raman spectra of TiO2-xNx nanocolloid with various Co concentrations.

Laser power is less than 10 mW. [22] All three samples exhibit characteristic

2+ rutile TiO2 peaks and the intensities of the peaks increase with Co

concentration...... 59

Figure 4.3 Overall Core level XPS spectra for selected samples as indicated in the Figure.

...... 62

Figure 4.4 Ti 2p XPS peaks for the CoCl2 / TiO2 and CoCl2 /TiON systems...... 63

Figure 4.5 N 1s XPS spectra for CoCl2 seeded TiO2 and TiO2-xNx...... 65

Figure 4.6 VB-XPS of TiO2 with different CoCl2 loading. For comparison, the VB-XPS

spectra of TiO2 and CoCl2 are also shown...... 66

Figure 4.7 Room temperature VB XPS spectra of CoCl2 / TiO2-xNx system with

2+ increased Co concentration. For comparison, the same spectra of TiO2-xNx

and CoCl2 are also included...... 67

Figure 4.8 TEM images of (a) CoCl2 / TiO2-xNx and (b) CoCl2 / TiO2 ...... 69

Figure 4.9 Reflectance Spectra of P25, cobalt doped TiO2 and TiON nanoparticles with

varying Co content as indicated in the legend. Comparing to P25, the cobalt

doped samples, even with a Co content as low as 0.025 g, exhibit a sharp

difference in the visible-light region. The features associated with Co are

more damped in the TiON matrix than those in the TiO2 matrix. The

measurements are done relative to MgO, which explains the >100%

reflectance of pure P25 TiO2...... 70

Figure 4.10 Photodecomposition of methylene blue under visible light (> 400 nm)

illumination with various photocatalysts...... 71

Table 5.1 Composition of the Prepared Fe2+ Doped Samples (mole % of Fe2+ in total

II composition of Fe + TiO2 or TiO2-xNx ) ...... 84

Figure 5.1 (a) Fe doped TiO2 core-level XPS spectra of all samples -VB-XPS spectra for

Fe doped TiO2 compared to TiO2 samples. (b) VB-XPS spectra for Fe doped

TiO2 compared to TiO2 and TiON samples. (c) Ti 2p core-level XPS spectra

for Fe doped TiO2. (d) Fe 2p core-level XPS spectra for Fe doped TiO2. (e)

O 1s core level XPS spectra for Fe doped TiO2. (f) N 1s core-level XPS

spectra for Fe doped TiO2...... 85

Figure 5.2 (a) Fe doped TiON core-level XPS spectra of all samples. (b) VB-XPS spectra

for Fe doped TiON compared to TiO2 and TiON. (c) Ti 2p core level XPS

spectra for Fe doped TiON. (d) Fe 2p3 core-level spectra for Fe doped TiON.

(e) O 1s core-level XPS spectra for Fe doped TiON. (f) N 1s core-level XPS

spectra for Fe doped TiON...... 86

Figure 5.3 (a) Raman spectra for FeCl2 doped TiO2 where the doping is done with FeCl2-4

H2O (concentrations correlating with Table 5.1 are indicated in the figure (b),

Raman spectra for nitrogen doped titanium oxynitride (from Ref. 1), (c)

Raman spectra for FeCl2 doped TiO2 where the doping is done initially with

anhydrous FeCl2 (concentrations correlating with Table 5.1 are given in the

figure)...... 93

Figure 5.4 TEM micrographs of FeCl2 doped TiO2. Figures (a) in broad (scale bar 0.5 µm)

and finer (200 nm) view demonstrate the apparent formation of crystallites

on the TiO2 surface and correspond to a doping level of 0.94 mole% (Table

5.1). Figure.(b) and (corresponds to a typical section showing a much more

pronounced agglomeration of the Fe in the TiO2 (scale bar 200 nm) and

correspond to a doping level of 3.67 mole% (Table 5.1)...... 96

Figure 5.5 (a) HRTEM image of Fe seeded TiO2 (1.87 mole %-Table 5.1) showing d

values, which are an average of 10 spacings. (b) Enlarged HRTEM image,

indicating a spacing of 0.35 nm, indicative of the TiO2 anatase 101 plane.

Scale bars are 5nm in (a) and –nm in (b)...... 97

Figure 5.6 EDX spectra for TiO2 – Fe (1.87 mole %). (a) Images 1 and 2 of two different

crystallites. (b) and (c) EDX spectra for the two regions showing virtually

identical titanium and oxygen concentrations. The Cu signal corresponds to

the TEM grid; adventitious carbon is also present. (d) Enlarged figure of the

EDX spectra showing the small iron concentration in the sample...... 99

Figure 5.7 TEM micrographs of FeCl2 doped TiON in two regions of the sample. Figure

(a) demonstrates the apparent formation of small iron particles on the TiO2

surface and corresponds to a doping level of 7.07 mole% (Table 5.1). Figure

(b) demonstrates a much more pronounced Fe-based particle formation on

the TiO2 surface and corresponds to a doping level of 37.9 mole% (Table

5.1). See also Figures 5 and 6. Scale bars are 200 nm...... 100

Figure 5.8 XRD spectra for (a) several FeCl2 seeded TiO2 and TiO2-xNx samples as

indicated in the inset of the figure, (b) expanded view of TiO2- FeCl2- 0.94

mole% (Table 5.I), and (c) XRD spectrum of DeGussa P25 with an 87%

anatase and 13% rutile crystalline phase [54]...... 102

2+ Figure 5.9 UV-Vis diffuse reflectance spectra for Fe doped (a) TiO2 and (b) TiO2-xNx

nanoparticles. The figure legends indicate the amount of FeCl2 used. For

comparison, UV-Vis diffuse reflectance spectra of commercial TiO2

photocatalyst, P25, and TiO2-xNx nanoparticles are also included...... 104

Figure 5.10 (a) Photocatalytic decomposition of methylene blue for iron chloride doped

TiO2 (seeded by anhydrous FeCl2). The legend shows the concentrations

corresponding with those in the first column of Table 5.1 and explained in

the caption of Table 5.1. (b) Photocatalytic decomposition of methylene blue

for iron chloride doped TiO2 (seeded by FeCl2 (4 H2O). The legend shows

the concentrations correlating with column 1 of Table 5.1. (c) Photocatalytic

decomposition of methylene blue by TiO2-xNx-FeCl2...... 105

Figure 5.11 VB-XPS spectrum of the products of the FeCl2 (0.2g = 37.9 mole % Fe)-

3+ TiO2-xNx interaction showing the formation of Fe whose XPS spectrum

peaks at 56 eV, Ti 3s, and remaining Fe2+. The indicated intensity for Fe3+

represents a lower bound...... 106

Figure 5.12 Correlation of XPS data obtained for Fe2+ seeded oxynitride with data

obtained for the Co2+ seeded oxynitride (Ref. 22). Note especially the peaks

at 17-17.5 eV for the cobalt system vs. the peaks at 15.5 to 16.5 eV for the

iron system and the broadening of the iron spectrum as discussed in the text.

...... 107

Figure 6.1 UV-vis diffuse reflectance spectra of Ga-TiO2-xNx NP samples...... 117

Figure 6.2 XRD patterns for Ga-TiO2-xNx samples with various Ti : Ga ratios sintered at

600 and 700 oC. The patterns are not smoothed to keep the original

appearance. The letter “A” in the figure stands for “anatase TiO2” peaks. . 119

Figure 6.3 (A)Full-scale (B) N1s XPS spectra of selected samples...... 121

Figure 6.4 Raman Spectra of the Ti : Ga = 4 : 1 sample sintered at 400, 500, 600, and 700

oC...... 123

Figure 6.5 TEM images of Ti : Ga = 4 : 1 samples sintered at (A) 400 oC, (B)500 oC (C)

600oC, and (D) 700 oC...... 124

Figure 6.6 Photocatalytic decomposition of methylene blue under visible-light

illumination. The legend is ranked according to the photoactivity from worst

to best. For comparison, the decomposition without catalyst and the

performance of commercial TiO2, P25 are also shown...... 125

Figure 7.1 From right to left, anhydrous FeCl3 (black), 1 : 1, 1 : 3, and 1 : 4 by weight

FeCl3 on CAB-O-SIL. The samples change in color from light orange to

bright yellow...... 131

Figure 7.2 Full-scale XPS spectra for Fe-CAB-O-SIL ratios as indicated in the figure. 132

Figure 7.3 De-convoluted Fe 3p XPS spectra for Fe-CAB-O-SIL samples with various Fe

to Cab ratios...... 133

Figure 7.4 Fe 3p XPS spectra for FeCl2 and FeCl3 ...... 134

Figure 7.5 High resolution XPS spectra of pure CAB-O-SIL at 50 to 62 eV...... 137

Figure 7.6 XPS spectra of Co 2p in a Co-TiO2 sample...... 137

Figure 7.7 Si XPS Binding Energy vs. Si Oxidation.[5] ...... 138

Figure 7.8 Right to left, comparison of bright yellow sample of Fe-CAB-O-SIL 1 : 4, Fe-

CAB-O-SIL 1 : 4 complexed with 1,10 phenanthroline, Fe2O3 over

phenanthroline under toluene, and powdered Fe2O3...... 140

Figure 7.9 UV-vis diffuse reflectance spectra for Fe-CAB-O-SIL (black) and the complex

of 1,10 phenanthroline with Fe-CAB-O-SIL(red)...... 141

Figure 7.10 FTIR of CAB-O-SIL and the Fe-CAB-O-SIL series samples...... 143

Figure 7.11 (a) side view before water droplet is introduced to the CAB-O-SIL surface (b)

side view after a water droplet has been introduced (c) front view after a

water droplet is introduced...... 144

Figure 7.12 XRD patterns of CAB-O-SIL and the Fe-CAB-O-SIL series samples...... 145

Figure 7.13 TEM images of (a) pure CAB-O-SIL (b) Fe Cab 1 : 1 (c) Fe Cab 1 : 3 (d) Fe

Cab 1 : 4...... 146

ACKNOWLEDGEMENTS

First, I would like to thank my advisor Prof. Clemens Burda for his guidance, encouragement, patience and support throughout my PhD study. He has a special way of enlightening me to think deeper and further. I am grateful to all the members in the Burda group. I learned a lot not only from senior students Yixin and Yu but also from younger students Tenny and Chi-Hung. And of course from my fellow students Baodong and

Keng-Chu. I also got help from the visiting students and undergraduate students.

Second, I would like to thank my committee members, Prof. Alfred B. Anderson,

Prof. Carlos E. Crespo Hernández, Prof. John E. Stuehr from CWRU and Prof. Jeffery

Dyck from John Carroll University for their time, guidance and help.

Third, I would like to thank our collaborators, Prof. James L. Gole in Georgia

Institute of Technology for providing samples and insights.

Also, I would like to thank my family for their care and support during my study in the US. In addition, I would like to thank Nikola Matic for his help as well.

Last but not least, I would like to thank the Chemistry Department of Case

Western Reserve University for having me as a graduate student.

LIST OF ABBREVIATIONS

CB conduction band

DFT density functional theory

DOS the density of state

EDX Energy-dispersive X-ray spectroscopy

Full Potential Linearized Augmented Plane Wave FLAPW Method

FTIR Fourier-transformed infrared

FWHM the full width at half-maximum

HOMO the highest occupied molecular orbital

HRTEM high-resolution transmission electron microscopy

IPCE the incident photo-to-current efficiency

LDA local density approximation

LUMO the lowest unoccupied molecular orbital

MB methylene blue

NHE normal hydrogen electrode

NP nanoparticle

PL photoluminescence

SEM scanning electron microscopy

TEM transmission electron microscopy

TOPO trioctylphosphonic oxide

XPS X-ray photoelectron spectroscopy

XRD X-ray diffraction

UV ultraviolet

UV-vis ultraviolet-visible

VB valence band

CHEMICAL DOPING OF METAL OXIDE NANOMATERIALS AND

CHARACTERIZATION OF THEIR PHYSICAL-CHEMICAL PROPERTIES

ABSTRACT

JUNWEI WANG

Energy and environment are vital to every aspect of our daily lives. It is paramount to have a new source of energy before the exhaustion of fossil fuels to prevent a setback in the standard of living in which we have grown accustomed to and to have emission of greenhouse gases under control to deflect tremendous global ecological tragedy. Photocatalysis using solar energy is considered by many as the most promising solution for both energy and environment. Among all the photocatalysts under research, titanium dioxide (TiO2) remains one of the most studied for decades. Reducing its bandgap via doping is crucial to harvest visible-light, which accounts for 48% of the total solar energy. In addition, it is very important to keep the particle size small to avoid recombination of the photo-generated electrons and holes.

In this dissertation, metal and / or non-metal elements have been used as dopants to modify the electronic structure of TiO2. Special synthesis techniques have been chosen so that the particle size is in the nano-regime. Various techniques such as X-ray diffraction patterns (XRD), Raman spectroscopy, UV-Vis diffuse reflectance spectroscopy (UV-DRS), Fourier transform infrared spectroscopy (FTIR), scanning electron microscopy (SEM), transmission electron microscopy (TEM) and X-ray photoelectron spectroscopy (XPS) have been used to obtain detailed information of the

16 crystal structure, light-absorbing property, organic remains, particle appearance, size distribution and elemental compositions, respectively. Based on the characterization results, we have successfully introduced doping levels into TiO2 to reduce its band gap and thus made the prepared materials visible-light reactive. The sizes of the products are controlled in nano-scale. For all the photocatalysts tested, the performance for the decomposition of methylene blue can be ranked as follows: Ga,N-TiO2 > N-TiO2 >

Co,N-TiO2 > Fe-TiO2 ≈ Fe,N-TiO2. Furthermore, there have been other important discoveries such as the observation of phenomena similar to that of nitrogen fixation in the Fe-TiO2 experiment and the observation of the surprising reducing properties of commercial material Cab-o-sil.

Chapter 1 Introduction

1.1 Motivation: Energy and Environment Crisis

The United States is not alone in the current energy crisis. Indeed, the tremendous global demand for oil, spurred by factors such as population growth and industrial development, has served to generate ever-rising gas prices. (See Figure 1.1 and

1.2). On January 2, 1986, the oil price was $25.56 per barrel. By July 14, 2008, the oil price reached a historical high of $145.16 per barrel. Despite a temporary drop in price of crude oil, recent trends suggest that the price may again reach and possibly exceed that of

2008. The price of gasoline, which is more relevant to daily life, demonstrates yet another concern: while the price of crude oil has not yet reached its historical high, the price of gasoline is almost once again at its highest cost in history. In a world that relies on vehicles for transportation and by extension, fuel, it is difficult to imagine what might happen if this perceived necessity would become a luxury. Should the price of gas reach

$10 per gallon, it would cost an estimated $132 just to fill the tank of an average compact car. According to statistics, a carrot travels 1838 miles on average before reaching a dinner table.[1] With the rising price of gasoline, eating something that was once as inane as a carrot might be considered a luxury. Indeed, such a scenario would create an extremely difficult economic climate, much more severe than ever before. The growing consumption of fossil fuels such as oil, coal, and gasoline has made a tremendous impact on the planet. While oil depletion is one of the main issues being addressed around the world, climate change may stand to be more jeopardizing to human survival. With the energy crisis and climate change clearly linked, the use of petroleum has increased carbon dioxide levels and has contributed to the change in weather around the world.

Frequent reports demonstrate the melting of glaciers in Antarctica. [2, 3] Such melting has already influenced people thousands of miles away near the equator, for example the island country Kiribati is planning to move the whole nation to Fiji[4] and people in

Thailand are experiencing an increase in floods per year.[5]It has also been suggested that the surge in hurricanes may also be a result of global climate change.[6]

Figure 1.1 Crude oil prices from January 2nd, 1986 to March 6th, 2012. Data from http://www.eia.gov/dnav/pet/hist/LeafHandler.ashx?n=PET&s=RWTC&f=A

4.5

4.0

3.5

3.0

2.5

2.0

1.5

1.0 Gasoline price / dollars price Gasoline 0.5 1986 1991 1996 2001 2006 2011 Year

Figure 1.2 Average gasoline prices in US from 1986 to 2011.

1.2 Potential Solutions: Photocatalysis

It is imperative for global cooperation to address these issues. Considering almost all the energy sources except nuclear energy can be traced back to solar energy, the solution will have to ultimately lie in the solar energy. Since carbon containing fossil fuels will cause the release of greenhouse gases, possible solutions could include converting the greenhouse gases into non-harmful existence or use carbon free energy. It is well-known that 70% of the earth is covered by water and when hydrogen is combusting, the only product will be water. If we can use solar energy to split water into hydrogen and oxygen and then make use of the energy that is released during hydrogen combustion with oxygen, we are not creating any pollution at all and we may even get purified water by doing so (see reaction 1-1).

using solar energy to split 2H O←→2H+ O (1-1) 2 combusting and releasing energy to use 22

In fact, researchers have already made this true.[7] The problem remaining is the low efficiency.[8] To improve the efficiency, we need to first understand how the solar energy split the water. The oceans are being illuminated all the time during daytime but little hydrogen is generated. Hence, we need something serving as a media for the solar energy to be utilized. And this media is usually the so-called photocatalyst. So far, most of the photocatalysts are semiconductor metal oxides, sulfides, or nitrides. These semiconductors have bandgaps between their conduction band and valence band. With the appropriate radiation, the electrons in the valence band can be ejected into the conduction band, leaving a positively charged hole in the valence band. The electrons and holes thus separated will serve as reductant and oxidant respectively to facilitate the

+ redox reactions. In the water splitting case, the electrons will reduce H , generating H2

- and the holes will oxidize the OH to get O2 (see Figure 1.3). However, in reality, the process can be much more complicate than it appears to be. The bandgap of the semiconductor needs to be in the right range to be able to use solar energy to generate electron hole separations; the electrons and holes have to transfer to the surface of the photocatalysts to be able to react with the reactants; the products, namely, hydrogen and oxygen, must be prevented from recombination before they are collected. Up to now, there is no readily available photocatablyst to fulfill the job of catalyzing the water splitting process efficiently enough for our daily use. Take the most investigated semiconductor, i.e., titanium dioxide, as an example, although it shows great potential due to its high activity, stability, non-toxicity and low cost, it has a band gap of 3.2 eV, which makes it capable of absorbing only radiations with wavelength shorter than 387 nm. If we looked at the energy distribution of sunlight by wavelength, as shown in Figure

1.4, we will notice that the energy contained in the radiations whose wavelength is shorter than 387 nm is less than 5% of the total solar energy. Therefore, actions are required to modify the bandgap so that the majority of the energy in the solar spectrum can be used.

Figure 1.3 Illustration of photocatalytic water splitting[9].

Figure 1.4 Solar light energy distribution by wavelength[10].

Apart from serving as catalyst to split water, another useful application for photocatalyst is to decompose the organic waste produced in our daily lives. Such waste can be either in air or in water and can be really harmful to ecology. For example, excess organic dyes discharged from dyeing and finishing industries into the rivers and lakes can cause serious problem not only by depleting the dissolved oxygen but also by releasing toxic compounds during the chemical or biological pathways. Photocatalysts have the capability to generate oxidative species under light illumination such as singlet oxygen or hydroxyl radicals which can destroy a wide range of tough organic contaminants.

Titanium dioxide has been used to decompose organic substances in both air and / or water.[11, 12] Figure 1.5 shows the suggested reaction pathways for photocatalytic decomposition of organic dyes. The process begins with the adsorption of the dyes onto the TiO2 surface. Upon light illumination, the electrons in the valence band of TiO2 will be ejected into the conduction band, leaving a positively charged hole in the valence band.

Should these electrons and holes transfer to the surface of TiO2, the holes will be able to react with water to generate hydroxyl radicals whereas the electrons may react with

23 oxygen to produce superoxide free radicals. The hydroxyl radicals and the free superoxide radicals may then attack the dyes. After some intermediate processes, the organic dyes are degraded to the inorganic minerals and the TiO2 will go back to its ground state to start another round of photocatalysis. Again, although TiO2 is proven to be a highly active photocatalyst, the problem is its bandgap is too big to absorb visible- light.

Figure 1.5 Suggested pathways for photocatalytic decomposition of organic dyes. [12]

1.3 Current research on metal oxides as photocatalysts

Suppose we have the bandgap of titanium dioxide adjusted to the suitable range.

There are still many obstacles that remain. Semiconductors are usually bulk materials that have low surface to volume ratios. Note that only the surface of the material is in contact with the reactants of the reaction that needs to be catalyzed. What this means is that the electrons and holes have to travel from the spot where they generate to the surface of the material. While traveling, it is very likely that the electrons will recombine with holes. As

24 a result, only a small portion of the photogenerated electrons and holes will be able to actually react with water to produce hydrogen and oxygen. In another word, the majority of energy carried by the excited electrons and holes is not able to pass to the water. To solve this issue, people are trying to make the material in the scale of nanometers.[13] By doing so, surface to volume ratio will be greatly increased, hence reducing the probability of electrons and holes recombination.

The third problem mentioned above is the reaction between the newly generated hydrogen and oxygen. Hydrogen has an explosion limit of 4% to 74%. The water splitting should produce hydrogen and oxygen in the ratio of 2 to 1. This means the percentage of hydrogen is 67%, right in the explosion limit. With constant light illumination, it is quite possible that the hydrogen will react with oxygen to generate water again. To avoid this from happening, it would be ideal to physically separate the hydrogen generating location from the oxygen generating location. This usually requires the application of external bias or the use of heterostructures to drive the electrons and holes to different locations. [7]

As previously discussed, it is necessary to reduce the band gap of TiO2. How exactly is that achieved? In general, two methods are roughly seen. They are metal doping and non-metal doping.[14] The optical property of the semiconductor is essentially determined by its electronic structure. Introducing new element may change the existing electronic structure hence may have the effect of reducing the band gap.

Umebayashi et al calculated the Density of States (DOS) of V-, Cr-, Mn-, Fe-, Co-, and

Ni-doped TiO2 by ab initio band calculations (see Figure 1.6). They found that an electron occupied level is formed when doped with the above mentioned metal elements

25 and the electrons are localized around each dopant. With the increase of the atomic number of the dopants, this localized level will shift lower. Specifically, the localized level for the Co-doped TiO2 is right at the top of the valence band of TiO2, while others like Mn, Cr tend to form midgap states. Currently, the major problem of metal doped

TiO2 is the recombination of the excited electrons and holes.[14] Scientists are still working to solve the issue.

Asahi et al have calculated the electronic band structures of non-metal doped

TiO2 using full potential linearized augmented plane wave (FLAPW) in the framework of the local density approximation (LDA).[15] They found that substitutional doping of N was the most effective because its p states contribute to the band-gap narrowing by mixing with O 2p states (See Figure 1.7). In addition, N doping was found more efficient than the molecularly existing species, e.g. NO and N2 dopants, which contribute to the bonding states below the O 2p valence bands and antibonding states deep in the band gap and are screened and hardly interact with the band states of TiO2. However, there are different opinions as to how the band structure changes by nitrogen doping. For example,

Di Valentin et al found that instead of mixing with the valence band, N2p localized states are just above the O2p states in the valence band.[16] Serpone et al argued that the visible-light absorption originates from oxygen defects formed during the process of TiO2 reduction. They proposed that TiO2 has been reduced during the nitrogen doping process and F-type color centers (a type of crystallographic defect in which an anionic vacancy in a crystal is filled by one or more electrons, depending on the charge of the missing ion in the crystal) are formed. These F-type centers together with the Ti3+ centers play the

26 dominant role of increasing the absorption in the visible-light region.[17] In addition to nitrogen, other non-metal elements have also been doped into TiO2 to obtain visible

Figure 1.6 DOS calculation for V, Cr, Mn, Fe, Co, or Ni doped TiO2. Gray solid lines: total DOS. Black solid lines: dopant’s DOS. The states are labelled (a) to (j). [18]

-light reactivity.[19] [20] So far, the shortage of the non-metal doped TiO2 is the low concentration of the dopant and the less control over the stability of the prepared samples.[14]

Figure 1.7 Total DOSs of doped TiO2 calculated by FLAPW.[15]

Throughout the years, scientists have developed a series of techniques to synthesize the TiO2 related samples. These techniques includes sol-gel, micelle and inverse micelle, hydrothermal, solvothermal, direct oxidation, chemical vapor deposition

(CVD), physical vapor deposition, electrodeposition, sonochemical method, microwave method etc. Details of these methods can be found in Xiaobo Chen and Samuel Mao’s recent review.[14]

1.4 About this dissertation

The Burda group is one of the first groups in the world to synthesize nano-scaled high concentration nitrogen doped titanium dioxide.[21] The purpose of this dissertation is to further perfect the synthetic techniques of nano-scaled Nitrogen doped TiO2 as well as metal doped TiO2. Various synthetic methods, including sol-gel, hydrothermal, ligand- assisted methods are used in this work. Various techniques such as X-ray diffraction patterns (XRD), Raman spectroscopy, UV-vis diffuse reflectance spectroscopy (UV-

DRS), Fourier transform infrared spectroscopy (FTIR), scanning electron microscopy

(SEM), transmission electron microscopy (TEM) and X-ray photoelectron spectroscopy

(XPS) etc. have been used to understand the surface physical and chemical properties of these prepared materials. We have successfully introduced doping levels to modify the band gap and thus made the prepared materials visible-light reactive (Chapter 2 to 6).

Most of the work have been published or submitted for publication. While collaborating with Professor James Gole at Georgia Institute of Technology, we noticed and investigated the unique properties of the reducing power of commercial product cab-o-sil and the work was recently submitted to PCCP for publication (Chapter 7).

1.5 References

1. Pirog, R., and Andrew Benjamin, Leopold Center for Sustainable Agriculture,, July 2003. 2. http://www.usatoday.com/weather/climate/globalwarming/2009-02-25- warming_N.htm. 3. http://news.sciencemag.org/sciencenow/2011/04/melting-antarctic-ice-causing- pe.html 4. REYES, D., NEWJERSEYNEWSROOM.COM, 2012. 5. http://www.worldweatherpost.com 6. Greene, C.H., the Christian Science Monitor, 2011. 7. Fujishima, A. and K. Honda, Nature, 1972. 238(5358), 37-38. 8. Chen, X.B., S.H. Shen, L.J. Guo, and S.S. Mao, Chemical Reviews, 2010. 110(11), 6503-6570. 9. http://newscenter.lbl.gov/wp-content/uploads/solar-water-splitting1.jpg. 10. http://www.incognitogb.com/learn_more.html. 11. Fujishima, A., T.N. Rao, and D.A. Tryk, Journal of Photochemistry and Photobiology C: Photochemistry Reviews, 2000. 1(1), 1-21. 12. Wu, C.-H. and J.-M. Chern, Industrial & Engineering Chemistry Research, 2006. 45(19), 6450-6457. 13. Burda, C., X.B. Chen, R. Narayanan, and M.A. El-Sayed, Chemical Reviews, 2005. 105(4), 1025-1102. 14. Chen, X. and S.S. Mao, Chemical Reviews, 2007. 107(7), 2891-2959. 15. Asahi, R., T. Morikawa, T. Ohwaki, K. Aoki, and Y. Taga, Science, 2001. 293(5528), 269-271. 16. Di Valentin, C., G. Pacchioni, and A. Selloni, Physical Review B, 2004. 70(8). 17. Kuznetsov, V.N. and N. Serpone, Journal of Physical Chemistry C, 2009. 113(34), 15110-15123. 18. Umebayashi, T., T. Yamaki, H. Itoh, and K. Asai, Journal of Physics and Chemistry of Solids, 2002. 63(10), 1909-1920. 19. Park, J.H., S. Kim, and A.J. Bard, Nano Letters, 2006. 6(1), 24-28. 20. Ksibi, M., S. Rossignol, J.M. Tatibouet, and C. Trapalis, Materials Letters, 2008. 62(26), 4204-4206. 21. Burda, C., Y. Lou, X. Chen, A.C.S. Samia, J. Stout, and J.L. Gole, Nano Letters, 2003. 3(8), 1049-1051.

Chapter 2 Hydrothermal Nitrogen-doped Titanium Dioxide with switchable surface

wettability

This work was published in Nanosacle, 2010, 2, 2257-2261

2.1 Introduction

Surface wettability is one of the most important properties of solid state materials relevant to practical applications in the field of self-cleaning,[1] anti-fogging,[2] anti- biofouling,[3] drug delivery,[4] and the of water and oil separation[5]. During the past several years, scientific interest in surface wetting has greatly increased.[2, 3, 6-10]

Generally, the surface wettability can be evaluated by measuring its water contact angle with water. Surfaces with water contact angles higher than 90º are categorized as hydrophobic whereas surfaces with contact angles lower than 90º are categorized as hydrophilic. In addition, superhydrophobic refers to surfaces that show water contact angles higher than 150o whereas superhydrophilic surfaces exhibit water contact angles of less than 10o. These surfaces attract much attention from the wetting community as a result of plentiful industrial and biological applications. Furthermore, “smart” surfaces that are able to switch between hydrophobic and hydrophilic have attracted even more attention because of their potential to enhance rapid water motion,[11] to improve micro- fluidic devices,[12] and to create smart membranes[13] and sensors.[14]

Titanium dioxide is one of the most popular semiconductor metal oxides currently under extensive investigation due to its stability, nontoxicity, cost effectiveness, and highly reactivity.[15-20] However its intrinsic disadvantage, i.e., having too large a bandgap, has to be addressed. Currently, a number of approaches have been used to solve the issue. For example, nitrogen doping was found to be one of the most promising

31 operandi.[21-30] There have been few reports detailing the synthesis of superhydrophobic TiO2,[19, 31-34] and a few studies of smart surfaces.[32, 35]

Nevertheless, to the best of our knowledge, no reports regarding the fabrication of visible-light-driven switchable hydrophobic to hydrophilic surfaces have been reported.

We report a hydrothermal synthesis of visible-light responsive anatase phase N-doped

TiO2 at low temperature (180 ºC, compared to 350 ºC, for the normal crystallization temperature of anatase TiO2 formed by the sol-gel method). The surface wettability of the material is switchable between hydrophobic and hydrophilic.

2.2 Experimental procedures

Synthesis

Nitrogen doped TiO2 nanoparticles were prepared using a hydrothermal method adapted from literature.[36] For a typical synthesis, 30.0 mL of 5.33 mM titanium tetrafluoride, 0.40 mL of 10% hydrofluoric acid and 1.50 mL of ethylenediamine (99%

Aldrich) were mixed in a Teflon-lined autoclave, which was then placed in an oven at

180 ºC for 6, 12, 18 and 24 hours to get four samples. The product was washed with deionized water and centrifuged at 6000 revolutions per minute for 15 minutes, twice. It was then dispersed in ~30 mL of deionized H2O and treated with ultrasonic wave for 30 min, after which the suspension was added to a wafer, drop by drop, in an 80 ºC oven to form a film.

Characterization

Structures of the samples were scanned by a Scintag X-1 Advanced X-ray powder diffractometer (XRD, 2º / min, Cu Kα radiation). The morphologies of these samples

32 were characterized by a Hitachi S4500 field-emission gun scanning electron microscope

(SEM). Optical properties were obtained by a Varian Cary Bio50 UV–vis spectrometer with a Barrelino remote diffuse reflection probe (reference material: MgO). Water contact angles were measured by analyzing high-definition photographs of the water droplets on sample surfaces. The visible-light source was obtained from a 150 W Xenon arc light with its UV light filtered by two paralleled 400 nm-cut-off filters. To induce hydrophilicity, the samples were illuminated with the filtered light for 6 hours. Various surface contact angles ranging from hydrophilic to hydrophobic were obtained during the

6 hour illumination. The hydrophobicity was recovered either by keeping the samples in the dark for 2 weeks or by keeping the samples in a 120 ºC oven over night.

Photocatalysis

The photocatalytic performance of the N-doped TiO2 nanoparticles was evaluated by monitoring the decomposition of methylene blue (MB) under visible-light irradiation.

The light source was a 150-W high-pressure Xenon arc lamp together with two 400 nm- cut-off filters. The change in the concentration of methylene blue (MB) was determined by measuring the absorbance of MB solution with a Varian Cary Bio50 UV–vis spectrometer. For comparison, the changes in the concentration of MB solution with no catalyst under dark and light conditions were also tested.

2.3 Results and discussion

Anatase 6 hrs 12 hrs Brookite 18 hrs 24 hrs

20 30 40 50 60 2 θ / degree

Figure 2.1 XRD patterns of N-doped TiO2 prepared by a hydrothermal method. Shown in the figure are samples with growing times ranging from 6 to 24 hours. In the bottom of the figure are the standard XRD peaks for anatase TiO2 (pink, JCPDS file NO. 21-1272) and brookite TiO2 (wine, JCPDS file NO.03-0380).

Figure 2.1 shows the XRD patterns of the N-doped TiO2 samples for different reaction time lengths. The structures of the prepared samples are found to be primarily anatase phase (body centered tetragonal structure, space group I41/ amd). A small fraction of brookite phase (primitive orthorhombic structure, space group Pbca) is also present.

There is no significant change in the intensities of the peaks among these samples, indicating that the sizes and crystallinities of the samples are similar. The peaks are relatively broad compared to those of the bulk material, which means that the crystal sizes of these samples are smaller than bulk samples. Compared to the standard XRD pattern of anatase TiO2 (JCPDS file no. 21-1272), the peaks at 2θ = 25.2º ( (101) reflection ) are clearly weakened whereas the peaks at 2θ = 53.8º and 55.0º ( (105) and

(211) reflections, respectively) are stronger (see Table 2.1 for the relative intensities of

34 the peaks). The difference in the intensities of the peaks compared to those of the standard sample indicates the existence of a preferred orientation for the growth of these samples, which should be the longitudinal axis of the nanorods depicted in Figure 2.2.

6 h 12 h

18 h 24h

Figure 2.2 SEM images of N-doped TiO2 samples after the indicated reaction time of 6, 12, 18 and 24 hours.

Figure 2.2 depicts the SEM images of the N-doped TiO2 samples. There is no significant difference between the morphologies of the nanorods produced over different time scales, which coincides with the similar XRD patterns of Figure 2.1. However, upon careful examination of these images, we can see that the 18-hour sample has the narrowest size distribution of the four samples. The average length of the nanorods is approximately 130 nm while their diameter is approximately 35 nm.

Table 2.1 Relative XRD peak intensities (at 2θ = 25.2º, 53.8º and 55.0º) for the prepared samples and standard anatase TiO2 (For comparison, the intensities of the peak at 2θ = 55.0º are set to be 1.0). Diffraction angle / 2θ 25.2º 53.8º 55.0º Standarda 6.4 1.0 1.0 6 h 3.1 1.2 1.0 12 h 4.1 1.0 1.0 18 h 3.9 1.1 1.0 24 h 3.8 1.1 1.0 a standard anatase XRD file (JCPDS file NO. 21-1272)

According to Lifshitz et al.,[37] diffusion controlled growth for a supersaturated solution can be divided into three distinct stages, nucleation, normal growth and competitive growth. The last stage, Ostwald ripening, may correspond to the dissolution of nanocrystallites that are smaller than a certain critical size and the transfer of their mass to nanocrystallites that are larger than a critical size.

Figure 2.3 depicts the UV-vis diffuse reflection spectra of the N-doped TiO2 prepared by the hydrothermal method for various time scales. The absorption in the visible region increases from growing times of 6 hours to 12 hours and continues to increase to 18 hours where it reaches a plateau. Based on these spectra, it can be concluded that the hydrothermal reaction takes about 18 hours to reach its equilibrium.

This is consistent with the SEM results (Figure 2.2). Two mechanisms have been suggested as the origin of the visible light absorption. First, bandgap narrowing is thought to result from N 2p-O 2p orbital mixing decreasing the energy increment to the conduction band of TiO2.[21, 26, 38] Second, oxygen defects formed in the reduction of

TiO2 may lead to visible light absorption.[39] It has been argued that TiO2 has been

36 reduced during the N-doping process, as F-type color centers are formed. The F-type color centers together with Ti3+ centers are suggested to play the dominant role in the visible-light-activity of TiO2 photocatalysts. More recently, researchers have concluded that both of the two mechanisms play a role.[40, 41]

The wettability of the N-doped TiO2 films was evaluated by contact angle measurement. All of the samples exhibit surface hydrophobicities. Shown in Figure 2.4 is a series of photographs of water droplets sitting on an N-doped TiO2 film. According to

the Cassie and Baxter relationships,[32, 42] the apparent contact angle θ f can be given using the following equation:

cosθθf=ff s cos wv − , (1)

where θw is the contact angle on a smooth surface (72º for anatase TiO2[43]), fs is the surface fraction of the solid and fv is the surface fraction of the air (fs + fv = 1). With the

large apparent contact angle as high as 141º, fv can be calculated (assuming θw is the same as that of anatase TiO2) to be as high as 83%. Air trapped in the films would appear to play the role of a cushion at the film-water interface and prevent the water droplet from penetrating into the porous regions of the film.

60 6 hrs

12 hrs 18 hrs 24 hrs 30 Reflectance / %

0 200 400 600 800 Wavelength / nm

Figure 2.3 UV-vis diffuse reflection spectra of N-doped TiO2 prepared by hydrothermal method. Shown in the figure are the curves for 4 different growing times, namely, 6, 12, 18 and 24 hours.

The hydrophobic nature of the N-doped TiO2 films transforms into a hydrophilic character when these samples are illuminated with visible-light. Shown in Table 2.2 are the contact angles for the samples before and after light illumination. The change from hydrophobic to hydrophilic character might be ascribed to a kinetically favored absorption of water molecules onto surface oxygen vacancies formed by the reaction of

Figure 2.4 Photographs of the water droplets on an N-doped TiO2 film with various surface wettabilities.

38 photogenerated holes and lattice oxygen.[10, 32] Adsorbed water can fill the porous regions along the nanorods, replacing trapped air. This can result in a hydrophilic surface.

This change in the wettability can be reversed when the samples are kept in the dark for two weeks. This behavior suggests that surface oxygen vacancies decay as the oxygen regeneration process is thermodynamically more favored.[32]

Table 2.2 Contact angle measurements for the N-doped TiO2 samples before and after visible-light illumination contact angles / º

before illumination after illumination

TiON-6 133 0a

TiON-12 133 0

TiON-18 140 0

TiON-24 135 0 a water droplet completely spread out.

The process of regenerating hydrophobicity can be shortened to 12 hours if the samples are kept in an oven at 120 ºC. To test the role of adsorbed water in the wettability changing process, we illuminated the hydrophobic surface in a vacuum system (10-3

Torr), and did not find any change in the surface hydrophobicity under the same light illumination as in the air-exposed case. We also did not find any change when the hydrophobic surface was kept in an argon environment while being illuminated. This again indicates that water is indispensable in the wettability transition process.

140 Vis Vis Vis 105

35 ο ο 120οC 120 C 120 C

Water contact angle / degree 0 0 1 2 3 Cycles

Figure 2.5 Contact angle measurements of the 24-hour N-TiO2 sample in 3 hydrophobic- hydrophilic cycles (in air-exposed system). Values shown are averages from three different measurements. It is noteworthy to point out that the same illumination on the surfaces that are kept in vacuum or argon environment did not result in the wettability change at all, indicating the important role of water in the wettability transition process.

Figure 2.5 shows the water contact angle measurement for the 24-hour N-TiO2 sample in 3 cycles of hydrophobic-hydrophilic transformation (in an air-exposed system).

The photocatalytic performance of the nitrogen doped TiO2 (24-hour sample) with various surface wettabilities is shown in Figure 2.6. As the surface becomes more hydrophilic, the photocatalytic activity is significantly enhanced. It is noteworthy that the decomposition processes of methylene blue for the 135º and 90º samples do not follow the same kinetics as those for the 60º and 0º samples. The reason is that with light illumination, the surfaces of the 135º and 90º samples continue to transform towards hydrophilicity, thus increasing the photocatalytic activity of the sample during the photocatalysis process.

100

no catalyst Degussa P25 TiON-135o TiON-90o 60 TiON-60o Percentage of MB leftPercentage of MB / % TiON-0o 0 1 2 3 4 5 Time / hour

Figure 2.6 Visible-light induced photocatalysis of methylene blue by nitrogen-doped TiO2 with various surface wettabilities. The numeric values of these surface contact angles are included in the figure legend. For comparison, the photocatalytic performance of Degussa P25 (a commercial available TiO2 product) under the same condition is also shown.

2.4 Conclusion

In conclusion, strongly hydrophobic N-doped TiO2 nanorods have been prepared using a hydrothermal method. These anatase TiO2 nanorods have a length of approximately 130 nm and diameter of approximately 35 nm. The optimum time for the nanorod growth was determined to be 18 hours in correspondence with both SEM and

UV-Vis diffuse reflective observations. The surfaces of the N-doped TiO2 samples are strongly hydrophobic with contact angles as high as 141º. These surfaces transform toward a hydrophilic character under visible-light illumination and reversibly to a strongly hydrophobic character after being kept in the dark for 2 weeks or after moderate sintering at 120 ºC. The photocatalytic activity of these nano-scaled catalysts is found to rely heavily on their surface wettability.

2.5 References

1 D. Quere, Reports on Progress in Physics, 2005, 68, 2495-2532. 2 P. Roach, N. J. Shirtcliffe and M. I. Newton, Soft Matter, 2008, 4, 224-240. 3 X. M. Li, D. Reinhoudt and M. Crego-Calama, Chemical Society Reviews, 2007, 36, 1350-1368. 4 M. L. Gou, X. Y. Li, M. Dai, C. Y. Gong, X. H. Wang, Y. Xie, H. X. Deng, L. J. Chen, X. Zhao, Z. Y. Qian and Y. Q. Wei, International Journal of Pharmaceutics, 2008, 359, 228-233. 5 L. Feng, Z. Y. Zhang, Z. H. Mai, Y. M. Ma, B. Q. Liu, L. Jiang and D. B. Zhu, Angewandte Chemie-International Edition, 2004, 43, 2012-2014. 6 T. Onda, S. Shibuichi, N. Satoh and K. Tsujii, Langmuir, 1996, 12, 2125-2127. 7 R. Wang, K. Hashimoto, A. Fujishima, M. Chikuni, E. Kojima, A. Kitamura, M. Shimohigoshi and T. Watanabe, Nature, 1997, 388, 431-432. 8 H. Y. Erbil, A. L. Demirel, Y. Avci and O. Mert, Science, 2003, 299, 1377-1380. 9 M. W. Denny, Science, 2008, 320, 886-886. 10 F. Xia, Y. Zhu, L. Feng and L. Jiang, Soft Matter, 2009, 5, 275-281. 11 S. Daniel, M. K. Chaudhury and J. C. Chen, Science, 2001, 291, 633-636. 12 D. J. Beebe, J. S. Moore, Q. Yu, R. H. Liu, M. L. Kraft, B. H. Jo and C. Devadoss, Proc. Natl. Acad. Sci. U. S. A., 2000, 97, 13488-13493. 13 G. Q. Zhai, Y. Lei, E. T. Kang and K. G. Neoh, Surface and Interface Analysis, 2004, 36, 1048-1051. 14 J. D. Jeyaprakash, S. Samuel, P. Ruther, H. P. Frerichs, M. Lehmann, O. Paul and J. Ruhe, Sensors and Actuators B-Chemical, 2005, 110, 218-224. 15 A. Fujishima and K. Honda, Nature, 1972, 238, 37-38. 16 M. R. Hoffmann, S. T. Martin, W. Y. Choi and D. W. Bahnemann, Chemical Reviews, 1995, 95, 69-96. 17 C. Burda, X. B. Chen, R. Narayanan and M. A. El-Sayed, Chemical Reviews, 2005, 105, 1025-1102. 18 I. Nakamura, N. Negishi, S. Kutsuna, T. Ihara, S. Sugihara and K. Takeuchi, Journal of Molecular Catalysis A: Chemical, 2000, 161, 205-212. 19 X. T. Zhang, M. Jin, Z. Y. Liu, D. A. Tryk, S. Nishimoto, T. Murakami and A. Fujishima, Journal of Physical Chemistry C, 2007, 111, 14521-14529. 20 Y. X. Zhao, X. F. Qiu and C. Burda, Chemistry of Materials, 2008, 20, 2629-2636. 21 R. Asahi, T. Morikawa, T. Ohwaki, K. Aoki and Y. Taga, Science, 2001, 293, 269-271. 22 S. U. M. Khan, M. Al-Shahry and W. B. Ingler, Science, 2002, 297, 2243-2245. 23 X. Chen and S. S. Mao, Chemical Reviews, 2007, 107, 2891-2959. 24 J. W. Wang, W. Zhu, Y. Q. Zhang and S. X. Liu, Journal of Physical Chemistry C, 2007, 111, 1010-1014. 25 F. Dong, W. R. Zhao and Z. B. Wu, Nanotechnology, 2008, 19, 365607.

26 D. N. Tafen, J. Wang, N. Q. Wu and J. P. Lewis, Applied Physics Letters, 2009, 94, 093101. 27 C. Burda, Y. B. Lou, X. B. Chen, A. C. S. Samia, J. Stout and J. L. Gole, Nano Letters, 2003, 3, 1049-1051. 28 X. B. Chen, Y. B. Lou, A. C. S. Samia, C. Burda and J. L. Gole, Advanced Functional Materials, 2005, 15, 41-49. 29 X. F. Qiu, Y. X. Zhao and C. Burda, Advanced Materials, 2007, 19, 3995-3999. 30 P. Romero- Gomez, V. Rico, A. Borras, A. Barranco, J. P. Espinos, J. Cotrino and A. R. Gonzalez-Elipe, The Journal of Physical Chemistry C, 2009, 113, 13341- 13351 31 W. T. Sun, S. Y. Zhou, P. Chen and L. M. Peng, Chemical Communications, 2008, 603-605. 32 X. J. Feng, J. Zhai and L. Jiang, Angewandte Chemie-International Edition, 2005, 44, 5115-5118. 33 X. H. Xu, Z. Z. Zhang and W. M. Liu, Colloids and Surfaces a-Physicochemical and Engineering Aspects, 2009, 341, 21-26. 34 X. T. Zhang, M. Jin, Z. Y. Liu, S. Nishimoto, H. Saito, T. Murakami and A. Fujishima, Langmuir, 2006, 22, 9477-9479. 35 H. S. Lim, D. Kwak, D. Y. Lee, S. G. Lee and K. Cho, Journal of the American Chemical Society, 2007, 129, 4128-4129. 36 H. G. Yang, C. H. Sun, S. Z. Qiao, J. Zou, G. Liu, S. C. Smith, H. M. Cheng and G. Q. Lu, Nature, 2008, 453, 638-642. 37 I. M. Lifshitz and V. V. Slyozov, Journal of Physics and Chemistry of Solids, 1961, 19, 35-50. 38 S. Sakthivel, M. Janczarek and H. Kisch, Journal of Physical Chemistry B, 2004, 108, 19384-19387. 39 V. N. Kuznetsov and N. Serpone, Journal of Physical Chemistry C, 2009, 113, 15110-15123. 40 P. Romero- Gomez, V. Rico, A. Borras, A. Barranco, J. P. Espinos, J. Cotrino and A. R. Gonzalez-Elipe, The Journal of Physical Chemistry C, 2009, 113, 13341- 13351. 41 J. Wang, D. N. Tafen, J. P. Lewis, Z. L. Hong, A. Manivannan, M. J. Zhi, M. Li and N. Q. Wu, Journal of the American Chemical Society, 2009, 131, 12290- 12297. 42 A. B. D. Cassie and S. Baxter, Transactions of the Faraday Society, 1944, 40, 546-551. 43 A. Fujishima, T. N. Rao and D. A. Tryk, Journal of Photochemistry and Photobiology C: Photochemistry Reviews, 2000, 1, 1-21.

Chapter 3 Modifying Titanium Dioxide with Tungsten

3.1 Introduction

As mentioned in Chapter 1, metal doping is one of the major methods to red shift the absorption edge. Of all the metals that have been attempted to dope into TiO2, tungsten is one that has attracted researchers’ attention but is underdeveloped.

Advantages of tungsten include its low cost, non-toxicity, and durability. In addition, tungsten oxide has a band gap of 2.7 eV, which is smaller than TiO2 and is able to absorb visible-light. In addition, the redox potentials of the conduction band and valence band of tungsten oxide are in such positions that if combined with TiO2, type II heterostructures

(with the composing semiconductors having staggering band alignment like shown in

Figure 3.1) that are able to separate the electrons and holes physically might be formed.

(See figure 3.1). In this chapter, tungsten precursors in the amount of monolayer and double layers are used in the synthesis in the hope of creating core-shell structure.

However, the results do not seem to be able to prove that such core-shell structures are formed. Nevertheless, it is clear that the prepared sample exhibit visible-light absorbing capability.

Figure 3.1 Bands alignment of TiO2 (anatase) and WO3 (monoclinic).

3.2 Experimental procedures

Single sol-gel method and double sol-gel methods are shown in scheme 3.1 and

3.2 respectively. All beginning materials were purchased to use without further purification. The TiO2 nanoparticle (NP) solution was prepared by dissolving 0.25 mL

Ti(OCH(CH3)2)4 in 4.75mL isopropanol and adding the resultant 5.0 mL solution under continuous stirring dropwise into 900 mL pH=2 water solution(pH was adjusted by concentrated HNO3).

The TiO2 nanotube was synthesized by anodic oxidation of Ti foil at 35 V in an electrolyte composed of 0.25 wt % NH4F and 0.75 wt % H2O in ethylene glycol for 20 h at room temperature. [26] Then the as-prepared nanotube samples were rinsed with H2O and ethanol and calcined at 500 oC for 1 hour. Then the sample was dipped into ~ 30 mL

H2O which contained WCl6 whose weight is about 10% of that of the TiO2 nanotube. The

o whole solution was stirred at ~ 35 C overnight to allow the WCl6 to fully hydrolyze.

o Finally, the WO3 loaded TiO2 nanotube was sintered at 200 C for 2 h.

Scheme 3.1 Procedures for the synthesis of TiO2 / WO3 heterostructures by single sol-gel method.

Scheme 3.2 Procedures for the synthesis of TiO2 / WO3 heterostructures by double sol- gel method.

3.3 Results and discussion

Figure 3.2 TEM images of various samples prepared by single sol-gel method. The uncalcined TiO2 is shown in negative.

Figure 3.2 shows the TEM images of uncalcined TiO2 and TiO2 with 1 and 2 layers WO3 coating. The uncalcined TiO2 remains to be amorphous while the heterostructures are clusters of nanometer sized crystallites. There is severe aggregation for the prepared samples. Based on the XRD result (Figure 3.3), the phase of TiO2 is anatase in both 1 layer and 2 layers samples. There is no WO3 in the crystal form identified in the XRD patterns, which means that WO3 exists in amorphous phase.

However, it is hard to tell whether these crystallites are core-shell structures or not with existing data.

A 1 layer TiO2 2 layers

A A A A

20 30 40 50 60 70 2θ / degree

Figure 3.3 XRD patterns for TiO2 and TiO2 with 1 and 2 layers WO3 coating. Letter “A” indicates the position of antase TiO2.

Figure 3.4 shows the optical properties of the samples made by single sol-gel method. The heterostructures exhibit different absorption spectra comparing to those of either pure TiO2 or pure WO3 (purchased from Sigma-Aldrich), which indicates a possible interaction between TiO2 and WO3 in the heterostructures. It is interesting to

120 1layer 2layers 100 TiO2 WO 80 3

Reflection/ % 40

200 300 400 500 600 700 800 Wavelength / nm

Figure 3.4 UV-Vis diffuse reflection spectra of the samples made by single sol-gel method.

48 note that in the range of 550 ~ 800 nm, the one layer sample comes closer to TiO2 while in comparison, the two layers sample is more like WO3.

Figure 3.5 TEM images of various samples prepared by double sol-gel method.

The TEM images of the samples prepared by double sol-gel method have been shown in Figure 3.5. The aggregation of the heterostructures is not as heavy as that of the samples prepared by the single sol-gel method. However, the size distribution is still not uniform enough for us to reach the conclusion that the size of the nanoparticles has gained considerably from 1 layer to 2 layers. Hence we are still not sure whether WO3 exists as shell or as dopant inside the particles.

140 1 layer 120 2 layers WO3

100 TiO2

40 Reflectance /% 20

200 300 400 500 600 700 800 wavelength / nm

Figure 3.6 UV-Vis diffuse reflection spectra of the samples made by double sol-gel method.

Figure 3.6 shows the absorption spectra of the samples made by double sol-gel method. Similar to the single sol-gel products, the heterostructures exhibit absorption spectra different than either pure TiO2 or pure WO3. However, the heterostructures made by double sol-gel method resembles TiO2 more than WO3, which is opposite to the situation of the heterostructures made by the single sol-gel method.

Figure 3.7 SEM image of WO3 loaded TiO2 nanotube in low (13.0K) and high (100K) magnification. The section view of TiO2 nanotube was also given (right) for comparison.

The microscopic appearance of the TiO2 nanotube with WO3 loading is shown in

Figure 3.7. The length of the tubes is in the order of micrometer and the diameter of the cross section is about 80 nm. The phase of the heterostructure is anatase with high crystallinity according to XRD result (Figure 3.8).

TiO2 nanotube

WO3 loaded TiO2 nanotube

TiO2/WO3

TiO2

20 30 40 50 60 70 2 θ / degree

Figure 3.8 XRD patterns for the TiO2 nanotube with and without WO3 loading. The dots in the figure correspond to the characterization peak position of metal titanium. All other peaks correspond to anatase TiO2.

120

100

40 P25 TiO2 nanotube with WO3 TiO nanotube Refelectance /% 20 2

200 300 400 500 600 700 800 Wavelength / nm

Figure 3.9 UV-Vis diffuse reflection spectra of TiO2 nanotubes with and without WO3 loading.

Figure 3.9 shows the UV-Vis diffuse reflection spectra of TiO2 nanotube with and without WO3 loading. XPS result (not shown) indicates that the WO3 content in the surface of the sample is about 6%. However, the loaded WO3 did not move the absorption edge of the nanotube towards the WO3, which would be very interesting to further investigate.

3.4 Conclusion

To summarize, nanometer sized TiO2 / WO3 heterostructures have been prepared using three methods. The appearance and optical properties of these heterostructures are quite different from each other. For example, Anatase TiO2 is the only crystal structure of the prepared samples for all three cases. WO3 is suspected to be in amorphous form.

There is clear evidence that these heterostructures are not simply mixture of TiO2 and

WO3. However, extensive effort in the future is required to better understand the interaction between the two metal oxides.

3.5 References

1. Vermaire, D. C.; Vanberge, P. C. Journal of Catalysis 1989, 116, 309-317. 2. Tennakone, K.; Ileperuma, O. A.; Bandara, J. M. S.; Kiridena, W. C. B. Semiconductor Science and Technology 1992, 7, 423-424. 3. Do, Y. R.; Lee, W.; Dwight, K.; Wold, A. Journal of Solid State Chemistry 1994, 108, 198-201. 4. Papp, J.; Soled, S.; Dwight, K.; Wold, A. Chemistry of Materials 1994, 6, 496-500. 5. Burrows, A.; Kiely, C. J.; Joyner, R. W.; Knozinger, H. K.; Lange, F. Catalysis Letters 1996, 39, 219-231. 6. Li, X. L.; Lou, T. J.; Sun, X. M.; Li, Y. D. Inorganic Chemistry 2004, 43, 5442-5449. 7. Zhi-Gang Zhao, M. M. Angewandte Chemie International Edition 2008, 47, 7051- 7055. 8. Blackman, C. S.; Parkin, I. P. Chemistry of Materials 2005, 17, 1583-1590. 9. Kuang, D.; Brillet, J.; Chen, P.; Takata, M.; Uchida, S.; Miura, H.; Sumioka, K.; Zakeeruddin, S. M.; Gratzel, M. Acs Nano 2008, 2, 1113-1116.

Chapter 4 A Comparative Study of Cobalt Ion Doping of TiO2 and TiO2-xNx at the

Nanoscale: Evidence for Room Temperature Conversion from

Anatase to Rutile Phase

This work was published in Nanoscale, 2010, 2, 1134-1140 4.1 Introduction

Nanoscale titania and its oxynitride (TiO2-xNx) have attracted considerable attention due to changes in their molecular electronic structure and the increase in surface/volume ratio at the nanoscale associated with heightened reactivity.[1-11]

Nanostructured titanium dioxides have been widely used as photovoltaics, electrochromics, sensors, and photocatalysts.[12-18] Further recent studies

[4,8,13,16,17,19,20] for sol-gel generated TiO2 of average particle size 10 nm [8,19,20] suggest the importance of structural porosity as it influences and enhances the rate and efficiency of the doping processes. It has been demonstrated that the introduction of small concentrations (less than 1%) of transition metal ions can be used to enhance the catalytic activity of TiO2.[13,17] This effect has been associated with a modification of the TiO2 bandgap. However, changes in morphology and redox potentials may also be affected at higher doping concentrations.

Here, we expand our studies on porous sol-gel generated TiO2 nanocolloids and their corresponding oxynitrides TiO2-xNx to evaluate the cumulative effects of higher concentration cobalt (Co2+) ion doping. Combinations of core level, valence band photoelectron, and reflectance spectroscopies are correlated with Raman spectroscopy to demonstrate that porous sol-gel generated anatase TiO2 and TiO2-xNx can be simultaneously doped with cobalt (Co2+) ions to introduce a spinel-like structure into

54 these lattices that is commensurate with a decrease in metal-oxygen binding energy relative to TiO2 and the oxynitride. Furthermore, the conversion of both the oxide and its oxynitride from the anatase to the rutile phase occurs at room temperature. TEM, Raman, and VB photoelectron spectra demonstrate that cobalt cluster formation is negligible even at seeding levels exceeding 25 weight % cobalt. The photocatalytic activity of these systems has also been studied, and despite strong visible light absorption, we find weaker visible-light reactivity relative to the TiO2-xNx based systems.

4.2 Experimental procedures

TiO2 and TiO2-xNx nanocolloids were prepared at room temperature using a previously published synthesis.[8,19,20] In all cases the preparation of the porous TiO2 nanocolloid was done under nitrogen using the 2 propanol / Ti[OCH(CH3)2]4-acetic acid

/ doubly ionized water synthesis combination. The average particle size of the present nanocolloids is 10 nm - 20 nm. The oxynitride samples were prepared using

2+ triethylamine (TEA) as nitrogen source for the nitridation of the TiO2.[8,19,20] Co ions were introduced using the direct seeding of cobalt chloride hexahydrate at room temperature during the nanoparticle synthesis.

X-ray photoelectron spectroscopy (PHI, ESCA-5600, Al Kα X-ray source, step size: 0.4 eV) was used to characterize both the core level and valence band (VB) electronic structures of the nitrogen doped and cobalt doped TiO2 nanoparticles. Spectra were collected under ultrahigh vacuum (5x10-8 Pa) without further cleaning steps since

Ar-ion sputtering or high temperature ramping can cause undesirable changes including the reduction or oxidation of the sample as well as changes in the composition of the nanoparticles.[21] The reported binding energies were calibrated with respect to the

55 carbon 1s core level energy centered at 284.6 eV.

The XPS data are correlated with recently obtained μ-Raman results obtained using a Mitutoyo Microscope and a SPEX Triplemate Spectrometer equipped with a

CCD [22]. The minimum Raman shift of this spectrometer is 200 cm-1 (anatase band of

-1 TiO2 at 144 cm is not observed). The 514.5 nm line of an Ar ion laser was used as the excitation source. The microscope is equipped with 10 x, 50 x, and 100 x objectives for focusing the laser light, coupled to the spectrometer through a fiber optic bundle. The light from the microscope was filtered by a 514.4 nm notch filter. The positions of the

Raman lines in a given spectrum were calibrated against the 546.0 nm emission line from a fluorescent light source. Comparisons of this approach to XRD analysis are considered in detail elsewhere.[22,23]

Transmission electron microscopy (TEM) analysis was performed at 100 kV using a JEOL 100C TEM and 200 kV using a Hitachi HT-2000 TEM.

Reflectance spectra were measured on a Cary 50 UV-visible spectrometer with a reflectance Barrelino attachment. Photocatalytic measurements are carried out in a self- built reactor that was recently published.[24] All measurements were carried out at room temperature.

4.3 Results

4.3.1 Cobalt Seeding of TiO2 and TiO2-xNx Nanocolloids

2+ Porous TiO2 and TiO2-xNx nanocolloids were seeded with cobalt (Co ) ions, introduced as the chloride. The prepared sols were subsequently dried in vacuum (≤ 10-2 torr). The CoCl2 forms hexahydrate coordination complexes which are pink in color.

However, after vacuum drying on TiO2 (≤ 3 hrs), the resulting material is deep blue,

56 corresponding closely to the anhydrous cobalt ions, and remains this color for several months under dry atmosphere. In contrast, the complexes remain deep purple (dihydrate) after vacuum drying for 12-24 hour periods and readily convert back to the hexahydrate in air. The weight percent of Co2+ (Table 4.1) introduced to the titanium-based systems exceeds that of previous workers and ranges from 2 to 27 % by weight. The weight % of

2+ Co ion in Table 4.1 is calculated for seeding both with CoCl2•6H2O and

Co(NO3)2•6H2O into a prescribed concentration of TiO2 nanocolloid solution. Column 1 in Table 4.1 provides the mass of hexahydrate disposed to a 5 ml solution of the TiO2 nanocolloid which contains 134 mg of the nanocolloid. This is converted in columns 2 and 3 to a weight percent of cobalt ions.

Table 4.1 Composition of the prepared samples Weight % Co2+

a Mass of Hexahydrate / g CoCl2•6H2O / TiO2 Co(NO3)2 • 6H2O / TiO2 0.0125 2.3% 1.9% 0.0250 4.4% 3.2% 0.0500 8.5% 7.0% 0.1000 15.6% 13.1% 0.2000 27% 23.2% a Corresponds to grams of hexahydrate added to oxide or oxynitride solution

2+ 4.3.2 Outline of Raman Spectroscopy for Co Seeded TiO2 and TiO2-xNx

Nanocolloids

A 440 TiO2-CoCl2-0.0125g 20000 TiO2-CoCl2-0.025g TiO2-CoCl2-0.05g 606 TiO -CoCl -0.1g 15000 2 2 240

10000 Eg 690 A 5000 1g Raman IntensityRaman / a.u. B2g 0 300 600 900 1200 Wavenumber / cm-1

4000 B TiO2-Co(NO3)2-0.2g TiO2-Co(NO3)2-0.05g 442 TiO2-Co(NO3)2-0.025g 3000 606

2000 244 380 688 1000 E A1g

Raman IntensityRaman / a.u. g

0 B2g 0 300 600 900 1200 Wavenumber / cm-1

Figure 4.1 Raman spectra for a TiO2 nanocolloid (<10 nm) prepared with various 2+ concentrations of Co using (A) CoCl2 and (B) Co(NO3)2 at room temperature. The Raman signal was obtained using a 1 μm spot size and a power of 25 mW or less.[19,22] The typical vibration modes of rutile TiO2 are indicted by solid vertical lines.

2+ The TiO2 nanocolloids, TiO2-xNx nanocolloids, and Co doped TiO2 and TiO2- xNx have been examined using Raman Spectroscopy. This work is outlined briefly and has been discussed in detail elsewhere.[22] The spectra of untreated TiO2 powders consist only of the anatase crystal phase as identified by the three Raman lines at 400 cm-

1, 515 cm-1, and 635 cm-1.[22,25]

The Raman spectrum of the triethylamine treated porous TiO2 nanocolloid demonstrates broad peaks corresponding to the anatase crystal structure. However, an additional feature appears at 550 cm-1 [8,19,20,22,25] that has been attributed to the presence of non-stoichiometric titanium oxynitride.[20,21,25,26] The Raman spectra [22]

2+ obtained on treatment of the porous TiO2 and TiO2-xNx nanocolloids with the Co ion are outlined, in part, in Figures 4.1 and 4.2. An intense Raman signal is obtained for the

2+ CoCl2 treated TiO2 nanocolloid which decreases as a function of increasing Co concentration, as the 440 cm-1 band redshifts and the seeded nanocolloid transforms to an amorphous structure. [22]

TiON-CoCl2-0.05g TiON-CoCl2-0.025g TiON-CoCl2-0.0125g 440 608

240 690 Raman IntensityRaman / a.u.

Eg A1g B2g 0 300 600 900 1200 -1 Wavenumber / cm

2+ Figure 4.2 Raman spectra of TiO2-xNx nanocolloid with various Co concentrations.

Laser power is less than 10 mW. [22] All three samples exhibit characteristic rutile TiO2 peaks and the intensities of the peaks increase with Co2+ concentration.

The Raman signal observed for the Co(NO3)2 treated TiO2 nanocolloid is weaker,

59 but increases with increasing Co2+ concentration. Similar results are also obtained for the CoCl2 doping of the nitrided, TiO2-xNx nanocolloid depicted in Figure 4.2. The

Raman signal strength again increases with increased Co2+ concentration, similar to the results obtained for Co(NO3)2 seeding, however, the observed spectra correspond to lower cobalt ion concentrations. Furthermore, the increase in the intensity of the Raman spectra is not linear with the Co concentration, damping at higher seeding levels.

An unexpected and significant result is obtained from these Raman studies as we compare the observed features for the anatase TiO2 nanocolloid with those features that

2+ characterize the Co systems in Figures 4.1 and 4.2. Although the initial TiO2 and

TiO2-xNx nanocolloids consist exclusively of the anatase crystal structure, [8,19,20] the

Co2+ doped systems have transformed almost completely at room temperature, exhibiting the more stable rutile structure [21,22,25] with vibrational bands at ~ 235 cm-1, 440 cm-1, and 605 cm-1. [27] This structural transformation is not unique to doping with the Co2+ ion, as preliminary results suggest a similar, albeit less pronounced, change for Ni2+

(NiCl2) doped TiO2 nanocolloids. [22]

The lines at approximately 380 and 690 cm-1 do not correspond to a structural phase of TiO2. However, it has been reported previously that the vibrations of spinel

2+ -1 (Co3O4) which are associated with tetrahedral Co sites result in a 383 cm line, [21]

-1 while the A1g phonon mode of spinel has been reported at 691 cm . [28] Thus, these phonon modes result from CoO sites in the TiO2 lattice, [22] similar to those sites in

Co3O4, formed during the metal ion seeding of the solution phase TiO2 nanocolloids. The redshift of the 440 cm-1 Raman line noted in Figure 4.1(a) is consistent with the growth of the spinel phase with increasing Co content and the reduction of the TiO2 rutile

60 fraction. The sensitivity of the Raman technique reveals this spinel formation.

The Raman studies have been assessed to evaluate the potential laser induced transformation of anatase to rutile TiO2. The Raman system was therefore operated at extremely low laser powers. [22] The CoCl2 / TiO2 spectra (Figure 4.1 (a)) signal a direct room temperature transformation from anatase to rutile TiO2 and the CoCl2 / TiO2-xNx spectra result, at most, in minor part, from a laser induced surface transformation. The apparent facile in-situ cobalt induced phase transformation at room temperature certainly is surprising as the stoichiometric conversion of anatase to rutile TiO2 requires up to 12 hours at 850 °C. [26,27] This is different from more recently reported doping studies at low Co concentrations.[29]

4.3.3 Core Level and Valence Band XPS Studies

Core level XPS studies elucidate the chemical environment of the different elements in the doped nanoparticles.[30] When the TiO2-xNx is formed via the nitridation of a TiO2 nanocolloid, the N 1s and Ti 2p core levels are the most important energy regions to be investigated for incorporation of nitrogen into the TiO2 lattice and the O 1s core level locations can play an important confirmatory role. Based on a previous study

[8,19,20] of these core levels for vacuum dried nitrogen doped titania nanoparticles, we have established the formation of NO sites, characterized by an N 1s binding energy peak centered at 401.3 eV ( extending from 397.4 to 403.7 eV). While core level XPS studies provide information on changes in the most tightly bound orbitals of a nitrogen doped and Co ion seeded TiO2 matrix, an evaluation of the VB structure tuning that

2+ accompanies the Co seeding of TiO2 and TiO2-xNx in concert with N doping can be obtained from VB XPS spectra.

4.3.3 Co2+ Core Level X-ray Photoelectron Spectroscopy

Overall X-ray photoelectron spectroscopic profiles for four representative Co seedings are depicted in Figure 4. 3. Two of the profiles correspond to the CoCl2

((0.0125, 0.100) Table I) seeded TiO2 nanocolloid. The third profile is that of Co(NO3)2

(0.050) seeded into TiO2 while the fourth profile is that of CoCl2 (0.025) seeded into

TiO2-xNx. The data presented for cobalt seeded TiO2 provide evidence for increased cobalt concentration, within a few nanometers of the particle surfaces, in the region near

780 eV. The cobalt peak intensities observed for the Co (NO3)2 and TiO2-xNx – CoCl2 systems are also consistent with the relative concentrations of cobalt ion. Only the TiO2- xNx - CoCl2 system displays a clear N 1s feature.

TiO2-CoCl2-0.0125g TiO2-CoCl2-0.05g 1sO

TiO2-Co(NO3)2-0.1g C 1s

TiON-CoCl2-0.025g Ti 2p c/s Co 2p Cl 2s Cl 2p N 1s

1000 800 600 400 200 0 Binding Energy / eV

Figure 4.3 Overall Core level XPS spectra for selected samples as indicated in the Figure.

We have examined, in more detail, four areas of these core level XPS spectra:the

Ti 2p region near 460 eV, the O 1s region near 530 eV, the N 1s region near 400 eV, and the Co2+ region near 780 eV. There are distinct differences between the binding energies associated with the three Co2+ seeded systems, corresponding primarily to a lowering of the binding energies in the TiO2 lattice, which the Raman data suggests results from the introduction of the Co2+ ion. The formation of a spinel-like structure is manifested by the red shift corresponding to the 440 cm-1 Raman line in Figure 4.1(a), and the growth of the 383 cm-1 line (Figures 4.1 and 4.2) associated with the cobalt oxide spinel.

TiO2-CoCl2-0.0125g TiO2-CoCl2-0.1g TiON-CoCl2-0.025g Ti 2p3/2

Ti 2p1/2 c/s

470 465 460 455 450 Binding Energy / eV

Figure 4.4 Ti 2p XPS peaks for the CoCl2 / TiO2 and CoCl2 /TiON systems.

2+ The Ti 2p region (Figure 4.4) for the Co -TiO2 system, whose Raman spectra correspond to those in Figure 4.1, demonstrates a red shift with increasing cobalt ion concentration. This shift extends from a Ti(+4) binding energy close to that measured for

the TiO2 nanocolloid (458.8 eV) [8,19,20] at the lowest concentrations, to a value close to 458.3 eV as the Co2+ concentration is increased by a factor of eight. The shift is consistent with that which should accompany the formation of the cobalt oxides in the

TiO2 lattice and parallels the conversion of anatase to rutile TiO2. [31] The Raman spectra demonstrate that this conversion is well on its way even at the lowest cobalt concentrations. The observed incorporation of Co and the appearance of spinel-like structure in the Raman spectrum for the seeded TiO2 lattice suggests that the XPS spectrum might also signal the transfer of charge from Co2+ to Ti4+, at increased co2+ concentration, and the subsequent partial formation of Co3+ and Ti3+.

The XPS spectra for the cobalt doped TiO2-xNx nanocolloid show a peak near

458.1 eV for the Ti 2p3/2 binding energy, weaker BE relative to that of the TiO2 nanocolloid (458.6).[8,20] The corresponding N 1s spectra of CoCl2-TiO2-xNx is depicted in Figure 4.5. As expected, no N 1s XPS spectra are observed for the CoCl2 seeded TiO2 nanocolloid. The XPS spectrum for the cobalt seeded TiO2-xNx nanocolloid also displays a peak at ~ 401.3 eV, due to nitric oxide (NO).[8,20] The typical binding energy for TiN is 397.2 eV[32,33] and Figure 4.5 demonstrates that neither TiN nor CoN has been formed.

TiON-CoCl -0.025g 2 TiO2-CoCl2-0.0125g N 1s c/s

410 405 400 395 390 Binding Energy (eV)

Figure 4.5 N 1s XPS spectra for CoCl2 seeded TiO2 and TiO2-xNx.

Rodriguez, et al.[31, 34] have carried out an XPS analysis of the interaction of

NO2 with several polycrystalline surfaces, observing a peak which they assign to for

NO2 at 404.5 eV[35,36], and to absorbed NO at 400-401.5 eV.[34,35] They also assign the 401 eV feature to NO.[8,20] Similar N1s features are observed when the TiO2 nanocolloid is seeded with Pd(NO3)2 to increase nitrogen uptake. [8,19,20]

VB XPS Spectra for the Co2+ doped systems can be compared to the corresponding room temperature spectra for the titanium oxide and titanium oxynitride, and that for the precursor CoCl2. Figure 4.6 summarizes the observed VB XPS spectra

2+ for CoCl2 seeded TiO2. With increase in the Co concentration they show a considerable interaction of Co with the TiO2 lattice. At the lowest concentrations for this study (Table 4.1) we observe a significant broadening of the O 2s region and a decrease in the peak binding energy from ~ 22.8 eV to 22 eV as the Co2+ concentration is doubled.

At higher cobalt concentrations, the XPS spectrum is dominated by the O 2s native cobalt oxide signal which continues to broaden to lower BE energy with increasing Co2+ concentration. Features associated with bulk cobalt,[37] and small cobalt clusters with a

Fermi level shifted by 1.8 - 1.9 eV toward higher binding energy[38] are absent.

TiO2-CoCl2-0.0125g CoCl2 TiO2-CoCl2-0.025g TiO2 TiO2-CoCl2-0.05g TiO2-CoCl2-0.5g c/s

30 20 10 0 Binding Energy (eV)

Figure 4.6 VB-XPS of TiO2 with different CoCl2 loading. For comparison, the VB-XPS spectra of TiO2 and CoCl2 are also shown.

The data in Figure 4.6, in conjunction with the Raman spectra, suggest the formation of Co3+-oxide as well as Co2+-oxide bonds commensurate with the formation of the cobalt oxide spinel. It is apparent that the porous nature of the nanocolloids allows the highly efficient diffusion of the metal ion seed to the TiO2 and TiO2-xNx lattice sites.

At the highest CoCl2 concentration a peak at 17 eV occurs, consistent with a partially substituted cobalt oxide within the TiO2 lattice. This feature, albeit slightly red shifted, is apparent in the spectrum for CoCl2 and may suggest the formation of an oxychloride.

This data suggests that it is not surprising that the Raman spectra for the TiO2 / CoCl2 system (Figure 4.1(a)) quench readily as a function of increased concentration as both the titanium oxide is weakened by substitution and the increased Co2+ seeding also perturbs the titanium dioxide crystal structure sufficiently to reduce it to a nearly amorphous state

(Figure 4.1(a)).

TiON-CoCl2-0.0125g TiON TiON-CoCl2-0.025g CoCl2 TiON-CoCl2-0.05g TiON-CoCl2-0.25g c/s

30 20 10 0 Binding Energy / eV

Figure 4.7 Room temperature VB XPS spectra of CoCl2 / TiO2-xNx system with increased 2+ Co concentration. For comparison, the same spectra of TiO2-xNx and CoCl2 are also included.

Figure 4.7 outlines the observed VB XPS spectra for the CoCl2 seeded TiO2-xNx nanocolloid. Here, the peak of the O 2s spectrum again shifts to a lower ~ 21 eV (vs.~

22 eV binding energy associated exclusively with the TiO2 nanocolloid). The O 2p region also demonstrates a lower peak binding energy, close to 5 eV. While we suggest that these decreases in binding energy are likely due to the Co replacement of Ti, they can also be associated with N 2s , O 2s and N 2p, O 2p orbital mixing catalyzed by the

67 introduction of cobalt.

At the lowest Co2+ concentrations for this study (compare with Table I) we observe a very significant broadening of the O 2s region such that the peak apparent for the oxynitride as well as the peaks observed in the TiO2 / CoCl2 system are strongly broadened and reduced. With increased Co2+ seeding, however, more clearly defined binding energy regions appear that correspond to the cobalt oxide. These regions, while at first slightly blue shifted relative to the oxynitride, subsequently transition to a comparable O 2s B.E. Based on the oxynitride VB XPS spectrum, it seems clear that the combination of Co2+ ions interacting with the oxynitride catalyzes this interaction.

While the Core level XPS spectra suggest that the oxynitride dominates the observed O

1s feature observed for this system, it seems apparent that the presence of Co2+ catalyzes the broadening and signals a strong synergistic interaction between Co2+ and the oxynitride that influences the nature of the observed Raman[22] and enhanced infrared spectra [22]. At the highest Co2+ concentrations (0.25) there is a considerably increased spectral broadening extending to the region between 15 and 12 eV. This is accompanied by a considerable lowering of the O 2p binding energy (to ~4 eV) to a level well below that of the oxynitride. For both the native oxynitride and the cobalt seeded systems, these features may result, in part, from the formation of NO sites in this region.[39]

Figure 4.8 TEM images of (a) CoCl2 / TiO2-xNx and (b) CoCl2 / TiO2

The TEM images presented in Figure 4.8 indicate the cobalt seeded TiO2 and

TiO2-xNx systems. The titania and titanium oxynitride particles are of average size 10 nm in these figures and the cobalt chloride molecules are intermingled among these nanoparticles.

120

P25 TiO2-CoCl2-0.025g 80 TiO2-CoCl2-0.1g

40 Reflectance / a.u.

200 400 600 800 Wavelength / nm

120

TiON-CoCl2-0.025g 80 TiON-CoCl2-0.1g

40 Reflectance / a.u.

200 400 600 800 Wavelength / nm

Figure 4.9 Reflectance Spectra of P25, cobalt doped TiO2 and TiON nanoparticles with varying Co content as indicated in the legend. Comparing to P25, the cobalt doped samples, even with a Co content as low as 0.025 g, exhibit a sharp difference in the visible-light region. The features associated with Co are more damped in the TiON matrix than those in the TiO2 matrix. The measurements are done relative to MgO, which explains the >100% reflectance of pure P25 TiO2.

The UV-Vis diffuse reflectance spectra of the cobalt doped TiO2 and TiON nanoparticles are shown in Figure 4.9. In comparison with P25, the commercial available

TiO2, these doped samples absorb significantly in the visible-light region, which should be ascribed mostly to the existence of cobalt. However, the features of these cobalt

70 absorptions are relatively weaker in TiON matrix than those in TiO2.

100 100

90 90

80 no catalyst 80 TiO2-CoCl2-0.025g TiO2-CoCl2-0.05g 70 70 TiON-CoCl2-0.025g TiON

Percentage of blue Methylene left / % 60 60 0 1 2 3 4 5 Time / hour

Figure 4.10 Photodecomposition of methylene blue under visible light (> 400 nm) illumination with various photocatalysts.

Shown in Figure 4.10 is the photocatalysis performance of the as-prepared nanoparticles. Although the Co-doped samples exhibit great absorption in the visible- light region, their photocatalytic activity is not as high as that of the TiON sample.

4.4. Discussion

The TiO2 and TiO2-xNx nanocolloids used in this study have been characterized previously using X-ray powder diffraction techniques (XRD) which, based on the

Debye- Scherrer equation [8, 19, 20] suggest that they have average diameters of 10 nm.[8] Broadened but regular XRD patterns are readily obtained for the anatase structured nanocolloids.[8,19,20] However, we have recently noted [22, 23] that it is difficult to obtain a room temperature Raman spectrum for these small nanocolloids without the sophisticated collection optics used in the present study,[8,19,20,22]

71 discussing the advantages and disadvantages of Raman and XRD characterization at the nanometer scale. [22,23]

In a system heated to elevated temperatures or a system in which doping is used to obtain the anatase to rutile phase transformation, the introduction of crystalline disordering with even moderate doping levels or increasing temperature can lead to significant broadening associated with disorder in the XRD spectrum. This renders mute the use of the Debye-Scherrer equation to evaluate crystallite size and a clear phase structure. The results which we have outlined previously [22,23] point to the difficulty of using XRD to characterize extremely small structures when a phase transformation and dopant seeding, in combination, are involved. However, in the application of Raman scattering, the two phases of TiO2 and its oxynitride have very distinct, widely spaced, fundamental vibrational frequencies as is also the case for the cobalt based spinel. This allows for the identification of the specific constituents and phases present in the various samples, and is the driving force for the correlations of the XPS and Raman data that we have made in this study.

While the role of dopants introduced into the TiO2 lattice is controversial,[40-43] the present study demonstrates that the cobalt cation, when introduced as a dopant, can induce a significant transformation of anatase to rutile TiO2 at room temperature. The

Raman spectra depicted in Figures 4.1, and 4.2 demonstrate a range of completeness for

-1 this transformation. The conversion of the 634 cm stretch mode of anatase TiO2 to the

-1 608 cm stretch mode of rutile TiO2 is readily completed and, we find little evidence for the broadening that might be expected to accompany a soft mode transformation.

Although the anatase phase of TiO2 can be readily produced at the 10 nm scale

72 using sol-gel techniques,[8,19,20] it is metastable with respect to the rutile phase.

Furthermore, the anatase phase reverts to the rutile phase at temperatures above

915 °C,[44] although the process is kinetically unfavorable at low temperatures. The transformation process, partially kinetically driven, can begin at temperatures between

700° and 800 °C.[23] Navrotsky and Kleppa [44] determined a negative enthalpy change for the anatase to rutile transformation, reporting that anatase is metastable with respect to rutile under all conditions of temperature and pressure. The JANAF Tables [45] demonstrate that the free energy of formation of rutile is always less than anatase making it the more stable structure. These factors, and the need to produce structures for the rutile phase of TiO2 which are comparable to those obtained for the sol-gel generated anatase phase suggests the importance of the cobalt ion dopant for the anatase to rutile conversion. The very small difference in the free energies of anatase and rutile TiO2 is the primary reason for their slow thermodynamic transformation. The introduction of cobalt ions would appear to drive the conversion, possibly by a catalyzed process. The modeling of this process will require considerable further study.

The Raman spectra in Figure 4.2 for the anatase phase show a virtually complete anatase to rutile transformation, the red shift associated with increasing Co content for only the 440 cm-1 band, being accompanied by the growth of the spinel phase (modes at

383 cm-1 and 690 cm-1). As more of the spinel phase forms, there is a decrease in the

-1 intensity of the 606 cm line, associated with the A1g vibration mode of the TiO2 rutile lattice. The conversion to rutile is consistent with the formation of the spinel structure within the TiO2 matrix, and the clear indication of a lowering of the O 2p and O 2s binding energies. By comparison, the transformation is clear but notably less effective in

73 the Ni2+ system and virtually absent in the Cu2+ system [22] but comparable in the higher spin Fe2+ system. Clearly, the exact mechanistic of these systems merits further study.

The evidence for an anatase-to-rutile conversion at the lowest Co concentrations may have exciting consequences. Hurrum, et al.[46] suggest that one can create a material of enhanced photocatalytic activity by controlling the degree of anatase to rutile transformation so as to form anatase/rutile interface regions.[22,23] The assessment of this exciting possibility for the cobalt ion seeded systems will, however, require further kinetic in situ Raman, XPS, and EPR spectroscopic studies. However, initial studies suggest that the photo-catalytic efficiency of these systems does not approach that of the oxynitride or TiO2.

By comparison, the formation of the oxynitride enhances the Raman spectrum. At

2+ the lowest concentrations of Co in the CoCl2 / TiO2-xNx system, the spinel seeded rutile spectrum is clearly observable, and the spinel signal increases with decreasing Co content, which leads to a stronger rutile signal.

5. Conclusion

2+ It is apparent that the introduction of Co into either a TiO2 or TiO2-xNx considerably lowers the activation energy for the anatase to rutile phase conversion.

Based on simple considerations, the activation energy of the reaction may be lowered by as much as a factor of four. It is also important to note that we are operating at much higher cobalt ion seeding levels than have typically been employed in these experiments.

It is not surprising that the cobalt ions have a significant influence on the electronic structure of the titania, leading to visible light absorbance and to the formation of cobalt oxide sites with a concomitant lowering of the binding energy verses the titanium oxides

74 and oxynitrides. The conduction band edges of titania are clearly affected by the presence of cobalt ions. Although the inherent visible light absorbance does not lead to a visible light absorbing photocatalyst comparable to the oxynitride, it is clear that catalysis has occurred, which suggests that the quenching of catalytic effects at low metal doping concentrations may be countered. Based on the Raman, TEM, and VB photoelectron data, it appears that the N-doped TiO2 nanoparticles are more efficient at taking up cobalt ions. We found no evidence for the formation of cobalt cluster sites even at the high seeding levels used in these experiments.

4.5 References

1. A. Fujishima and K. Honda, Nature, 1972, 238, 37-38. 2. R. Asahi, T. Morikawa, T. Ohwaki, K. Aoki and Y. Taga, Science, 2001, 293, 269-271. 3. J. L. Gole and M. G. White, Journal of Catalysis, 2001, 204, 249-252. 4. C. Burda, Y. B. Lou, X. B. Chen, A. C. S. Samia, J. Stout and J. L. Gole, Nano Letters, 2003, 3, 1049-1051. 5. J. L. Gole, B. D. Shinall, A. V. Iretskii, M. G. White, W. B. Carter and A. S. Erickson, Chemphyschem, 2003, 4, 1016-1021. 6. X. B. Chen and C. Burda, Journal of Physical Chemistry B, 2004, 108, 15446- 15449. 7. C. Burda, X. B. Chen, R. Narayanan and M. A. El-Sayed, Chemical Reviews, 2005, 105, 1025-1102. 8. X. B. Chen, Y. B. Lou, A. C. S. Samia, C. Burda and J. L. Gole, Advanced Functional Materials, 2005, 15, 41-49. 9. Y. Liu, X. Chen, J. Li and C. Burda, Chemosphere, 2005, 61, 11-18. 10. X. F. Qiu and C. Burda, Chemical Physics, 2007, 339, 1-10. 11. X. B. Chen and C. Burda, Journal of the American Chemical Society, 2008, 130, 5018-5019. 12. M. R. Hoffmann, S. T. Martin, W. Y. Choi and D. W. Bahnemann, Chemical Reviews, 1995, 95, 69-96. 13. M. Anpo, H. Yamashita, Y. Ichihashi, Y. Fujii and M. Honda, Journal of Physical Chemistry B, 1997, 101, 2632-2636. 14. A. Mills and S. LeHunte, Journal of Photochemistry and Photobiology a- Chemistry, 1997, 108, 1-35. 15. M. Gratzel, Nature, 2001, 414, 338-344. 16. S. Kumar, A. G. Fedorov and J. L. Gole, Applied Catalysis B-Environmental, 2005, 57, 93-107. 17. L. Xiao, J. L. Zhang, Y. Cong, B. Z. Tian, F. Chen and M. Anpo, Catalysis Letters, 2006, 111, 207-211. 18. T. Gerfin, M. Graezel and L. Walder, in Progress in Inorganic Chemistry, ed. D. K. Kenneth, 2007, pp. 345-393. 19. J. L. Gole, J. D. Stout, C. Burda, Y. B. Lou and X. B. Chen, Journal of Physical Chemistry B, 2004, 108, 1230-1240. 20. S. M. Prokes, J. L. Gole, X. B. Chen, C. Burda and W. E. Carlos, Advanced Functional Materials, 2005, 15, 161-167. 21. C. F. Windisch, G. J. Exarhos and R. R. Owings, Journal of Applied Physics, 2004, 95, 5435-5442. 22. J. L. Gole, S. M. Prokes and O. J. Glembocki, Journal of Physical Chemistry C, 2008, 112, 1782-1788.

23. A. Ogden, J. A. Corno, J. I. Hong, A. Fedorov and J. L. Gole, Journal of Physics and Chemistry of Solids, 2008, 69, 2898-2906. 24. X. B. Chen, S. M. Halasz, E. C. Giles, J. V. Mankus, J. C. Johnson and C. Burda, Journal of Chemical Education, 2006, 83, 265-267. 25. A. Brevet, F. Fabreguette, L. Imhoff, M. C. M. de Lucas, O. Heintz, L. Saviot, M. Sacilotti and S. Bourgeois, Surface & Coatings Technology, 2002, 151, 36-41. 26. E. Gyorgy, A. P. del Pino, P. Serra and J. L. Morenza, Applied Surface Science, 2002, 186, 130-134. 27. M. E. Straumanis, T. Ejima and W. J. James, Acta Crystallographica, 1961, 14, 493-497. 28. M. J. Escudero, T. Rodrigo, L. Mendoza, M. Cassir and L. Daza, Journal of Power Sources, 2005, 140, 81-87. 29. J. Choi, H. Park and M. R. Hoffmann, Journal of Physical Chemistry C, 2010, 114, 783-792. 30. R. Sanjines, H. Tang, H. Berger, F. Gozzo, G. Margaritondo and F. Levy, Journal of Applied Physics, 1994, 75, 2945-2951. 31. J. A. Rodriguez, T. Jirsak, J. Dvorak, S. Sambasivan and D. Fischer, Journal of Physical Chemistry B, 2000, 104, 319-328. 32. N. C. Saha and H. G. Tompkins, Journal of Applied Physics, 1992, 72, 3072-3079. 33. E. Gyorgy, A. P. del Pino, P. Serra and J. L. Morenza, Surface & Coatings Technology, 2003, 173, 265-270. 34. T. Jirsak, J. Dvorak and J. A. Rodriguez, Surface Science, 1999, 436, L683-L690. 35. W. M. R. C.D. Wagner, L.E. Davis, J.R. Moulder, G.E. Muilenberg Handbook of X-ray Photoelectron Spectroscopy and E. P. published by Perkin-Elmer Corp., MN, USA, 1979. 36. L. A. Delouise and N. Winograd, Surface Science, 1985, 159, 199-213. 37. M. V. B.V. Crist Handbook of Monochromatic XPS Spectra (Elements and Native Oxides) published by XPS International LLC, CA, USA, 2004. 38. G. Peto, G. Molnar, G. Bogdanyi and L. Guczi, Catalysis Letters, 1994, 26, 383- 392. 39. K. Irokawa, S. Ito, T. Kioka and H. Miki, Surface Science, 1999, 435, 297-301. 40. C. N. R. Rao, A. Turner and J. M. Honig, Journal of Physics and Chemistry of Solids, 1959, 11, 173-175. 41. S. Robert D and P. Joseph A, Journal of the American Ceramic Society, 1965, 48, 391-398. 42. S. Ram, R. Larry and H. D. Burtron, Journal of the American Ceramic Society, 1990, 73, 3528-3530. 43. P. F. Becher and M. V. Swain, Journal of the American Ceramic Society, 1992, 75, 493-502. 44. C. Burda and M. A. El-Sayed, Pure and Applied Chemistry, 2000, 72, 165-177.

45. J. T. T. D.R. Stull, ” Joint Army-Navy-Air Force-ARPA-NASA Thermochemical Working Group, 1996. 46. D. C. Hurum, A. G. Agrios, K. A. Gray, T. Rajh and M. C. Thurnauer, The Journal of Physical Chemistry B, 2003, 107, 4545-4549.

Chapter 5 Facile Charge Transfer, Phase Transformation, and the Analog of

Nitrogen Fixation in Heavily Iron Ion Doped Titanium Oxide

and Oxynitride Nanocolloids

This work is submitted for publication.

5.1 Introduction

Recent studies [1-7] of sol-gel generated TiO2 of average particle size 10 nm [2,

4-7] suggest the importance of structural porosity as it influences and enhances the rate and efficiency of doping processes. The introduction of small concentrations (< 1 %) of transition metal ions is known to influence the catalytic activity of TiO2 [4]. The result of the doping has been attributed to a modification of the TiO2 bandgap. However, at higher doping concentrations, modified morphologies and redox potentials that characterize nanophase systems can produce unique phase changes [8] and result in facile oxidation state changes.

Nanoscale TiO2 and its oxynitride, TiO2-xNx, have attracted significant attention primarily due to changes in their molecular electronic structure and their heightened activity [1,2,8-16] due, in part, to an enhanced surface/volume ratio. Titania-based nanocolloids have been used as photovoltaics, electrochromics, sensors, and photocatalysts[3-5,17-20]. Here, we expand our previous studies of porous sol-gel generated TiO2 nanocolloid and its corresponding oxynitride, TiO2-xNx, to evaluate the cumulative effects of additional high iron (Fe2+ ) ion doping. We find that this high level of transition metal seeding results in several novel and unexpected effects. Combinations of core level and valence band photoelectron spectroscopy are correlated with the results

79 of DFT calculations to demonstrate (1) not only a facile charge transfer from Fe2+ to Ti4+, producing Ti3+ and Fe3+, but also the subsequent formation of Ti2+ via the process Fe2+ +

Ti3+ to Fe3+ + Ti2+ associated with the formation of Ti(II) at the surface of the titania- based nanoparticles and suggest the possible conversion of nitrogen to its oxides in a process analogous to nitrogen fixation, (2) the surprising room temperature facile conversion of the anatase forms of TiO2 and TiO2-xNx, generated in the sol-gel preparation, to the rutile form in contrast to the usually observed minimum temperature of 550ºC, and (3) the manifestation of an oscillating catalytic activity. Raman spectroscopy, in direct analogy to results previously obtained for Co2+ ion doping,21-23

2+ suggests that the doping of TiO2 and TiO2-xNx with iron (Fe ) ions introduces an iron- based spinel-like structure to these nanocolloids commensurate with a decrease in the metal-oxygen binding energy relative to TiO2 and its oxynitride. The photocatalytic activity of these systems is comparable to that of the oxynitride, TiO2-xNx, and appears to demonstrate the possibility of multiple “sweet” zones for catalytic activity as a function of changing Fe2+ concentration. We first consider the XPS and DFT results as they apply facile charge transfer. We then consider Raman spectroscopy as it demonstrates phase transformation. TEM and XRD analyses, which correlate, are considered. We conclude with a discussion of DRS optical spectroscopy and an assessment of photocatalytic behavior relative to TiO2 and TiO2-xNx. We follow this with a Discussion and Conclusion.

5.2 Experimental procedures

TiO2 and TiO2-xNx nanocolloids were prepared at room temperature using a previously published synthesis.[2,6,7] In all cases, the preparation of the porous TiO2 nanocolloid was done under nitrogen using the 2 propanol/Ti[OCH(CH3)2]4-acetic

80 acid/doubly deionized water synthesis combination. The average particle size of the present nanocolloids is 10 nm - 20 nm. The oxynitride samples were prepared using

2+ triethylamine (TEA) as the nitrogen source for the nitridation of the TiO2.[2,6,7] Fe ions were introduced using the direct seeding of ferrous chloride tetrahydrate or anhydrous FeCl2 at room temperature following the TiO2 nanoparticle synthesis.

The TiO2 nanocolloids can be stored under refrigeration for extended periods; however, at room temperature they will form a semitransparent gel within few days. At the lowest Fe2+ doping levels, we find that the gelation process is virtually unimpeded.

However, as the Fe2+ concentration is increased, the time period for gelation increases precipitously. The process of nitridation is carried out on the TiO2 nanocolloids, which have previously been seeded with Fe2+ ions. Here also the presence of Fe2+ ions impedes

(1) the rate at which the oxynitride forms and (2) the degree of precipitation inherent to the oxynitride [24].

X-ray photoelectron spectroscopy (PHI, ESCA-5600, Al Kα X-ray source, step size: 0.4 eV) was used to characterize both the core level and valence band (VB) electronic structures of the nitrogen doped and iron doped TiO2 nanoparticles. Spectra were collected under ultrahigh vacuum (5x10-8Pa) without further cleaning steps since

Ar-ion sputtering or high temperature ramping can cause undesirable changes including the reduction or oxidation of the sample as well as changes in the composition of the nanoparticles.[1,12] The reported binding energies were calibrated with respect to the carbon 1s core level energy centered at 284.6 eV.

The XPS data are correlated with μ-Raman results obtained using a Mitutoyo

Microscope and a SPEX Triplemate Spectrometer equipped with a CCD [21]. The

81 minimum Raman shift that can be detected with this spectrometer is 200 cm-1 (anatase

-1 band of TiO2 at 144 cm is not observed). The 514.5 nm line of an Ar ion laser was used as the excitation source. The microscope is equipped with 10 x, 50 x, and 100 x objectives for focusing the laser light, coupled to the spectrometer through a fiber optic bundle. The spot size was kept on the order of 100 μm in order to reduce the power density. The light from the microscope was filtered by a 514.5 nm notch filter. The positions of the Raman lines in a given spectrum were calibrated against the 546.0 nm emission line from a fluorescent light source. Comparisons of this approach to XRD analysis are considered in detail elsewhere.[21,26]

Transmission electron microscopy (TEM) analysis was performed at 60 kV using a Phillips CM 12 TEM. The HRTEM study was performed on a Zeiss LIBRA 200FE at

200 kV.

Reflectance spectra were measured on a Cary 50 UV-visible spectrometer with a reflectance Barrelino attachment. Photocatalytic measurements are carried out in a self- built reactor whose description was recently published [27]. All measurements were carried out at room temperature.

Calculations to study the binding energies of the Ti4+, and Ti2+ states [28] of model compounds were carried out at the density functional theory (DFT) level with the

B3LYP exchange correlation functional [29] and the DZVP2 basis set [30] following our prior work on interpreting the XPS spectrum of Cabosil [31]. The structures of a variety of TiO2 nanoclusters have been reported previously [32,33]. Here we compare the 2s and

2p binding energies. Additional data on the 3s and 3p binding energies and, the O 1s

82 binding energies are discussed elsewhere [28]. All calculations were done with the

Gaussian program system [34].

5.3 Results

XPS Spectroscopy

Figures 5.1(a)-5.1(f) and 5.2(a)-5.2(f) depict the core level and valence band

2+ 2+ photoelectron spectra obtained for Fe doped TiO2 and TiO2-xNx respectively. The Fe doping levels (mole %) are defined in Table 5.1. Figures 5.1(a) and 5.2(a) present the

2+ full-scale core level XPS spectra for doped Fe TiO2 and TiO2-xNx. These spectra and the higher resolution spectra, which follow, are generated at considerably higher Fe2+ ion doping concentrations than previous transition metal ion doping studies [4].

Figure 5.1(b) compares the valence band (VB) photoelectron spectra for several

2+ Fe ion doping levels versus the VB spectra for TiO2 and TiO2-xNx. The VB XPS spectra for TiO2 consist of two oxide related peaks near 22-24 eV that is generally associated with the O 2s region and the transition metal-oxygen bond [35-38]. The region between 3 and 9 eV, attributed to the O 2p orbitals in pure TiO2, is very sensitive to the Ti-Ti and

Ti-O distances.[39,40] Upon formation of the oxynitride, the valence band edge [35] shifts toward lower binding energy as a result of significant electronic structure changes

[41]. The O 2p orbital is also dramatically altered as clear features associated with pi nonbonding (~5-6 eV) and sigma bonding (~8 eV) are not discerned and a broad feature at 5-6 eV is observed.[39,40] With increased nitrogen doping, we produce an effective bandgap narrowing of ~1eV, most likely due to a mixing of N 2p and O 2p orbitals [42].

No additional peaks emerge in the bandgap.[35,42]

Table 5.1 Composition of the Prepared Fe2+ Doped Samples (mole % of Fe2+ in total II composition of Fe + TiO2 or TiO2-xNx )

a 2+ Corresponding mass of Mass of Tetrahydrate /g Fe /TiO2 anhydrous FeCl2 /g

0.00312 0.94% 0.002 0.00625 1.9% 0.004 0.0125 3.7% 0.008 0.025 7.1% 0.016 0.050 13.2% 0.032 0.100 23.3% 0.064 0.200 37.9% 0.128 a Corresponds to grams of tetrahydrate added to titanium oxide or titanium oxynitride

solution

TiO2 FeCl2 0.1 TiO FeCl 0.05 A 2 2 B FeCl2 TiO2 FeCl2 0.025

TiO2 FeCl2 0.1 TiO2 FeCl2 0.0125 TiO2

TiO2 FeCl2 0.05 TiO2 FeCl2 0.1 c/s

c/s TiO FeCl 0.05 TiO2 FeCl2 0.025 2 2 TiO2 FeCl2 0.025 TiO2 FeCl2 0.0125 TiO2 FeCl2 0.0125

TiON

1000 800 600 400 200 0 30 20 10 0 Binding Energy (eV) Binding Energy (eV) TiO FeCl 0.1 2 2 TiO2 FeCl2 0.1 TiO FeCl 0.05 C D Fe 2p3/2 2 2 TiO2 FeCl2 0.05 TiO FeCl 0.025 TiO FeCl 0.1 2 2 Ti 2p3/2 2 2 TiO2 FeCl2 0.025 TiO FeCl 0.0125 2 2 TiO2 FeCl2 0.0125 Ti 2p1/2 TiO2 FeCl2 0.05 TiO FeCl 0.1 c/s

c/s 2 2

TiO2 FeCl2 0.05 TiO2 FeCl2 0.025

TiO2 FeCl2 0.025 TiO2 FeCl2 0.0125

TiO2 FeCl2 0.0125 470 465 460 455 450 730 720 710 700 690 Binding Energy (eV) Binding Energy (eV)

TiO FeCl 0.1 2 2 TiO2 FeCl2 0.1 E TiO FeCl 0.05 F O 1s 2 2 TiO2 FeCl2 0.05 TiO FeCl 0.025 2 2 TiO2 FeCl2 0.025 TiO FeCl 0.0125 N 1s 2 2 TiO2 FeCl2 0.1 TiO2 FeCl2 0.0125

TiO2 FeCl2 0.1 TiO2 FeCl2 0.05 c/s c/s

TiO2 FeCl2 0.05 TiO2 FeCl2 0.025

TiO2 FeCl2 0.025 TiO2 FeCl2 0.0125

TiO2 FeCl2 0.0125

540 530 520 410 405 400 395 390 Binding Energy (eV) Binding Energy (eV)

Figure 5.1 (a) Fe doped TiO2 core-level XPS spectra of all samples -VB-XPS spectra for Fe doped TiO2 compared to TiO2 samples. (b) VB-XPS spectra for Fe doped TiO2 compared to TiO2 and TiON samples. (c) Ti 2p core-level XPS spectra for Fe doped TiO2. (d) Fe 2p core-level XPS spectra for Fe doped TiO2. (e) O 1s core level XPS spectra for Fe doped TiO2. (f) N 1s core-level XPS spectra for Fe doped TiO2.

TiON FeCl 0.2 CoCl TiO 2 A 2 2 B TiON FeCl 0.05 2 VB-TiON FeCl2 0.2 VB-TiON FeCl2 0.05 TiON FeCl 0.025 2 VB-TiON FeCl2 0.025 TiON CoCl TiON FeCl2 0.2 2

TiO2

VB-TiON FeCl 0.2 c/s TiON FeCl2 0.05 2 c/s

VB-TiON FeCl2 0.05 TiON FeCl 0.025 2 VB-TiON FeCl2 0.025 TiON

1000 800 600 400 200 0 30 20 10 0 Binding Energy (eV) Binding Energy (eV)

TiON FeCl 0.2 TiON FeCl 0.2 2 C D 2 TiON FeCl 0.05 TiON FeCl 0.05 2 Fe 2p3 2 TiON FeCl 0.025 TiON FeCl2 0.025 2 Ti 2p3/2 Ti 2p1/2 TiON FeCl2 0.2

TiON FeCl2 0.2 c/s c/s TiON FeCl 0.05 TiON FeCl2 0.05 2

TiON FeCl 0.025 TiON FeCl2 0.025 2

470 465 460 455 450 730 720 710 700 690 Binding Energy (eV) Binding Energy (eV)

TiON FeCl2 0.2 TiON FeCl2 0.2 TiON FeCl 0.05 E F TiON FeCl 0.05 2 N 1s 2 TiON FeCl 0.025 TiON FeCl 0.025 2 O 1s 2

TiON FeCl2 0.2

TiON FeCl2 0.2 c/s c/s

TiON FeCl2 0.05

TiON FeCl2 0.025

550 540 530 520 510 410 405 400 395 390 Binding Energy (eV) Binding Energy (eV)

Figure 5.2 (a) Fe doped TiON core-level XPS spectra of all samples. (b) VB-XPS spectra for Fe doped TiON compared to TiO2 and TiON. (c) Ti 2p core level XPS spectra for Fe doped TiON. (d) Fe 2p3 core-level spectra for Fe doped TiON. (e) O 1s core-level XPS spectra for Fe doped TiON. (f) N 1s core-level XPS spectra for Fe doped TiON.

2+ The photoelectron spectra for the Fe doped TiO2 nanocolloid are shifted to considerably lower binding energies. However, at higher Fe2+ doping levels, the spectra show a shift to higher binding energy, while still not approaching the binding energy for the O 2s and O 2p levels of TiO2. Similar characteristic shifts (Table 5.2) are observed for the Ti 2p core level XPS spectra depicted in Figure 5.1(c). While first displaying a shift to lower binding energies for the lower concentrations, at higher concentrations the shift is to higher binding energy. All of the observed spectra lie at lower binding energy relative to the (2p1/2-2p1) peak for the undoped TiO2 nanocolloid at 458.8 eV[2]. The observed spectra strongly suggest that doping with Fe2+ results in a charge transfer that converts the surface localized Ti4+ centers to Ti3+ with a concomitant effective oxidation of the Fe2+ to Fe3+ (Table 5.2). In addition, while the shift at lower concentrations suggests the transformation of Ti4+ to Ti3+ at a lower binding energy, the process of charge transfer appears to go further. The shift to higher binding energies for the higher

Fe2+ seeding concentrations is attributed to the process Fe2+ + Ti3+ to Fe3+ + Ti2+ . In fact, our DFT calculations [28] summarized in Table 5.3 demonstrate that the binding energy of formal Ti2+ centers can exceed that of Ti3+ and Ti4+. The calculations support the assignment of a shift to higher binding energy in Figures 5.1(b,c) and Figure 5.2(b,c) to a partial conversion to Ti2+ at the surface of the seeded nanoparticle. Consistent also with these observations is (1) a significant increase in binding energy (from 710 to ~712 eV.) for the Fe 2p XPS peak (0.1-23.3 mole %-Table 5.1) associated with the highest Fe2+ doping level depicted in Figure 5.1(d) compared to the corresponding peak in the XPS spectrum of FeCl2,[38] (2) a possible slight decrease in the binding energy which is also appears to be corroborated by XAS data43.

Table 5.2 2s and 2p Core level XPS Binding Energies as Function of Fe2+ Doping Level (eV) - for several mole fractions (peak intensities for 2p3/2 (2p3) and 2p1/2 (2p1) are indicated)

Mole % Fe (II) Ti Fe O N

3.67 – TiO2 457.8(2p3), 463.5(2p1) - 530.2 -

7.07 – TiO2 458.2(2p3), 464.0(2p1) - 530.4 -

13.2 – TiO2 458.8 (3), 464.7 (1) 711.2 530.6 -

23.3 – TiO2 458.9 (3), 464.7 (1) 711.8 530.8 - 7.07 – TiON 458.0 (3), 463.4 (1) 710.6 530.0 400.2 13.2 – TiON 457.4 (3), 462.9 (1) 710.1 529.5 400.2 37.9 – TiON 458.2 (3), 464.5 (1) 711.4 531.2 401.3

Further, the associated increase in the O 1s core level XPS binding energy with increasing iron ion concentration in Figure 5.1(e) is consistent also with observations on

2+ both TiO2 and the oxynitride [2] and in recent studies of Co ion doping of TiO2 and its oxynitride.[23] Again, we observe a decrease in binding energy for the lower Fe2+ seeding concentrations and a corresponding shift to higher binding energy for the higher Fe2+ concentrations. We suggest that this shift is directly related to the Ti3+ to Ti2+ conversion.

This is further suggested by our DFT calculations discussed in detail elsewhere [28] and summarized in Table 5.3. These calculations clearly indicate that the binding energies of the nominally (+2) TiO and Ti(OH)2 sites, especially TiO, can exceed those of the nominally (+4) , TiO2 based sites. This strongly suggests that a process that first converts

Ti4+ to Ti3+, producing a known decrease in the Ti binding energy [2] can readily be followed by a further oxidation-reduction in which Fe2+ reduces Ti3+ to Ti2+ with a

88 commensurate increase in binding energy. Hence, the resulting XPS spectra in Figures

5.1(b,c) and 5.2(b,c) show an expected shift to lower binding energy at low Fe2+ seeding concentrations that subsequently increases to a higher binding energy at the higher Fe2+ seeding levels.

Table 5.3 Calculated 2s and 2p Binding Energies in eV.

Oxid. Molecule 2s 2p 2p 2p State x y z 1TiO II -544.91 -455.92a - -455.29 1 Ti(OH)2 II -542.51 -453.60 -453.28 -453.06 1 TiO2 IV -541.96 -452.70 -452.83 -452.84 1 TiO2(H2O) IV -541.23 -451.93 -452.07 -452.11 1 TiO(OH)2 IV -543.16 -454.02 -453.97 -453.84 1 Ti(OH)4 IV -543.65 -454.42 -454.40 -454.40 1 Ti2O4 C2h IV -542.86 -453.55 -453.72 -453.70 1 a Ti2O4 C3v III -543.56 -454.35 - -454.39 1 a Ti2O4 C3v IV -543.19 -454.05 - -453.82 1 Ti3O6 Cs III -544.53 -455.34 -455.36 -455.28 1 Ti3O6 Cs IV -543.12 -453.99 -453.96 -453.77 a Degenerate value for 2px and 2py.

Figure 5.2(b) compares the valence band photoelectron spectra for several Fe2+ ion doping levels verses the valence band spectra for TiO2 and TiO2-xNx .We find that the effects of iron ion seeding are much more pronounced for the oxynitride system. The VB

XPS spectra for the lowest Fe2+ ion doping levels are shifted to considerably lower binding energies than the oxynitride. At higher Fe2+ doping levels, the spectra show a

89 profound change as a new peak appears in the spectrum at approximately 15.5 to 16.5 eV and the titanium peak is broadened and shifted to higher binding energy consistent with the conversion from Ti3+ to Ti2+. The O 2s feature is considerably broadened, and the feature associated with a mixed N 2p – O 2p valency is skewed to a significantly lower binding energy than the corresponding O 2p peak in TiO2. The appearance of the feature at 15.5 to 16.5 eV is consistent with the iron catalyzed formation of NOx species [44] on the oxynitride surface. The location of this peak may indicate the formation of the higher order nitrogen oxides with x>1. In complement the N 1s XPS spectra in Figure 5.2(f) display a clear trend to increasing binding energy with increasing Fe2+ concentration, again indicative of the formation of the higher oxides of NOx. The process displayed for the VB and N 1s core level XPS spectra would appear to be analogous to that accompanying nitrogen fixation.[45] By contrast, the data in Figure 5.1(f) indicates, as expected, the absence of a nitrogen signal. Again the Ti 2p core level XPS spectra depicted in Figure 5.2(c) display a complex pattern accompanying the doping of TiO2-xNx.

The spectra first shift to lower binding energy with increasing concentration, followed by an increase in the binding energy to a peak at ~ 458.0 eV (vs. 458.6 for the oxynitride

[2]). This is accompanied by an attendant decrease in spectral intensity. A similar behavior is observed for the O 1s XPS spectrum in Figure 5.2(e). These observations suggest the much stronger interaction of the iron ions with the titanium oxynitride at the surface of the titanium oxynitride nanoparticles. The Fe 2p XPS spectrum in Figure 5.2(d)

(0.2(37.9 mole%)Table 5.1) shows a similar shift to higher binding energy (~712 eV) relative to FeCl2 at 710.2 eV.[46] This is also corroborated by XAS spectra.

The most significant comparison between the XPS and XAS data suggests that the

XPS spectra, as expected, probe the surface of the doped titania and its oxynitride whereas the XAS spectra probe somewhat further into the sample. The analysis of these

XAS spectra can be quite complicated as Fargus et al [43] have indicated and the interpretation of this data to extract binding energies tenuous. Therefore, they are not discussed in detail in this monograph. However, the XAS spectra are accompanied by significant structure extending from the oxygen peaks to higher binding energy. This structure is consistent with the presence of residual water [47] complexed with Fe2+ at the sub-surface of the iron seeded titania and its oxynitride. We suggest that this is a result of the incomplete drying of those regions that lie below the surface and are formed from the interaction of the FeCl2 tetrahydrate.

The Fe2+ doping of the titania nanocolloid and its oxynitride displays the manifestation of a process which appears to produce progressively higher oxides of nitrogen with the introduction of higher concentrations of iron ions to the oxynitride surface. This is in analogy to the process of nitrogen fixation [45].

Figure 5.3 (a) Raman spectra for FeCl2 doped TiO2 where the doping is done with FeCl2- 4 H2O (concentrations correlating with Table 5.1 are indicated in the figure (b), Raman spectra for nitrogen doped titanium oxynitride (from Ref. 1), (c) Raman spectra for FeCl2 doped TiO2 where the doping is done initially with anhydrous FeCl2 (concentrations correlating with Table 5.1 are given in the figure).

The N 1s XPS nitrogen spectrum of Figure 2(f) for iron doped TiO2 at the highest iron concentrations displays the manifestation of a considerable shift to higher binding energy with increased Fe2+ concentration. This is indicative of the formation of the nitrogen oxides [2].

Raman Spectroscopy and Phase Transformation

Using Raman spectroscopy, we have established that Co2+ ion seeding introduces a spinel-like structure into both the TiO2 and TiO2-xNx anatase nanocolloids. This leads to the ready conversion of both the oxide and oxynitride from the anatase to the rutile form

93 at room temperature, in contrast to temperatures in excess of 700°C.[48] The assessment of this conversion process is more difficult with Fe2+ ion doping because the hematite, magnetite, wuistite, and maghemite iron oxide spinels all have significant and distinct

Raman spectra [49] below 700 cm-1. Raman spectra obtained for the oxynitride doped with FeCl2-4H2O and anhydrous FeCl2 are shown in Figures 3(a) and 3(c) respectively.

These samples were first prepared by doping the TiO2 nanocolloid with iron and subsequently carrying out the nitridation process. Therefore, in view of the water base [1-

3] for both systems, it is not surprising that the two spectra in Figures 3(a,c) are very similar. The anatase oxynitride demonstrates three broad features [21,49] centered at 400,

515, and 635 cm-1 with an additional feature at ~550 cm-1[6,7,20,21] attributed to the presence of non-stoichiometric titanium oxyinitride.[7,25,48,50,51] In contrast, the Co2+ treated system [21-23] displays features at ~250, 380, 430,606, and 690 cm-1, where the

-1 features at 250, 430, and 606 cm are associated with the rutile phase [25] of TiO2-xNx and those at 380 [24] and 690 cm-1 [52] are associated with the cobalt oxide (spinel).

While the assessment of the change from anatase to rutile TiO2-xNx is more difficult for the iron-based system, it is apparent from the data presented in Figures 3(a) and (c) that a conversion from the anatase to rutile forms of the oxynitride has occurred. In Figure

5.3(a), although the spectra for the four lowest FeCl2 concentrations show mostly an anatase structure, the spectrum at the highest FeCl2 concentration (0.1- 23.3 mole %), shows two main broad peaks, one at 431 cm-1 and the other at 614 cm-1, which are exactly in the range expected for rutile TiO2 nanoparticles, as reported by T. Mazza et al.[53] These authors found that the Raman rutile Eg and A1g modes range between 430

-1 -1 -1 cm and 445 cm for Eg and the A1g mode is excited at frequencies between 608 cm

94 and 615 cm-1, depending on particle size. Additionally, one might note the total absence of the 515 cm-1 peak, indicating that no anatase phase is present. Similar results are also evident in Figure 5.3(c). Features in the 250 and 430 cm-1 region are clearly present accompanied by a significant broadening of the Ti-O stretching mode to lower frequency.

Note also that the 613 cm-1 line is also evident for all concentrations shown in Figure

5.3(c) except for 0.0063, again confirming the transformation to the rutile phase in this system. There is also a clear increase in spectral intensity with increased iron doping, followed by a significant quenching of the spectra at higher doping concentrations. A similar behavior has been observed for the cobalt, nickel, and manganese treated systems.[24]

Transmission Electron Microscopy

2+ Figures 5.4(a)-(d) depict the results of the Fe seeding of TiO2 at the lowest concentration used in these experiments, 1.9 mole% Fe2+, and at a concentration corresponding to 3.7 mole%. There is a clear change in the character of the system over this range.

At the lowest concentrations we find the formation of what appear to be TiO2 crystallites which, in conjunction with the following XRD analyses, we assign to a slightly distorted anatase phase. At the lowest iron seeding concentrations the XRD analyses suggests a partial conversion to the rutile structure of TiO2. This can be readily inferred by comparison with the XRD pattern for DeGussa P25 for a 13% rutile by composition sample of TiO2 (Figure 5.8(c)) [47]. However, upon the introduction of iron

95 or its hexahydrate, the bulk crystalline nature of the sample is largely lost as an amorphous structure dominates.

(a) TiO2-Fe(II)-0.94

(b) TiO2-Fe(II)- 3.67

Figure 5.4 TEM micrographs of FeCl2 doped TiO2. Figures (a) in broad (scale bar 0.5 µm) and finer (200 nm) view demonstrate the apparent formation of crystallites on the TiO2 surface and correspond to a doping level of 0.94 mole% (Table 5.1). Figure.(b) and (corresponds to a typical section showing a much more pronounced agglomeration of the Fe in the TiO2 (scale bar 200 nm) and correspond to a doping level of 3.67 mole% (Table 5.1).

A B

Figure 5.5 (a) HRTEM image of Fe seeded TiO2 (1.87 mole %-Table 5.1) showing d values, which are an average of 10 spacings. (b) Enlarged HRTEM image, indicating a spacing of 0.35 nm, indicative of the TiO2 anatase 101 plane. Scale bars are 5nm in (a) and –nm in (b).

Figure 5.5 depicts HRTEM micrographs obtained for the lowest (1.9 mole %)

Fe2+ doping levels as indicated in Figure 5.4(a). The HRTEM image in Figure 5.5(a) and the enlarged image in Figure 5.5(b) indicate a “d spacing” consistent with the anatase crystalline form of TiO2. The EDX data presented in Figures 6(a)-(d) indicate that the crystallites are composed of titanium and oxygen and suggest the uniformity of the crystalline composition. Figure 5.6(d), an enlargement of the EDX spectrum, demonstrates iron peaks consistent with the small seeding level of this sample.

5.6 (a)

5.6 (b)

5.6 (c)

5.6 (d)

Figure 5.6 EDX spectra for TiO2 – Fe (1.87 mole %). (a) Images 1 and 2 of two different crystallites. (b) and (c) EDX spectra for the two regions showing virtually identical titanium and oxygen concentrations. The Cu signal corresponds to the TEM grid; adventitious carbon is also present. (d) Enlarged figure of the EDX spectra showing the small iron concentration in the sample.

At the higher concentrations (e.g., 3.7 mole %), we observe the manifestation of a

2+ stronger interaction with the anatase TiO2 nanocolloid leading to an interweaving of Fe

99 particles with the nanocolloids. The titania particles are of average 10 nm size [2] and the

(darker) Fe2+ ions are intermingled among aggregates of these nanoparticles. The change in the mode of interaction of Fe2+ ions with the nanocolloids correlates well with the rate and degree of physical gelation [5] associated with the doping process both as Fe2+ is introduced to the TiO2 and in the process of oxynitride formation in these systems. The introduction of Fe2+ at significant concentrations likely leads to the conversion of the anatase/rutile crystalline structure for the surface-based doped titanias versus a more amorphous structure created in the subsurface and bulk.

(a) TiON-Fe-0.025 (b) TiON-Fe-0.2

Figure 5.7 TEM micrographs of FeCl2 doped TiON in two regions of the sample. Figure (a) demonstrates the apparent formation of small iron particles on the TiO2 surface and corresponds to a doping level of 7.07 mole% (Table 5.1). Figure (b) demonstrates a much more pronounced Fe-based particle formation on the TiO2 surface and corresponds to a doping level of 37.9 mole% (Table 5.1). See also Figures 5 and 6. Scale bars are 200 nm.

2+ TEM images for the Fe seeding of TiO2-xNx are depicted in Figures 5.7(a,b).

2+ Here the Fe concentration levels are considerably higher than those for the TiO2 samples and correspond to a doping of 7.1 mole% (0.025) and 37.9 mole % (0.2) respectively. The distribution of iron in these samples appears very uniform although a

100 few larger crystallites are apparent in the highly doped sample. This result would also appear to suggest the extremely slow gelation of these samples.

XRD Analysis

Figure 5.8(a) indicates XRD spectra for four of the samples studied. For the TiO2 samples, the data suggest that the introduction of 0.002 and 0.008 grams of anhydrous

2+ FeCl2 (1.9 and 3.7 mole % Fe ) perturbs the anatase structure to the extent that the spectrum is weak, those peaks at lower angle are slightly shifted, and the higher angle diffraction peaks are greatly reduced. The enlarged spectrum for the 1.87 % doped sample also indicates the degree of perturbation and suggests the onset of perturbed rutile features. This is consistent with the slightly distorted crystallites observed in the TEM spectrum (Figures 4 (a) and (b)). For the TiO2-xNx samples formed from FeCl2 • 4H2O,

(Fe2+, 7.07 and 37.9 % ) anatase features may still be apparent, however, the weak XRD pattern might also correspond to a highly perturbed rutile structure. The heavily iron ion seeded TiO2-xNx sample (0.2g = 37.9 %) results in a spectrum which appears to be more dominated by the perturbed rutile structure of the oxynitride whose spectrum also decays at higher scattering angles, however, the spectrum is notably more complex. The spectrum appears to demonstrate the conversion from a pure anatase to a highly perturbed anatase/rutile amorphous structure. The enlarged spectrum (Figure 5.8(b)) provides an indication of a conversion from the anatase to the rutile phase. However, the weak XRD patterns also suggest that the samples have been largely converted to an amorphous structure as shown in the comparison of the results in Figure 5.8(b) with those for

DeGussa P25 in Figure 5.8(c).

101

A TiON-FeCl2-37.9% TiON-FeCl -7.1% A 2 TiO -FeCl -3.7% A 2 2 TiO2-FeCl2-0.94% R black drop line: Standard anatase red drop line: Standard rutile R Intensity /a.u.

20 30 40 50 60 70 θ 2 / degree

A TiO2-FeCl2-0.002 B black drop line : standard anatase R red drop line: standard rutile A R A A A R R R

R Intensity/ a.u.

20 30 40 50 60 70 2θ / degree

Figure 5.8 XRD spectra for (a) several FeCl2 seeded TiO2 and TiO2-xNx samples as indicated in the inset of the figure, (b) expanded view of TiO2- FeCl2- 0.94 mole% (Table 5.I), and (c) XRD spectrum of DeGussa P25 with an 87% anatase and 13% rutile crystalline phase [54].

102

Optical Spectroscopy

2+ UV-Vis diffuse reflectance spectra for the Fe doped TiO2 and TiO2-xNx nanoparticles are depicted in Figures 9 (a) and (b). By comparison with Degussa P25, a commercially availableTiO2, the absorption edges of these doped samples extend to the visible region and the magnitude of the absorption increases with the amount of FeCl2 dopant added to the titanium oxide and oxynitride. However, the degree of absorption does not approach that of the Co2+ seeded systems in our previous work,[22] in agreement with the recent work published by Hoffmann’s group (see Figure 5.6 in the reference).[55] The extended absorption of visible light corresponds to “bandgap narrowing” of the matrix or mid-band gap states. The mechanism for this red-shifted light absorption is not unequivocally established. Some researchers emphasize that visible light absorption of TiO2-derived materials results from the formation of oxygen vacancies and the reduction of Ti4+ to Ti3+,[56] others suggest that the dopant plays the dominant role of bandgap engineering.[9,57,58] We suggest that the bandgap narrowing may be due to both effects [59,60]. The difference DRS data 61 indicate that the absorption in the

< 500 nm region arises from additional dopant bands in the bandgap while absorption in the region >500 nm may result from the oxygen vacancies and defects.

103

120 120 A B

80 80

40 P25 40 TiO2-FeCl2-0.002 TiO2-FeCl2-0.004 N-TiO2 P25 Reflectance / % Reflectance / % TiO2-FeCl2-0.032 TiO2-FeCl2-0.080 TiON-FeCl2-0.025 TiON-FeCl2-0.200 0 0 200 400 600 800 200 400 600 800 Wavelength / nm Wavelength / nm

2+ Figure 5.9 UV-Vis diffuse reflectance spectra for Fe doped (a) TiO2 and (b) TiO2-xNx nanoparticles. The figure legends indicate the amount of FeCl2 used. For comparison, UV-Vis diffuse reflectance spectra of commercial TiO2 photocatalyst, P25, and TiO2-xNx nanoparticles are also included.

Photocatalytic Behavior

Depicted in Figure 5.10 is the photocatalysis performance of the as-prepared nanoparticles. The iron-doped samples exhibit absorption in the visible-light region.

Further, their photocatalytic activity is comparable to that of the TiO2-xNx sample [2].

However, probably of most significance is the data in Figures 5.10(a) and 10(b) which demonstrate clear oscillations in the photocatalytic decomposition of methylene blue.

That is, the photocatalytic activity of the iron–doped TiO2 does not decrease monotonically with increased iron concentration but, in fact, shows a decrease followed subsequently by an increase in photocatalytic activity. This reversal would suggest that there is more than one maximum in the catalytic activity as a function of concentration.

This observation leads us to examine the potential role of Fe3+ in the photocatalysis. Here,

104 we refer to the bulk dominance of Fe3+ vs. Fe2+ as a promoter of the anatase to rutile phase transformation [62].

102 100 A 100 B 98 98 96 96 94 0.002

92 0.004 94 0.008 90 0.016 TiO FeCl 0.0125 0.032 92 2 2 TiO FeCl 0.025 88 0.080 2 2 Percentage of MB leftPercentage of MB / %

Percentage left of MB / % 90 TiO2 FeCl2 0.05 86 TiO2 FeCl2 0.1 84 88 0 1 2 3 4 5 6 0 1 2 3 4 5 6 Time / hour Time / hour

102 100 C TiON-FeCl2-0.025 98 TiON-FeCl2-0.2 96 94

92 90 88 Percentage of MB leftPercentage of MB / % 86 0 1 2 3 4 5 6 Time / hour

Figure 5.10 (a) Photocatalytic decomposition of methylene blue for iron chloride doped TiO2 (seeded by anhydrous FeCl2). The legend shows the concentrations corresponding with those in the first column of Table 5.1 and explained in the caption of Table 5.1. (b) Photocatalytic decomposition of methylene blue for iron chloride doped TiO2 (seeded by FeCl2 (4 H2O). The legend shows the concentrations correlating with column 1 of Table 5.1. (c) Photocatalytic decomposition of methylene blue by TiO2-xNx-FeCl2.

105

Fe 3p Fe2+ Ti 3s Fe3+

Intensity / a.u.

65 60 55 50 Binding Energy / eV

Figure 5.11 VB-XPS spectrum of the products of the FeCl2 (0.2g = 37.9 mole % Fe)- 3+ TiO2-xNx interaction showing the formation of Fe whose XPS spectrum peaks at 56 eV, Ti 3s, and remaining Fe2+. The indicated intensity for Fe3+ represents a lower bound.

Figure 5.11 in fact indicates the formation of Fe3+ in these systems through the

VB-XPS spectrum at 56 eV [63]. It has been suggested [62] that mixed anatase/rutile phases may lead to improved catalytic activity due to the enhancement of electron transfer. We suggest that these mechanisms may be operative in the present study.

5.4 Discussion

From all of the data presented, it should be apparent that Fe2+ ions doped into

TiO2 or TiO2-xNx nanoparticles undergo a substantial interaction, producing significant transformations at the nanoscale. Among the most intriguing of these interactions is that monitored primarily by XPS on the oxynitride surface. This is exemplified by the VB

XPS data in Figure 5.2(b) and the N 1s core level data in Figure 5.2(f). These XPS spectra are consistent with the formation of NOx sites with increasing oxygen content whose formation correlates with an increased iron ion doping concentration. The data obtained for the Fe2+ seeded oxynitride correlates with data previously obtained for

106

TiO2

CoCl2

TiO2-CoCl2-0.5

VB-TiON FeCl 0.2 c/s 2

VB-TiON FeCl2 0.05

VB-TiON FeCl2 0.025 TiON

30 20 10 0 Binding Energy (eV)

Figure 5.12 Correlation of XPS data obtained for Fe2+ seeded oxynitride with data obtained for the Co2+ seeded oxynitride (Ref. 22). Note especially the peaks at 17-17.5 eV for the cobalt system vs. the peaks at 15.5 to 16.5 eV for the iron system and the broadening of the iron spectrum as discussed in the text.

2+ 2+ seeding with Co ion and CoCl2 depicted in Figure 5.12. At the highest Co

2+ concentration (55.7 mole% Co (ex: Table1))[22] introduced to TiO2 a peak at ~17-17.5 eV is observed, consistent with a partially substituted cobalt oxide. At the highest Co2+ concentrations (38.6 mole% Co2+) introduced to the oxynitride [22] there is a considerable increase in spectral broadening extending to the regions 20-30 eV and 15-12 eV, accompanied by the considerable lowering of the O 2p binding energy (to ~ 4 eV) to an energy well below that of the oxynitride. It is apparent that this decrease in the O 2p binding energy is also characteristic of the Fe2+ doped sample depicted in the VB XPS spectrum of Figure 5.2(b). However, the behavior of the Fe2+ and Co2+ doped systems is quite distinct in the 20-12 eV region. For the oxynitride, this is especially true in the 15 to

12 eV region where NOx sites are monitored [43]. Figure 5.12 clearly demonstrates the distinction between the 15.5-16.5 eV broad peak observed for the Fe2+ (0.2g-37.9

107 molel%- Table 5.1) doped oxynitride system and the broad, virtually structureless, feature observed in the Co2+ doped system. Further, Figure 5.12 demonstrates that the 17-17.5 eV

2+ feature observed most clearly in the Co doped TiO2 nanocolloid and for CoCl2[22] is

2+ distinct from the broad peak observed upon seeding with Fe . The FeCl2 spectrum in

Figure 5.2(b) clearly indicates that the broad feature in the range 12.5-16.5 eV is correlated at least in part with FeCl2, however, the feature is broader and correlates closely with the shift to higher binding energy of the N 1s core level XPS spectrum depicted in Figure 5.2(f). The observations suggest that these complementary spectral features are also, in part, associated with the formation [46] of nitrogen oxides with increasing oxygen content, paralleling the increase of Fe2+ ion doping concentrations.

The process whereby the 12.5-16.5 eV peak forms and the shift of the N 1s core level spectra in Figure 5.2(f) bears a distinct resemblance to the process of nitrogen fixation.[46] The nature of this process and the activity of the Fe2+ ion at the nanoscale may also be demonstrated in the doping of TiO2 as the VB XPS spectrum of Figure 5.1(b) demonstrates a clear peak near 12 eV in the region for the known VB XPS spectrum of

NO.[44,64-68] Further, the N 1s core level XPS spectrum, at the highest Fe2+ concentrations, displays a feature which might be associated with nitrogen formation on the TiO2 surface. Alternatively, the feature in the 12.5-16.5 eV region of the VB XPS spectrum can be consistent with the appearance of spectra associated with the formation of significant N 2s sites within the TiO2-xNx framework.[69]

The data in Figure 5.1 (b,c) and 2(b,c) in conjunction with our DFT calculations suggest that the double transfer process Ti4+ + Fe2+ to Ti3+ + Fe3+ followed by Fe2+ + Ti3+ to Fe2+ + Ti2+ has occurred. These observations suggest a highly facile charge transfer in

108 these highly Fe2+ doped systems counter to that previously observed. This may strongly influence the photocatalytic properties of the nanoparticles.

The TiO2 and TiO2-xNx nanocolloids used in this study have been characterized previously using X-ray powder diffraction techniques (XRD) which, based on the Debye-

Scherrer equation,[2,6,7] suggest that they have average diameters of 10 nm.[2]

Broadened but regular XRD patterns are readily obtained for the anatase structured nanocolloids [2]. However, we have recently demonstrated that it is difficult to obtain a room temperature Raman spectrum for these small nanocolloids without the sophisticated collection optics used in the present study [2,6,7,21]. The use of Raman spectroscopy in this study is crucial to our final conclusions. It would be difficult to assess the transformation from anatase to rutile TiO2 and TiO2-xNx without the clear indication of the ~250 cm-1 band observed in Figures 3(a) and (c) the previous work of Massa et al.[53] and the comparison with the results obtained for the Co2+ doped system.[21,22] The large number of potential iron oxide spinel features below 700 cm-1 can lead to the significant spectral broadening observed in Figures 3(a) and (c) vs. the far simpler Co2+ doped system [22].

The data presented in Figures 1 and 2 suggest that while the XPS studies probe

2+ the surface of the Fe doped TiO2 and TiO2-xNx nanoparticles. Corroborating XAS spectral studies probe somewhat deeper into the doped oxide and oxynitride even though the XAS spectra are generated by electron vs. fluorescence analysis. This observation is suggested by the N 1s core level nitridation data, clearly apparent in Figure 5.2(f). We have been unable to generate this data using an XAS probe, suggesting that the nitridation occurs at the surface of the TiO2 and TiO2-xNx nanoparticles. Further, the O 1s

109 core level XPS and XAS spectra are distinct with the latter indicating the presence of water that is absent for the surface of the vacuum dried nanoparticles.

5.5 Conclusion

2+ Thus, we conclude that the Fe doping of porous sol-gel generated TiO2 and

TiO2-xNx at the nanoscale profoundly affects the electronic structure and morphology of these oxides. We have observed (1) a ready and unexpected oxidation state conversion on the oxide and its oxynitride surface, (2) conversion of the anatase phase to that of rutile at least at the surface of TiO2 nanocolloids at room temperature in contrast to the generally observed conversion at 550°C, (3) oscillating and efficient photocatalytic activity at notably higher iron concentrations than those observed by previous researchers, and (4) what appears to be an analog of nitrogen fixation.

110

5.6 References

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46. C.D. Wagner, W. H. Riggs, L. E. Davis, J. F. Moulder, and G. E. Mullenberg, Handbook of Photoelectron Spectroscopy, Perkin-Elmer Corp., Eden Prairie, MN, USA. 47. See for example. J.E. Huhey, E.A. Kiefer, R.L. Kiefer, Inorganic Chemistry: Principles and Applications, 4th Edition, Harper Collins College Publishers. 48. (a) E. Gyorgy, A.P. del Pino, P. Serra, J.L. Morenza, Appl. Surf. Sci. 2002, 186, 130-134; (b) A. Ogden, A. Fedorov, J-il Hong, and J.L. Gole, Phys. and Chem. of Solids, 2008, 69, 2898 and references therein. 49. See for example. (a) D.L.A. deFaria, S.V. Silva, M.T. Olivera, J. Raman Spectro. 1998, 28, 873-878; (b) I. Chamritski, G. Burns, J. Phys. Chem. 2005, 109(B), 4965-4968. 50. A. Brevet, F. Fabreguette, L. Imhoff, M.C.M. deLucas, O. Heintz, L. Saviot, M. Sacilotti, S. Bourgeois, Surf. & Coatings Tech. 2002, 151, 36-41. 51. M.E. Straumanis, T. Ejima, W.J. James. ActaCrystallographica 1961, 14, 493- 497. (a) M.J. Escuder, T. Rodrigo, L. Mendoza, M. Cassir, L. Daza, J. Power Sources 2005, 140, 81- 87. (b) J. Choi, H. Park, M.F. Hoffmann, J. Phys. Chem. C 2010, 114, 783-792. 52. J. Choi, H. Park and M. R. Hoffmann, J. Phys. Chem. C, 2010, 114, 783-792. 53. T. Mazza, E. Barborini, P. Piseri, P. Milani, D. Cattaneo, A. Li Bassi and C.E. Bottani, Phys. Rev. B, 2007, 75, 045416 54. Mark White private communication 55. V. N. Kuznetsov and N. Serpone, J. Phys. Chem. C, 2009, 113, 15110-15123. 56. D. N. Tafen, J. Wang, N. Q. Wu and J. P. Lewis, Appl. Phys. Lett., 2009, 94, 093101. 57. S. Sakthivel, M. Janczarek and H. Kisch, J. Phys. Chem. B, 2004, 108, 19384- 19387. 58. P. Romero- Gomez, V. Rico, A. Borras, A. Barranco, J. P. Espinos, J. Cotrino and A. R. Gonzalez-Elipe, J. Phys. Chem. C, 2009, 113, 13341-13351. 59. J. Wang, D. N. Tafen, J. P. Lewis, Z. L. Hong, A. Manivannan, M. J. Zhi, M. Li and N. Q. Wu, J. Amer. Chem. Soc., 2009, 131, 12290-12297. 60. Q. Xiaofeng, Y. Zhao, C. Burda, Adv. Mater. (Weinheim, Germany) 2007, 19 (22), 3995-3999. 61. J.A. Rodriguez, T. Jirsak, J. Dvorak, S. Sambasivan, D. Fischer, J. Phys. Chem. B 2000, 104, 319-328. 62. D.A.H. Hanaor and C.C. Sorrell, J.Mater. Sci., 46, 855-74 (2011) 63. P. Hays, Applied Surface Science, 254, 2441-49 (2008) 64. E. Gyorgy, A.P. del Pino, P. Serra, J.L., Morenza,rface and Coatings Technology 2003, 173,265. 65. T. Jirsak, J. T. Dvorak, J.A. Rodriguez, Surface Science 1999, 436, L683-L690 66. C.D. Wagner, L.E. Davis, J.R. Moulder, G.E. Muilenberg Handbook of X-ray. 67. Photoelectron Spectroscopy, published by Perkin-Elmer. E.P., MN, USA 1979. 68. L.A. Delouise, N. Winograd, Surface Science 1985, 159, 199-213 69. The N 2s features are expected to lie at lower binding energy than do those features associated with O 2s.

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Chapter 6 Synthesis and Characterization of Nano-scaled Gallium, Nitrogen Co-

doped Titanium Dioxide with High Nitrogen Concentration

6.1 Introduction

The utilization of solar energy to solve energy deficiency and environmental pollution has attracted the attention of scientists in fields such as material chemistry,[1, 2] material physics, [3, 4] physical chemistry[5] and material engineering[6, 7] for years.

Although there are many ways to harvest solar energy, using semiconductors is considered one of the best solutions.[8] Among the semiconductors, titanium dioxide

(TiO2) has been the focus of researchers due to the fact that it is active, cost effective, highly abundant in nature, and stable under light illumination.[9, 10] Nevertheless, there is one well-known key disadvantage of TiO2 preventing the full usage of the solar energy, which leads to absence of activity in the visible-light region due to its large bandgap.[11]

Without modification, TiO2 can only harvest the UV-light part of the solar spectrum, which constitutes only about 5% of total solar energy, leaving the majority of solar energy unused. To solve this issue, scientists have doped new elements into the TiO2 lattice in the goal of reducing its bandgap so that the visible-light that constitutes about

50% of the solar spectrum can be utilized.[7, 9, 11, 12] Of all the elements that have been attempted to be doped into TiO2, nitrogen appears to be one the most effective dopants and has aroused extensive studies among researchers.[7, 9, 13-15] Now, the bandgap of the TiO2 has been successfully reduced to absorb visible-light, which is a significant step towards utilizing the visible-light region of the solar energy. However, the concentration of the doping nitrogen is usually rather low and tends to get even lower when sintering the photocatalyst to make it more crystalline at higher temperature. Additionally,

114 sintering of the photocatalyst at high temperatures will cause another problem, which is particle agglomeration that decreases the surface to volume ratio hence impacting the overall efficiency of the photocatalyst. In this study, we attempt to introduce a codopant of Gallium to solve the problems mentioned above. Particles of Ga-TiO2-xNx are found to remain in the nanoscale without aggregation even after sintering at 700 oC for 2 hours while the concentration of nitrogen is found to be much higher compared to previous reports.

6.2 Experimental procedures

All starting materials were used without further purification. The Ga, N co-doped

TiO2 was synthesized as follows: with magnetic stirring, 4.86 mL of Titanium (IV) isopropoxide (Aldrich, 97%) was added dropwise to 100 mL of deionized water to form a white suspension. 6.85 mL concentrated HNO3 was then added to the stirring suspension slowly. After the suspension became clear, a certain amount of Ga(NO3)3 (crystalline,

99.9% trace metals basis ordered from Aldrich) was added to the solution so that the molar ratio between Ga and Ti is 1, 2, 3, 4 specifically. 100 mL 10% (weight percentage) aqueous ammonia solution was added to form a precipitate. This precipitate was then washed with deionized water three times and dried in the open air before it was sintered

o with Ar / NH3 flow at 400, 500, 600 and 700 C for 2 hours. To clarify, the NH3 was brought by Ar from a flask of 1.5 L aqueous ammonia solution. The equivalent flow rate of NH3 was calculated to be ~300 mL per minute based on the result of effluent titration.

X-ray Photoelectron Spectroscopy spectra were obtained using a PHI-

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Western Reserve University). UV-vis Diffuse Reflectance Spectra were acquired on a

Varian Cary Bio50 UV–vis spectrometer with a Barrelino remote diffuse reflection probe using MgO as reference. X-ray Diffraction patterns of the prepared samples were obtained on a Scintag X-1 Advanced X-ray powder diffractometer (XRD, 4º / min, Cu

Kα radiation). TEM analysis was performed using a transmission electron microscope

(TEM, JEOL 1200CX) with an accelerating voltage of 80 kV.

The photocatalytic activity of the prepared samples was evaluated by measuring the decomposition of methylene blue (MB) under visible-light irradiation. The light source used was a 150-W high-pressure Xenon arc lamp with a 400 nm long pass filter.

The concentration change of methylene blue was monitored by measuring the absorbance of the MB solution with a Varian Cary Bio50 UV–vis spectrometer at 664 nm. For a typical photodecomposition experiment, 1.0 mg photocatalyst is mixed with 2.5 mL methylene blue solution in a standard quartz cuvette (10.0 mm path length, ~3.5 mL) together with a micro magnetic stir bar (Fisherbrand, 7 mm length, 2.4 mm diameter).

The cuvette was then kept in the dark with continuous stirring until there is no change in absorbance of the solution ( to make sure that physical adsorption on the catalyst will not play a role in reducing the opticle absorption). The starting concentration of the methylene blue was set as 20 ppm so that after physical adsorption in the dark, the absorbance of the methylene blue solution would be within the range of 1.0 ~ 2.0.

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6.3 Results

100 100

TiO2

TiO2 o 400 C 400 oC o

50 o 500 C 50 600 C 500 oC o 700 C o 600 C o Reflectance / %

Reflectance / % 700 C

Ti : Ga = 1 : 1 Ti :Ga = 2 : 1

0 0 300 400 500 600 700 800 300 400 500 600 700 Wavelength / nm Wavelength / nm

100 100

TiO 2 TiO2

700 oC

o o

50 o 600 C 50 o 400 C 600 C 400 C o o 500 C 700 oC 500 C Reflectance / % Reflectance / a.u.

Ti : Ga = 3 : 1 Ti : Ga =4 : 1

0 0 300 400 500 600 700 300 400 500 600 700 Wavelength / nm Wavelength / nm

Figure 6.1 UV-vis diffuse reflectance spectra of Ga-TiO2-xNx NP samples.

Figure 6.1 shows the UV-vis diffuse reflectance spectra of the prepared powder samples. Compared to pure TiO2, there is clearly absorption in the visible-light region, especially for the 600 and 700 oC sintered samples. The absorption edge goes up to 550 nm (2.3 eV). Ga2O3 has a huge bandgap of more than 4.0 eV[16, 17] which corresponds to wavelengths of less than 310 nm. Therefore, the visible-light absorption cannot be caused due to Ga2O3 or TiO2 alone. Nitrogen doped titania NPs exhibit visible-light absorption for samples sintered at below 500 oC. Furthermore, the visible-light absorption

117 usually decreases as the sintering temperature increases.[12, 14] However, in the current project, the visible-light absorption for samples sintered at 600 and 700 oC are stronger than the samples sintered at 400 and 500 oC (except for the case of Ti : Ga = 4 : 1 sample, in which the absorptions in visible-light region are equally intense). Moreover, the 700 oC sintered samples show no decrease at all in the visible light specturm compared to

o samples sintered at 600 C. The significant difference between the current Ga-TiO2-xNx and normal nitrogen-doped TiO2 indicates that the existence of Ga raised the barrier for both nitrogen incorporation as well as nitrogen loss out of the titania matrix. It is noteworthy that with increasing content of Ga, the barrier becomes more and more evident. The detailed structure and composition that led to this difference will be discussed in combination with additional characterization techniques later.

Figure 6.2 shows the X-ray diffraction patterns of the Ga-TiO2-xNx samples with various Ti : Ga ratios sintered at 600 and 700 oC. The samples for 400 and 500 oC are all amorphous with the exception of the Ti : Ga = 4 : 1 and hence are not shown here. The

Ti : Ga = 4 : 1 samples show clear crystal structures of anatase TiO2, although there is a wide bulge centered at around 2θ = 30ο which was hard to assign since it is neither any phase of TiO2 nor any of the major Ga2O3 or GaN crystal phases. A broad diffraction band is usually a reflection of an amorphous material. However, amorphous TiO2 usually shows similar broad appearance in XRD patterns but centered at 2θ < 30o [12, 18-20].

This might indicate the existence of a new structure within the Ga-TiO2-xNx system. By comparing the XRD patterns of samples with various Ti : Ga ratios, it can be noted that as the percentage of Ga in the sample increases, the difficulty in which anatase TiO2 forms also increases. Normally, anatase TiO2 can be obtained by sintering the TiO2

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Ti : Ga = 1 : 1 Ti : Ga = 2 : 1 600 OC

600 OC A A

700 OC Intensity / a.u. O Intensity / a.u. A 700 C A A A A A

20 30 40 50 60 70 20 30 40 50 60 70 2 θ / degree 2 θ / degree

A Ti : Ga = 4 : 1 Ti : Ga = 3 : 1 A 400 oC A O A 600 C A o A A 500 C A A

A 600 oC Intensity / a.u. Intensity / a.u. 700 OC 700 oC

20 30 40 50 60 70 20 30 40 50 60 2 θ / degree 2θ / degree

Figure 6.2 XRD patterns for Ga-TiO2-xNx samples with various Ti : Ga ratios sintered at 600 and 700 oC. The patterns are not smoothed to keep the original appearance. The letter “A” in the figure stands for “anatase TiO2” peaks.

o precursor at 400 to 500 C or lower.[12, 14, 21] And rutile phase TiO2 usually begins to form at 600 oC.[21] Together with the conclusion from UV-vis diffuse reflectance spectra, we can see that Gallium not only elevates the threshold of nitrogen binding and desorbing off the system but also increases the temperature for anatase TiO2 formation. Accordingly, the formation of rutile TiO2 is also deferred to higher temperature (not even seen in the current work). It must be noted that nitrogen also has the effect of keeping anatase TiO2

119 from transforming to rutile TiO2 as well.[12, 15] The size of the particles can be estimated using the Scherrer equation,

Kλ D = (1) βθcos where D is the mean size of the crystallites, K is the shape factor whose value is 0.94 for dimensionless shapes, λ is the X-ray wavelength (0.1548 nm for Cu Kα radiation), β is the full width at half maximum (FWHM) in radians, and θ is the Bragg angle of the corresponding peak. Due to the larger signal to noise ratios for the Ti : Ga = 1 : 1 and 2 :

1 samples, we only estimated the Ti : Ga = 3 : 1 and 4 : 1 samples and the results are shown in Table 6.1. The sizes of the samples, based on this estimation, are around 10 nm.

For both the Ti : Ga = 3 : 1 and 4 : 1 samples, there are slight but noticeable increases in the size when the sintering temperature increases from 600 oC to 700 oC. Within the same sintering temperature, the 4 : 1 samples are bigger than the 3 : 1 samples, although the margin is miniscule.

Table 6.1 Particle size calculated based on Scherrer equation.

Sample FWHM / radian 2θ / degree D / nm

Ti : Ga = 3 : 1, 600 oC 0.0200 25.3 7.4

Ti : Ga = 3 : 1, 700 oC 0.0185 25.3 8.1

Ti : Ga = 4: 1, 600 oC 0.0174 25.3 8.5

Ti : Ga = 4 : 1, 700 oC 0.0114 25.3 10.7

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Ti : Ga = 4 : 1 700oC o

Ti : Ga = 3 : 1 700 C O1s A o Ti : Ga = 2 : 1 600 C Ti2p Ga LMM Ga Ti KVV Ti N1s Ga3d Ga3p Ga3s

C1s Ti3p

1000 800 600 400 200 0 Binding Energy / eV

Ti : Ga = 4 : 1, 700 oC B Ti : Ga = 3 : 1, 700 oC Ti : Ga = 2 : 1, 600 oC experiment 2, 800 oC

408 404 400 396 392 Binding energy / eV

Figure 6.3 (A)Full-scale (B) N1s XPS spectra of selected samples.

To obtain information of the composition of the samples, XPS spectra were measured for selected samples and results are shown in Figure 6.3. As we can see in

Figure 6.3a, the major composing elements, i.e., Ti, O, Ga, N, are all present in the full- scale XPS spectra. There is significant amounts of nitrogen incorporated in the sample.

To have a more detailed understanding of how much nitrogen is incorporated, we performed high resolution XPS and the percentage of nitrogen is listed in Table 6.2. As

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Palgrave et al. pointed out recently[22], it is usually a problem that the concentration of nitrogen in the TiO2-xNx system is generally low. We developed N-TiO2 with high nitrogen concentrations of up to several percent. [13, 23] Now with the help of Gallium, the nitrogen content is even higher as indicated in Table 6.2. Notably, the binding energy of N1s decreases with increasing concentration of Gallium, which implies that the N1s is

“being reduced” as more gallium is incorporated.

Table 6.2 Percentage and Binding Energies of N1s in the samples based on XPS analysis.

Sample N1s Binding Energy / eV N1s percentage

Ti : Ga = 4 : 1, 700 oC 397.4 24%

Ti : Ga = 3 : 1, 700 oC 396.8 15%

Ti : Ga = 2 : 1, 600 oC 396.7 19%

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Ti : Ga = 4 : 1

700 oC

600 oC

500 oC

400 oC

0 1000 2000 3000

Raman shift / cm-1

Figure 6.4 Raman Spectra of the Ti : Ga = 4 : 1 sample sintered at 400, 500, 600, and 700 oC.

To further understand the structure of the samples, Raman spectra of the Ti : Ga =

4 : 1 samples were measured and shown in Figure 6.4. The peaks that appeared at 142,

-1 397, 518 and 640 cm correspond to anatase TiO2 fundamental phonon modes of Eg,

B1g, A1g+B1g, and Eg respectively.[24] There should be another peak at ~197 cm-1 for the Eg mode as well, which can be barely seen even for the 700 oC sintered sample. The reason is two-fold: on one hand, this peak is intrinsically low in intensity; on the other hand, the crystallinity of the samples is poor. The Raman spectra lead to the conclusion that the host crystal structures of the 600 and 700 oC sintered samples are mainly anatase

TiO2. No other peaks were observed, which excludes the possibility of the existence of crystalized Ga2O3 or GaN. These findings are in accordance with the XRD results.

123

Nevertheless, there is still the possibility that other forms of metal oxides or nitrides exist in these amorphous phases.

Figure 6.5 TEM images of Ti : Ga = 4 : 1 samples sintered at (A) 400 oC, (B)500 oC (C) 600oC, and (D) 700 oC.

Shown in Figure 6.5 are TEM images for selected samples. It is evident that the shapes of the particles are not regular but close to spheres. Most of the samples show particles separating well from each other, with the exception of the 400 oC sintered

124 sample. For sol-gel method synthesized particles, the agglomeration phenomenon is not uncommon as there is no surface ligand in between to keep the particles from aggregating.

Agglomeration is undesirable due to the fact that the surface to volume ratio might be affected. The 400 oC sintered sample shows the most severe aggregation, which may be related to its existence as the only sample among the four shown in Figure 6.5 that is in amorphous phase according to the XRD results. The average particle sizes are measured

(15 particles are randomly picked to measure using free version of ImageJ) to be 10 nm

(with a standard deviation of 2 nm) in average for the 500 oC sintered sample and 12 nm

(with a standard deviation of 4 nm) for the 700 oC sample. These size measurements correspond well with the calculated estimate based on the Scherrer equation (1).

105

90 P25 no catalyst 75

60 Ti:Ga = 4:1, 700oC Ti:Ga = 1:1, 700oC o 45 Ti:Ga = 3:1, 700 C Ti:Ga = 2:1, 700oC Ti:Ga = 2:1, 600oC 30 o Ti:Ga = 3:1, 600 C Percentage left of MB (%) Ti:Ga = 1:1, 600oC 15 Ti:Ga = 4:1, 600oC 0 1 2 3 4 Time / h

Figure 6.6 Photocatalytic decomposition of methylene blue under visible-light illumination. The legend is ranked according to the photoactivity from worst to best. For comparison, the decomposition without catalyst and the performance of commercial TiO2, P25 are also shown.

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The photocatalytic performances of the 600 and 700 oC samples are evaluated and results are shown in Figure 6.6. In general, the catalytic performances of these samples as photocatalysts to decompose methylene blue are not too different. There is also no linear trend with the percentage of gallium. However, it is clear that the 600 oC sample outperformed the 700 oC samples slightly, suggesting that crystallinity did not help the catalytic performance of Ga-TiO2-xNx. Considering that there is no drastic change in the

UV-vis absorption properties for the 600 and 700 oC sintered samples (Figure 6.1), a possible explanations as to why the 600 oC samples outperformed the 700 oC sample may lie in the slightly increased particle size when the sintering temperature increased.

Overall, the photocatalytic performance of these samples are better than the hydrothermally synthesized samples in previous work [25].

6.4 Conclusion

In summation, we used gallium as a co-dopant for nitrogen doped TiO2 to help stabilize the existence of nitrogen as well as to increase the overall threshold of particle aggregation. Preliminary results revealed that compared to usual nitrogen doped TiO2, the samples in this work exhibited much more nitrogen concentration and this concentration do not decrease as temperature goes up as high as 700 oC. Furthermore, unlike normal cases of TiO2 NPs in which particles agglomerate upon sintering, there is no heavy aggregation observed in the TEM images. The introduction of the codopant gallium shows the effect of increasing the barrier of (i) anatase TiO2 formation and (ii) anatase

TiO2 transferring to rutile TiO2. The photocatalytic decomposition of methylene blue under visible-light is found to be faster compared to recently made N-doped TiO2 samples.[25]

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6.5 References

1. Yan, X., C. Zou, X. Gao, and W. Gao, J. Mater. Chem., 2012. 22(12). 2. Fan, K., J. Chen, F. Yang, and T. Peng, J. Mater. Chem., 2012. 22(11), 4681-4686. 3. Gambhire, A.B., M.K. Lande, B.R. Arbad, S.B. Rathod, R.S. Gholap, and K.R. Patil, Mater. Chem. Phys., 2011. 125(3), 807-812. 4. Liu, C.-J., T.-Y. Yang, C.-H. Wang, C.-C. Chien, S.-T. Chen, C.-L. Wang, W.-H. Leng, Y. Hwu, H.-M. Lin, Y.-C. Lee, C.-L. Cheng, J.H. Je, and G. Margaritondo, Mater. Chem. Phys., 2009. 117(1), 74-79. 5. Ma, X., Y. Wu, Y. Lu, J. Xu, Y. Wang, and Y. Zhu, J. Phys. Chem. C, 2011. 115(34), 16963-16969. 6. Liu, X., C. Liang, H. Wang, X. Yang, L. Lu, and X. Wang, Materials Science and Engineering: A, 2002. 326(2), 235-239. 7. Li, D., H. Haneda, S. Hishita, and N. Ohashi, Materials Science and Engineering: B, 2005. 117(1), 67-75. 8. Chen, X.B., S.H. Shen, L.J. Guo, and S.S. Mao, Chem. Rev., 2010. 110(11), 6503- 6570. 9. Chen, X. and S.S. Mao, Chem. Rev., 2007. 107(7), 2891-2959. 10. Fujishima, A. and K. Honda, Nature, 1972. 238(5358), 37-+. 11. Asahi, R., T. Morikawa, T. Ohwaki, K. Aoki, and Y. Taga, Science, 2001. 293(5528), 269-271. 12. Wang, J., W. Zhu, Y. Zhang, and S. Liu, J. Phys. Chem. C, 2006. 111(2), 1010- 1014. 13. Burda, C., Y. Lou, X. Chen, A.C.S. Samia, J. Stout, and J.L. Gole, Nano Lett., 2003. 3(8), 1049-1051. 14. Zhao, Y., X. Qiu, and C. Burda, Chem. Mater., 2008. 20(8), 2629-2636. 15. Li, Q. and J.K. Shang, J. Am. Ceram. Soc., 2008. 91(10), 3167-3172. 16. Mohammadi, M., M. Ghorbani, M. Cordero-Cabrera, and D. Fray, J. Mater. Sci., 2007. 42(13), 4976-4986. 17. Ueda, N., H. Hosono, R. Waseda, and H. Kawazoe, Appl. Phys. Lett., 1997. 70(26), 3561-3563. 18. Nakamura, M., T. Aoki, and Y. Hatanaka, Vacuum, 2000. 59(2–3), 506-513. 19. Oskam, G. and F.D.P. Poot, Journal of Sol-Gel Science and Technology, 2006. 37(3), 157-160. 20. Li, Y., G. Zhao, X. Zhou, L. Pan, and Y. Ren, Journal of Sol-Gel Science and Technology, 2010. 56(1), 61-65. 21. Sadeghzadeh Attar, A., M. Sasani Ghamsari, F. Hajiesmaeilbaigi, S. Mirdamadi, K. Katagiri, and K. Koumoto, J. Mater. Sci., 2008. 43(17), 5924-5929. 22. Palgrave, R.G., D.J. Payne, and R.G. Egdell, J. Mater. Chem., 2009. 19(44). 23. Qiu, X.F. and C. Burda, Chem. Phys., 2007. 339(1-3), 1-10.

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24. Zhang, W.F., Y.L. He, M.S. Zhang, Z. Yin, and Q. Chen, Journal of Physics D- Applied Physics, 2000. 33(8), 912-916. 25. Wang, J., B. Mao, J.L. Gole, and C. Burda, Nanoscale, 2010. 2(10).

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Chapter 7 Interactive Iron-Silicon Oxidation/Reduction Processes on Fumed Silica

This work is submitted for publication

7.1 Introduction

We recently reported on the surprising oxidation state of fumed silica and the nature of water binding to the silicon oxides and hydroxides.[1] X-ray photoelectron spectroscopy data was used to suggest that the average oxidation state of silicon, at least at the surface of fumed silica, was +1 in contrast to the assumed value of +4. This data

® suggested a less hydrophilic character for CAB-O-SIL (commercial fumed SiO2) than the oxides of silicon in an average formal oxidation state of +3 or +4. This result was further supported by FTIR analysis and elemental analysis. Molecular electronic structure calculations1, as they describe the binding of water to the silicon oxides and hydroxides, demonstrate the likelihood of a slow initial hydration of the +1 and +2 oxidation states of silicon followed by the more rapid formation of the +3 oxidation state. The nature of the

+1 oxidation state and the nature of water binding to the +1 and +4 oxidation states can influence the observed isotherms for water uptake and can have important implications for oxidation/reduction processes and the expanded use of such silica-based materials[2,3] in catalytic applications. There may be a role for such oxidation state changes in microelectronics where gradients from Si0 to Si4+play a role and in silicon-based laser systems.[4] The existence of a +1 oxidation state[1,5] on the surface of fumed silica therefore has potential important implications.

As opposed to the relative stabilities of surfaces with different mixtures of Si1+ to

Si4+ we have obtained evidence for distinct Si1+ sites. The existence of the Si1+ sites implies the potential for an active oxidation / reduction chemistry at the surface of silica.

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In this study, we detail experiments to demonstrate that these lower oxidation states of silica undergo charge transfer with supported metal ions in an oxidation/reduction coupled reaction. We combine X-ray and Valence band photoelectron (XPS, UPS) spectroscopy, diffuse reflectance (DRS) and FTIR spectroscopy, Color Analysis, X-ray diffraction (XRD), and Transmission Electron microscopy (TEM) to provide convincing evidence to demonstrate that Fe3+ metal ions in contact with the Si1+ and Si2+ ions are partially reduced at room temperature to form Fe2+ ion while the Si ions are oxidized to higher oxidation states. Further, we obtain similar results for the Cu2+/Cu1+ couple.

7.2 Experimental procedures

The CAB-O-SIL® powder used in these experiments was obtained from Aldrich and has a surface area of 390 ± 40 m2/g. The fumed silica in methanol was exposed in a vacuum hood to Fe3+ and Cu2+ as the chlorides. The subsequent samples were then vacuum dried at ~ 10-2 Torr for a period close to two hours. These are the samples that have been characterized.

XPS data for the Fe3+-(CAB-O-SIL) and Cu2+-(CAB-O-SIL) were obtained using a PHI-VERSAPROBE XPS scanning microprobe, as a focused, highly monochromatic

Western Reserve University) and XPSpeak41 (free version downloaded from http://public.wsu.edu/~scudiero/). DRS spectra were determined by a Varian Cary Bio50

UV–vis spectrometer with a Barrelino remote diffuse reflection probe using MgO as reference (CWRU). FTIR spectra were obtained using a Thermo Nexus 870 FTIR spectrometer with an attenuated total reflection (ATR) accessory (CWRU). XRD analysis

130 of the prepared samples was performed on a Scintag X-1 Advanced X-ray powder diffractometer (XRD, 2º / min, Cu Kα radiation, CWRU). TEM analysis was performed using a transmission electron microscope (TEM, JEOL 1200CX) with an accelerating voltage of 80 kV (CWRU).

7.3 Results

Interaction of Fe3+ with fumed silica-X-ray photoelectron data

Figure 7.1 From right to left, anhydrous FeCl3 (black), 1 : 1, 1 : 3, and 1 : 4 by weight FeCl3 on CAB-O-SIL. The samples change in color from light orange to bright yellow.

Figure 7.1 depicts the treated Fe3+-(CAB-O-SIL) samples after vacuum drying.

Here the anhydrous FeCl3 is seen as a dark black powder and the Fe-Cab samples are shown in various ratios by weight from 1 : 1 to 1 : 4. It is these samples that we have characterized by XPS.

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Fe Cab 1 : 4 Fe Cab 1 : 3 Fe Cab 1 : 1 O 1sO

Cl2p Fe 2p Fe Si2/32p Si2s C1s

Cl2s Fe 3p Fe C /C S

1000 800 600 400 200 0 Binding Energy / eV

Figure 7.2 Full-scale XPS spectra for Fe-CAB-O-SIL ratios as indicated in the figure.

Figure 7.2 depicts the full XPS spectra for the three Fe-Cab samples studied. The existence of Fe, Cl, and Si, is clearly apparent in the spectra. However, in order to better understand the interrelationship between the Fe3+ deposition and its transformation on the fumed silica surface which acts as a support for this metal ion, higher resolution XPS spectra are necessary. These higher resolution spectra provide more detailed information on the chemical environment of the metal ion. Figure 7.3 depicts the XPS spectra obtained for the Fe 3p transition. It is well known that a stronger band for Fe lies in the

2p region, however, for quantitative analysis of the Fe2+/Fe3+ ratio the satellite peaks in the Fe 2p region add considerable complexity to the analysis. Although weaker than Fe

2p, the Fe 3p region has been used to quantitatively analyze the Fe2+ / Fe3+ ratios in the pioneering work of Mekki et al.[6] and the validity of approach has been further

132 investigated by Yamashita et al.[7-9] and Paparazzo[10]. A detailed set of fitting parameters has been given to achieve the best fit by Yamashita group[7-9]. The Mekki et al.6 and Yamashita et al.[7-9] research groups identified a peak at ~54 eV to be Fe2+ and a peak at ~56 eV to be Fe3+. Figure 7.4 demonstrates the similar positions for which we have obtained XPS spectra for Fe2+and Fe3+.

Fe Cab 1 : 1 Fe Cab 1 to 3

Intensity/ a.u. Intensity/ a.u.

62 60 58 56 54 52 62 60 58 56 54 52 Binding Energy / eV Binding Energy / eV

Fe Cab 1 : 4

Intensity/ a.u.

62 60 58 56 54 52 Binding Energy / eV

Figure 7.3 De-convoluted Fe 3p XPS spectra for Fe-CAB-O-SIL samples with various Fe to Cab ratios.

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FeCl2

FeCl3

FeCl2 FeCl3

64 62 60 58 56 54 52 50 48 Binding energy / eV

Figure 7.4 Fe 3p XPS spectra for FeCl2 and FeCl3

Table 7.1 lists the de-convoluted peaks used to fit the Fe 3p XPS spectra for the

Fe-fumed silica samples given in Figure 7.2 for various Fe to Cab ratios. Clearly, the larger the ratio of fumed silica in the sample the larger is the percentage of Fe2+ in the sample. This indicates that CAB-O-SIL reduces the Fe3+. This reducing property of silicon is consistent with our earlier work on the oxidation states of Si in CAB-O-SIL[5] and with our recent theoretical calculations[1].Table 7.2 shows the Si 2p XPS peak positions for the Fe-CAB-O-SIL samples. After the incorporation of FeCl3, the silicon 2p

XPS binding energy shifts toward higher binding energy (BE) significantly for all three samples, which implies the occurrence of Si oxidation. There is a weak trend to a decreasing value of the binding energy when the relative amount of CAB-O-SIL in the

134 sample increases. This seems reasonable as the more CAB-O-SIL in the sample, on average, the less oxidation will occur on average at silicon.

The fit of the experimental data given in Figure 7.2 suggests a third feature at ~58 eV. To better understand the origin of this higher binding energy component, the XPS spectra for FeCl2 and FeCl3 (Figure 7.4) were taken. Interestingly, both FeCl3 and FeCl2 show a similar broad feature in the same 58 eV region (Figure 7.4), although there is a clear difference in the position of the major peaks for FeCl3 and FeCl2. We must question whether this feature arises from the chloride or iron ions. To answer this question, we further tested the Fe 3p XPS of FeSO4 and found the same shoulder at ~58 eV. Because no chloride peaks have been reported at 58 eV, it is unlikely that this ion is responsible for the 58 eV feature. To rule out an artifact associated with CAB-O-SIL itself, we analyzed the same region to obtain the spectrum shown in Figure 7.5.

Table 7.1 XPS de-convoluted peaks for Fe Cab samples using Gaussian functions Temp. center / Height / Sample Area/% Width / eV R2 Assignment eV a.u.

X* 49.4 57.1 2.9 116.9 Fe Cab 1 : 1 Fe3+ 46.7 55.8 1.7 189.6 0.996 Fe2+ 3.9 54.1 0.8 35.6 X 36.1 57.6 3.6 75.5 Fe Cab 1 : 3 Fe3+ 54.8 55.7 2.2 183.6 0.995 Fe2+ 9.1 54.4 1.4 47.1 X 20.2 58.3 2.8 40.5 Fe Cab 1 : 4 Fe3+ 65.3 55.9 2.3 159.3 0.997 Fe2+ 14.4 54.4 1.5 56.1 * X indicates the shoulder that corresponds to the tail caused by electron exchange interaction effects and electron correlation effects.[11]

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Table 7.2 XPS results of the Silicon 2p in the Fe-CAB-O-SIL samples Single peak fit results Sample 2 position / eV R pure Cab* 100.3 - Fe Cab 1 : 1 103.2 0.9980 Fe Cab 1 : 3 103.0 0.9969 Fe Cab 1 : 4 102.8 0.9974 * value from previous work[5]

Clearly, there is no peak at ~58 eV. Therefore, we can now safely reach the conclusion that the feature indicated at ~58 eV does come from iron. Here, we must ask whether this is a real shoulder or simply an artifact of the fitting procedure in concert with the asymmetrical tail intrinsic to many of the first row transitional metals.

Paparazzo[10] has carefully noted in discussing the studies of Yamashita et al.[7,8], that the XPS 2p and 3p spectra of the 1st row transitional metal paramagnetic ions all possess a complex line shape that originates from both electron exchange interaction effects and electron correlation effects. Electron exchange interaction effects are responsible for the so-called “multiplet splitting” states whereas electron correlation effects produce “shake- up” and/or “shake-off” states. As a result, the XPS spectra always contain a broad leading signal that is usually accompanied by an asymmetric tail to the higher binding energy

(BE) side of the peak region. This may include some high-BE satellite signals as well.

This is corroborated by the Co 2p XPS spectra shown in Figure 7.6. Furthermore, if we translate the Fe 3p XPS spectrum of FeCl2 so that it overlaps with that of FeCl3, we find that the high binding energy tails of these two are almost identical, suggesting that it is less likely that these features emanate from a specific chemical environment of Iron. With these considerations, we suggest that the higher binding energy feature at ~58 eV results from an asymmetrical tail. That being said, we focus on the remaining two de-convoluted

136 peaks. The lower binding energy corresponds to the lower oxidation states. As there are only two major oxidation states for iron ions, namely Fe2+ and Fe3+, it is reasonable, especially in conjunction with Figure 7.4, to ascribe the lower BE de-convoluted peak as

Fe2+ and the higher BE de-convoluted peak as Fe3+. The peak positions correspond well to the database value.[11]

Pure Cab at Fe 3p region

Intensity/ a.u.

62 60 58 56 54 52 Binding Energy / eV Figure 7.5 High resolution XPS spectra of pure CAB-O-SIL at 50 to 62 eV.

Co 2p in Co-TiO2

810 800 790 780 770 Binding Energy / eV

Figure 7.6 XPS spectra of Co 2p in a Co-TiO2 sample.

Table 7.2 reports the Si 2p binding energies for these three Fe-Cab samples. The sample most rich in silica shows the lowest binding energy (102.8 eV); whereas, the sample richest in Fe shows the highest binding energy (103.2 eV). A silica sample having

137 no Fe was reported to show a binding energy of 100.3 eV. These data suggest that the iron ions influence the Si 2p binding energies. To understand these results better, one may use literature XPS data of non-equilibrium silica that show Si oxidations states of 0 to 4, Figure 7.7 (reference 5 and citations therein).

4 3.5 y = 1.0399x - 103.9 3 2.5 2 1.5

Si Si State Oxidation 1 0.5 0 99 100 101 102 103 104 Si 2p Binding Energy, eV

Figure 7.7 Si XPS Binding Energy vs. Si Oxidation.[5]

These data may be correlated to show the relationship between observed Si 2p binding energy and the Si oxidation state. This correlation shows that the Fe-Cab 1 : 4 sample demonstrates a Si oxidation state of near 3; whereas, the Fe-Cab 1 : 1 samples shows a Si oxidation state of 3.4. The Fe-Cab 1 : 3 shows an average oxidation state of

3.2. The undecorated CAB-O-SIL shows an average Si oxidation state near 1. Thus, one may interpret the Si XPS data to show that the Fe3+ ions interact with the Si lower valent ions to undergo redox so that Fe2+ ions are produced as the Si ions are oxidized. When silica is in great excess, the average oxidation of the Si is lower (~3) compared to a value of 3.4 when equal amounts of Fe and Si are present initially.

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Surface Reaction Model.

A simple model was proposed to simulate the interaction between the iron species and the Cab-o-Sil surface. The aqueous source solution of FeCl3 can be assumed to show

3+ a variety of Fe species to include: FeCl3 nH2O; Fe(OH)mCl3-m nH2O, m=1-3;

FeO(OH)mCl1-m nH2O, m = 0-1. The Cab-o-Sil shows Si species terminated by silanol groups, hydrated silanol groups, siloxane bridging oxygens, and hydrated siloxane bridging oxygens. We have proposed a diverse distribution of Si ion oxidation states to include zero-valent Si up to tetravalent Si as first discussed by us previously.[1]

This model may be used to estimate the fraction of the total iron ions that come in contact with the Cab-o-Sil surface by assuming a value of 1.3 OH per nm2 for the silanol density and by assuming that one Fe3+ reacts with three surface OH groups upon initially contacting the surface. The assumed silanol density of 1.3 OH/nm2 is less than the customary number[12] of 2-4 OH/nm2 since a dry MeOH solvent is used and we expect this solvent to reduce the silanol density found on the CAB-O-SIL. Next, we assume that each of these surface ion species is reduced from 3+ to 2+ by this surface interaction between the Fe and Si species. As an example of this calculation, consider the 1/1 sample having a mass of 1 gram which shows 0.5 g of FeCl3 and 0.5 g of SiO2. The number of moles of Fe ions is 3.083 mmol whereas the number of mole of SiO2 is 8.33 mmol.

The number of surface SiOH is (0.5 g x 390 m2/g x 1.3 OH per nm2=) 2.6 x 1020

OH groups and the number of surface Fe species initially sequestered is 1/3 of this number or 8.6 x 1019 surface Fe species. If these surface Si species is 3+ or less, then all of these Fe species can be reduced from 3+ to 2+. Thus, the number of Fe2+ is 8.6 x 1019.

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The fraction of total Fe species in the 2+ state is 0.86 x 1020/1.86 x 1021 = 4.6%. The observed fraction of Fe in 2+ state is 9.7%. This model was used to estimate the fraction of Fe2+ in the other two Fe-Cab samples (Table 7.3). The predicted Fe2+ amounts for the other two samples show close agreement with the observed Fe2+ from XPS measurements.

This simple model suggests that the Fe ions in contact with the surface Si species are capable of undergoing redox to produce the observed divalent Fe species.

Table 7.3 Predicted and Observed Fe(II) and Fe(III) amounts on the Cab-O-Sil Surface. Model Predictions Observed from XPS SiO /FeCl , mass 2 3 % Fe (II) % Fe (III) % Fe (II) % Fe (III) ratio 1 4.6% 95% 9.7% 92.3% 3 13.9% 86% 14.2% 86% 4 18.6% 81% 18.1% 82%

Complexation with 1,10 Phenanthroline

Figure 7.8 Right to left, comparison of bright yellow sample of Fe-CAB-O-SIL 1 : 4, Fe- CAB-O-SIL 1 : 4 complexed with 1,10 phenanthroline, Fe2O3 over phenanthroline under toluene, and powdered Fe2O3.

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1,10 Phenanthroline is known to bind selectively to Fe2+, forming a red complex13.

In contrast, Fe3+, which is usually brown in color, does not complex with phenantroline.

Figure 7.8 shows the results of an experiment where both Fe2O3 and the product of introducing Fe3+ to fumed silica are exposed to 1,10 phenanthroline under toluene. Figure

7.9 displays the corresponding DRS spectra which clearly show the manifestation of the complex of phenanthroline and Fe2+, demonstrating the reduction of the Fe3+ originally placed on the surface of the CAB-O-SIL sample.

1,10-phenanthroline complex Fe Cab 1 : 4

Reflectance / a.u.

200 300 400 500 600 700 800 Wavelength / nm

Figure 7.9 UV-vis diffuse reflectance spectra for Fe-CAB-O-SIL (black) and the complex of 1,10 phenanthroline with Fe-CAB-O-SIL(red).

FTIR Spectroscopy

Figure 7.10 shows the FTIR of the CAB-O-SIL and the Fe-CAB-O-SIL series samples. The broad peak around 3000~3600 cm-1 together with the peak at about 1610

141 cm-1 corresponds to surface adsorbed water14. At first glance, this would seem to suggest that the surface of the CAB-O-SIL is somewhat hydrophobic or at least that the surface of the iron doped CAB-O-SIL samples has become more hydrophilic. However, a simple test of water contact angle[15], the result of which is shown in Figure 7.11, indicates that the surface of CAB-O-SIL is actually hydrophilic. However, it is to be noted that the +1 and +2 silicon sub-oxides uptake water and form surface hydroxyl groups at a considerably slower rate than does the surface consisting primarily of +3 and +4 oxidation states.[1] Thus the weak IR absorption band for water, especially for the untreated CAB-O-SIL sample, is not surprising. However, the seeming contradiction between the IR data and the water contact angle experiments suggests the need for a closer examination of the IR data. In comparison with our previous results[1], although the shapes of the IR peaks are quite similar, there is a shift in the position of the Si-O stretching peak from 1105 cm-1 in the previous work to 1074 cm-1 in the current study.

By contrast, there is only a very small shift observed for the Si-O-Si bending peak which is located at 812 cm-1 in our previous report1 and 810 cm-1 in the current study. It is noteworthy that there is a sequential change in the frequency of the IR peaks with increasing iron concentration. As shown in Table 7.4, the Si-O-Si bending peak decreases in frequency from 810 cm-1 to 796 cm-1 whereas the Si-O stretching mode decreases from

1074 cm-1 to 1049 cm-1.

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1049 Fe Cab 1 : 1 Fe Cab 1 : 2 1057 Fe Cab 1 : 4 Cab 1061

1074

796

Absorbance a.u. / 800 1608 802 810 3600 3000 2400 1800 1200 600 Wavenumber / cm-1

Figure 7.10 FTIR of CAB-O-SIL and the Fe-CAB-O-SIL series samples.

Table 7.4 Detailed analyses of IR peaks.

Si-O-Si bending difference Si-O stretching difference sample / cm-1 / cm-1 / cm-1 / cm-1 pure cab-o-sil 810 1074 8 13 Fe Cab 1 : 4 802 1061 2 4 Fe Cab 1 : 2 800 1057 4 8 Fe Cab 1 : 1 796 1049

143

Figure 7.11 (a) side view before water droplet is introduced to the CAB-O-SIL surface (b) side view after a water droplet has been introduced (c) front view after a water droplet is introduced.

The introduction of iron onto SiO2 results in a perturbation to both the Si-O stretching mode and the Si-O-Si bending mode, which is consistent with a significant interaction with the silicon site. Up to a monolayer coverage, the more iron in the system, the stronger its influence.

XRD Analysis

Figure 7.12 depicts the XRD patterns of the CAB-O-SIL and Fe-CAB-O-SIL samples. Apart from broad half peak from 20o to 50o, which is stipulated to be characteristic of amorphous silicon oxides,[16] there are no visible peaks in the entire 20o to 70o range. This may exclude the possibility of the existence of any crystalized silicon

144 oxide or iron oxide, at least at the surface of the fumed silica. In addition, the intensity of the broad half peak for the pure CAB-O-SIL sample is stronger than those for the Fe-

CAB-O-SIL samples.

CAB-O-SIL Fe Cab 1 : 1 Fe Cab 1 : 3 Fe Cab 1 : 4

Intensity / a.u. / Intensity

20 30 40 50 60 70

2 Theta / degree

Figure 7.12 XRD patterns of CAB-O-SIL and the Fe-CAB-O-SIL series samples.

TEM Analysis

Figure 7.13 catalogues TEM images of the CAB-O-SIL and Fe-Cab samples. The average observed particles sizes are determined to be 12.8 ± 3.3, 13.8 ± 2.5, 12.9 ± 3.3,

13.8 ± 3.6 nm for pure CAB-O-SIL, Fe Cab 1 : 1, Fe Cab 1 : 3, and Fe Cab 1 : 4, respectively. These determinations were made by measuring randomly selected particles within each image. It is reasonable to conclude that there is no major change after the introduction of Fe into the CAB-O-SIL system. This is consistent with the XRD result that there are no crystallites of iron formed. The shapes of the particles composing the

TEM images are dominantly irregular spheres or spheroids.

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Figure 7.13 TEM images of (a) pure CAB-O-SIL (b) Fe Cab 1 : 1 (c) Fe Cab 1 : 3 (d) Fe Cab 1 : 4.

7.4 Conclusion

We have investigated the changing properties of CAB-O-SIL before and after it is doped with Fe3+ ions. The size of the particles composing both the original sample and the iron treated CAB-O-SIL is determined to be around 13 nm based on the TEM results.

There is no obvious change in particle size after doping. The valence state of Si in fumed

146 silica, originally found to be +1, was found to change from +1 to +3 to +4 by XPS results.

In concert with the change in the deposited iron oxidation state from +3 to +2, we are lead to the conclusion that the surprising oxidation state of Si1+ indeed exhibits its reducing power once doped with an oxidant like FeCl3. It appears that the more CAB-O-

SIL is present in the sample (ratio of Cab to Fe3+), the higher will be the ratio of Fe2+/Fe3+ in the product. In other words, by tuning the percentage of the composing CAB-O-SIL, we will be able to control the ratio of Fe2+/Fe3+ in a sample. Interestingly, the ratio obtained from deconvolution of XPS spectra matches well with the calculated estimation.

This may be of great value in the design and synthesis of iron-related catalysts. The significance of this result is that only those Fe ions in contact with the Si ions participated in this redox reaction and thus we are able to exercise control of the oxidation state of the supported iron.

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7.5 References

1. Wang, T. H.; Gole, J. L.; White, M. G.; Watkins, C.; Street, S. C.; Fang, Z. T.; Dixon, D. A. Chemical Physics Letters 2011, 501, 159-165. 2. Prokes, S. M.; Carlos, W. E.; Seals, L.; Lewis, S.; Gole,J. L.; Materials Letter 2002, 54, 85. 3. Gole, J. L.; White, M. G., Journal of Catalysis 2001, 204,249. 4. Pavesi, L.; Gaponenko, S.; Dal Negro, I. (Eds.) Towards the First Silicon Laser, NATO Science Series II, Mathematics, Physics, and Chemistry, vol.93, Kluwer Academic Publisher, Dordrecht, 2003. 5. Gole, J. L.; Shinall, B. D.; Iretskii, A. V.; White, M. G.; Carter, W. B.; Erickson, A. S. Chemphyschem 2003, 4, 1016-1021. 6. Mekki, A.; Holland, D.; McConville, C. F.; Salim, M. Journal of Non-Crystalline Solids 1996, 208, 267-276. 7. Yamashita, T.; Hayes, P. Journal of Electron Spectroscopy and Related Phenomena 2006, 152, 6-11. 8. Yamashita, T.; Hayes, P. Journal of Electron Spectroscopy and Related Phenomena 2006, 154, 41-42. 9. Yamashita, T.; Hayes, P. Applied Surface Science 2008, 254, 2441-2449. 10. Paparazzo, E. Journal of Electron Spectroscopy and Related Phenomena 2006, 154, 38-40. 11. http://lasurface.com/database/elementxps.php 12. Handbook of Fillers for Plastics, H. S. Katz and J. V. Milewksi, 1987, P. 300. 13. Noorjaha, M.; Kumari, V. D.; Subrahmanyam, M.; Panda, L.; Applied Catalysis B: Environmental 2005, 57, 291-298. 14. Davydov, V. Y.; Kiselev, A. V.; and Zhuravlev, L. T. Study of the surface and bulk hydroxyl groups of silica by infrared spectra and D20 exchange. Transactions of the Faraday Society 1964, 69, 2254-2264. 15. Wang, J.; Mao, B.; Gole, J. L.; Burda, C.; Nanoscale, 2010, 2(10), 2257-2261. 16. Liu, K.; Feng Q.; Yang Y.; Zhang G.; Ou, L.; Lu, Y.; Journal of Non-Crystalline Solids 2007 353, 1534-1539. 17. Moulder, F. J.; Stickle, F. W.; Sobol, E. P.; and Bomben, D. K.; Handbook of X- ray Photoelectron Spectroscopy, Perkin-Elmer Corporation, Minnesota, 1992.

148

Chapter 8 Summary and Outlook

In the effort to render visible-light reactivity, we attempted to modify the electronic band structure of Titanium Dioxide with either metal or non-metal elements, or both. The results demonstrate different light absorbing properties for different methods: while metal dopants such as tungsten, iron and cobalt tend to move the whole absorption edge towards visible-light region, nitrogen-doped TiO2 always exhibit a shoulder-like absorption peak in the visible-light region. Most of the samples are successfully controlled in the scale of nanometers, although particle aggregation is observed. For all the photocatalysts tested, the performance for the decomposition of methylene blue can be ranked as follows: Ga,N-TiO2 > N-TiO2 > Co,N-TiO2 > Fe-TiO2 ≈ Fe,N-TiO2.

Nevertheless, the underlying mechanism as to why such ranking is observed is not clear.

Understanding the relationship between the absorption and the photocatalytic reactivity remains elusive. The source of this problem is two-fold. First, theoretical calculation is far from being able to get the exact result of experiments. The band gap calculated by the ab initio method is not even close in the case of anatase TiO2. Second, there are other factors such as crystallinity, surface defects, particles shapes and sizes that play a role in determining the outcome of photocatalytic activity. It is still a challenge to control each of these factors precisely. Therefore, future efforts should be focused on improving both theory and practice. Specifically, it would be beneficial to calculate the change of band gap when the dopant concentration changes. Also, in the case of co-doping, we can calculate the relationship between band gap and the relative position of the two dopants within the unit cell. In addition, calculating the free energy change for the reaction of

149 nitrogen release with and without a second dopant can provide us useful information as to what type of co-dopant should be used. As far as practice is concerned, future work should be focused on understanding the lifetime of the excitons and its relation with the properties such as crystallinity, the content of the dopant, the light absorption or the particle sizes, etc. Correlating the photocatalytic activity with the exiton lifetime would also be desired.

Finally, it is worth mentioning that there is already ongoing research of utilizing solar energy to split water with bioengineered microorganism. In the event that the right solid state material cannot be found in time, biomaterials may stand to serve as a commendable alternative.

150

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