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CHAPTER 1
INTRODUCTION
1.1 Literature review
Titanium is the world’s fourth most abundant metal and ninth most abundant element. It
was discovered in 1791 in England by Reverend William Gregor, who recognised the
presence of a new element in ilmenite [1]. It was then rediscovered in rutile ore several
years later by a German chemist, Heinrich Klaporth who named it after Titans,
mythological first sons of the goddess Ge (earth in Greek mythology). Titanium has the
symbol Ti, with atomic number 22, and atomic weight 47.90. It is placed in the fourth
group of the periodic table, and its chemistry shows similarities to that of silicon and
zirconium. The outer electronic arrangement is 3d24s2, and the principal valence state is
IV; III, II states are also known, but are less stable. The element burns in air when
heated to give the oxide, TiO2. Titanium is not found in its elemental state, it occurs
mainly in minerals like rutile, ilmenite, leucoxene, anatase, brookite, perovskite and
spene. It is also found in titanates and many iron ores. The metal has been detected in
meteorites and stars. In fact, samples brought back from the moon by Apollo 17
contained 12.1 % TiO2. The primary source and the most stable form of titanium
dioxide is rutile ore. It was discovered in Spain by Werner in 1803. Its name is derived
from the Latin rutilus, red because of the deep colour observed in some specimens when
the transmitted light is viewed. Rutile is one of three main polymorphs of titanium
dioxide (TiO2), the other polymorphs being; anatase and brookite [2, 3]. Brookite was
discovered in 1825 by A. Levy and was named after an English mineralogist, H. J.
Brooke. In 1801 anatase was named by R. J. Hauy from the Greek word ‘anatasis’
meaning extension, due to its longer vertical axis compared to that of rutile. In all three
forms, titanium (Ti4+) atoms are co-ordinated to six oxygen (O2-) atoms, forming TiO6
octahedra [4]. All three forms differ only in the arrangement of these octahedra. The
anatase structure, is made up of corner (vertices) sharing octahedral (figure 1.1a)
resulting in a tetragonal structure.
In rutile the octahedra share edges to give a tetragonal structure (figure 1.1b) and in
brookite both edges and corners are shared to give an orthorhombic structure (figure
1.1c)[5].
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(a) Anatase
(b) Rutile
(c)-Brookite
Figure 1.1 Crystalline structures of titanium dioxide (a)-rutile, (b)-anatase, (c)brookite [6]
Figure 1.2 more clearly shows the unit cell structures of the rutile and anatase TiO2 [7].
Both these structures can be described in terms of chains of TiO6 octahedra. The two
crystal structures differ in the distortion of each octahedron and by the assembly pattern
of the octahedra chains. In rutile, the octahedron shows a slight orthorhombic distortion;
in anatase, the octahedron is significantly distorted so that its symmetry is lower than
orthorhombic. The Ti-Ti distances in anatase are larger, whereas the Ti-O distances are
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shorter than those in rutile. In the rutile structure, each octahedron is in contact with 10
neighbour octahedrons (two sharing edge oxygen pairs and eight sharing corner oxygen
atoms), while, in the anatase structure, each octahedron is in contact with eight
neighbours (four sharing an edge and four sharing a corner). These differences in lattice
structures cause different mass densities and electronic band structures between the two
forms of TiO2.
Figure 1.2 Lattice structures of rutile and anatase TiO2
Titanium dioxide is an n-type semiconductor [8] that has a band gap of 3.2 eV for
anatase,[9-12] 3.0 eV for rutile [13], and ~3.2 eV for brookite[14]. Titanium dioxide
(TiO2) is the most widely investigated photocatalyst due to its strong oxidative
properties, low cost, non-toxicity, chemical and thermal stability [15]. Anatase and
rutile are the most researched polymorphs. Their properties are summarised in Table.1
and Table 1.2.
In the past few decades there have been several exciting breakthroughs with respect to
titanium dioxide. The first major breakthrough was in 1972 when Fujishima and Honda
reported the photoelectrochemical splitting of water (2H2O → 2H2 + O2) using a TiO2
anode and a Pt counter electrode. Titanium dioxide first showed promise for the
remediation of environmental pollutants in 1977 when Frank and Bard investigated the
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reduction of CN- in water [16]. This led into an increasingly well researched area of
TiO2 because of the potential implications for environmental water and air purification
utilising solar energy [17]. In 1997 Wang et al reported TiO2 surfaces with excellent
anti-fogging and self-cleaning abilities which were attributed to the super hydrophilic
attributes of the TiO2 surfaces [18].
Table 1.1 Physical and structural properties of anatase and rutile
Property
Anatase
Molecular weight ( g/mol) 79.88
Rutile
79.88
Melting point (°C)
1825
Boiling Point (°C)
2500 ~ 3000 2500 ~ 3000
Specific gravity
3.9
4.0
Light absorption (nm)
< 390
< 415
Mohr’s Hardness
5.5
6.5-7.0
Refractive index
2.55
2.75
Dielectric constant
31
114
Crystal structure
Tetragonal
Tetragonal
Lattice constants (Å)
a = 3.78
a = 4.59
c = 9.52
c = 2.96
Density (g/cm )
3.79
4.13
Ti–O bond length (Å)
1.94 (4)
1.95 (4)
1.98 (2)
1.98 (2)
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Nano sized titanium dioxide was employed to excellent use in an efficient solar cell, the
dye sensitised solar cell (DSSC) as reported by Graetzel and O’Regan in 1991[19].
1.2 Anatase to rutile transformation
The anatase to rutile phase transformation in TiO2 is the most studied area of scientific
and technological interest [20]. The anatase to rutile transformation (ART) is kinetically
defined and the reaction rate is determined by parameters such as particle shape/size
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[21], purity [22], source effects [23], atmosphere [24] and reaction conditions [25]. It is
agreed that the mechanism for phase transformation of titania is one of nucleation and
growth [26]. Anatase nanocrystals coarsen, grow and then transform to rutile only when
a critical size is reached [27]. Therefore, phase transformation is dominated by effects
such as defect concentration [28], grain boundary concentration [29], and particle
packing [30].
Table 1.2 X-Ray Data on TiO2 modifications (Clark, 1968: 268)
Titania
Space group
Z
Phase
Cell parameters A
a
Anatase
C4h= C4/ amc
8
3.784
Brookite
D2h= Pbca
8
9.51
Rutile
D4h= P42/mnm
2
4.593
b
5.44
Ti-O
c
( A)b
9.515
1.934 (4), 1.98 (2)
4.593
1.84-2.03
2.959
1.94(4), 1.98(2)
Rutile is the thermodynamically most stable phase, while anatase and brookite are both
metastable, transferring to rutile under heat treatment at temperatures typically ranging
from 873 to 973 K [2]. Anatase is widely regarded as the most photocatalytically active
of the three crystalline structures [31-33]. The generally accepted theory of phase
transformation is that two Ti–O bonds break in the anatase structure, allowing
rearrangement of the Ti–O octahedra, which leads to a smaller volume, forming a dense
rutile phase [34]. The removal of oxygen ions, which generate lattice vacancies,
accelerates the transformation. The transition follows first order kinetics; with activation
energy of ~ 418 k J mol-1[22]. The breaking of these bonds can be affected by a number
of factors, including the addition of dopants, synthesis method and thermal treatment.
1.3 Electronic Structure of a Semiconductor
When molecular orbitals are formed from two atoms, each type of atomic orbital gives
rise to two molecular orbitals. When N atoms are used, N molecular orbitals are formed.
In solids, N is very large, resulting in a large number of orbitals [35]. The overlap of a
large number of orbitals leads to molecular orbitals that are closely spaced in energy and
so form a virtually continuous band (figure 1.3) [36]. The overlap of the lowest
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unoccupied molecular orbitals (LUMO) results in the formation of a conduction band
and a valence band is formed from overlapping the highest occupied molecular orbitals
(HOMO). The band separation is known as the band gap (Eg), a region devoid of
energy levels. From the illustration shown in figure 1.3, the reduction in the band gap
size with the formation of bands can clearly be seen.
Figure 1.3 Change in the electronic structure of a semiconductor compound with
increasing number of monomer units [37]
If a band is formed from the molecular overlap of s orbitals it is called an s band and
likewise, an overlap of p orbitals, forms a p band and an overlap of d orbitals will give a
d band. Typically, p orbitals have higher energy than s orbitals, resulting in a band gap.
However, if the s and p bands are of similar energy, then the two bands overlap [38]. In
titanium dioxide, the valence band consists of oxygen 2 p orbitals and the conduction
band is made up from the titanium 3 d orbitals [12]. The band gap of TiO2 anatase is 3.2
eV and rutile is 3.0 eV corresponding to an absorbance threshold, λ = 388 and 415 nm
respectively. Figure 1.4 shows the various band positions of different semiconductors.
For a semiconductor to be capable of producing hydroxyl radicals, the potential of the
valence band must be greater than the potential of OH•. From figure 1.4 it can be seen
that ZnO, TiO2, WO3 and SnO2 have the potential to produce hydroxyl radicals.
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Figure 1.4 Band positions of various semiconductors and relevant redox couples
Titanium dioxide is regarded as an n-type semiconductor due to the presence of oxygen
vacancies in the lattice. These vacancies are formed upon the release of two electrons
and molecular oxygen leaving a positive (+2) oxide ion vacancy [2]. When electrons of
energy lower than the conduction band are present, the result is an n-type semiconductor
(figure 1.5a). Alternatively if a material is added with fewer electrons than the host,
positive holes are added above the valence band resulting in a p-type semiconductor
(figure 1.5b).
(a)
(b)
Figure 1.5 Semiconductors, n-type (a), p-type (b)
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1.4 A brief study of TiO2 as photo catalyst
Photo catalytic degradation of organic compounds on semiconductor surfaces is an
important area of current research with respect to both fundamental understanding [39]
and potential practical applications [40]. Several simple oxide and sulfide
semiconductors have band-gap energies sufficient for promoting or catalysing a wide
range of chemical reactions of environmental interest. The primary criterion for good
semiconductor photocatalysts for organic compound degradation is that the redox
potential of the H2O/·OH (OH- = ·OH + e-; Eº= -2.8 V) couple lies within the band gap
domain of the material (photocatalysts) and that the material is stable over prolonged
periods of time. Among the semiconductors TiO2 (Eg = 3.2 eV), WO3 (Eg = 2.8 eV),
SrTiO3 (Eg= 3.2 eV), a-Fe2O3 (Eg = 3.1 eV), ZnO (Eg = 3.2 eV), and ZnS (Eg = 3.6 eV)
[41], TiO2 is the most promising, low-cost, robust photocatalyst and extensively
investigated material. Semiconductor photocatalysis with a primary focus on TiO2 as a
durable photocatalyst has been applied to a variety of problems of environmental
interest in addition to water and air purification, for the destruction of microorganisms
[42], for the inactivation of cancer ce11s [43], for the photo splitting of water to produce
hydrogen gas [44], and for the clean-up of oil spills [45].
In the light of solid-state physics, semiconductors (and insulators) are defined as
solids in which at 0 K (and without excitations) the uppermost band of occupied
electron energy states is completely filled up with electrons. It is well-known from
solid-state physics that electrical conduction in solids occurs only via electrons in
partially-filled bands, so conduction in pure semiconductors occurs only when electrons
have been excited--thermally, optically, etc.--into higher unfilled bands.
At room temperature, a proportion (generally very small, but not negligible) of
electrons in a semiconductor have been thermally excited from the "valence band," the
band filled at 0 K, to the "conduction band," the next higher band. The ease with which
electrons can be excited from the valence band to the conduction band depends on the
energy gap between the bands, and it is the size of this energy band gap that serves as an
arbitrary
dividing
line
between
semiconductors
and
insulators.
Generally,
semiconductors are defined as materials with band gap less than 4 eV at room
temperature. Semiconductor electronic structures are characterized by a filled valence
band (VB), and an empty conduction band (CB). When a photon with energy of hv
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matches or exceeds the band gap energy (Eg) of the semiconductor, an electron (ecb-), is
promoted from the valence band, into the conduction band, leaving a hole (hvb+) behind.
Excited state conduction-band electrons and valence-band holes can react with electron
donors and electron acceptors adsorbed on the semiconductor surface or within the
surrounding electrical double layer of the charged particles or recombine and dissipate
the input energy as heat; get trapped in metastable surface states. The above process in
the photocatalysis is illustrated in Figure 1.6.
Figure 1.6 Schematic photo excitation in a solid followed by de-excitation events [7]
Once excitation occurs across the band gap there should be a sufficient lifetime (in the
nanosecond regime) for the created electron-hole pair to undergo charge transfer to
adsorbed species on the semiconductor surface from solution or gas phase contact. If the
semiconductor remains intact and the charge transfer to the adsorbed species is
continuous and exothermic the process is termed heterogeneous photocatalysis [46]. The
initial process for heterogeneous photocatalysis of organic and inorganic compounds by
semiconductors is the generation of electron-hole pairs in the semiconductor particles.
Upon excitation, the fate of the separated electron and hole can follow several
pathways. Recombination of the separated electron and hole can occur at the surface
(pathway A) in the volume of the semiconductor particle (pathway B) or with the
release of heat. The photo induced electron/hole can migrate to the semiconductor
surface. At the surface the semiconductor can donate electrons to reduce an electron
acceptor (usually oxygen in an aerated solution) (pathway C); in turn, a hole can
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migrate to the surface where an electron from a donor species can combine with the
surface hole oxidizing the donor species (pathway D). The electron transfer process is
more efficient if the species are spreadsorbed on the surface [47]. The probability and
rate of the charge transfer processes for electrons and holes depends upon the respective
positions of the band edges for the conduction and valence bands and the redox
potential levels of the adsorb ate species.
1.5 TiO2 nanostructured materials as semiconductor photocatalysts
Titanium (Ti) is a light, strong, lustrous, and corrosion-resistant metal. Titanium dioxide
(TiO2) is the most commonly used compound of titanium. Since its commercial
production in the early twentieth century, TiO2 has been widely used as a pigment in
sunscreens, paints, ointments, and toothpaste. It is also used in cements, gemstones, as
an optical opacifier in paper, and a strengthening agent in graphite composite fishing
rods and golf clubs. TiO2 powder is chemically inert, stable under sunlight, and is very
opaque. This allows it to impart a pure and brilliant white colour to the brown or gray
chemicals that form the majority of household plastics. However, in 1972, Fujishima
and Honda discovered the phenomenon of photocatalytic splitting of water on a TiO2
electrode under ultraviolet (UV) light. Since then, enormous efforts have been devoted
to the research of TiO2, which has led to many promising applications such as
photocatalysis, photovoltaic, photoelectrochromics and sensors. Some potential
applications of TiO2 are mentioned below.
1.6 Photocatalytic Applications
As has been pointed out by Heller, the development of semiconductor photo
electrochemistry during the 1970 and 1980s has greatly assisted the development of
photocatalysis [47, 48]. In particular, it turned out that TiO2 is excellent for
photocatalytically breaking down organic compounds. For example, if one puts
catalytically active TiO2 powder into a shallow pool of polluted water and allows it to
be illuminated with sunlight, the water will gradually become purified. Ever since 1977,
when Frank and Bard first examined the possibilities of using TiO2 to decompose
cyanide in water [16], there has been increasing interest in environmental applications.
These authors quite correctly pointed out the implications of their result for the field of
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environmental purification. Their prediction has indeed been borne out, as evidenced by
the extensive global efforts in this area [49]. One of the most important aspects of
environmental photocatalysis is the availability of a material such as titanium dioxide,
which is close to being an ideal photocatalyst in several respects. For example, it is
relatively inexpensive, highly stable chemically, and the photogenerated holes are
highly oxidizing. In addition, photogenerated electrons are reducing enough to produce
superoxide from dioxygen. The commonly studied principle of the photocatalysis
reaction mechanisms is given in Figure 1.7.
SEMICONDUCTOR TiO2
Figure 1.7 The principle of TiO2 photocatalysis [50]
When photons with energies larger than the band gap of TiO2 are absorbed, electrons
are excited from the valence band to the conduction band, creating electron-hole pairs
[50]. These photogenerated charge carriers can undergo recombination, become trapped
in metastable states, or can migrate to the surface of the TiO2, where they can react with
adsorbed molecules. In an aqueous environment saturated with air, the photogenerated
electrons and holes participate in reacting with dissolved molecular oxygen, surface
hydroxyl groups, and adsorbed water molecules to form hydroxyl and superoxide
radicals.
The energy band diagram for TiO2 in pH 7 solution is shown in Fig. 1.8.
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Figure 1.8 The energy band diagram for TiO2
As shown, the redox potential for photo generated holes is + 2.53V [41] versus the
standard hydrogen electrode (SHE). After reaction with water, these holes can produce
hydroxyl radicals (OH), whose redox potential is only slightly decreased. Both are more
positive than that for ozone. The redox potential for conduction band electrons is
−0.52V, which is in principle negative enough to evolve hydrogen from water, but the
electrons can become trapped and lose some of their reducing power, as shown.
However, even after trapping, a significant number are still able to reduce dioxygen to
superoxide O2−., or to hydrogen peroxide H2O2. Depending upon the exact conditions,
the holes, OH radicals, O2−., H2O2 and O2, all can play important role in the
photocatalytic reaction mechanisms.
Although the detailed mechanism of TiO2
photocatalysis reactions differs from one pollutant to another, it has been widely
recognized that superoxide and, specifically, •OH hydroxyl radicals act as active
reagents in the degradation of organic compounds. These radicals are formed by
scavenging of the electron–hole pair by molecular oxygen and water, through the
following sequences [51]:
TiO2 + hν → e− + h+
O2 + e− →O2 − •
H2O + h+ → •OH + H+
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•OH + •OH → H2O2
H2O2 + O2−•→ •OH + OH− +O2
Following a similar mechanism, the TiO2 photocatalyst can reduce CO2 to CH3OH or
produce H2 in a controlled environment (inert or oxygen free). However, the
photocatalytic activity of a semiconductor widely depends on its: (i) light absorption
properties, (ii) surface reduction and oxidation rates by the electron and hole, and (iii)
electron-hole recombination rates. On the other hand, the following three factors
pertaining to the band structure of semiconductors have the greatest effect on the
photocatalytic reactions: (i) band gap energy, (ii) position of the lowest point in the
conduction band, and (iii) position of the highest point in the valence band. In
photocatalytic reactions, the band gap energy principally determines which light
wavelength is the most effective, and the position of the highest point in the valence
band is the main determinant of the oxidative decomposition power of the photocatalyst.
Naturally occurring TiO2 has three polymorphs, i.e. anatase, brookite, and rutile.
Although all three types of polymorphs are expressed using the same chemical formula
(TiO2), their structures are different. Rutile is thermodynamically the most stable phase,
although the anatase phase forms at lower temperatures. Despite the fact that the band
gap values are 3.0 eV for the rutile and 3.2 eV for the anatase phases, both absorb only
UV rays. Hence both of these phases show photocatalytic activity whereas the brookite
phase does not. Characteristics of the rutile phase seem more suitable for use as a
photocatalyst because the rutile phase can absorb light of a wider range. But, the anatase
phase exhibits higher photocatalytic activity [52]. The most prominent reasons are
attributed to the difference in the energy structure between the two phase types and the
surface area. In both phases, the position of the valence band is deep, and the resulting
positive holes show sufficient oxidative power. However, the conduction band in the
anatase phase is closer to the negative position than in the rutile phase. Therefore, the
reducing power of the anatase phase is stronger than that of the rutile phase. Usually,
the anatase crystal phase formed at lower temperatures, show higher surface areas
compared to the rutile phase. A larger surface area with a constant surface density of
adsorbents leads to faster surface photocatalytic reaction rates. In this sense, the higher
the specific surface area, the higher the photocatalytic activity that one can expect.
Therefore, the anatase phase exhibits higher overall photocatalytic activity compared to
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the rutile phase. However, Hermann et al. [53] hypothesized that the anatase and rutile
mixed phases showed better photocatalytic activity because any kind of solid–solid
interface is a key structural feature that facilitates the charge separation to hinder
recombination. The interface also acts as an active site to initiate the catalytic activity
and enhance photocatalytic efficiency [54]. In addition, TiO2 nanomaterials with high
crystallinity show superior photocatalytic activity. High temperature treatment usually
improves the crystallinity of TiO2 nanomaterials, which can induce the aggregation of
small nanoparticles and decrease the surface area [55]. Therefore, it is very difficult to
predict the photocatalytic activities from the physical properties of TiO2 nanomaterials.
Optimal conditions are sought by taking into account all these considerations, which
may vary from case to case.
1.7 Other Potential Applications
In addition to the photocatalysis and photovoltaic applications, TiO2 nanomaterials can
also be used in biomedical applications [56], functionalized hybrid materials [57], and in
sensors and nanocomposites [58, 59]. The one dimensional TiO2 nanotubes are
promising for reinforcement due to their unique physical properties, i.e. large surface to
volume ratio, low cost and better biocompatibility compared with carbon nanotubes
[60]. These materials can be used as fillers for many applications such as a radio
pacifier in bone cement, a solid plasticizer of poly (ethylene oxide) for lithium batteries,
a dye in a conjugated polymer for photoelectrochemical photoconductive agents, and as
a photocatalyst in a photodegradable TiO2-polystyrene nanocomposite film [61, 62].
Moreover, TiO2 nanocrystalline films have been widely studied as sensors for various
gases. For instance, Varghese et al. observed that TiO2 nanotubes were excellent roomtemperature hydrogen sensors not only with a high sensitivity but also with ability for
self-cleaning after environmental contamination [63]. TiO2 nanomaterials also have
been widely explored as electro chromic devices, such as electro chromic windows and
displays. Electrochromism can be defined as the ability of a material to undergo colour
change upon oxidation or reduction. Electro chromic devices are able to vary their
throughput of visible light and solar radiation upon electrical charging and discharging
using a low voltage.
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1.8 Water splitting using semiconductor photocatalysis
The importance of hydrogen energy has recently been recognized again, especially for
use in fuel cells with due concern for environmental issues. However, industrially
hydrogen is produced by steam reforming of hydrocarbons such as methane.
Dihydrogen has to be produced from water using natural energy if one is concerned
about environmental issues. Therefore, water splitting using a photocatalyst is a
challenging reaction because it is one of the most important reactions for solving energy
and environmental problems. Water splitting using photocatalysts is not new. Water
splitting has been studied in the fields of catalysis, electrochemistry, photochemistry,
organic and inorganic chemistry, etc., for about 30 years since the Honda–Fujishima
effect was reported using a TiO2 semiconductor electrode. However, the number of
reported photocatalysts that are able to decompose water into H2 and O2 in a
stoichiometric amount with a reasonable activity has been limited. The only active
photocatalysts at present are oxide materials. Many people believe that water splitting is
impossible at a practical level. Against such a background, various types of new
photocatalysts for water splitting have been developed and this research field is taking a
new turn.
H2O cannot be photo decomposed on clean TiO2 surfaces, even though TiO2 can be
effectively photo excited under band-gap irradiation. Figure 1.9 illustrates the band edge
positions of TiO2 relative to the electrochemical potentials of the H2/H2O and O2/H2O
couple [64].
Figure 1.9 Potential energy diagram for the H2/H2O and O2/H2O redox couples
relative to the band-edge positions for TiO2.
15
According to this electron energy diagram, water photolysis is energetically favourable.
However, due to the presence of a large over potential for the evolution of H2 and O2 on
the TiO2 surface, TiO2 alone becomes inactive. Sustained photodecomposition of water
has been achieved under conditions where photogenerated electrons and holes are
separated for maximum photoreaction yield.
1.9 Limitation of TiO2 as an efficient photocatalyst and the modification of TiO2
There are some important drawbacks that severely limit the application of titanium
dioxide photocatalyst as a general tool either to degrade organic pollutants in the gas or
liquid phase or to perform useful transformations of organic compounds [65, 66]. One
of the most important limitations is the lack of TiO2 photocatalytic activity with visible
light [67, 68].
Figure 1.10 Solar spectrum at sea level with the sun at zenith with the absorption
region of TiO2
The reason for this is that the anatase form of TiO2 is a wide band gap semiconductor
with a band gap of 3.2 eV in most media, corresponding to an onset of the optical
absorption band at light photocatalytic activity, since approximately 5% of the solar
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light energy (figure 1.10) can be absorbed by TiO2. The above comments explain the
continued interest in improving the photocatalytic efficiency of TiO2.
1. 10 Modified Titania systems for improved photocatalytic activity
Two general strategies have been developed to increase the photocatalytic activity of
TiO2, namely the use of an organic dye as photosensitizer or doping TiO2 with metallic
and non-metallic elements [69-71]. The first route (using an organic dye that absorbs
visible light) has worked very well under conditions where oxygen/air is excluded and
the degradation of the dye is minimized by the efficient quenching of the dye oxidation
state with an appropriate electrolyte [72, 73]. Otherwise, particularly the dye becomes
rapidly mineralized and the photocatalytic system loses its response towards visible
light in the presence of oxygen. The mechanism of TiO2 dye sensitization has been
determined using time resolved sub nanosecond laser flash photolysis techniques [74].
In dye sensitization, the most relevant points are the absorption spectrum of the dye in
the visible region and the energy of the electron in the excited electronic state of the
dye, which has to be high enough to be transferred to the semiconductor conduction
band.
A second chemical route to promote titania photoresponse into the visible spectra is the
doping of TiO2 material either with metallic or non-metallic elements [75]. In this case,
doping introduces occupied or unoccupied orbitals in the band gap region leading to
negative or positive doping, respectively. A summary of some novel preparations of UV
and visible light responsive titania photocatalysts developed over the last few years has
been recently compiled [76].The limitations of titania semiconductor as a photocatalyst
for a particular use can be surmounted by modifying the surface of the semiconductor in
the following ways:
1.10.1 Metal semiconductor modification
Among the metal doped titania, Pt doping has recently attracted a great deal of attention
due to promising improvement in photo oxidation rate especially in gas phase. It has
been found that Pt–TiO2 improves the photo oxidation rate of ethanol, acetaldehyde and
acetone in the gaseous phase [77]. Li et al [78] reported a further development of
mesoporous titania as photocatalyst. These authors have prepared a photocatalyst
constituted by mesoporous titania embedding gold nanoparticles (Au/TiO2) whose
preparation requires P-123 as structure directing agent, a mixture of TiCl4 and Ti(OBu)4
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and AuCl3 as the source of gold using ethanol as solvent. The gel is cast on a Petri dish
to form a thin layer that is subsequently aged at 373 K to form a homogeneous
mesostructured nanocomposite. Calcinations at 350°C in air remove the template while
inducing crystallization of TiO2 and formation of gold nanoparticles. TiO2 has also been
doped with other metals, including V, Cr, Mn, Fe and Ni. The presence of the dopant
was found to cause large shift in the absorption band of titanium dioxide towards the
visible region. However, there are contradictory reports, particularly in the case of metal
doping, describing either an increase or a decrease of the photocatalytic activity [79].
This controversy arises in part from the difficulty to establish valid comparisons
between the photocatalytic activity of various solids testing different probe molecules
and employing inconsistent parameters. Also, the doping procedure and the nature of
the resulting material are very often not well defined and, most probably, controversial
results can be obtained depending on the way in which the metal has been introduced. It
also depends on the final concentration of the dopant. Thus, it has often been reported
that there is an optimum doping level to achieve the maximum efficiency and beyond
this point a decrease in photocatalytic activity is again observed [80]. Nevertheless, a
generalised consensus has been reached with regards to the inappropriateness of metal
doping as valid solution to enhance the photocatalytic activity of TiO2.
1.10.2 Composite semiconductors
Figure 1.11 Photoexcitation in composite semiconductor photocatalyst.
The efficiency of a photocatalytic process can be increased by coupled semiconductor
photocatalysts which provide an interesting way to increase the charge separation, and
extending the energy range of photoexcitation for the system. Figure 1.11 illustrates
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geometrically and energetically the photoexcitation process for the composite (coupled)
semiconductor photocatalyst CdS-TiO2. The energy of the excitation light is too small
to directly excite the TiO2 portion of the photocatalyst, but it is large enough to excite
an electron from the valence band across the band gap of CdS (Eg = 2.5 eV) to the
conduction band. The hole produced in the CdS valence band from the excitation
process remains in the CdS particle while the electron transfers to the conduction band
of the TiO2 particle. The electron transfer from CdS to TiO2 increases the charge
separation and efficiency of the photocatalytic process. The separated electron and hole
are then free to undergo electron transfer with adsorbates on the surface. The quantum
yield for the reduction of methylviologen drastically increased and approached an
optimum value of 1 when the concentration of TiO2 was increased in a CdS-TiO2
system [81]. The coupling of semiconductors with the appropriate energy levels can
produce a more efficient photocatalyst via better charge separation.
1.10.3 Surface sensitization
Figure 1.12 Excitation steps using dye molecule sensitizer.
The efficiency of the excitation process of wide band-gap semiconductor photocatalyst
(TiO2) can also be increased by surface sensitization of a chemisorbed or physisorbed
dyes. The photosensitization process can also expand the wavelength range of excitation
for the photocatalyst through excitation of the sensitizer followed by charge transfer to
the semiconductor. Erythrosin B, 3 thionine, and analogs of Ru (bpy) 3
2+
are some of
the common dyes which are used as sensitizers. Figure 1.12 illustrates the excitation and
charge injection steps involved for the regenerative dye sensitizer surface process.
Excitation of an electron in the dye molecule occurs to either the singlet or triplet
excited state of the molecule. If the oxidative energy level of the excited state of the dye
19
molecule with respect to the conduction band energy level of the semiconductor is
favourable (i.e. more negative), then the dye molecule can transfer the electron to the
conduction band of the semiconductor. The surface acts as a quencher accepting an
electron from the excited dye molecule. The electron in turn can be transferred to reduce
an organic acceptor molecule adsorbed on the surface.
1.10.4 Transition metal doping
An interesting area of semiconductor modification is the study of the influence of
dissolved transition metal impurity ions on the photocatalytic properties of TiO2. The
advantage of transition metal doping species is the improved trapping of electrons to
inhibit electron-hole recombination during light illumination. Different metals have
been employed to tune the electronic structure of TiO2-based material either by the ion
implantation method or by a wet chemical method [82]. The photocatalytic reactivity of
metal-doped TiO2 depends on many factors, including the dopant concentration, the
energy level pattern of the dopants within the TiO2 lattice, their d-electronic
configuration, and the distribution of the dopants. Enhanced photocatalytic activity has
been reported for Fe (III), Mo (V), Ru (III), Re (V), V (IV), and Rh (III) metals doped
TiO2 at 0.5% atomic ratio. It has been reported that the metal ions physically implanted
have produce significant changes in the TiO2 material. However, metal doping can
result in thermal instability and increased carrier trapping.
1.10.5 Non-metal doped TiO2 for photocatalysis
The elements used as the dopants in TiO2 cover almost the whole periodic table and is
summarized in Figure 1.13. Recent theoretical and experimental studies have shown
that the desired band gap narrowing of TiO2 can also be achieved by using a nonmetal,
although there is controversy regarding the origin of the resulting band gap narrowing.
Asahi and co-workers calculated the electronic band structures of anatase TiO2 with
different substitutional dopants, including C, N, F, P, or S and found that the
substitutional doping of N was the most effective for band gap narrowing. This is
because the p states of N mix with the 2p states of O, while the molecularly existing
species, e.g., NO and N2 dopants, give rise to the bonding states below the O 2p valence
20
bands, while the antibonding states deep within the band gap hardly interact with the
band states of TiO2 [83].
Figure 1.13 The elements used as the dopants in TiO2
Di Valentin et al. reported that the N 2p orbital formed a localized state just above the
top of the O 2p valence band, both for the anatase and rutile phases. However, these two
crystal systems showed the opposite photoresponce. In anatase, these dopant states
caused a red shift of the absorption band edge toward the visible region, while in rutile,
an overall “blue shift” was observed. This experimental evidence confirmed that Ndoped TiO2 formed N induced midgap levels slightly above the oxygen 2p valence band
[84]. Nakano et al. reported that in N-doped TiO2, deep levels located at approximately
1.18 and 2.48 eV below the conduction band were attributed to the O vacancy state and
a band gap narrowing was achieved by mixing with the O 2p valence band, respectively
[85]. N-doped TiO2 normally has a colour from white to yellow or even light gray, and
the onset of the absorption spectra red shifts to longer wavelengths. Table 1.3 shows
different methods used to prepare N-doped and S-doped titania. In N-doped TiO2
nonmaterial, the band gap absorption onset shifted 600 nm from 380 nm for the
undoped TiO2, extending the absorption up to 600 nm. The optical absorption of Ndoped TiO2 in the visible light region was primarily located between 400 and 500 nm,
while that for oxygen deficient TiO2 was mainly above 500 nm. N-F-co-doped TiO2
prepared by spray pyrolysis was shown to absorb light up to 550 nm in the visible
spectrum [102, 103]. Livraghi et al. recently found that N-doped TiO2 contained single
atom nitrogen impurity centres, localized in the band gap of the oxide, which were
responsible for the visible light absorption with promotion of electrons from the band
gap localized states to the conduction band [104].
21
Table 1.3 Various methods of preparation of N and S doped titania
N-Doped titania
Reference
By hydrolysis of TTIP in a water/amine mixture and the post 86, 87
treatment of the TiO2 sol with amines
Directly from a Ti-bipyridine complex
88
By ball milling of TiO2 in a NH3 water solution
89
By heating TiO2 under NH3 flux at 500-600 °C
90, 91
By calcination of the hydrolysis product ofTi(SO4)2
92
with ammonia as precipitation
By decomposition of gas-phase TiCl4 with
93
an atmosphere microwave plasma torch
By sputtering/ion-implanting techniques with
94, 95
nitrogen or N2+ gas flux.
S-Doped Titania
Reference
By mixing TTIP with ethanol containing thiourea
96, 97
By heating sulfide powder
98, 99
10. By using sputtering or ion-implanting
100, 101
techniques with S+ ion flux
1.11 Synthetic Methods for TiO2 Nanostructures: Various synthetic methods used for
the preparation of TiO2 Nanostructures are as follows1.11.1 Sol-gel methods:
The sol-gel method is a versatile process which is widely used for synthesizing various
oxide materials [105-107]. This synthetic method generally allows control of the
texture, the chemical, and the morphological properties of the solid. This method also
has several advantages over other methods, such as allowing impregnation or co
precipitation, which can be used to introduce dopants. The major advantages of the solgel technique includes molecular scale mixing, high purity of the precursors, and
homogeneity of the sol-gel products with a high purity of physical, morphological, and
chemical properties. In a typical sol-gel process, a colloidal suspension, or a sol, is
22
formed from the hydrolysis and polymerization reactions of the precursors, which are
usually inorganic metal salts or metal organic compounds such as metal alkoxides.
When a metal precursor alkoxide/salt reacts with water, nucleophilic substitution occurs
as shown in Figure 1.14 For the coordinative saturated metals such as the metal
alkoxides, hydrolysis and condensation both occur by nucleophilic substitution (SN)
mechanisms, in which coordination and proton transfer are involved and followed by
removal of either alcohol or water. Complete polymerization and loss of solvent leads to
the liquid sol transforming into a solid gel phase.
Figure 1.14 Hydrolysis and condensation pathways of metal alkoxides
Thin films can be produced on a piece of substrate by spin-coating or dip-coating. A
wet gel will form when the sol is cast into a mold, and the wet gel is converted into a
dense ceramic upon further drying and heat treatment. A highly porous and extremely
low-density material, called an aerogel, is obtained if the solvent in a wet gel is removed
under supercritical conditions. Figure 1.15 shows the conventional scheme to produce
aerogels. This method involves hydrolysis of the metal alkoxide with water and a
catalyst, i.e. an acid or a base, condensation into macromolecules, forming a colloidal
sol and subsequently three-dimensional network, solvent exchange to remove water by
alcohol, and then drying the wet gel using a supercritical fluid to produce the aerogel.
Aerogels can be used for advanced applications including electrochemical devices, thin
coatings, composite biomaterials, catalysts, ceramics, and heat and electric insulation
23
devices due to the aerogels having unique morphological and chemical properties.
Ceramic fibbers can be drawn from the sol when the viscosity is adjusted into the proper
range. Since the precursors react with water quickly and tend to precipitate, many
methods have been used to control the reaction rate in order to obtain desired
nanostructures. However, in binary metal systems, the main difficulty is to control the
hydrolysis rate of the precursors. Modifying the precursors with a suitable complexing
agent is one way to alter the hydrolysis rate, while generating in-situ water (instead of
externally added water) is another possible way to control the hydrolysis rate. In
practice, acetic acid is often used to modify the reactivity of metal alkoxides by
replacing the alkoxy groups bonded to the metal by acetate groups, forming M-OAc
complexes and alcohols.
Metal alkoxide
Water+ acid (or base)
+alcohol
Hydrolysis
Aging
Calcinatio
Figure 1.15- Conventional scheme to produce aerogels
Moreover, acetic acid (HO-Ac) can react with the alcoholic solvent and intermediates,
(produced by modification reaction) to form water via an esterification reaction [108]
However, the properties of the sol-gel products depend on the precursors, processing
temperature, catalyst, solvents, and solvent removal process [109]
1.11.1.1 Mechanism of Sol-Gel Synthesis
The sol-gel process involves hydrolysis and condensation of the metal alkoxide
followed by heat treatment at elevated temperatures which induce polymerisation,
24
producing a metal oxide network [110]. In general, transition metals have low
electronegativity’s and their oxidation state is frequently lower than their coordination
number in an oxide network. Therefore, coordination expansion occurs spontaneously
upon reaction with water or other nucleophilic reagents to achieve their preferred
coordination [111]. Metal alkoxides are in general very reactive due to the presence of
highly electronegative OR groups (hard-π donors) that stabilise the metal in its highest
oxidation state and render it very susceptible to nucleophilic attack. The lower
electronegativity of transition metals causes them to be more electrophilic and thus less
stable toward hydrolysis, condensation and other nucleophilic reactions. Controlling the
conditions can be difficult but successful control of the reaction conditions has the
potential to produce materials of consistent size, shape and structure [112-114].
1.11.1.2 Mechanism of hydrolysis
As stated previously the sol-gel reaction involves hydrolysis of the metal alkoxide
followed by condensation. Hydrolysis of titanium alkoxides occurs through a
nucleophilic substitution (SN) reaction. When a nucleophile, such as water is introduced
to titanium alkoxide a rapid exothermic reaction proceeds. The nucleophilic addition
(AN) of water involves a proton from the attacking nucleophile (water) being
transferred to the alkoxide group. The protonated species is then removed as either
alcohol or water (scheme 1.) [107, 115].
HOH+ RO M → HO M + ROH
Scheme 1- Hydrolysis reaction
The nucleophilic substitution reaction that occurs during hydrolysis can be described as
follows [115]:
1.
Nucleophilic addition of the H2O onto the positively charged metal atom.
2.
Proton transfer, within the transition state from the entering molecule to the
leaving alkoxy group.
1.11.1.3 Mechanism of condensation
Condensation reactions complete the sol-gel process. Condensation can proceed through
either alcoxolation or through oxolation. In both processes an oxo bridge is formed
25
between the metals (M–O–M) but the leaving group differs. During alcoxolation, two
partially hydrolysed metal alkoxide molecules combine and an oxo bridge is formed
between the two metals with alcohol departing as the leaving group (scheme 2) [109,
115].
Scheme 2 Alcoxolation reaction
In oxolation (scheme 3) two partially hydrolysed metal alkides combine to form an oxo
bridge between the metal centres but water is the leaving group.
Scheme 3 Oxolation reaction
Condensation reactions ultimately polymerised to form large chains of molecules.
Solvent used also influence the reaction kinetics of sol-gel reactions. Because if a
suitable solvent is chosen; it may be preferentially hydrolysed over the alkoxide ligands.
This allows for greater control over the reaction kinetics through the correct use of
solvent [116]. Hydrolysis and condensation kinetics are also affected by the organic
ligand. The hydrolysis rate of titanium alkoxides decreases with increasing alkyl chain
length [117, 118]. This is mainly due to the steric effect which is expected for an
associative SN reaction mechanism
1.11.1.4 Role of the catalyst
Acid and base catalysts can have a strong influence over hydrolysis and condensation
rates as well as the structural properties of the final product. The addition of an acid will
26
result in the preferential protonation of the negatively charged alkoxide group (scheme
4). Protonation of the alkoxide group produces good leaving groups which enhance the
reaction kinetics, eliminating any proton transfer that occurs in the transition state and
thus allowing hydrolysis to go to completion upon the addition of water.
H
M
+
OR + H → + M
O
R
Scheme 4- Acid catalyst reaction
In alkaline conditions, hydroxo ligands are deprotonated, producing strong
nucelophiles:
L–OH +: B → L–O- + BH+
where L = M or H and B = OH- or NH3. The hydrolysis rate of Ti (OBu) 4 was less in
basic conditions than in acidic or neutral conditions. It was postulated that the
nucleophilic addition of OH- may reduce the partial charge of the titanium metal centre
(δ (Ti)), making the hydrolysis reaction less favourable [109, 119]. While hydrolysis
kinetics is retarded in basic conditions, the condensation kinetics of metal alkoxides is
enhanced.
1.11.2 Hydrothermal Method
The hydrothermal method is widely used for the production of small particles in the
ceramics industry. Usually, this process is conducted in steel pressure vessels called
autoclaves under controlled temperature and/or pressure with the reactions occurring in
aqueous solution. The temperature can be elevated above the boiling point of water,
reaching the pressure of vapour saturation. The temperature and the amount of solution
added to the autoclave largely determine the internal pressure produced. The
hydrothermal method has been widely used to prepare TiO2 nanotubes after the
pioneering work by Kasuga et al. in 1998 [120]. Generally, several grams of TiO2
powders are added to concentrated NaOH solution and held at 20-110oC for 20 h in an
autoclave. TiO2 nanotubes were obtained after the products were washed with a dilute
HCl aqueous solution and distilled water. When the raw TiO2 material was treated with
27
NaOH aqueous solution, some of the Ti-O-Ti bonds were broken and new Ti-O-Na and
Ti-OH bonds were formed. New Ti-O-Ti bonds were formed after the Ti-O-Na and TiOH bonds reacted with acid and water. The bond distance from one Ti atom to the next
Ti atom on the surface was found to decrease. This resulted in the folding of the sheets
and connection between the ends of the sheets, resulting in the formation of a tube
structure. In this mechanism, the TiO2 nanotubes were formed in the stage of the acid
treatment following the alkali treatment. The hydrothermal method is also currently
being used to produce many new hybrid materials with new morphologies and
improved crystallinity.
1.11. 3 Micelle and Inverse Micelle Methods
Micelles and inverse micelles are commonly employed to synthesize TiO2
nanomaterials. A statistical experimental design method was conducted by Kim et al. to
optimize experimental conditions for the preparation of TiO2 nanoparticles [121]. The
values of H2O/surfactant, H2O/titanium precursor, ammonia concentration, feed rate,
and reaction temperature were significant parameters in controlling TiO2 nanoparticle
size and size distribution.
1.11.4 Sol Method: The sol method here refers to the nonhydrolytic sol-gel processes
andusually involves the reaction of titanium chloride with a variety of different oxygen
donor molecules, e.g., a metal alkoxide or organic ether [122-125]
1.11.5 Solvothermal Method: The solvothermal method is almost identical to the
hydrothermal method except that the solvent used here is nonaqueous.
1.11.6 Direct Oxidation Method: TiO2 nanomaterials can be obtained by oxidation of
titanium metal using oxidants or under anodization.
1.11.7 Chemical Vapour Deposition: Vapour deposition refers to any process in which
materials in a vapour state are condensed to form a solid-phase material. Thick
crystalline TiO2 films with grain sizes below 30 nm as well as TiO2 nanoparticles with
28
sizes below 10 nm can be prepared by pyrolysis of TTIP in a mixed helium/oxygen
atmosphere, using liquid precursor delivery [126].
1.12 Characterization Methods for TiO2 Nanostructures
1.12.1 XRD: XRD gives the following information:
i)
The nature of materials such as amorphous or crystalline
ii)
Types of structure and phases
iii)
Degree of crystallinity and thermal stability
iv)
In case of metal doping in a particular structure it can tell us whether the metal
atom has been incorporated into the framework by measuring the unit cell volume.
The powdered diffraction pattern of titania semiconductors are generally recorded
between 2θ values of 10-90 0. XRD patterns of nano-TiO2 and microTiO2 in rutile and
anatase phases are shown in Figure 1.16 and Figure 1.17, respectively [127]. In Figure
1.16, XRD patterns exhibited strong diffraction peaks at 27°, 36° and 55° indicating
TiO2 in the rutile phase. On the other hand, in Figure 1.17, XRD patterns exhibited
strong diffraction peaks at 25° and 48° indicating TiO2 in the anatase phase. All peaks
are in good agreement with the standard spectrum (JCPDS no.: 88-1175 and 84-1286).
Figure 1.16 X-ray diffraction of rutile TiO2 (a) micropowders and (b)
nanopowders.
29
Figure 1.17 X-ray diffraction of anatase TiO2 (a) micropowders and (b)
nanopowders.
These results suggested that the nano-TiO2 powder is composed of irregular
polycrystalline. Amorphous revealed a broad pattern with low intensity; however, the
effect of the amorphous materials on the broadening of the XRD patterns of nanosized
TiO2 is negligible. The low-angle X-ray diffraction patterns of the mesostructured
titanium dioxide thin films before and after calcinations (573 or 673 K for 4 h) are
shown in Figure 1.18 [128].
Figure 1.18 Low-angle X-ray diffraction patterns of a mesostructured titanium
dioxide thin film before (––) and after (- - - -) calcination (573 K for 4 h). The insets
show expanded views of the (110) and (200) regions for each.
30
Table 1.4 2 Theta values and hkl planes of anatase, rutile and brookite phase of
titania.
PHASE
2-THETA
ANATASE
25.30
(101)
0
37.8
(004)
540
(105)
25.40
(120)
30.80
(121)
40.20
(022)
0
(110)
0
(101)
0
(211)
BROOKITE
RUTILE
hkl PLANE
27.3
35.9
54.1
JCPDS
84-1286
29-1360
86-1175
The XRD of the as-synthesised films display 3 peaks in the low-angle range with d
(interatomic plane) spacing corresponding to 9.59 nm, 5.66 nm and 4.93 nm.
1.12.2 Raman Spectroscopy
Raman spectroscopy is used in condensed matter physics and chemistry to study
vibrational, rotational, and other low-frequency modes in a system. It relies on inelastic
scattering, or Raman scattering of monochromatic light, usually from a laser. The laser
light interacts with phonons or other excitations in the system, resulting in the energy of
the laser photons being shifted up or down. The shift in energy gives information about
the phonon modes in the system. Infrared spectroscopy yields similar, but
complementary information.
31
Figure 1.19 Raman spectra of anatase, rutile and brookite phases of titania
The three crystalline phases of titania; anatase, rutile and brookite, exhibit wellseparated and distinct Raman activities (figure 1.19) following 1064 nm excitation
which allows unambiguous identification of phase composition [129].
Knowledge of
the Raman scattering cross-sections allows analysis of both crystallite size and phase
compositions by comparing the relative intensities of isolated bands such as those
appearing at 320 cm-1 (brookite), 608 cm-1 (rutile) and 519 cm-1 (anatase)[130, 131].
1.12.3 I.R.Study
The Diffuse Reflectance Fourier-Transform Infrared (DRFT-IR) spectroscopy helps to
identify characteristic bands of the mesoporous framework. For example let us consider
the FT-IR spectra of the F-doped TiO2 sol particles shown in Fig. 1.20 [132]. The TiO2
sol particles showed the main bands at 400–700 cm–1, which were attributed to Ti-O
stretching and Ti-O-Ti bridging stretching modes. The small peak at 889.0 cm–1 was
attributed to Ti-F vibration [133]. The peak at 1384.0 cm–1 was produced by NO3–1. The
stronger peak at 1626.0 cm–1 was attributed to bending vibrations of O-H and N-H. The
IR spectra of the sol sample dried at 383 K revealed that Ti-O, Ti-F, N-H, Ti-OH, H-OH groups existed in the as-prepared sample by sol-gel-hydrothermal method.
32
Figure 1.20 FT-IR spectra of the F-doped TiO2 sol particles
The forming of Ti-F indicated that F atoms were incorporated into the TiO2 crystal
lattice.
1.12.4 UV-Vis Spectroscopy
Many molecules absorb ultraviolet (UV) or visible light. The absorption of UV or
visible radiation is caused by the excitation of outer electrons, from their ground state to
an excited state. This technique is based on the reflection of light in the ultra violet,
visible and near infrared region by a powdered sample. The DRS study (Fig. 1.21) gives
valuable information on the coordination of the transition metal such as Ti. Titania
samples have in common an intense UV absorption band with a maximum in the range
220–320 nm due to charge transfer from oxygen to titanium (IV) [134]. The position of
this band is affected by the coordination geometry around the titanium atom and by the
presence of adsorbents. More precisely, the bands in the region 210– 240 nm are
attributed to oxygen to tetrahedral Ti (IV) [135], whereas the bands at higher
wavelength (λ>240 nm) are due to octahedral Ti (IV) sites [136]. For TiO2 in the form
of anatase the transition occurs in 315 nm region. Hence Diffuse Reflectance UVVisible Spectroscopy is a widely available technique and is a very useful tool for
characterization TiO2 structure.
33
Figure 1.21 DRS of F-doped titania
The absorption spectra of pure undoped TiO2, pure F-doped TiO2 and F- doped TiO2 sol
prepared by the sol-gel-hydrothermal method and dried in a rotatory evaporator and
dried in vacuum at 333 K was shown in Fig.1.21A [132]. It was found that sample FTiO2 could cause a new absorption band in the visible range of 400–600 nm apart from
the fundamental absorption edge of TiO2, which was located in the UV region at about
385 nm (shown in Fig. 1.21C); whereas, pure F-TiO2 did not lead to any significant
shift in optical absorption of TiO2 (shown in Fig. 1.21B).
1.12.5 X-ray Photoelectron Spectroscopy (XPS)
XPS is a quantitative spectroscopic technique that measures the elemental composition,
empirical formula, and the chemical state and electronic states of the elements that exist
within a material. XPS spectra are obtained by irradiating a material with a beam of
aluminium or magnesium X-rays while simultaneously measuring the kinetic energy
(KE) and number of electrons that escape from the top 1 to 10 nm of the material being
analysed. Because the energy of a particular X-ray wavelength equals a known quantity,
we can determine the electron binding energy (BE) of each of the emitted electrons by
using an equation that is based on the work of Ernest Rutherford (1914):
Ebinding = Ephoton - Ekinetic – Φ; where E binding is the energy of the electron
emitted from one electron configuration within the atom, E photon is the energy of the
34
X-ray photons being used, E kinetic is the kinetic energy of the emitted electron as
measured by the instrument, and Φ is the work function of the spectrometer.
Fig. 1.22 XPS survey spectra of the 4%Fe doped TiO2 films grown on LAO
substrate. [137].
Fig. 1.23 XPS spectra of (a) N-doped mesoporous titania (1:1), (b) N-doped
mesoporous titania (1:2) and (c) N-doped mesoporous titania (1:3) [138].
Figure 1.22 and 1.23 gives the XPS spectra of Fe-doped and N-doped titania.
35
1.12.6 Thermal Analysis
Thermal analysis is an important technique for the characterization of solid materials, as
it provides information on the physical and chemical changes involving endothermic
and exothermic processes, temperatures for phase transitions, melting points and
crystallization, and the weight loss when the temperature is increased. TGA can
determine: (1) moisture/liquid content and the presence of volatile species, (2)
decomposition temperatures, and (3) the rate of degradation.
Figure 1.24 gives the TGA data of Er-doped luminescent TiO2 nanoparticles [139]. The
results indicate there were about 4% weight loss from room temperature to 773 K and
little weight loss after that. The weight loss from room temperature to 373 K is almost
certainly from removal of water, while it is suspected that the weight loss above 373 K
is due to removal of water, and perhaps hydroxyl groups.
Figure 1.24: TGA data obtained from 3% Er-doped TiO2
Full decomposition of hydroxyl groups requires a heat treatment around 1273 K;
however, it is seen that little weight loss between 773 K and 1073 K, indicating that
only trace structural water is present at 773 K since most would likely be removed by
1073 K.
1.12.7 N2 Physisorption
The textural characterization, such as surface area, pore volume and the pore size
distribution of the titania samples can be obtained by N2 physisorption studies. Prior to
N2 physisorption, the samples have to be degassed at 473 K under vacuum for 2 hours.
From the N2 adsorption isotherms, the specific surface area can be calculated. The
36
mesopore volume (VBJH), the average pore diameters (Dpore), and the pore size
distributions can be estimated by the Barret–Joyner– Halenda (BJH) method applied to
the desorption branch of the isotherm.
BET Surface Area
When a highly dispersed solid is exposed in a closed space to a gas or vapour at some
definite pressure, the solid begins to adsorb the gas. Adsorption on porous materials
proceeds via monolayer adsorption followed by multilayer adsorption and capillary
condensation. At a low relative pressure, a monolayer of adsorbent is absorbed on the
surface like non porous materials. The amount of adsorbent is used to calculate the
surface area. Using the Brunauer, Emmett and Teller (BET) theory, one can calculate
the specific surface area of a solid sample from the number of monolayer gas molecules
required to cover the solid surface, and the cross-sectional area of the gas molecule
being adsorbed.
BJH Pore-Size Distribution and Pore Volume
According to the IUPAC definitions, microporous, mesoporous and macroporous
materials exhibit pore diameters of less than 2 nm, in the range of 2-50 nm and above
50 nm, respectively [140]. During a physisorption process, beyond a monolayer
formation, the amount of adsorbed gas gradually increases as the relative pressure
increases, and then at higher pressures, the amount of gas adsorbed increases steeply
due to capillary condensation in the mesopores. Once these pores are filled, the
adsorption isotherm is complete. The capillary condensation is believed to be
proportional to the equilibrium gas pressure and the size of the pores inside the solid.
The Barrett, Joyner and Halenda (BJH) computational method allows the determination
of pore sizes from the equilibrium gas pressures.
Hysteresis Loops
After an adsorption isotherm, desorption isotherm can be generated by withdrawing the
adsorbed gas by reducing the pressure. However, capillary condensation and capillary
evaporations do not take place at the same pressure, resulting in hysteresis loops. The
resulting hysteresis loops can be mechanistically related to particular pore shapes.
37
Figure 1.25 shows the classification of hysteresis loops denoted as H1, H2, H3 and H4
by IUPAC. Type H1 and H2 loops were obtained from agglomerated spherical particles
and corpuscular systems, respectively, while type H3 and H4 were obtained from slitshaped pores or plate-like particles.
Figure 1.25 Hysteresis loops
Figure 1.26 BET surface area of Phosphated titania [141]
Figure 1.26 gives the BET surface area of phosphate titania which demonstrates that the
structure of pure titanium dioxide collapses after calcination at 973 K. In contrast,
hysteresis observed in the case of phosphated titanium dioxide samples is indicative of
porous structure. The specific surface area (as BET) increases with increasing
phosphate content at a given calcination temperature
38
1.12.8 Scanning electron microscopy (SEM)
Manfred von Ardenne pioneered the scanning electron microscope (SEM) and built his
universal electron microscope in the 1930s. In the SEM, a very fine beam of electrons
with energies up to several tens keV is focused on the surface of a specimen, and is
scanned across it in a parallel pattern. The intensity of emission of secondary and
backscattered electrons is very sensitive to the angle at which the electron beam strikes
the surface of the sample. The emitted electron current is collected and amplified. The
magnification produced by the SEM is the ratio between the dimension of the final
image display and the field scanning on the specimen. Usually, the magnification range
of the SEM is between 10 to 222,000 times, and the resolution is between 4-10 nm
(figure 1.27).
Figure 1.27: SEM images of Ni-doped TiO2 spheres: (a) side and (b) top views.
SEM images of Co-doped TiO2 spheres: (c) side and (d) top views [142].
Generally, the TEM resolution is about an order of magnitude greater than the SEM
resolution, however, because the SEM image relies on surface processes rather than
transmission, it is able to image bulk samples and has a much greater depth of view, and
so can produce images that are a good representation of the overall 3D structure of the
sample.
SEM images of the Fe-doped TiO2 agglomerates are presented in Fig. 1.27.
39
1.12.9 Transmission Electron Microscopy (TEM)
In 1931, Knoll and Ruska built the first electron microscope prototype. In 1938, Eli
Franklin Burton built the first practical electron microscope and in 1939, Siemens
produced the first commercial TEM. However, the TEM was not used for material
studies until 40 years ago, when the thin-foil preparation technique was developed.
More improvements of the TEM technique in the 1990s provided 0.1 nm of resolution,
which made TEM an indispensable analysis technique for studying materials in the
micron or nano size region.
TEM bright field images of TiO2 nanopowders in anatase phases are shown in Figure
(1.28a), respectively [127].
Figure 1.26: Images of anatase phase. (a) TEM image of micro-TiO2 powder; (b)
TEM image of nano-TiO2 powder; (c) SAED pattern of nano-TiO2 powder and (d)
HRTEM image of nano-TiO2 powder.
40
It can be estimated that the particle size of powders in Figure (1.28 b) are nanoscale
with the grain size less than 10 nm. The corresponding selected area electron diffraction
(SAED) patterns of nano- TiO2 powder anatase phases are shown in Figure (1.28c),
respectively. SAED patterns of nano-TiO2 powders in anatase phase (Figure (1.28c)
shows that the brightness and intensity of polymorphic ring is weak, so they are poorly
crystallized and partly amorphous. The crystallinity of nano-TiO2 powders can also be
observed by phase-contrast images or Moire patterns. Figure (1.28) show crystal lattice
planes of nano-TiO2 in rutile and anatase phases, respectively. It is seen that, anatase
give many crystal lattice planes with d-spacing of 0.313 nm for the plane (101).
1.12.10 Photoluminescence:
Photoluminescence (PL) emission spectra have been widely used to investigate the
efficiency of charge carrier trapping, migration, and transfer in order to understand the
fate of electron/hole pairs in semiconductor particles, since PL emission results from the
recombination of free carriers. Figure 1.29a [143] shows the room-temperature PL
spectra for C-doped and undoped TiO2 using the excitation light of 280 nm UV light. It
can also be seen that the PL intensity of the undoped TiO2 is much higher than that in
the spectra of C-doped samples, indicating that carbon doping can effectively inhibit
excited electron and hole recombination. The effective inhibition of excited electron and
hole recombination due to carbon doping can be explained by the following two
pathways. First, the surface hydroxyl groups can be trapped by the holes to form
hydroxyl radicals which can suppress electron-hole recombination, thus the PL
intensity. Second, theoretical and experimental investigation indicated that carbon
doping favours the formation of oxygen vacancies. The electron is trapped by the
oxygen vacancy, while the hole is trapped by the doped carbon, which can decrease the
PL intensity. However, no PL signal is observed for undoped TiO2, while a strong PL
peak is excited for C-doped TiO2 under the excitation of 514.5 nm visible light (Figure
1.29b). The energy of the excitation visible light used is not sufficient enough to
promote electronic transitions from the VB to conduction band (CB) for undoped TiO2
according to the UV-vis DRS and VB XPS. As a result, no electron/hole pairs can be
generated to give PL signals for undoped TiO2 under 514.5 nm light irradiation. As for
C-doped TiO2, electron/hole pairs can be generated and recombine radiatively to give
41
broad and strong PL signals. This can be ascribed to the carbon doping which reduces
the band gap and increases visible light absorption.
Figure 1.29: PL spectra of undoped TiO2 and C-doped TiO2 under UV-280 nm (a)
and visible light-514.5 nm (b).
1.12.11 EPR Study
This technique provides the information of oxidation state of the metal atom in the
mesopores and the position of the metal ion in the framework
1.12.12 Photocatalysis of dyes with titania:
The effluents, gaseous or liquid produced by some of our industries are harmful to the
health and general well-being of man. When undesirable substances are present in liquid
effluents, it can be disastrous as their presence pose severe threat to the immediate
recipients. Wastewaters from various industries, factories, laboratories, etc. are serious
problems to the environment. The discharged wastes containing dyes are toxic to
microorganisms, aquatic life and human beings [144]. These deleterious effects of
chemicals on the earth ecosystems are a cause for serious concern. Several of these
chemicals such as azo dyes, herbicides, and pesticides are actually present in rivers and
lakes, and are in part suspected of being endocrine-disrupting chemicals (EDCs) [145,
146]. Konstantinou and Albanis [147] reported that textile dyes and other industrial
dyestuffs constitute one of the largest groups of organic compounds that represent an
increasing environmental danger. About 1–20% of the total world production of dyes is
42
lost during the dying process and is released in the textile effluents [148]. Photocatalysis
may be termed as a photoinduced reaction which is accelerated by the presence of a
catalyst [37]. These types of reactions are activated by absorption of a photon with
sufficient energy (equals or higher than the band-gap energy (Eg) of the catalyst). The
absorption leads to a charge separation due to promotion of an electron (e−) from the
valence band of the semiconductor catalyst to the conduction band, thus generating a
hole (h+) in the valence band. The recombination of the electron and the hole must be
prevented as much as possible if a photocatalyzed reaction must be favoured. The
ultimate goal of the process is to have a reaction between the activated electrons with an
oxidant to produce a reduced product, and also a reaction between the generated holes
with a reductant to produce an oxidized product. The photogenerated electrons could
reduce the dye or react with electron acceptors such as O2 adsorbed on the Ti (III)surface or dissolved in water, reducing it to superoxide radical anion O2−• . The
photogenerated holes can oxidize the organic molecule to form R+, or react with OH− or
H2O oxidizing them into OH• radicals. Together with other highly oxidant species
(peroxide radicals) they are reported to be responsible for the heterogeneous TiO2
photodecomposition of organic substrates as dyes. The resulting •OH radical, being a
very strong oxidizing agent (standard redox potential +2.8 V) can oxidize most azo dyes
to the mineral end-products. According to this, the relevant reactions at the
semiconductor surface causing the degradation of dyes can be expressed as follows
[149]:
TiO2 +hv (UV) → TiO2 (eCB− +hVB+)
(1)
TiO2 (hVB+) + H2O → TiO2 +H+ +OH•
(2)
TiO2 (hVB+) + OH−→ TiO2 +OH•
(3)
−
TiO2 (eCB ) + O2→ TiO2 +O2
−•
−•
(4)
•
(5)
Dye + OH → degradation products
(6)
Dye + hVB+→ oxidation products
(7)
Dye + eCB−→ reduction products
(8)
O2 + H+→ HO2
•
where hv is photon energy required to excite the semiconductor electron from the
valence band (VB) region to conduction band (CB) region.
43
Table 1.5 Photocatalysis of different dyes with titania:
Dye degraded
Type of
Reference
Catalyst
no
Methylene Violet(MV), cationic red (X-GRL)
Ag–TiO2
150
Acid orange 7 (AO7)
Ag–TiO2
151
Sirius Gelb GC (SG-GC)
Ag–TiO2
152
0
Rhodamine B (RB)
Ag –TiO2
153
nanosol
Acid red B (ARB), reactive red (K-2G), cationic red (X-
Ag–AgBr–
GRL reactive brilliant red (X-3B )
TiO2
Methylene blue (MB)
Au–TiO2
155
Fast Green FCF, Patent Blue VF
TiO2
156
Orange G
Sn-TiO2
157
BRL
K-TiO2
158
Methyl orange
Pt-TiO2
159
Acid Red B
Ce-TiO2
160
Orange II
Zn-TiO2
161
Bromocresol purple
TiO2
162
methyl orange (MO), thymol blue (TB), and bromocresol M-TiO2 , (M =
green (BG)
154
163
Ni, Cu, Zn)
1.13. Specific aims of this dissertation
Since the discovery of photoinduced splitting of water on single crystal TiO2 electrodes,
the technology of semiconductor-based photocatalysis has shown potential application
in various areas such as sewage disposal, air purification, antibiosis and detoxification,
glass antifogging and self-cleaning, reclamation of precious metal ions, and
development of solar energy. TiO2 is a chemically stable, nontoxic, biocompatible,
inexpensive material with very high dielectric constant and interesting photocatalytic
activities. It is a wide-gap semiconductor, and depending on its chemical composition, it
shows a large range of electrical conductivity. However, the performance of TiO2
44
nanomaterials strongly relies on their crystallinity, crystallite size, crystal structure,
specific surface area, thermal stability and quantum efficiency. Bulk modifications such
as cation or anion doping have been found very effective to improve the properties of
TiO2 to enhance their performance. Several different processing techniques have been
used for the preparation of TiO2-based nanostructures, such as template techniques,
hydrothermal processes, and soft chemical processes. However, each of these methods
has limitations. The main objective of this project was to develop a synthesis route to
produce superior quality TiO2 based nanomaterials by introducing Zr, Fe, Co, N, S, and
P as doping materials.
The following have been identified as the specific objectives of this project:
I. Synthesis of the following nanomaterials using sol-gel technique and hydrothermal
technique.
A. Non-metal doped titania
(i)
S-doped TiO2
(ii)
P- doped TiO2
(iii)
N- doped TiO2
B. Metal doped titania
(i)
Fe- doped TiO2
(ii)
Zr- doped TiO2
(iii)
Co- doped TiO2
II. Characterization of the synthesized nanomaterials using different physical and
chemical techniques in order to,
a.
Study the surface characteristics, morphology, composition and electronic
properties.
b.
Investigate the effects of different operating variables and doping agents (ions)
on the properties.
III. Evaluation of the performance of selective nanomaterials as photocatalysts for the
degradation of different dyes as well as some pesticides.
45
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