**2. TiO2**

These species are very effective oxidant agents, and they are capable of degrading recalcitrant compounds due to the high potential of oxidation. These ROS are produced on the

promising material because it is inexpensive, nontoxic, chemically and thermally stable, abundant, and environmentally friendly. However, this semiconductor only will produce these species when it receives a minimum energy amount, called the bandgap energy (*Eg*

which is capable to remove an electron from the valence band (VB) and transfer it to the

is supplied by photons with frequency in the ultraviolet light region, which possess wave-

Despite the ability of photocatalytic processes to degrade several compounds by the use of hydroxyl radicals, their use is still not widespread, with scarce industrial applications. This is due to some inherent features of the catalysts employed, such as the recombination of photogenerated charges, which reduces the formation of radicals and, consequently, the efficiency of photocatalytic degradation. Therefore, plenty of studies have been performed to overcome this drawback, like surface modifications that allow the capture of the generated electrons and avoid the recombination. Furthermore, the bandgap energy required for the formation of the

source that has wavelengths in the ultraviolet region (≈390 nm). Therefore, only 5% of sunlight can be used in photocatalysis that applies this semiconductor, which makes the process more expensive. In this way, superficial modification through metal and nonmetal doping is funda-

Among the possible dopants, rare earths have been investigated for their ability to increase photocatalytic activity, possibly by reducing bandgap energy due to the introduction of orbitals between the conduction and valence bands, generating impurity energy levels in the semiconductor elements. These states are generated from the 4f level, which are electron deficient. Another hypothesis for this increase in contaminant degradability is that the adsorption of these lanthanides on the surface of the semiconductors generates an imbalance of surface charges, which can produce surface defects and vacancies of oxygen and titanium. These two propositions lead to states that serve as electron scavengers and reduce the recombination of photogenerated charges, increasing the probability of •OH formation. Moreover, the lanthanides adsorbed on the surface may act as traps to water molecule and –OH (hydroxyl anion) groups, which increase their density on the photocatalyst surface and can promote more intense formation of hydroxyl radicals. Another important feature related to the use of these compounds is that they serve as a Lewis base, which could concentrate the contaminants dispersed in aqueous medium on the semiconductor surface enhancing the electron transfer for the direct degradation of the contaminant or increasing the probability of interac-

been the same, with two ways most discussed. Rare earths can be included in the TiO2

tice by direct linking or substitution producing a≡Ti–O–Ln–O–Ti≡arrangement, which cause distortions/defects in the lattice due to the mismatch of ionic radius of Ln3+ and Ti4+. On the

), which is considered the most

is equal to 3.2 eV, restricts this catalyst to the use of a light

proposed by different researchers have not

lat-

, this minimum energy

),

semiconductors' surface, such as titanium dioxide (TiO2

lengths below 400 nm [1–6].

82 Photocatalysts - Applications and Attributes

electron/vacancy pair, which for TiO2

mental to overcome these drawbacks [7–10].

tion between the molecules and the radicals formed [11–15].

The mechanisms of lanthanides doping on TiO2

conduction band (CB), thus creating the electron-hole pair. For TiO2

The crystalline structures of titanium dioxide are anatase, rutile, and brookite, which is difficult to be synthesized in laboratory. Therefore, the first two lattices are the most prominent in researches. All crystal lattices are composed of TiO<sup>6</sup> octahedra, but with distinct connections between these structures, since anatase has four connections between the edges, brookite has three connections, and rutile has only two. This structure confers to rutile a greater thermodynamic stability among the other polymorphs, as established by the third Pauling rule [31–33].

There is a great interest in the application of titanium dioxide as a photocatalyst because this compound is nontoxic, economical, chemically inert, and photostable to corrosion, besides high thermal stability, intense photocatalytic activity, and strong oxidant power. However, not all crystalline structures have the same efficiency in the absorption of light for catalysis; for example, although rutile is the polymorph thermodynamically more stable, its photocatalytic activity is lower than anatase, possibly due to the high temperature required for its preparation, which causes an increase in particle size, a high rate of electron/vacancy recombination that reduces the number of hydroxyl groups on the surface, and lower electron mobility in relation to anatase [34–37].

However, the rutile phase has a bandgap equal to 3.0 eV, having the ability to absorb radiation in the visible light spectrum, while the anatase has a bandgap of approximately 3.2 eV. Therefore, the interaction of these two crystalline structures often has better photocatalytic results than when they are purely applied, which is possibly due to the visible light absorption by the rutile, which serves as a photosensitizer for the anatase phase. An example of this phase interaction is the commercial P25® titanium dioxide from Degussa Evonik, which has a composition of 80% anatase and 20% rutile, exhibiting high photocatalytic activity [32, 38, 39].

**3. Lanthanides**

4fn6s2

[52, 53].

or [Xe] 4fn5d1

3+ (Ln3+) [47–51].

6s2

properties do not undergo sudden changes [48, 54, 55].

**4. Surface modification by lanthanides**

used, because these incorporate levels 3d in the TiO2

The rare earths (RE) are a group of elements with physicochemical properties that are very similar, which is constituted by lanthanides (lanthanum to lutetium), scandium, and yttrium, totalizing 17 elements of the periodic table. This series can be separated into light rare earths that comprise low atomic mass elements (lanthanum to europium) and heavy rare earths,

These elements have shown great potential for industrial applications because of their unique magnetic, optical, and/or redox properties, with important uses in the field of catalysis, high temperature superconductors, hybrid cars, permanent magnets, nuclear magnetic resonance, rechargeable batteries, manufacture of glass and ceramic materials, shift reagents, etc. The electronic configuration of these chemical elements assure these properties due to the distribution of the electrons with the complete levels until 5, such as xenon, but with the level *f*, it was incomplete and protected by orbitals 6 *s* and/or 5*d*, except for of the scandium and yttrium. Thus, the representation of the electronic configuration is described as [Xe]

The orbitals 4*f* that were partially completed are responsible for the optical properties of rare earths, since different arrangements of these orbitals generate different levels of energy, allowing the absorption in a broad spectrum of light radiation, which can vary from ultraviolet to visible. Moreover, the luminescence of some rare earth ions arises from the *f-f* electronic transitions within their partially filled 4*f* orbitals that can occur by electric or magnetic dipole

Another advantage of this configuration is that these orbitals are sterically shielded from the surrounding microenvironment by the filled 5s and 5p orbitals, meaning that there are almost no perturbations of these transitions by other bonding elements, which assure that the optical

The surface modification of semiconductors that occurs by the addition of tiny and controlled impurities is denominated doping, and this process is applied to achieve distinct electronic properties from those raw materials. Metal doping has a main feature, the bandgap reduction, which optimizes the ability of a semiconductor to absorb wavelengths with less energy, reduce recombination of the vacancy/electron pairs, and modify the adsorptive capacity of the surface of this photocatalyst. Bandgap reduction occurs possibly due to the formation of energy levels between the conduction and valence bands, such as when transition metals are

electronic levels occupied near to the conduction band. However, visible light absorption by

for lanthanides (Ln), which usually have an oxidation state equal to

lattice, which leads to the formation of

Lanthanides Effects on TiO2 Photocatalysts http://dx.doi.org/10.5772/intechopen.80906 85

which are constituted from lutetium to gadolinium, besides yttrium [19, 46].

Regarding to the electronic structure, this catalyst can be considered a semiconductor with indirect bandgap for anatase phase, meaning that the valence and conduction bands of this lattice do not have the same maximum and minimum momentum, respectively. This condition leads to a recombination of the electron/vacancy pair that only occurs if assisted by a phonon, resulting in a longer lifetime of this electron generated. Furthermore, TiO2 is an n-type semiconductor; in other words it has a greater density of electrons than vacancies produced by oxygen, because they are compensated by the presence of Ti3+, which brings the Fermi level closer to the conduction band [40–42].

Bandgap energy for the structure with the highest photocatalytic activity is approximately 3.2 eV, and the valence band, formed mainly by the 2p orbital of the oxygen, has an energy level of approximately 2.60 V, which is more positive than the oxidation potential of water to hydroxyl radicals (*E0* = 2.27 V). In the case of the conduction band, which is constituted mainly by titanium 3d orbitals, the energy level is equal to −0.51 V, with a value that is negatively greater than the oxygen reduction potential of the superoxide radical (*E0* = −0.33 V). As depicted in **Figure 1**, with these conditions titanium dioxide can degrade organic compounds through activation by UV light, which makes it a prominent catalyst [13, 32, 43–45].

**Figure 1.** Redox potential of conduction and valence bands and the radical formation.
