+ TiO<sup>2</sup> (2)

concentra-

#

P-25 on

) on the pesticide reactants'

that is shown in **Figure 3**, when only a small amount of TiO<sup>2</sup>

neous reaction of photodegradation. However, at higher TiO<sup>2</sup>

be explained by the fact that suspended particles of TiO<sup>2</sup>

of the TiO<sup>2</sup>

increase in TiO<sup>2</sup>

tion (Eq. (2)):

where TiO<sup>2</sup>

compound [17, 26].

initial concentration (*Co*

studies [17].

concentration.

250 Titanium Dioxide - Material for a Sustainable Environment

TiO<sup>2</sup>

is the deactivated form of the catalyst [24].

According to several laboratory studies that have conducted, the use of Langmuir-Hinshelwood kinetics model and first-order rate equations provided reasonable simulations to the observed photocatalysis process of various organic pollutants over illuminated TiO<sup>2</sup> [5, 6, 13, 17]. Langmuir-Hinshelwood model is described by the following relationship:

**Figure 4.** Effect of initial concentration of selected OPPs on photocatalytic degradation rate (pesticide's concentration, 5–60 mg L−1; catalyst's concentration, 100 mg L−1; pH<sup>o</sup> , 6.06, λUV, 365 nm, in the presence of methanol).

$$r = -\frac{\mathbf{d} \cdot \mathbf{C}}{\mathbf{d}t} = \frac{k\_\gamma K \, \mathbf{C}}{1 + K \, \mathbf{C}} \tag{3}$$

employing different initial concentrations of the oxidant (not shown). This level of concen-

Photocatalytic Degradation of Selected Organophosphorus Pesticides Using Titanium Dioxide…

Blank experiments with the oxidant without the catalyst or UV light (not shown) proved a negligible reduction of OPPs studied. Experiments with UV light and the oxidant, without the catalyst, were performed too, are presented by the dotted lines of **Figure 5**, and showed that

added into the system. The kinetic parameters shown in **Table 3** demonstrated 18.20- and

is able to cause photooxidation of the pesticides tested. **Figure 5** illustrates the photo-

**(mg L−1 h−1) (×10−6 M h−1) (mg−1 L) (×10−6 M−1) (h−1) (h)**


system.

http://dx.doi.org/10.5772/intechopen.72193

(100 mg L−1) system, in the presence of

O2 was 253

, 6.06;

is in accordance with relative published studies [17].

A significant enhancement on degradation efficiency was observed when 5 mM of H<sup>2</sup>

**Pesticides** *R***<sup>2</sup>** *k<sup>r</sup> K kobs t1/2*

Azinphos methyl 0.9076 0.4569 1.4398 0.0894 0.2816 0.2293 3.02 Azinphos ethyl 0.8663 0.4574 1.3242 0.1320 0.3822 0.2378 2.91 Dimethoate 0.9612 0.0321 0.1402 0.0823 0.3592 0.0184 37.67 Disulfoton 0.9254 0.4553 1.6594 0.0789 0.2875 0.2587 2.68 Fenthion 0.9704 0.1289 0.4632 0.1203 0.4323 0.0535 12.96

**Figure 5.** Effect of the addition of hydrogen peroxide on the photocatalytic degradation of dimethoate and fenthion (pesticide's concentration, 10 mg L−1; catalyst's concentration, 100 mg L−1; oxidant's concentration, 5 mM; pH<sup>o</sup>

catalytic decomposition of fenthion and dimethoate in the UV-TiO<sup>2</sup>

**Table 2.** Photocatalytic kinetic parameters of selected pesticides in UV-TiO<sup>2</sup>

methanol, using a radiant UV energy of 14.5 mW cm−2.

λUV, 365 nm, in the presence of methanol).

tration of H<sup>2</sup>

H2 O2 O2

where *r* is the rate of reaction (in mg L−1 h−1), *C* is the concentration at any time t during degradation (in mg L−1), *t* is the irradiation time (in h), *kr* is the limiting rate constant of reaction at maximum coverage under the employed experimental conditions (in mg L−1 h−1), and finally, *K* is the equilibrium constant for adsorption of the substrate onto catalyst TiO<sup>2</sup> particles (in mg−1 L) [6]. The linear transformation of Eq. 3 is Eq. (4), which is used in the bibliography for demonstrating the linearity of data when plotted as the inverse rate versus inverse concentration (as in the inset of **Figure 4**, previously presented):

$$
\begin{pmatrix} \frac{1}{\mathcal{T}} \end{pmatrix} = \begin{pmatrix} \frac{1}{k\_r} \end{pmatrix} + \begin{pmatrix} \frac{1}{k\_r K} \end{pmatrix} \begin{pmatrix} \frac{1}{C} \end{pmatrix} \tag{4}
$$

Integration of Eq. (3) yields Eq. (5):

$$
\ln\left(\frac{\mathbb{C}\_s}{\mathbb{C}}\right) + K\left(\mathbb{C}\_o - \mathbb{C}\right) = k\_r K \, t \tag{5}
$$

When the initial concentration of the pesticide *Co* is very small (in millimolar solution as in the case), Eq. (5) is transformed into Eqs. (6) and (7):

$$\ln\left(\frac{C\_s}{C}\right) = k\_r K t = K\_{obs} t \tag{6}$$

$$\mathbf{C}\_t = \mathbf{C}\_o e^{\mathbf{x}\_{ho}^\prime t} \tag{7}$$

where *Ct* represents the concentration at time *t*, *Co* represents the initial concentration, and *Kobs* is the rate constant. In all cases, reduction process followed an apparent first-order rate reaction, and the calculations were performed using Eq. (7). The kinetic parameters calculated by Eqs. (4) and (7) are shown in **Table 2**. In this point, it must be mentioned that these constants reflect to the experimental conditions which have been used and are only useful for comparison between reactants that have been oxidized using the same catalyst and illumination source.

#### **3.3. Photocatalytic degradation of OPPs in UV-TiO2 -H2 O2 system**

The addition of an oxidant into a semiconductor suspension has been proven to enhance the photodegradation of a variety of organic pollutants among which pesticides are included [8, 17, 27, 28]. In current study, the addition of hydrogen peroxide (H<sup>2</sup> O2 ) was evaluated only for the cases of dimethoate and fenthion as these compounds proved to be more resistant in the previous set of photocatalytic experiments. The initial concentration (5 mM) of H2 O2 was chosen according to the results obtained in previously conducted experiments employing different initial concentrations of the oxidant (not shown). This level of concentration of H<sup>2</sup> O2 is in accordance with relative published studies [17].

*r* = −\_\_\_\_\_\_\_\_\_\_\_\_\_\_ <sup>d</sup> *<sup>C</sup>*

252 Titanium Dioxide - Material for a Sustainable Environment

radation (in mg L−1), *t* is the irradiation time (in h), *kr*

tion (as in the inset of **Figure 4**, previously presented):

(\_\_<sup>1</sup>

Integration of Eq. (3) yields Eq. (5):

ln(

*l*n(

where *Ct*

tion source.

H2 O2

case), Eq. (5) is transformed into Eqs. (6) and (7):

represents the concentration at time *t*, *Co*

**3.3. Photocatalytic degradation of OPPs in UV-TiO2**

<sup>d</sup>*<sup>t</sup>* <sup>=</sup> *kr <sup>K</sup> <sup>C</sup>* \_\_\_\_\_\_\_\_\_\_

where *r* is the rate of reaction (in mg L−1 h−1), *C* is the concentration at any time t during deg-

maximum coverage under the employed experimental conditions (in mg L−1 h−1), and finally,

mg−1 L) [6]. The linear transformation of Eq. 3 is Eq. (4), which is used in the bibliography for demonstrating the linearity of data when plotted as the inverse rate versus inverse concentra-

When the initial concentration of the pesticide *Co* is very small (in millimolar solution as in the

*Ct* = *Co e*‐*Kobs <sup>t</sup>* (7)

*Kobs* is the rate constant. In all cases, reduction process followed an apparent first-order rate reaction, and the calculations were performed using Eq. (7). The kinetic parameters calculated by Eqs. (4) and (7) are shown in **Table 2**. In this point, it must be mentioned that these constants reflect to the experimental conditions which have been used and are only useful for comparison between reactants that have been oxidized using the same catalyst and illumina-

The addition of an oxidant into a semiconductor suspension has been proven to enhance the photodegradation of a variety of organic pollutants among which pesticides are included

only for the cases of dimethoate and fenthion as these compounds proved to be more resistant in the previous set of photocatalytic experiments. The initial concentration (5 mM) of

was chosen according to the results obtained in previously conducted experiments

[8, 17, 27, 28]. In current study, the addition of hydrogen peroxide (H<sup>2</sup>

**-H2 O2**

 **system**

\_\_\_1 *kr <sup>K</sup>*) ( \_\_1

*K* is the equilibrium constant for adsorption of the substrate onto catalyst TiO<sup>2</sup>

*<sup>r</sup>*) <sup>=</sup> ( \_\_1 *kr*) <sup>+</sup> (

*C*\_\_\_*o*

*C*\_\_\_*o*

<sup>1</sup> <sup>+</sup> *<sup>K</sup> <sup>C</sup>* (3)

is the limiting rate constant of reaction at

*<sup>C</sup>*) (4)

*<sup>C</sup>*) + *K* (*Co* − *C*) = *kr K t* (5)

*<sup>C</sup>*) = *kr K t* = *Kobs t* (6)

represents the initial concentration, and

O2

) was evaluated

particles (in

Blank experiments with the oxidant without the catalyst or UV light (not shown) proved a negligible reduction of OPPs studied. Experiments with UV light and the oxidant, without the catalyst, were performed too, are presented by the dotted lines of **Figure 5**, and showed that H2 O2 is able to cause photooxidation of the pesticides tested. **Figure 5** illustrates the photocatalytic decomposition of fenthion and dimethoate in the UV-TiO<sup>2</sup> -H<sup>2</sup> O2 system.

A significant enhancement on degradation efficiency was observed when 5 mM of H<sup>2</sup> O2 was added into the system. The kinetic parameters shown in **Table 3** demonstrated 18.20- and


**Table 2.** Photocatalytic kinetic parameters of selected pesticides in UV-TiO<sup>2</sup> (100 mg L−1) system, in the presence of methanol, using a radiant UV energy of 14.5 mW cm−2.

**Figure 5.** Effect of the addition of hydrogen peroxide on the photocatalytic degradation of dimethoate and fenthion (pesticide's concentration, 10 mg L−1; catalyst's concentration, 100 mg L−1; oxidant's concentration, 5 mM; pH<sup>o</sup> , 6.06; λUV, 365 nm, in the presence of methanol).


**Table 3.** Photocatalytic kinetic parameters of selected pesticides in UV-TiO<sup>2</sup> (100 mg L−1)-Η<sup>2</sup> Ο2 (5 mM) system, in the presence of methanol, using a radiant UV energy of 14.5 mW cm−2.

4.04-fold increase in rate constants, kobs, compared to the results obtained in UV-TiO<sup>2</sup> systems for fenthion and dimethoate, respectively.

The magnitude of enhancement is in agreement with other studies concerning phosphoruscontaining pesticides [27, 28]. This observation is explained by the fact that H<sup>2</sup> O2 can act as an alternative electron acceptor to oxygen (Eq. (8)) that is a thermodynamically more favorable reaction than oxygen reduction [29]. This should consequently promote the charge separation and accelerate the heterogeneous photocatalysis [30]. At the same time, hydroxyl radicals able to oxidize organic pollutants, such as pesticides, are generated by either the reduction of H<sup>2</sup> O2 at the conductance band [29] or the acceptance of an electron from superoxide again (chain reactions) (Eq. (9)) [31]. As a consequence, and regardless which conductance band reaction overrules, additional •OH oxidizing species may be produced resulting in the increase of the oxidizing power of the system [30]:

$$\text{e}\_{\text{co}}\text{–}+\text{H}\_{2}\text{O}\_{2}\rightarrow\text{OH}^{-}+\text{OH}^{\*}\tag{8}$$

*3.4.1. TOC content*

As shown in **Figure 6**, in the presence of TiO<sup>2</sup>

**Figure 6.** Pesticide and TOC reduction and evolution of sulfate (SO4

and ammonium (NH4

+

time (pesticide's concentration level, 10 mg L−1; TiO<sup>2</sup>

and under UV irradiation, the TOC concentra-

http://dx.doi.org/10.5772/intechopen.72193

255

2−), phosphate (PO4

) ions originating from photocatalytic degradation of selected pesticides as a function of irradiation

O2

, 100 mg L−1; H<sup>2</sup>

3−), nitrite (NO<sup>2</sup>

, 5 mM; λUV, 365 nm, in the presence of methanol).

−

), nitrate (NO3

− ),

tion of the insecticides exhibited a constant decrease with time, reaching TOC reduction that ranged between 58 and 100% (total mineralization) after the end of illumination time. In addition, it can be seen that the rates of TOC reduction during irradiation of solutions of azinphos methyl, azinphos ethyl, and disulfoton were higher compared to those of dimethoate and fenthion, suggesting that former ones exhibit greater susceptibility to photocatalysis than the

Photocatalytic Degradation of Selected Organophosphorus Pesticides Using Titanium Dioxide…

$$\rm O\_2^- + H\_2O\_2 \rightarrow OH^- + OH^\* + O\_2 \tag{9}$$

Experimental data obtained proved that the addition of H<sup>2</sup> O2 in the aqueous suspensions of 100 mg L−1 TiO<sup>2</sup> /10 mg L−1 OPPs was more effective on the photocatalytic oxidation of fenthion (100% decomposition in 5 h) rather than in the case of dimethoate (95% decomposition in 36 h). The *t 1/2* in the presence of both TiO<sup>2</sup> and H<sup>2</sup> O2 was 0.71 h (42.6 min) for fenthion and 9.32 h for dimethoate.

#### **3.4. Mineralization studies**

The general stoichiometric reaction proposed for the photocatalytic oxidation of the studied pesticides that leads to the conversion of all of their carbon atoms to gaseous CO<sup>2</sup> and the heteroatoms into inorganic anions at their highest oxidation states which remain in the solution is described by the following reaction (Eq. (10)). Consequently, in order to assess the extent of mineralization during photocatalysis of selected organic pollutants, TOC measurements were carried out along with determination of released inorganic anions containing the heteroatoms of the selected organics (**Figure 6**):

$$\begin{aligned} \mathbf{C}\_{\times} \mathbf{H}\_{\psi} \mathbf{N}\_{\theta} \mathbf{O}\_{\omega} \mathbf{P}\_{z} \mathbf{S}\_{\phi} + \left[ (2\mathbf{x} + 3\boldsymbol{\Theta} + 4\mathbf{z} + 4\boldsymbol{\phi} + 1 - \boldsymbol{\omega})/2 \right] \mathbf{O}\_{z} \boldsymbol{\Theta} & \quad \mathbf{x} \mathbf{CO}\_{z} + \boldsymbol{\theta} \mathbf{NO}\_{3}^{-} \\ &+ \mathbf{z} \mathbf{PO}\_{4}{\cdot}^{3-} + \boldsymbol{\phi} \mathbf{SO}\_{4}{\cdot}^{2-} + \mathbf{H}\_{z} \mathbf{O} + (\boldsymbol{\psi} - 2) \mathbf{H}^{\*} \end{aligned} \tag{10}$$
