Any reaction or process that enhances the rate of conversion of Fe(III) back to Fe(II) will in principle accelerate the rate of the Fenton reaction. In the electro-Fenton reaction, this is accomplished electrochemically. Chemically, the Fenton reaction can be efficiently catalyzed by certain types of organic molecules, especially benzoquinones or dihydroxybenzene (DHB) derivatives. The catalytic influence of DHBs on the Fenton reaction was originally reported by Hamilton and coworkers (Hamilton et al., 1966a, 1966b). Because DHBs such as catechol or hydroquinone and their analogs are common initial intermediates in the degradation of aromatic molecules, their presence in the reaction medium can result in efficient DHB-catalyzed redox cycling of Fe3+ back o Fe2+ (Chen & Pignatello, 1997; Nogueira et al., 2005). Indeed, in several cases, it has been shown that the addition of catechol or catechol derivatives can enhance the rate and the overall mineralization efficiency of Fenton reactions (Aguiar et al., 2007; Zanta et al., 2010). In addition, DHB-catalyzed redox cycling of iron may be important in fungal degradation of lignin and several possible mechanisms have been suggested (Aguiar et al., 2007). The third method of accelerating the Fenton reaction is via irradiation with ultraviolet light, generally known as the photo-assisted Fenton or photo-Fenton reaction, which will be the primary focus of the remainder of this chapter.

#### **4. The Photo-Fenton reaction**

#### **4.1 General features of the Photo-Fenton reaction**

About two decades ago, it was found that the irradiation of Fenton reaction systems with UV/Visible light strongly accelerated the rate of degradation of a variety of pollutants (Huston & Pignatello, 1999; Ruppert et al., 1993). This behavior upon irradiation is due principally to the photochemical reduction of Fe(III) back to Fe(II), for which the net reaction can be written as:

$$\mathrm{Fe^{3+}\_{aq} + H\_2O \xrightarrow{hv} Fe^{2+}\_{aq} + HO^{\bullet} + H^{\bullet}} \tag{4}$$

More detailed studies of the pH dependence of the photo-Fenton reaction have shown that the optimum pH range is ca. pH 3. The reason for this pH dependence becomes clear when one examines the speciation of Fe(III) as a function of pH (Figure 1) and the absorption spectra of the relevant Fe(III) species (Figure 2). At pH < 2, the dominant species is Organic Pollutants Ten Years After 274 the Stockholm Convention – Environmental and Analytical Update

formation of oxalic acid, H2C2O4, a poisonous and persistent oxidation product of many degradation reactions. Since Fenton reactions are typically run at an initial pH of about 3 and oxalic acid is a relatively strong acid (with a first pKa of 1.4), the accumulation of oxalic acid causes further acidification of the reaction mixture (to ca. pH 2) as the reaction proceeds. In addition, Fe(III) is very efficiently chelated by the oxalate anion. This prevents the reduction of Fe(III) back to Fe(II) and hence the complete mineralization of the organic matter. A still unresolved question in the mechanism of the Fenton degradation of organic material is the relative importance of other potential oxidizing species besides the hydroxyl radical. The stoichiometry of Fenton degradation reactions is complex and, in addition to Fe2+/Fe3+ and hydrogen peroxide, can involve the participation of the hydroperoxyl radical,

, iron(IV) or ferryl, FeO2+, dissolved molecular oxygen, organic hydroperoxides, and other intermediates formed during the degradation. Despite these potential limitations, the conventional Fenton reaction has been widely used for the treatment of effluents (Benitez et

Any reaction or process that enhances the rate of conversion of Fe(III) back to Fe(II) will in principle accelerate the rate of the Fenton reaction. In the electro-Fenton reaction, this is accomplished electrochemically. Chemically, the Fenton reaction can be efficiently catalyzed by certain types of organic molecules, especially benzoquinones or dihydroxybenzene (DHB) derivatives. The catalytic influence of DHBs on the Fenton reaction was originally reported by Hamilton and coworkers (Hamilton et al., 1966a, 1966b). Because DHBs such as catechol or hydroquinone and their analogs are common initial intermediates in the degradation of aromatic molecules, their presence in the reaction medium can result in efficient DHB-catalyzed redox cycling of Fe3+ back o Fe2+ (Chen & Pignatello, 1997; Nogueira et al., 2005). Indeed, in several cases, it has been shown that the addition of catechol or catechol derivatives can enhance the rate and the overall mineralization efficiency of Fenton reactions (Aguiar et al., 2007; Zanta et al., 2010). In addition, DHB-catalyzed redox cycling of iron may be important in fungal degradation of lignin and several possible mechanisms have been suggested (Aguiar et al., 2007). The third method of accelerating the Fenton reaction is via irradiation with ultraviolet light, generally known as the photo-assisted Fenton or photo-Fenton reaction, which will be the primary focus of the remainder of this

About two decades ago, it was found that the irradiation of Fenton reaction systems with UV/Visible light strongly accelerated the rate of degradation of a variety of pollutants (Huston & Pignatello, 1999; Ruppert et al., 1993). This behavior upon irradiation is due principally to the photochemical reduction of Fe(III) back to Fe(II), for which the net reaction

⎯→*v* Fe2+

More detailed studies of the pH dependence of the photo-Fenton reaction have shown that the optimum pH range is ca. pH 3. The reason for this pH dependence becomes clear when one examines the speciation of Fe(III) as a function of pH (Figure 1) and the absorption spectra of the relevant Fe(III) species (Figure 2). At pH < 2, the dominant species is

aq + HO•

+ H+ (4)

aq + H2O ⎯h

HOO•

chapter.

can be written as:

al., 1999; Pignatello et al., 2006).

**4. The Photo-Fenton reaction**

Fe3+

**4.1 General features of the Photo-Fenton reaction**

hexaquoiron(III), Fe(H2O)63+ [or simply Fe3+ for convenience], which absorbs weakly in the ultraviolet above 300 nm. At pH > 3, freshly prepared solutions of Fe(III) are supersaturated with respect to formation of colloidal iron hydroxide, Fe(OH)3 and prone to precipitation of hydrated iron oxides upon standing for a prolonged period. At pH 3, however, the predominant Fe(III) species present in aqueous solution is Fe(H2O)5(OH)2+ [or simply Fe(OH)2+], which absorbs throughout much of the ultraviolet spectral region (Martyanov et al., 1997).

Fig. 1. Speciation of 0.5 mM Fe(III) between pH 1 and 4 at 25○C and an ionic strength of 0.1. Mole fractions of each species were calculated with the public domain program Hydra (Puigdomenech, 2010), employing the equilibrium constants for complexation supplied with the program (excluding insoluble iron species).

Studies of the photochemistry of Fe(OH)2+ have shown (Pozdnyakov et al., 2000) that Fe(OH)2+ undergoes a relatively efficient photoreaction to produce Fe(II) and the hydroxyl radical:

$$\text{Fe(OH)2\*} \xrightarrow{\text{hv}} \text{Fe}^{2\*} + \text{HO}^{\bullet} \tag{5}$$

Thus, irradiation of the Fenton reaction not only regenerates Fe(II), the crucial catalytic species in the Fenton reaction (reaction 2), but also produces an additional hydroxyl radical, the species responsible for the degradation of organic material. As a consequence of these two effects, the photo-Fenton process is faster than the conventional thermal Fenton process. Moreover, since Fe(II) is regenerated by light with decomposition of water (equations 4-5) rather than H2O2 (reaction 3), the photo-Fenton process consumes less H2O2 and requires only catalytic amounts of Fe(II).

The photo-Fenton reaction has several operational and environmental advantages. The classes of organic compounds that are susceptible to photodegradation via the Fenton reaction are rather well known (Bigda, 1995). The photo-Fenton process produces no new

Fundamental Mechanistic Studies of

al., 2007).

of the less-reactive Cl2

the Photo-Fenton Reaction for the Degradation of Organic Pollutants 277

Fig. 3. Degradation of 12 mM phenol in the presence of 200 mM H2O2 and 0.5 mM Fe2+. Fenton oxidation at pHinitial 3.0 in the presence (curve a, ■) and absence (curve b, ) of 0.5 M NaCl. Photo-Fenton degradation (400 W medium pressure Hg vapor lamp – incident photon intensities were 1.2 x 10-4 Ein/s) at pHinitial 3.0 with no pH control in the absence (curve e, ▲) or presence (curve c, ●) of 0.5 M NaCl; or with pH 3.0 maintained throughout the reaction in the absence (curve f, ∆) or presence of 0.5 M NaCl (curve d, ○). For details, see (Machulek et

In the presence of chloride ion, the extent of mineralization in the photo-Fenton reaction (Figure 3, curve c) is comparable to that observed in the analogous thermal Fenton reaction carried out in the dark in the presence or absence of chloride ion (Figure 3, curves a and b). Although we (Moraes et al., 2004a, 2004b) and others (De Laat & Le, 2006; Kiwi et al., 2000; Maciel et al., 2004; Pignatello, 1992) had ascribed this inhibition to the preferential formation

of photo-Fenton reactions required a fuller understanding of the mechanistic details of the inhibition. For this purpose, we (Machulek et al., 2006) used nanosecond laser flash photolysis to investigate the influence of added chloride ion on the photocatalytic step that converts Fe(III) back to Fe(II) (equation 5), deliberately omitting H2O2 from the reaction mixture to prevent the thermal Fenton reaction. Although direct spectroscopic detection of

absorbs at 340 nm (**ε**340nm *=* 8000 M-1cm-1), can be readily detected upon excitation of aqueous solutions of iron(III) at acidic pH in the presence of added sodium chloride at 355 nm with the third harmonic of a Nd-YAG laser. Differential absorption spectra (Figure 4) and kinetic

of the laser pulse (5 ns), and its subsequent decay via mixed first and second order kinetics. The net decrease in absorption at longer times, relative to that prior to the laser pulse, reflects the conversion of Fe(III) to Fe(II), which does not absorb in this spectral region.

the hydroxyl radical has proved elusive (Marin et al., 2011), the Cl2

traces (insert, Figure 4) showed fast formation of the Cl2

•- radical anion instead of the desired hydroxyl radical, optimization

•**-** radical anion, which

•- radical anion, within the lifetime

pollutants and requires only small quantities of iron salt. At the end of the reaction, if necessary, the residual Fe(III) can be precipitated as iron hydroxide by increasing the pH. Any residual hydrogen peroxide that is not consumed in the process will spontaneously decompose into water and molecular oxygen and is thus a "clean" reagent in itself. These features make homogeneous photo-Fenton based AOPs the leading candidate for costefficient, environmental friendly treatment of industrial effluents on a small to moderate scale (Pignatello et al., 2006). An early example of an industrial-scale application of the photo-Fenton process was the decontamination of 500 L batches of an industrial effluent containing 2,4-dimethylaniline in a photochemical reactor fitted with a 10 kW medium pressure mercury lamp (Oliveros et al., 1997).

(a) Solid curve: pH 1.5 with perchloric acid, 90% Fe(H2O)6 3+ or Fe3+, see Figure 1; (b) Dashed Curve: pH 2.5 with perchloric acid, ca. 50:50 FeOH2+:Fe3+, see Figure 1; (c) Dotted Curve: pH 1.5 plus added 500 mmol L-1 NaCl, 30% FeCl2+ and 65% FeCl2+.

Fig. 2. Absorption spectra of 0.43 mM Fe(III) perchlorate under three typical conditions, expressed as the apparent extinction coefficient, ε, defined as the observed absorbance divided by the total Fe(III) concentration.

#### **4.2 Inhibition of the Photo-Fenton reaction by chloride ion**

Our interest in the fundamental mechanistic aspects of the photo-Fenton reaction arose from the necessity to optimize the degradation of organic material in effluents or wastewaters containing high concentrations of chloride ion. With such effluents, such as petroleum wastewaters from offshore marine environments or residues from organochloride pesticide production, the photo-Fenton reaction is strongly inhibited. Thus, although the photo-Fenton reaction results in essentially complete mineralization of phenol in the absence of chloride ion (Figure 3, curve e), in the presence of the chloride ion the mineralization process (indicated by the reduction in the amount of total organic carbon, TOC) stops after only partial decomposition of the organic material (Figure 3, curve c).

Organic Pollutants Ten Years After 276 the Stockholm Convention – Environmental and Analytical Update

pollutants and requires only small quantities of iron salt. At the end of the reaction, if necessary, the residual Fe(III) can be precipitated as iron hydroxide by increasing the pH. Any residual hydrogen peroxide that is not consumed in the process will spontaneously decompose into water and molecular oxygen and is thus a "clean" reagent in itself. These features make homogeneous photo-Fenton based AOPs the leading candidate for costefficient, environmental friendly treatment of industrial effluents on a small to moderate scale (Pignatello et al., 2006). An early example of an industrial-scale application of the photo-Fenton process was the decontamination of 500 L batches of an industrial effluent containing 2,4-dimethylaniline in a photochemical reactor fitted with a 10 kW medium

pressure mercury lamp (Oliveros et al., 1997).

divided by the total Fe(III) concentration.

(a) Solid curve: pH 1.5 with perchloric acid, 90% Fe(H2O)6 3+ or Fe3+, see Figure 1; (b) Dashed Curve: pH 2.5 with perchloric acid, ca. 50:50 FeOH2+:Fe3+, see Figure 1; (c) Dotted Curve: pH 1.5 plus added 500 mmol L-1 NaCl, 30% FeCl2+ and 65% FeCl2+.

**4.2 Inhibition of the Photo-Fenton reaction by chloride ion**

partial decomposition of the organic material (Figure 3, curve c).

Fig. 2. Absorption spectra of 0.43 mM Fe(III) perchlorate under three typical conditions, expressed as the apparent extinction coefficient, ε, defined as the observed absorbance

Our interest in the fundamental mechanistic aspects of the photo-Fenton reaction arose from the necessity to optimize the degradation of organic material in effluents or wastewaters containing high concentrations of chloride ion. With such effluents, such as petroleum wastewaters from offshore marine environments or residues from organochloride pesticide production, the photo-Fenton reaction is strongly inhibited. Thus, although the photo-Fenton reaction results in essentially complete mineralization of phenol in the absence of chloride ion (Figure 3, curve e), in the presence of the chloride ion the mineralization process (indicated by the reduction in the amount of total organic carbon, TOC) stops after only

Fig. 3. Degradation of 12 mM phenol in the presence of 200 mM H2O2 and 0.5 mM Fe2+. Fenton oxidation at pHinitial 3.0 in the presence (curve a, ■) and absence (curve b, ) of 0.5 M NaCl. Photo-Fenton degradation (400 W medium pressure Hg vapor lamp – incident photon intensities were 1.2 x 10-4 Ein/s) at pHinitial 3.0 with no pH control in the absence (curve e, ▲) or presence (curve c, ●) of 0.5 M NaCl; or with pH 3.0 maintained throughout the reaction in the absence (curve f, ∆) or presence of 0.5 M NaCl (curve d, ○). For details, see (Machulek et al., 2007).

In the presence of chloride ion, the extent of mineralization in the photo-Fenton reaction (Figure 3, curve c) is comparable to that observed in the analogous thermal Fenton reaction carried out in the dark in the presence or absence of chloride ion (Figure 3, curves a and b). Although we (Moraes et al., 2004a, 2004b) and others (De Laat & Le, 2006; Kiwi et al., 2000; Maciel et al., 2004; Pignatello, 1992) had ascribed this inhibition to the preferential formation of the less-reactive Cl2 •- radical anion instead of the desired hydroxyl radical, optimization of photo-Fenton reactions required a fuller understanding of the mechanistic details of the inhibition. For this purpose, we (Machulek et al., 2006) used nanosecond laser flash photolysis to investigate the influence of added chloride ion on the photocatalytic step that converts Fe(III) back to Fe(II) (equation 5), deliberately omitting H2O2 from the reaction mixture to prevent the thermal Fenton reaction. Although direct spectroscopic detection of the hydroxyl radical has proved elusive (Marin et al., 2011), the Cl2 •**-** radical anion, which absorbs at 340 nm (**ε**340nm *=* 8000 M-1cm-1), can be readily detected upon excitation of aqueous solutions of iron(III) at acidic pH in the presence of added sodium chloride at 355 nm with the third harmonic of a Nd-YAG laser. Differential absorption spectra (Figure 4) and kinetic traces (insert, Figure 4) showed fast formation of the Cl2 •- radical anion, within the lifetime of the laser pulse (5 ns), and its subsequent decay via mixed first and second order kinetics. The net decrease in absorption at longer times, relative to that prior to the laser pulse, reflects the conversion of Fe(III) to Fe(II), which does not absorb in this spectral region.

Fundamental Mechanistic Studies of

Fe(OH)2+ (equation 5).

Fe(Cl)2+ ⎯h

Cl• + Cl-

Cl2

*Speciation Equilibria (I = 1)*

*Reactions of Iron Species*

*Reactions of Reactive Chlorine Species*

and Fe3+ and NaCl concentrations.

4 Fe2+ + Cl2

the Photo-Fenton Reaction for the Degradation of Organic Pollutants 279

wide range of pH and concentration of Fe(III) and chloride ion and these data used to develop an explicit mechanistic model for the initial photoinduced processes involved in the photo-Fenton reaction in the presence of chloride ion (Machulek et al., 2006). As indicated in Figure 5, in the presence of chloride ion, the dominant species present at pH < 2.5 is no

Moreover, FeCl2+ not only absorbs further out in the ultraviolet than Fe(OH)2+ (Figure 2), but it also undergoes photolysis (equation 6) with a quantum yield higher than that of

→ Cl2

•- were then obtained over a

•- (equation

•- took into

•-

⎯→*v* Fe2+ + Cl• (6)

•- → 2 Cl- + Cl2 (8)

•- (7)

Laser flash photolysis data for the production and decay of Cl2

longer Fe3+, but rather FeCl2+ and FeCl2+ (coordinated waters being omitted).

The photoproduced chlorine atoms rapidly react with chloride ions to form Cl2

and whose second order decay component is due to the disproportionation reaction:

account the initial photochemical formation of the hydroxyl radical from Fe(OH)2+ (equation 5) and of chlorine atoms from FeCl2+ (equation 6) and the subsequent reactions of these via

•- + Cl2

The model employed to fit the experimentally observed decay curves for Cl2

the set of elementary steps outlined in Table 1 (Machulek et al., 2006).

3 Fe2+ + Cl• → Fe3+ + Cl− *k3* = 5.9 x 109

6 Cl• + H2O2 → HO2• + Cl− + H+ *k6* = 1 x 109 7 Cl2•− + Cl2•− → 2Cl− + Cl2 *k7* = 2.8 x 109 8 Cl2•− + H2O2 → HO2• + 2Cl− + H+ *k8* = 1.4 x 105 9 Cl2•− + HO2• → 2Cl− + H+ + O2 *k9* = 3.1 x 109

12 Cl2•− + Cl• → Cl− + Cl2 *k12* = 1.4 x 109

**No. Reaction** *k* **(s-1 or mol-1 L s-1)**

1 Fe3+ + Cl− ⇄ FeCl2+ *k1* = 4.79 x 1010, *k-1* = 1 x 1010 2 Fe3+ + 2 Cl− ⇄ FeCl2+ *k2* = 6.31 x 1010, *k-2* = 1 x 1010

•− → Fe3+ + 2Cl− *k4* = 5 x 106

5 Cl• + Cl− ⇄ Cl2•− *k5* = 7.8 x 109, *k-5* = 5.7 x 104

10 Cl− + HO• ⇄ HOCl•− *k10* = 4.2 x 109, *k*-*10* = 6.1 x 109 11 HOCl•− + H+ ⇄ H2O + Cl• *k11* = 2.4 x 1010, *k*-*11* = 1.8 x 105

Table 1. Set of reactions required to simulate the kinetics for production and decay of Cl2

upon laser flash photolysis of Fe3+ in the presence of chloride ion over a wide range of pH

7), the species actually detected in the laser flash photolysis experiments:

Fig. 4. Laser flash photolysis at 355 nm of 0.50 mM Fe(III) in the presence of 0.5 M NaCl at pH 1. Transient absorption spectra of Cl2 •- immediately (●) and 15 μs (○) after the laser pulse. The insert shows the kinetics of Cl2 • disappearance monitored at λ= 340 nm.

Fig. 5. Speciation of 0.5 mM Fe(III) between pH 1 and 4 at 25○C and an ionic strength of 0.2 in the presence of 0.1 M NaCl. Mole fractions of each species were calculated with the public domain program Hydra (Puigdomenech, 2010), employing the equilibrium constants for complexation supplied with the program (excluding insoluble iron species).

Organic Pollutants Ten Years After 278 the Stockholm Convention – Environmental and Analytical Update

Fig. 4. Laser flash photolysis at 355 nm of 0.50 mM Fe(III) in the presence of 0.5 M NaCl at

Fig. 5. Speciation of 0.5 mM Fe(III) between pH 1 and 4 at 25○C and an ionic strength of 0.2 in the presence of 0.1 M NaCl. Mole fractions of each species were calculated with the public domain program Hydra (Puigdomenech, 2010), employing the equilibrium constants for

complexation supplied with the program (excluding insoluble iron species).

•-

•- immediately (●) and 15 μs (○) after the laser

disappearance monitored at λ= 340 nm.

pH 1. Transient absorption spectra of Cl2

pulse. The insert shows the kinetics of Cl2

Laser flash photolysis data for the production and decay of Cl2 •- were then obtained over a wide range of pH and concentration of Fe(III) and chloride ion and these data used to develop an explicit mechanistic model for the initial photoinduced processes involved in the photo-Fenton reaction in the presence of chloride ion (Machulek et al., 2006). As indicated in Figure 5, in the presence of chloride ion, the dominant species present at pH < 2.5 is no longer Fe3+, but rather FeCl2+ and FeCl2+ (coordinated waters being omitted).

Moreover, FeCl2+ not only absorbs further out in the ultraviolet than Fe(OH)2+ (Figure 2), but it also undergoes photolysis (equation 6) with a quantum yield higher than that of Fe(OH)2+ (equation 5).

$$\text{Fe(Cl)2\*} \xrightarrow{hv} \text{Fe2\*} + \text{Cl}^\* \tag{6}$$

The photoproduced chlorine atoms rapidly react with chloride ions to form Cl2 •- (equation 7), the species actually detected in the laser flash photolysis experiments:

$$\mathsf{Cl^{\bullet}} + \mathsf{Cl^{\cdot}} \to \mathsf{Cl\_{2}^{\bullet}} \tag{7}$$

and whose second order decay component is due to the disproportionation reaction:

$$\rm Cl\_2^{\bullet\_\*} + Cl\_2^{\bullet\_\*} \to 2\ Cl\_1 + Cl\_2 \tag{8}$$

The model employed to fit the experimentally observed decay curves for Cl2 •- took into account the initial photochemical formation of the hydroxyl radical from Fe(OH)2+ (equation 5) and of chlorine atoms from FeCl2+ (equation 6) and the subsequent reactions of these via the set of elementary steps outlined in Table 1 (Machulek et al., 2006).

Table 1. Set of reactions required to simulate the kinetics for production and decay of Cl2 • upon laser flash photolysis of Fe3+ in the presence of chloride ion over a wide range of pH and Fe3+ and NaCl concentrations.

Fundamental Mechanistic Studies of

and f in Figure 3).

and decay of Cl2

listed in Table 2.

the Photo-Fenton Reaction for the Degradation of Organic Pollutants 281

al., 2007; Moraes et al., 2004a, 2004b; Pignatello, 1992) will occur at moderate chloride ion concentrations [>0.03 *M* NaCl for aliphatic hydrocarbons (Moraes et al., 2004a); >0.2 *M* NaCl for phenols] when the pH of the medium reaches pH 2.0 or below. On the other hand, this sequence of events clearly indicates that it should be possible to circumvent the inhibition by chloride ion by simply maintaining the pH at or slightly above pH 3 throughout the degradation process (see Figure 5). This is nicely illustrated by curve d in Figure 3, which shows that pH control does indeed permit nearly complete mineralization of phenol in the presence of chloride ion (Machulek et al., 2006, 2007). Although pH control also enhances the rate of the photo-Fenton reaction in the absence of chloride ion by optimizing the concentration of Fe(OH)2+ (see Figure 1), the effect is much less dramatic (compare curves e

Although the mechanism outlined in Table 1 nicely rationalizes the kinetics of formation

Fe(II) in the same system in the absence of H2O2 under steady-state irradiation. For this purpose, one must include additional kinetic steps that are too slow to be relevant on the time scale of the laser flash photolysis experiments. Experimentally (Machulek et al., 2009b) it is found that the concentration of Fe2+ eventually reaches a plateau value at long times, more quickly in the presence of chloride ion than in its absence (Figure 6). Simulation of this behavior required the inclusion of additional kinetic steps, in particular for the back reactions that result in reoxidation of Fe2+. These additional kinetic steps are

Fig. 6. The accumulation of ferrous ion during irradiation (350 nm) of 1.0 mM Fe(III) at: (A) pH 1.0 in the presence of 0.75 mol L-1 chloride ion; (B) at pH 3 in the absence of chloride

ion; and (C) at pH 3 in the presence of 1.8 mmol L-1 sulfate ions.

•- on a fast time scale, it does not permit prediction of the accumulation of

The initial concentrations of HO• and C1• produced at t = 0 by the laser pulse were calculated from the relationship (equation 9):

$$\left[\left.X^{\bullet}\right]\_{t=0} = \left(\mathbf{I}\_{\text{Laser}}\right) \left(\frac{\mathbf{A}\_{\text{Fe(\%)}}}{\mathbf{A}\_{\text{T}}}\right) \left(\mathbf{1} - \mathbf{1}0^{A\_{\text{T}}}\right) \left(\Phi\_{\text{X}}\right) \tag{9}$$

where X = OH or Cl and ΦHO = 0.21 and ΦCl = 0.47 are, respectively, the quantum yields for production of Fe(I1) from Fe(OH)2+ or Fe(Cl)2+ at 355 nm. AFe(X) is the absorbance of either Fe(OH)2+ or Fe(Cl)2+ and **AT** the total absorbance of the solution, both at the laser excitation wavelength of 355 nm (in an effective optical path length of 0.5 cm). The ratio AFe(X)/AT was calculated from the initial concentrations of the species Fe(Cl)2+, Fe(Cl)2+, Fe3+, Fe(OH)2+, Fe(OH)2+ and [Fe2(H2O)4(OH)2]4+, together with molar absorption coefficients at 355 nm taken from the literature absorption spectra (Byrne & Kester, 1978, 1981). For each experimental condition (pH, [Cl-], [Fe(III)]), the relative concentrations of the iron(III) species were calculated with the public domain speciation program Hydra (Puigdomenech, 2010), employing the set of standard equilibrium constants for complexation supplied with the program. The incident laser intensity (Ilaser) was estimated by fitting the transient absorbance of Cl2 •- at time zero at pH 1.0 in the presence of 0.5 M NaCl, where competitive photolysis of Fe(OH)2+ is negligible. The concentrations of HO• and Cl• at t = 0 calculated from equation 9 served as the initial conditions for numerical solution of the set of reactions and rate constants listed in Table 1. This kinetic model provided a quantitative fit of the observed transient decay curves (like that shown in the inset of Figure 4) over the entire range of pH and concentrations of chloride ion and Fe(III) investigated.

Having established the basic mechanistic scheme for the photoinduced steps of the photo-Fenton reaction, one can then use it to infer the course of a typical photo-Fenton degradation in the presence of chloride ion. At the beginning of the photo-Fenton reaction, when the pH is still ca. 3.0, the concentration of Fe(OH)2+ exceeds that of FeCl2+ or FeCl2+, even in the presence of relatively high concentrations of chloride ion (Figure 5). The pH-dependent scavenging of HO• by chloride ion (equations 10-11):

$$\text{HO}^\* \text{+ Cl} \leftrightharpoons \text{HOCl}^\* \text{:} \tag{10}$$

$$\rm H^{+} + \rm HCOCl^{\*} \rightleftharpoons \rm H\_{2}OCl^{\*} \rightarrow \rm H\_{2}O + Cl^{\*} \tag{11}$$

is also still a relatively inefficient process, so that the photochemical formation of HO• should predominate. However, as the photo-Fenton process proceeds, partial degradation of the organic material causes the pH of the medium to fall to ca. pH 2.0, where photolysis of FeCl2+ dominates over photolysis of Fe(OH)2+ and where the chloride ion efficiently converts any HO• formed in the system into the intrinsically much less reactive Cl2 •- radical anion (Bacardit et al., 2007; Buxton et al., 1999; De Laat et al., 2004; George & Chovelon, 2002; Kiwi et al., 2000; McElroy, 1990; Moraes, et al., 2004a; Nadtochenko & Kiwi, 1998; Pignatello, 1992; Soler et al., 2009; Truong et al., 2004; Yu, 2004; Yu et al., 2004; Yu & Barker, 2003a, 2003b; Zapata et al., 2009). As a result, virtually complete inhibition of the photo-Fenton degradation of typical organic substrates (Kiwi & Nadtochenko, 2000; Machulek et Organic Pollutants Ten Years After 280 the Stockholm Convention – Environmental and Analytical Update

A I 1 10 A

*X t X*•

where X = OH or Cl and ΦHO = 0.21 and ΦCl = 0.47 are, respectively, the quantum yields for production of Fe(I1) from Fe(OH)2+ or Fe(Cl)2+ at 355 nm. AFe(X) is the absorbance of either Fe(OH)2+ or Fe(Cl)2+ and **AT** the total absorbance of the solution, both at the laser excitation wavelength of 355 nm (in an effective optical path length of 0.5 cm). The ratio AFe(X)/AT was calculated from the initial concentrations of the species Fe(Cl)2+, Fe(Cl)2+, Fe3+, Fe(OH)2+, Fe(OH)2+ and [Fe2(H2O)4(OH)2]4+, together with molar absorption coefficients at 355 nm taken from the literature absorption spectra (Byrne & Kester, 1978, 1981). For each experimental condition (pH, [Cl-], [Fe(III)]), the relative concentrations of the iron(III) species were calculated with the public domain speciation program Hydra (Puigdomenech, 2010), employing the set of standard equilibrium constants for complexation supplied with the program. The incident laser intensity (Ilaser) was estimated

NaCl, where competitive photolysis of Fe(OH)2+ is negligible. The concentrations of HO• and Cl• at t = 0 calculated from equation 9 served as the initial conditions for numerical solution of the set of reactions and rate constants listed in Table 1. This kinetic model provided a quantitative fit of the observed transient decay curves (like that shown in the inset of Figure 4) over the entire range of pH and concentrations of chloride ion and

Having established the basic mechanistic scheme for the photoinduced steps of the photo-Fenton reaction, one can then use it to infer the course of a typical photo-Fenton degradation in the presence of chloride ion. At the beginning of the photo-Fenton reaction, when the pH is still ca. 3.0, the concentration of Fe(OH)2+ exceeds that of FeCl2+ or FeCl2+, even in the presence of relatively high concentrations of chloride ion (Figure 5). The pH-dependent

 H+ + HOCl•- ⇄ H2OCl• → H2O + Cl• (11) is also still a relatively inefficient process, so that the photochemical formation of HO• should predominate. However, as the photo-Fenton process proceeds, partial degradation of the organic material causes the pH of the medium to fall to ca. pH 2.0, where photolysis of FeCl2+ dominates over photolysis of Fe(OH)2+ and where the chloride ion efficiently

anion (Bacardit et al., 2007; Buxton et al., 1999; De Laat et al., 2004; George & Chovelon, 2002; Kiwi et al., 2000; McElroy, 1990; Moraes, et al., 2004a; Nadtochenko & Kiwi, 1998; Pignatello, 1992; Soler et al., 2009; Truong et al., 2004; Yu, 2004; Yu et al., 2004; Yu & Barker, 2003a, 2003b; Zapata et al., 2009). As a result, virtually complete inhibition of the photo-Fenton degradation of typical organic substrates (Kiwi & Nadtochenko, 2000; Machulek et

formed in the system into the intrinsically much less reactive Cl2

by chloride ion (equations 10-11):

Laser 0 T

( ) ( ) ( ) Fe(X)

*AT*

= − Φ (9)

produced at t = 0 by the laser pulse were

•- at time zero at pH 1.0 in the presence of 0.5 M

+ Cl- ⇄ HOCl•- (10)

•- radical

The initial concentrations of HO• and C1•

calculated from the relationship (equation 9):

by fitting the transient absorbance of Cl2

HO•

Fe(III) investigated.

scavenging of HO•

converts any HO•

=

al., 2007; Moraes et al., 2004a, 2004b; Pignatello, 1992) will occur at moderate chloride ion concentrations [>0.03 *M* NaCl for aliphatic hydrocarbons (Moraes et al., 2004a); >0.2 *M* NaCl for phenols] when the pH of the medium reaches pH 2.0 or below. On the other hand, this sequence of events clearly indicates that it should be possible to circumvent the inhibition by chloride ion by simply maintaining the pH at or slightly above pH 3 throughout the degradation process (see Figure 5). This is nicely illustrated by curve d in Figure 3, which shows that pH control does indeed permit nearly complete mineralization of phenol in the presence of chloride ion (Machulek et al., 2006, 2007). Although pH control also enhances the rate of the photo-Fenton reaction in the absence of chloride ion by optimizing the concentration of Fe(OH)2+ (see Figure 1), the effect is much less dramatic (compare curves e and f in Figure 3).

Although the mechanism outlined in Table 1 nicely rationalizes the kinetics of formation and decay of Cl2 •- on a fast time scale, it does not permit prediction of the accumulation of Fe(II) in the same system in the absence of H2O2 under steady-state irradiation. For this purpose, one must include additional kinetic steps that are too slow to be relevant on the time scale of the laser flash photolysis experiments. Experimentally (Machulek et al., 2009b) it is found that the concentration of Fe2+ eventually reaches a plateau value at long times, more quickly in the presence of chloride ion than in its absence (Figure 6). Simulation of this behavior required the inclusion of additional kinetic steps, in particular for the back reactions that result in reoxidation of Fe2+. These additional kinetic steps are listed in Table 2.

Fig. 6. The accumulation of ferrous ion during irradiation (350 nm) of 1.0 mM Fe(III) at: (A) pH 1.0 in the presence of 0.75 mol L-1 chloride ion; (B) at pH 3 in the absence of chloride ion; and (C) at pH 3 in the presence of 1.8 mmol L-1 sulfate ions.

Fundamental Mechanistic Studies of

*Speciation Equilibria (I = 1)*

*Reactions of Sulfate Ions*

*Reactions of Iron Species*

7 SO4

2009b).

the Photo-Fenton Reaction for the Degradation of Organic Pollutants 283

1 Fe3+ + SO42- ⇄ FeSO4+ *k1* = 2.09 x 1012, *k-1* = 1 x 1010 (I=1) 2 Fe3+ + 2SO42- ⇄ Fe(SO4)2- *k2* = 1.95 x 1013, *k-2* = 1 x 1010 (I=1) 3 Fe2+ + SO42- ⇄ FeSO4 *k3* = 1.55 x 1011, *k-3* = 1 x 1010 (I=1)

4 H+ + SO42- ⇄ HSO4- *k4* = 2.8 x 1011, *k-4* = 1 x 1010 (I=1)

Table 3. Set of additional ground state reactions required to fit the accumulation of Fe2+ during the steady-state irradiation of Fe3+ in the presence of sulfate ion (Machulek et al.,

Fig. 7. Speciation of 1.0 mM Fe(III) between pH 1 and 4 at 25○C and an ionic strength of 0.2 in the presence of 1.8 mM sulfate. Mole fractions of each species were calculated with the public domain program Hydra (Puigdomenech, 2010), employing the equilibrium constants

for complexation supplied with the program (excluding the insoluble iron species).

•− + OH− → SO42− + HO• *k7* = 1.4 x 107

**No. Reaction** *k* **(s-1 or mol-1 L s-1)**

5 HSO4− + HO• → SO4•− + H2O *k5* = 3.5 x 105 6 SO4•−+ H2O→ H+ + SO42− + HO• *k6* = 6.6 x 102

8 SO4•−+H2O2→SO42−+ H+ + HO2• *k8* = 1.2 x 107 9 SO4•− + HO2• → SO42− + H+ + O2 *k9* = 3.5 x 109 10 SO4•− + SO4•−→ S2O82− *k10* = 2.7 x 108

11 Fe2+ + SO4•−→ Fe3+ + SO42− *k11* = 3.0 x 108

Table 2. Set of additional ground state reactions required to fit the accumulation of Fe2+ during the steady-state irradiation of Fe3+ in the presence and absence of chloride ion (Machulek et al., 2009b)

#### **4.3 Inhibition of the Photo-Fenton reaction by sulfate ion**

Also indicated in Figure 6 is the inhibition of the formation of Fe2+ in the presence of a relatively low concentration of added sulfate ions, based on the rate constants for complexation of Fe(III) by sulfate in Table 3. Sulfate ion is often present in photo-Fenton reactions as a result of the addition of iron in the form of readily available sulfate salts (De Laat & Le, 2005). Sulfate ion complexes strongly with Fe(III) over a wide pH range (Figure 7) and the quantum yield of production of Fe2+ plus a sulfate anion radical from photolysis of Fe(SO4)+ is only about 0.05 (Benkelberg & Warneck, 1995). Iron phosphate is even more photoinert than the iron sulfate complex (Benkelberg & Warneck, 1995; Lee et al., 2003) and should be an even more powerful inhibitor of the photo-Fenton reaction.

#### **4.4 Catalysis of the Photo-Fenton reaction by complexation of Fe(III)**

The efficiency of the photo-Fenton process can be further enhanced by using organic carboxylic acids to complex Fe(III) (Pignatello et al., 2006). A particularly important example is provided by oxalic acid. Thus, unlike the thermal Fenton reaction, in which oxalic acid is a recalcitrant intermediate, in the photo-Fenton reaction it can act as a catalyst. Thus, ferrioxalate complexes can absorb light as far out as 570 nm, i.e., well into the visible region of the spectrum. Moreover, upon irradiation, they decompose efficiently (quantum yields of the order of unity) to Fe(II) and CO2. The net result is that, in the presence of oxalate, the photo-Fenton reaction is intrinsically more efficient, can be induced by a wider range of wavelengths of light, and results in the mineralization of the oxalate ion. Thus, for example, in a municipal water treatment system, Kim & Vogelpohl (1998) found that, with UV irradiation, the photo-Fenton process was at least 30% more energy efficient in the presence of oxalate than in its absence. Clearly, the sensitivity of the ferrioxalate-catalyzed photo-Fenton process to both UV and visible light makes it particularly attractive for applications in which the sun is employed as the radiation source (Silva et al., 2010; Trovó & Nogueira, 2011).

Organic Pollutants Ten Years After 282 the Stockholm Convention – Environmental and Analytical Update

• *k5* = 2.7 x 107

**No. Reaction** *k* **(s-1 or mol-1 L s-1)**

4 HO• + HO• → H2O2 *k4* = 6.0 x 109

8 Fe2+ + HO• → Fe3+ + OH− *k8* = 4.3 x 108 9 Fe2+ + H2O2 → Fe3+ + HO• + OH− *k9* = 63 10 Fe3+ + HO2• → Fe2+ + O2 + H+ *k10* = 1 x 106 11 Fe3+ + H2O2 → Fe2+ + HO2• + H+ *k11* = 0.01

**4.3 Inhibition of the Photo-Fenton reaction by sulfate ion**

should be an even more powerful inhibitor of the photo-Fenton reaction.

**4.4 Catalysis of the Photo-Fenton reaction by complexation of Fe(III)**

employed as the radiation source (Silva et al., 2010; Trovó & Nogueira, 2011).

• + H2O2 → HO• + O2 + H2O *k7* = 0.5

1 Fe3+ + H2O ⇄ FeOH2+ + H+ *k1* = 1.91 x 107, *k-1* = 1 x 1010 2 Fe3+ + 2H2O ⇄ Fe(OH)2+ + 2H+ *k2* = 3.39 x 103, *k-2* = 1 x 1010 3 2Fe3+ + 2H2O ⇄ [Fe2(OH)2]4+ + 2H+ *k3* = 1.12 x 107, *k-3* = 1 x 1010

• → H2O2 + O2 *k6* = 8.3 x 105

Table 2. Set of additional ground state reactions required to fit the accumulation of Fe2+ during the steady-state irradiation of Fe3+ in the presence and absence of chloride ion

Also indicated in Figure 6 is the inhibition of the formation of Fe2+ in the presence of a relatively low concentration of added sulfate ions, based on the rate constants for complexation of Fe(III) by sulfate in Table 3. Sulfate ion is often present in photo-Fenton reactions as a result of the addition of iron in the form of readily available sulfate salts (De Laat & Le, 2005). Sulfate ion complexes strongly with Fe(III) over a wide pH range (Figure 7) and the quantum yield of production of Fe2+ plus a sulfate anion radical from photolysis of Fe(SO4)+ is only about 0.05 (Benkelberg & Warneck, 1995). Iron phosphate is even more photoinert than the iron sulfate complex (Benkelberg & Warneck, 1995; Lee et al., 2003) and

The efficiency of the photo-Fenton process can be further enhanced by using organic carboxylic acids to complex Fe(III) (Pignatello et al., 2006). A particularly important example is provided by oxalic acid. Thus, unlike the thermal Fenton reaction, in which oxalic acid is a recalcitrant intermediate, in the photo-Fenton reaction it can act as a catalyst. Thus, ferrioxalate complexes can absorb light as far out as 570 nm, i.e., well into the visible region of the spectrum. Moreover, upon irradiation, they decompose efficiently (quantum yields of the order of unity) to Fe(II) and CO2. The net result is that, in the presence of oxalate, the photo-Fenton reaction is intrinsically more efficient, can be induced by a wider range of wavelengths of light, and results in the mineralization of the oxalate ion. Thus, for example, in a municipal water treatment system, Kim & Vogelpohl (1998) found that, with UV irradiation, the photo-Fenton process was at least 30% more energy efficient in the presence of oxalate than in its absence. Clearly, the sensitivity of the ferrioxalate-catalyzed photo-Fenton process to both UV and visible light makes it particularly attractive for applications in which the sun is

*Speciation Equilibria (I = 1)*

6 HO2• + HO2

*Reactions of Iron Species*

(Machulek et al., 2009b)

7 HO2

*Reactions of Reactive Oxygen Radicals*

5 HO• + H2O2 → H2O + HO2

Table 3. Set of additional ground state reactions required to fit the accumulation of Fe2+ during the steady-state irradiation of Fe3+ in the presence of sulfate ion (Machulek et al., 2009b).

Fig. 7. Speciation of 1.0 mM Fe(III) between pH 1 and 4 at 25○C and an ionic strength of 0.2 in the presence of 1.8 mM sulfate. Mole fractions of each species were calculated with the public domain program Hydra (Puigdomenech, 2010), employing the equilibrium constants for complexation supplied with the program (excluding the insoluble iron species).

Fundamental Mechanistic Studies of

**6. Conclusions**

concurrent thermal Fenton reaction.

**7. Acknowledgment**

at UFMS.

absorbed light energy is dissipated as heat in the solution.

the Photo-Fenton Reaction for the Degradation of Organic Pollutants 285

the reactor walls and maximize the total amount of light absorbed by the system. Reactor design must also take into account the strong interdependence between light absorption, mass transport and reaction kinetics in solutions with high optical densities and optimize the balance between the concurrent photochemical and thermal contributions to the Fenton reaction. Although the relatively long irradiation times required for the complete mineralization of organic contaminants tend to compromise the economic viability of standalone photo-Fenton processes, they can be viable as a pre-treatment to reduce the toxicity of the effluent to levels compatible with other remediation technologies such as conventional biological treatment. Hence, the design of photo-Fenton reactors should also take into account the necessity of facile integration with other effluent decontamination technologies. A final point that has only recently been considered in a more systematic context is the intensification of Fenton processes by increasing the temperature (Zazo et al., 2011), particularly attractive in the case of photo-Fenton reactions where a substantial part of the

Of the currently known advanced oxidation processes, the Fenton reaction and the photo-Fenton reaction appear to be the most promising for practical industrial applications on a moderate scale. The fundamental photomechanistic aspects of the photo-Fenton reaction have been clarified by studies of the photochemistry of Fe(IIII) in the absence of hydrogen peroxide in order to avoid the complications due to the competing Fenton reaction. The Fe(III) species primarily responsible for the initial photochemical generation of Fe(II) plus a hydroxyl radical is Fe(OH)2+. Inhibition of the photo-Fenton by ions such as chloride or sulfate can be understood by considering the competitive pH-dependent complexation of Fe(III) by these ions, the competitive absorption and photochemistry of these complexes and/or the subsequent chemical reactions that can convert the hydroxyl radical into less reactive species. Adequate mechanisms, together with the relevant rate constants, are now available for the inhibition of the net photoconversion of Fe(III) to Fe(II) by both chloride and sulfate, allowing quantitative or semi-quantitative simulation of the experimentally observed inhibitory effects. As predicted by the mechanistic models, the inhibitory effect of chloride ion on real photo-Fenton degradations can be readily circumvented by simply maintaining the medium pH at ca. pH 3 throughout the reaction. Areas where additional progress would be welcome include an understanding of the role of temperature effects on the Fenton and photo-Fenton reactions and new strategies for further enhancing the efficiency of the photo-Fenton reaction. Improvements in photoreactor design are necessary in order to optimize the contribution of the photoinduced processes relative to the

The authors acknowledge the Brazilian funding agencies CAPES and CNPq for financial and fellowship support. F.H.Q. is associated with NAP-PhotoTech, the USP Research Consortium for Photochemical Technology, and INCT-Catalysis. A.M. Jr. and J.E.F. Moraes are affiliated with INCT-EMA. A.M.Jr. thanks FUNDECT for financial support of the work

#### **4.5 Concurrent Photo-Fenton and Fenton reactions**

A final point that should be emphasized is that, under typical photo-Fenton reaction conditions, both the thermal Fenton reaction and the photo-Fenton reactions occur concurrently in the reaction mixture. Light is almost always the limiting reagent in photochemical reactions and photo-Fenton reactions are often conducted in reactors in which the solution is recirculated through the photoreactor from an external reservoir. Thus, at any given moment, the photo-Fenton reaction will be operative only in the portion of the total reaction mixture that is contained in the photoreactor *per se*, while the thermal Fenton reaction can proceed throughout the entire reaction volume. Indeed, comparison of the degradation of phenol under various conditions (Figure 3) strongly suggests that the initial phase of the reaction is dominated by the thermal reaction, with the inflection at about 80 min where the curves diverge indicating the point at which the photo-Fenton reaction becomes dominant.

Particularly interesting is the apparent insensitivity of the thermal Fenton degradation of phenol to the presence of added chloride ion, as indicated by the similarity of curves a and b of Figure 3. At the initial pH of 3 of the reaction mixture, the hydroxyl radical is inefficiently scavenged by chloride and produces, as the initial products of phenol degradation, the DHBs catechol and hydroquinone. Thus, unlike the photo-Fenton reaction, which depends on the intermediacy of the hydroxyl radical, the thermal degradation of phenol shifts to the DHB-catalyzed reaction mentioned above for the Fenton reaction until all of the DHBs have themselves been degraded (Chen and Pignatello, 1997). At that point, most of the iron(III) is rendered redox inert by complexation with oxalic acid and/or other recalcitrant aliphatic acids and the Fenton reaction ceases (Zanta et al., 2010).

#### **5. Areas for future development**

Despite significant advances in the last decade, the widespread successful application of photochemical technology for the treatment and decontamination of industrial residues and wastewaters is still not a reality. Research chemists and process engineers must work together to solve several important practical problems that limit the technological applications of the photo-Fenton reaction on a large scale (Braun et al., 1991). Since the overall effective quantum efficiency can be relatively low, long irradiation times must be employed in order to achieve total mineralization of the organic components of effluents. Electrical power consumption can then easily represent 60% of the total operating expense of a lamp-based photochemical reactor. Additional costs associated with non-solar photochemical reactors include cooling, maintenance and depreciation of the lamps and UV radiation shielding and protection for workers. One of the obvious economic benefits of the use of solar photochemical reactors is the elimination of costs associated with installation and maintenance of lamps and with electrical power consumption (Moraes et al., 2004b; Nascimento et al., 2007). In parallel with the development of solar reactors or hybrid lamp/solar reactors, an important area is the development of photocatalysts that operate effectively under visible or solar irradiation (Anpo, 2000).

Particularly in larger scale reactors, operational problems such as the formation of deposits on the reactor walls that block the incident radiation must be addressed. Optimization strategies (Braun et al., 1991, 1993; Cassano et al., 1995; Nascimento et al., 1994; Oliveros et al., 1998) must be applied to the design of reactors that homogenize the incidence of light on the reactor walls and maximize the total amount of light absorbed by the system. Reactor design must also take into account the strong interdependence between light absorption, mass transport and reaction kinetics in solutions with high optical densities and optimize the balance between the concurrent photochemical and thermal contributions to the Fenton reaction. Although the relatively long irradiation times required for the complete mineralization of organic contaminants tend to compromise the economic viability of standalone photo-Fenton processes, they can be viable as a pre-treatment to reduce the toxicity of the effluent to levels compatible with other remediation technologies such as conventional biological treatment. Hence, the design of photo-Fenton reactors should also take into account the necessity of facile integration with other effluent decontamination technologies. A final point that has only recently been considered in a more systematic context is the intensification of Fenton processes by increasing the temperature (Zazo et al., 2011), particularly attractive in the case of photo-Fenton reactions where a substantial part of the absorbed light energy is dissipated as heat in the solution.

#### **6. Conclusions**

Organic Pollutants Ten Years After 284 the Stockholm Convention – Environmental and Analytical Update

A final point that should be emphasized is that, under typical photo-Fenton reaction conditions, both the thermal Fenton reaction and the photo-Fenton reactions occur concurrently in the reaction mixture. Light is almost always the limiting reagent in photochemical reactions and photo-Fenton reactions are often conducted in reactors in which the solution is recirculated through the photoreactor from an external reservoir. Thus, at any given moment, the photo-Fenton reaction will be operative only in the portion of the total reaction mixture that is contained in the photoreactor *per se*, while the thermal Fenton reaction can proceed throughout the entire reaction volume. Indeed, comparison of the degradation of phenol under various conditions (Figure 3) strongly suggests that the initial phase of the reaction is dominated by the thermal reaction, with the inflection at about 80 min where the curves diverge indicating the point at which the photo-Fenton reaction

Particularly interesting is the apparent insensitivity of the thermal Fenton degradation of phenol to the presence of added chloride ion, as indicated by the similarity of curves a and b of Figure 3. At the initial pH of 3 of the reaction mixture, the hydroxyl radical is inefficiently scavenged by chloride and produces, as the initial products of phenol degradation, the DHBs catechol and hydroquinone. Thus, unlike the photo-Fenton reaction, which depends on the intermediacy of the hydroxyl radical, the thermal degradation of phenol shifts to the DHB-catalyzed reaction mentioned above for the Fenton reaction until all of the DHBs have themselves been degraded (Chen and Pignatello, 1997). At that point, most of the iron(III) is rendered redox inert by complexation with oxalic acid and/or other recalcitrant aliphatic

Despite significant advances in the last decade, the widespread successful application of photochemical technology for the treatment and decontamination of industrial residues and wastewaters is still not a reality. Research chemists and process engineers must work together to solve several important practical problems that limit the technological applications of the photo-Fenton reaction on a large scale (Braun et al., 1991). Since the overall effective quantum efficiency can be relatively low, long irradiation times must be employed in order to achieve total mineralization of the organic components of effluents. Electrical power consumption can then easily represent 60% of the total operating expense of a lamp-based photochemical reactor. Additional costs associated with non-solar photochemical reactors include cooling, maintenance and depreciation of the lamps and UV radiation shielding and protection for workers. One of the obvious economic benefits of the use of solar photochemical reactors is the elimination of costs associated with installation and maintenance of lamps and with electrical power consumption (Moraes et al., 2004b; Nascimento et al., 2007). In parallel with the development of solar reactors or hybrid lamp/solar reactors, an important area is the development of photocatalysts that operate

Particularly in larger scale reactors, operational problems such as the formation of deposits on the reactor walls that block the incident radiation must be addressed. Optimization strategies (Braun et al., 1991, 1993; Cassano et al., 1995; Nascimento et al., 1994; Oliveros et al., 1998) must be applied to the design of reactors that homogenize the incidence of light on

**4.5 Concurrent Photo-Fenton and Fenton reactions**

acids and the Fenton reaction ceases (Zanta et al., 2010).

effectively under visible or solar irradiation (Anpo, 2000).

**5. Areas for future development**

becomes dominant.

Of the currently known advanced oxidation processes, the Fenton reaction and the photo-Fenton reaction appear to be the most promising for practical industrial applications on a moderate scale. The fundamental photomechanistic aspects of the photo-Fenton reaction have been clarified by studies of the photochemistry of Fe(IIII) in the absence of hydrogen peroxide in order to avoid the complications due to the competing Fenton reaction. The Fe(III) species primarily responsible for the initial photochemical generation of Fe(II) plus a hydroxyl radical is Fe(OH)2+. Inhibition of the photo-Fenton by ions such as chloride or sulfate can be understood by considering the competitive pH-dependent complexation of Fe(III) by these ions, the competitive absorption and photochemistry of these complexes and/or the subsequent chemical reactions that can convert the hydroxyl radical into less reactive species. Adequate mechanisms, together with the relevant rate constants, are now available for the inhibition of the net photoconversion of Fe(III) to Fe(II) by both chloride and sulfate, allowing quantitative or semi-quantitative simulation of the experimentally observed inhibitory effects. As predicted by the mechanistic models, the inhibitory effect of chloride ion on real photo-Fenton degradations can be readily circumvented by simply maintaining the medium pH at ca. pH 3 throughout the reaction. Areas where additional progress would be welcome include an understanding of the role of temperature effects on the Fenton and photo-Fenton reactions and new strategies for further enhancing the efficiency of the photo-Fenton reaction. Improvements in photoreactor design are necessary in order to optimize the contribution of the photoinduced processes relative to the concurrent thermal Fenton reaction.

#### **7. Acknowledgment**

The authors acknowledge the Brazilian funding agencies CAPES and CNPq for financial and fellowship support. F.H.Q. is associated with NAP-PhotoTech, the USP Research Consortium for Photochemical Technology, and INCT-Catalysis. A.M. Jr. and J.E.F. Moraes are affiliated with INCT-EMA. A.M.Jr. thanks FUNDECT for financial support of the work at UFMS.

Fundamental Mechanistic Studies of

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**12**

*1India 2Germany*

**Photocatalytic Degradation of Organic**

Malik Mohibbul Haque1, Detlef Bahnemann2 and Mohammad Muneer1\*

A wide variety of organic pollutants are introduced into the water system from various sources such as industrial effluents, agricultural runoff and chemical spills (Muszkat et al., 1994; Cohen et al., 1986). Their toxicity, stability to natural decomposition and persistence in the environment has been the cause of much concern to the societies and regulation

Development of appropriate methods for the degradation of contaminated drinking, ground, surface waters, wastewaters containing toxic or nonbiodegradable compounds is necessary. Among many processes proposed and/or being developed for the destruction of the organic contaminants, biodegradation has received the greatest attention. However, many organic chemicals, especially which are toxic or refractory, are not amendable to microbial degradation. Researcher showed their interest and started the intensive studies on heterogeneous photocatalysis, after the discovery of the photo-induced splitting of water on

Semiconductor particles have been found to act as heterogeneous photocatalysts in a number of environmentally important reactions (Blake, 2001; Pirkanniemi, & Sillanpää, 2002; Gaya & Abdullah, 2008).  Semiconductor photocatalysts have been shown to be useful as carbon dioxide (Irvine et al., 1990) and nitrogen (Khan & Rao, 1991) fixatives and for the decomposition of O3 (Ohtani et al., 1992), destruction of microorganisms such as bacteria (Matsunaga & Okochi, 1995; Zhang et al., 1994; Dunlop et al.,

**1. Introduction**

authorities around the world (Dowd et al., 1998).

TiO2 electrodes (Fujishima and Honda, 1972).

**Pollutants: Mechanisms and Kinetics**

*2Institut fuer Technische Chemie, Leibniz Universität Hannover,*

*1Department of Chemistry, Aligarh Muslim University,*

### **Photocatalytic Degradation of Organic Pollutants: Mechanisms and Kinetics**

Malik Mohibbul Haque1, Detlef Bahnemann2 and Mohammad Muneer1\* *1Department of Chemistry, Aligarh Muslim University, 2Institut fuer Technische Chemie, Leibniz Universität Hannover, 1India 2Germany*

#### **1. Introduction**

Organic Pollutants Ten Years After 292 the Stockholm Convention – Environmental and Analytical Update

Zanta, C.L.P.S., Friedrich, L.C., Machulek Jr., A, Higa, K.M. & Quina, F.H. (2010). Surfactant

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0920-5861

degradation by a catechol-driven Fenton reaction. *Journal of Hazardous Materials*,

Evaluation of operational parameters involved in solar photo-Fenton degradation of a commercial pesticide mixture. *Catalysis Today*, Vol. 144, No. 1-2, pp. 94-99, ISSN

Fenton process by increasing the temperature. *Industrial & Engineering Chemistry*

A wide variety of organic pollutants are introduced into the water system from various sources such as industrial effluents, agricultural runoff and chemical spills (Muszkat et al., 1994; Cohen et al., 1986). Their toxicity, stability to natural decomposition and persistence in the environment has been the cause of much concern to the societies and regulation authorities around the world (Dowd et al., 1998).

Development of appropriate methods for the degradation of contaminated drinking, ground, surface waters, wastewaters containing toxic or nonbiodegradable compounds is necessary. Among many processes proposed and/or being developed for the destruction of the organic contaminants, biodegradation has received the greatest attention. However, many organic chemicals, especially which are toxic or refractory, are not amendable to microbial degradation. Researcher showed their interest and started the intensive studies on heterogeneous photocatalysis, after the discovery of the photo-induced splitting of water on TiO2 electrodes (Fujishima and Honda, 1972).

Semiconductor particles have been found to act as heterogeneous photocatalysts in a number of environmentally important reactions (Blake, 2001; Pirkanniemi, & Sillanpää, 2002; Gaya & Abdullah, 2008).  Semiconductor photocatalysts have been shown to be useful as carbon dioxide (Irvine et al., 1990) and nitrogen (Khan & Rao, 1991) fixatives and for the decomposition of O3 (Ohtani et al., 1992), destruction of microorganisms such as bacteria (Matsunaga & Okochi, 1995; Zhang et al., 1994; Dunlop et al.,

Photocatalytic Degradation of Organic Pollutants: Mechanisms and Kinetics 295

If charge separation is maintained, the electron and hole may migrate to the catalyst surface where they participate in redox reactions with sorbed species. Specially, h+vb may react with

TiO2+ h ν → e cb + h+vb (1)

O2 + e cb → O2 (2)

 (H2O ⇔ H+ + OH-)ads + hvb+ → •OH + H+ (3) It has been suggested that the hydroxyl radical (•OH) and superoxide radical anions (O2

are the primary oxidizing species in the photocatalytic oxidation processes. These oxidative reactions would results in the degradation of the pollutants as shown in the following

For oxidation reactions to occur, the VB must have a higher oxidation potential than the material under consideration. The redox potential of the VB and the CB for different semiconductors varies between +4.0 and -1.5 volts versus Normal Hydrogen Electrode (NHE) respectively. The VB and CB energies of the TiO2 are estimated to be +3.1 and -0.1 volts, respectively, which means that its band gap energy is 3.2 eV and therefore absorbs in the near UV light (λ<387 nm). Many organic compounds have a potential above that of the TiO2 valence band and therefore can be oxidized. In contrast, fewer organic compounds can be reduced since a smaller number of them have a potential below that of the TiO2 conduction band.

Due to non-toxic, easily available, inexpensive, biologically and chemically inert and stable

to photo and chemical corrosion, TiO2 is used in various fields as shown in Figure 2.

to generate superoxide radical anion (O2 ), as indicated in the following equations 1-3;

to produce the hydroxyl radical and e cb is picked up by oxygen

R + •OH → R'• + H2O (4)

R + h+ → R•+ → degradation products (5)

•- )

surface-bound H2O or OH -

equations 4-5;

**1.2 Use of semiconductor (TiO2) in various fields**

Fig. 2. Application of TiO2 in various fields**.**

2002), viruses (Lee et al., 1997), for the killing of malignant cancer-cells (Kubota et al., 1994), for the photo splitting of water (Pleskov & Krotova, 1993; Grätzel, 1981), for the cleanup of oil spills (Gerisher & Heller, 1992; Berry & Mueller, 1994), and for the control of the quality of meat freshness in food industry (Funazaki et al., 1995). Photocatalytic semiconductor films have been studied for the purposes of laboratory scale up of some of the above applications (Dunlop et al., 2002; Byrne et al., 2002) and, in the case of TiO2 films, for their photo-induced wettability properties (Wang et al., 1998, 1999; Sakai et al., 1998; Yu et al., 2002). This latter phenomenon, termed superhydrophilicity, is being explored for applications in the development of self-cleaning and antifogging surfaces.

Our research group at the Department of Chemistry, Aligarh Muslim University, Aligarh, India in collaboration with the Institut fuer Technische Chemie, Leibniz Universität Hannover, Hannover, Germany, have been actively involved in studying the photocatalytic degradation of variety of priority organic pollutants in aqueous suspensions.

#### **1.1 Mechanism of photooxidation process**

The acceleration of a chemical transformation by the presence of a catalyst with light is called photocatalysis. The catalyst may accelerate the photoreaction by interaction with the substrate in its ground or excited state and/or with a primary photoproduct, depending upon the mechanism of the photoreaction and itself remaining unaltered at the end of each catalytic cycle. Heterogeneous photocatalysis is a process in which two active phases solid and liquid are present. The solid phase is a catalyst, usually a semiconductor. The molecular orbital of semiconductors has a band structure. The bands of interest in photocatalysis are the populated valence band (VB) and it's largely vacant conduction band (CB), which is commonly characterized by band gap energy (Ebg). The semiconductors may be photoexcited to form electron-donor sites (reducing sites) and electron-acceptor sites (oxidising sites), providing great scope for redox reaction. When the semiconductor is illuminated with light (hν) of greater energy than that of the band gap, an electron is promoted from the VB to the CB leaving a positive hole in the valence band and an electron in the conduction band as illustrated in Figure 1.

Fig. 1. Photoexcitation of semiconductor leading to charge separation / oxidation / reduction sites.

If charge separation is maintained, the electron and hole may migrate to the catalyst surface where they participate in redox reactions with sorbed species. Specially, h+vb may react with surface-bound H2O or OH to produce the hydroxyl radical and e cb is picked up by oxygen to generate superoxide radical anion (O2 ), as indicated in the following equations 1-3;

$$\rm TiO\_2 + h \, v \to e^- \, e \flat + h^\* \, v \, \tag{1}$$

Organic Pollutants Ten Years After 294 the Stockholm Convention – Environmental and Analytical Update

2002), viruses (Lee et al., 1997), for the killing of malignant cancer-cells (Kubota et al., 1994), for the photo splitting of water (Pleskov & Krotova, 1993; Grätzel, 1981), for the cleanup of oil spills (Gerisher & Heller, 1992; Berry & Mueller, 1994), and for the control of the quality of meat freshness in food industry (Funazaki et al., 1995). Photocatalytic semiconductor films have been studied for the purposes of laboratory scale up of some of the above applications (Dunlop et al., 2002; Byrne et al., 2002) and, in the case of TiO2 films, for their photo-induced wettability properties (Wang et al., 1998, 1999; Sakai et al., 1998; Yu et al., 2002). This latter phenomenon, termed superhydrophilicity, is being explored for

Our research group at the Department of Chemistry, Aligarh Muslim University, Aligarh, India in collaboration with the Institut fuer Technische Chemie, Leibniz Universität Hannover, Hannover, Germany, have been actively involved in studying the photocatalytic

The acceleration of a chemical transformation by the presence of a catalyst with light is called photocatalysis. The catalyst may accelerate the photoreaction by interaction with the substrate in its ground or excited state and/or with a primary photoproduct, depending upon the mechanism of the photoreaction and itself remaining unaltered at the end of each catalytic cycle. Heterogeneous photocatalysis is a process in which two active phases solid and liquid are present. The solid phase is a catalyst, usually a semiconductor. The molecular orbital of semiconductors has a band structure. The bands of interest in photocatalysis are the populated valence band (VB) and it's largely vacant conduction band (CB), which is commonly characterized by band gap energy (Ebg). The semiconductors may be photoexcited to form electron-donor sites (reducing sites) and electron-acceptor sites (oxidising sites), providing great scope for redox reaction. When the semiconductor is illuminated with light (hν) of greater energy than that of the band gap, an electron is promoted from the VB to the CB leaving a positive hole in the valence band and an electron

Recombination

**D**

Fig. 1. Photoexcitation of semiconductor leading to charge separation / oxidation /

**D**

Photooxidation

**A**

Photoreduction

**A**

Recombination

applications in the development of self-cleaning and antifogging surfaces.

degradation of variety of priority organic pollutants in aqueous suspensions.

**1.1 Mechanism of photooxidation process**

in the conduction band as illustrated in Figure 1.

hν > Ebg

Energy

reduction sites.

CB

VB

Ebg

$$\text{O}\_2 + \text{e}^{\bullet} \text{c}^{\bullet} \rightarrow \text{O}\_2^{\bullet} \text{7} \tag{2}$$

$$\text{(H}\_2\text{O} \leftrightarrow \text{H}\_\* + \text{OH}\text{)}\_{\text{ads}} + \text{h}\_{\text{vb}\*} \rightarrow \text{.OH} + \text{H}\_\* \tag{3}$$

It has been suggested that the hydroxyl radical (•OH) and superoxide radical anions (O2 •- ) are the primary oxidizing species in the photocatalytic oxidation processes. These oxidative reactions would results in the degradation of the pollutants as shown in the following equations 4-5;

$$\text{R} + \text{\textasciicircum} \text{OH} \rightarrow \text{R}'\\\text{\textasciicircum} \text{R}' + \text{H}\_2\text{O} \tag{4}$$

$$\mathbf{R} + \mathbf{h}^\* \to \mathbf{R}^{\*\*} \to \text{degradation products} \tag{5}$$

For oxidation reactions to occur, the VB must have a higher oxidation potential than the material under consideration. The redox potential of the VB and the CB for different semiconductors varies between +4.0 and -1.5 volts versus Normal Hydrogen Electrode (NHE) respectively. The VB and CB energies of the TiO2 are estimated to be +3.1 and -0.1 volts, respectively, which means that its band gap energy is 3.2 eV and therefore absorbs in the near UV light (λ<387 nm). Many organic compounds have a potential above that of the TiO2 valence band and therefore can be oxidized. In contrast, fewer organic compounds can be reduced since a smaller number of them have a potential below that of the TiO2 conduction band.

#### **1.2 Use of semiconductor (TiO2) in various fields**

Due to non-toxic, easily available, inexpensive, biologically and chemically inert and stable to photo and chemical corrosion, TiO2 is used in various fields as shown in Figure 2.

Fig. 2. Application of TiO2 in various fields**.**

Photocatalytic Degradation of Organic Pollutants: Mechanisms and Kinetics 297

11 Acephate 2005; Atiqur Rahman et al., J. Adv.

12 Dichlorvos and Phosphamidon 2005; Atiqur Rahman & Muneer,

13 Propham, Propachlor and Tebuthiuron 2005; Muneer et al., Chemosphere,

17 Thiram 2005; Haque & Muneer, Indian J.

18 Picloram, Dicamba and Floumeturon 2005; Atiqur Rahman & Muneer, J.

19 2,4-dichlorophenoxy Acetic Acid 2004; Singh & Muneer Res. Chem.

20 Maleic Hydrazide 2004; Singh et al., J. Adv. Oxid.

22 Dimethyl Terephthalate 2003; Atiqur Rahman et al., Res.

24 Bromacil 2003; Singh et al., Photochem.

23 Isoproturon 2003; Haque & Muneer, J. Environ.

25 Diphenamid 2003; Atiqur Rahman et al., J. Adv.

26 Benzidine and 1,2-diphenyl hydrazine 2002; Muneer et al., Chemosphere,

27 Terbacil and 2,4,5-tribromoimidazole 2002; Muneer & Bahnemann, Appl.

28 1,2-diethyphthalate 2001; Muneer et al., J. Photochem.

29 Diuron 1999; Muneer et al., Res. Chem.

14 Dichlone, 2-amino-5-chloropyridine,

15 2,2′-dinitro biphenyl , N,N′-dimethyl-4 nitroso aniline , 4-dimethyl amino benzaldehyde, phthalaldehyde and

16 Indole-3-acetic acid and Indole-3-butyric

tetramethyl benzoquinone

21 Methoxychlor, Chlorothalonil and

Disulfoton

Chart 1.

acid

acid

benzoyl peroxide and 3-chloro perbenzoic

Oxid. Technol., 9, 1.

Desalination, 181, 161.

Intermed., 31, 807.

General, 289, 224.

Mat., 120, 219.

Chem. Technol., 12, 68.

Intermed., 30, 317.

Intermed., 30, 663.

Manag., 69, 169.

49, 193.

Chem. Intermed., 29, 35.

Photobiol. Sci., 2, 151.

Oxid. Technol., 6, 100.

Catal. B: Environ., 36, 95.

Intermed., 25, 667.

Photobiol., A: Chem., 143, 213.

Tech., 7, 184.

2005; Qamar et al., Res. Chem.

2005; Muneer et al., Appl. Catal. A:

2005; Qamar & Muneer, J. Hazard.

Environ. Sci. & Health, B40, 257.

2004; Muneer et al., Res. Chem.

61, 457.

The process of photocatalysis is also widely being contributed to various sub-discipline of Chemistry as shown in Figure 3.

Fig. 3**.** Application of photocatalysis in various sub-divisions of Chemistry.

#### **1.3 Our research focus**

We have studied the photocatalysed degradation of a large variety of organic pollutants under different reaction conditions to determine the detailed degradation kinetics and product identification in few selected systems for better mechanistic understanding. The different class of organic pollutants studied by our research group are shown below in chart 1.

Chart 1.

Organic Pollutants Ten Years After 296 the Stockholm Convention – Environmental and Analytical Update

The process of photocatalysis is also widely being contributed to various sub-discipline of

We have studied the photocatalysed degradation of a large variety of organic pollutants under different reaction conditions to determine the detailed degradation kinetics and product identification in few selected systems for better mechanistic understanding. The different class

Intermed., DOI 10.1007/s11164-

Photoenergy, article ID 197346

2007; Singh et al., J. Mol. Catal.

2007; Abu Tariq et al., J. Mol. Catal.

011-0299-6

3, 512,

261, 169.

142, 374.

142, 425.

A: Chem., 264, 66.

A: Chem., 265, 231.

Technol., 40, 4765.

Manag., 80, 99.

Fig. 3**.** Application of photocatalysis in various sub-divisions of Chemistry.

of organic pollutants studied by our research group are shown below in chart 1.

1 Dinoterb 2011; Dar et al., Res. Chem.

2 Fenoprop and Dichloroprop-P 2010; Faisal et al., Adv. Sci. Lett.

3 Acetamiprid 2010; Khan et al., Desalination,

4 Glyphosate 2008; Muneer & Boxall, Int. J.

5 4-chlorophenoxyacetic acid 2007; Singh et al., J. Hazard. Mat.,

7 Uracil and 5-bromouracil 2007; Singh et al., J. Hazard. Mat.,

9 Chlorotoluron 2006; Haque et al., Environ. Sci.

10 Trichlopyr and Daminozid 2006; Qamar et al., J. Environ.

**S. No. Compound Studied Reference**

Chemistry as shown in Figure 3.

**1.3 Our research focus**

6 Phenoxyacetic acid and 2,4,5-

8 4-bromoaniline, 3-nitroaniline,

and diphenylamine

trichlorophenoxyacetic acid

pentachlorophenol, 1,2,3-trichlorobenzene

Photocatalytic Degradation of Organic Pollutants: Mechanisms and Kinetics 299

platform under sunlight with continuous stirring and purging of molecular oxygen. Samples (10 mL) were collected before and at regular intervals during the illumination and

The photomineralization of the pesticide was measured using Total Organic Carbon analyzer (Shimadzu TOC 5000 A). The main principle of TOC analyzer involves the use of carrier gas (oxygen), which is flow-regulated (150 ml / min) and allows to flow through the total carbon (TC) combustion tube, which is packed with catalyst, and kept at 6800 C. When the sample enters the TC combustion tube, TC in the sample is oxidized to carbon dioxide. The carrier gas containing the combustion products from the TC combustion tube flows through the inorganic carbon (IC) reaction vessel, dehumidifier, halogen scrubber and finally reaches the sample cell of the nondispersive infrared (NDIR) detector which measures the carbon dioxide content. The output signal (analog) of the NDIR detector is displayed as peaks. The peak areas are measured and processed by the data processing unit. Since the peak areas are proportional to the total carbon concentration, the total carbon in a sample may be easily determined from the calibration curve prepared using standard solution of known carbon content. Total carbon is the sum of TOC (Total Organic Carbon)

The photodegradation of the pesticide was measured using UV-Vis spectrophotometry or

Intermediate product formed during the photooxidation process was characterized by monitoring the reaction as a function of time using GC/MS analysis technique. For GC/MS analysis a Shimadzu Gas Chromatograph and Mass Spectrometer (GCMS-QP 5050) equipped with a 25m CP SIL 19 CB (d=0.25mm) capillary column, operating temperature programmed (2200C for 40 min at the rate of 100C min-1) in splitless mode injection volume

Irradiation of an aqueous suspension of desired organic pollutants in the presence of TiO2 lead to decrease in absorption intensity and depletion in TOC content as a function of

As a representative example Figure 5 shows the change in absorption intensity and depletion in TOC as a function of time on irradiation of an aqueous solution of isoproturon (0.5 mM, 250 ml) in the presence and absence of photocatalyst (Degussa P25, 1 gL-1) by the "Pyrex" filtered output of a 125 W medium Pressure mercury lamp (radiant flux 4.860

**3. Photocatalysis of organic pollutants under different conditions**

analyzed after centrifugation.

and IC (Inorganic Carbon).

HPLC analysis techniques.

**3.1 In the presence of TiO2**

time.

mW/cm2).

**2.2 Photodegradation of the pesticide**

**2.3 Characterization of intermediate products**

(1.0 μL) with helium as a carrier gas was used.

**2.1 Photomineralization of organic pollutants**

**2. Analysis**

The following text describes the results of the photocatalysed degradation of different pollutants under variety of conditions for degradation kinetics and product identification.

#### **1.4 Procedure for conducting the degradation experiments**

Stock solutions of the pollutants containing the desired concentration were prepared in double distilled water. An immersion well photochemical reactor made of Pyrex glass equipped with a magnetic stirring bar, water circulating jacket and an opening for supply of molecular oxygen was used. A simplified diagram of the reactor system is shown in Figure 4.

Fig. 4. Simplified diagram of Photochemical Reaction Vessel.

For irradiation experiment an aqueous solution of the pollutants with desired concentration is taken into the photoreactor and required amount of photocatalyst was added. Then the solution was stirred and bubbled with molecular / atmospheric oxygen for at least 15 minutes in the dark to allow equilibration of the system so that the loss of compound due to adsorption can be taken into account. The zero time reading was obtained from blank solution kept in the dark in the presence of TiO2 and oxygen but otherwise treated similarly to the irradiated solution. The suspensions were continuously purged with molecular / atmospheric oxygen throughout each experiment. Irradiations were carried out using medium pressure mercury lamp. The light intensity was measured using UV-light intensity detector. IR-radiation and short-wavelength UV-radiation were eliminated by water circulating Pyrex glass jacket. Samples (10 mL) were collected before and at regular intervals during irradiation and analyzed after centrifugation.

The sunlight experiments were carried out between 9:00 A.M to 2:30 P.M. on a sunny day. The light intensity was measured using UV-light intensity detector (Lutron UV-340), which was found to be in the range of 0.370 to 0.480 mW/cm2. Reactions were carried out in the same reaction vessel as described above. Aqueous solution of desired concentration of the model compound containing required amount of photocatalyst was taken and stirred for 15 min. in the dark in presence of oxygen for equilibration. The solution was then placed on flat platform under sunlight with continuous stirring and purging of molecular oxygen. Samples (10 mL) were collected before and at regular intervals during the illumination and analyzed after centrifugation.

#### **2. Analysis**

Organic Pollutants Ten Years After 298 the Stockholm Convention – Environmental and Analytical Update

The following text describes the results of the photocatalysed degradation of different pollutants under variety of conditions for degradation kinetics and product identification.

Stock solutions of the pollutants containing the desired concentration were prepared in double distilled water. An immersion well photochemical reactor made of Pyrex glass equipped with a magnetic stirring bar, water circulating jacket and an opening for supply of molecular oxygen was used. A simplified diagram of the reactor system is shown in

For irradiation experiment an aqueous solution of the pollutants with desired concentration is taken into the photoreactor and required amount of photocatalyst was added. Then the solution was stirred and bubbled with molecular / atmospheric oxygen for at least 15 minutes in the dark to allow equilibration of the system so that the loss of compound due to adsorption can be taken into account. The zero time reading was obtained from blank solution kept in the dark in the presence of TiO2 and oxygen but otherwise treated similarly to the irradiated solution. The suspensions were continuously purged with molecular / atmospheric oxygen throughout each experiment. Irradiations were carried out using medium pressure mercury lamp. The light intensity was measured using UV-light intensity detector. IR-radiation and short-wavelength UV-radiation were eliminated by water circulating Pyrex glass jacket. Samples (10 mL) were collected before and at regular intervals

The sunlight experiments were carried out between 9:00 A.M to 2:30 P.M. on a sunny day. The light intensity was measured using UV-light intensity detector (Lutron UV-340), which was found to be in the range of 0.370 to 0.480 mW/cm2. Reactions were carried out in the same reaction vessel as described above. Aqueous solution of desired concentration of the model compound containing required amount of photocatalyst was taken and stirred for 15 min. in the dark in presence of oxygen for equilibration. The solution was then placed on flat

**1.4 Procedure for conducting the degradation experiments**

Fig. 4. Simplified diagram of Photochemical Reaction Vessel.

during irradiation and analyzed after centrifugation.

Figure 4.

#### **2.1 Photomineralization of organic pollutants**

The photomineralization of the pesticide was measured using Total Organic Carbon analyzer (Shimadzu TOC 5000 A). The main principle of TOC analyzer involves the use of carrier gas (oxygen), which is flow-regulated (150 ml / min) and allows to flow through the total carbon (TC) combustion tube, which is packed with catalyst, and kept at 6800 C. When the sample enters the TC combustion tube, TC in the sample is oxidized to carbon dioxide. The carrier gas containing the combustion products from the TC combustion tube flows through the inorganic carbon (IC) reaction vessel, dehumidifier, halogen scrubber and finally reaches the sample cell of the nondispersive infrared (NDIR) detector which measures the carbon dioxide content. The output signal (analog) of the NDIR detector is displayed as peaks. The peak areas are measured and processed by the data processing unit. Since the peak areas are proportional to the total carbon concentration, the total carbon in a sample may be easily determined from the calibration curve prepared using standard solution of known carbon content. Total carbon is the sum of TOC (Total Organic Carbon) and IC (Inorganic Carbon).

#### **2.2 Photodegradation of the pesticide**

The photodegradation of the pesticide was measured using UV-Vis spectrophotometry or HPLC analysis techniques.

#### **2.3 Characterization of intermediate products**

Intermediate product formed during the photooxidation process was characterized by monitoring the reaction as a function of time using GC/MS analysis technique. For GC/MS analysis a Shimadzu Gas Chromatograph and Mass Spectrometer (GCMS-QP 5050) equipped with a 25m CP SIL 19 CB (d=0.25mm) capillary column, operating temperature programmed (2200C for 40 min at the rate of 100C min-1) in splitless mode injection volume (1.0 μL) with helium as a carrier gas was used.

#### **3. Photocatalysis of organic pollutants under different conditions**

#### **3.1 In the presence of TiO2**

Irradiation of an aqueous suspension of desired organic pollutants in the presence of TiO2 lead to decrease in absorption intensity and depletion in TOC content as a function of time.

As a representative example Figure 5 shows the change in absorption intensity and depletion in TOC as a function of time on irradiation of an aqueous solution of isoproturon (0.5 mM, 250 ml) in the presence and absence of photocatalyst (Degussa P25, 1 gL-1) by the "Pyrex" filtered output of a 125 W medium Pressure mercury lamp (radiant flux 4.860 mW/cm2).

Photocatalytic Degradation of Organic Pollutants: Mechanisms and Kinetics 301

TOC = Total Organic Carbon, k = rate constant, c = concentration of the pollutant, n = order

Fig. 6. Plot showing the linear regression curve fit for the natural logarithm of absorbance of

Titanium dioxide is known to be the semiconductor with the highest photocatalytic activity and stable in aqueous solution. Several reviews have been written, regarding the mechanistic and kinetic details as well as the influence of experimental parameters. It has been demonstrated that degradation by photocatalysis can be more efficient than by other

We have tested the photocatalytic activity of four different commercially available TiO2 powders (namely Degussa P25, Sachtleben Hombikat UV100, Milenium Inorganic PC500 and Travancore Titanium Product, India) on the degradation kinetics of the pollutants. In most of the cases it has been observed that the degradation of pollutants under investigation proceed much more rapidly in the presence of Degussa P25 as compared with other TiO2 samples (Muneer and Bahnemann, 2001; Haque & Muneer, 2003; Bahnemann et al., 2007) as shown in Figure 7 (isoproturon). The better photocatalytic activity of Degussa P25 has also been reported by Pizzaro et al., in 2005. While in some cases Hombikat UV 100 is found to be better than Degussa P25 for the degradation of benzidine, eosine yellowish, and Remazol brilliant blue R (Muneer et al., 2002; Saquib and Muneer, 2002, 2003) as shown in Figure 8 (eosine yellowish). In an earlier study Lindner et al. 1995, showed that Hombikat UV100 was almost four times more effective than P25 when dichloroacetic acid was used as

the pesticide concentration of isoproturon against irradiation time for the first order

of reaction.

reaction.

the model pollutant.

**3.2 Comparison of different photocatalysts**

wet-oxidation technique (Weichrebe and Vogelpohl, 1995).

Fig. 5. Depletion in TOC and change in absorption intensity as a function of irradiation time for an aqueous solution of isoproturon in the presence and absence of the photocatalyst. Experimental conditions**:** 0.5 mM isoproturon, V=250 mL, photocatalyst: TiO2 (Degussa P25, 1 gL-1), immersion well photoreactor, 125 W medium pressure Hg lamp, absorbance was followed at 238 nm after 75% dilution, cont. O2 purging and stirring, irradiation time = 80 min.

It could be seen that 94.8% degradation and 78.45 % mineralization of the isoproturon takes place after 80 min of illumination. On the other hand no observable loss of the compound was found when the irradiations were carried out in the absence of the photocatalysts (Haque & Muneer, 2003).

Both the mineralization (depletion of TOC Vs irradiation time) and decomposition (decrease in absorption intensity/concentration Vs irradiation time) curves can be fitted reasonably well by an exponential decay curve suggesting the first order kinetics. Figure 6 shows the linear regression curve fit for the natural logarithm of the degradation of isoproturon Vs irradiation time for first order reaction. For each experiment, the degradation rate constant was calculated from the plot of the natural logarithm of the TOC depletion or absorbance / concentration of the pesticide as a function of irradiation time (the concentration of the pollutant was calculated by taking the absorbance of the pollutant at its λmax of standard concentration and then by plotting the graph of absorbance Vs standard concentration of the pollutant). The degradation rate for the mineralization and for the decomposition of the pollutants for first order reaction was calculated using formula given below;

$$\text{-d[TOC]}/\text{dt} = \text{kcf}^n\tag{6}$$

$$-\mathbf{d}[\mathbf{c}]/\mathbf{dt} = \mathbf{k}\mathbf{c}^n\tag{7}$$

Organic Pollutants Ten Years After 300 the Stockholm Convention – Environmental and Analytical Update

Isoproturon (hν) TOC (hν) TOC (TiO2 / hν) Isoproturon (TiO2 / hν)

 **A / A0**

min.

**TOC / TOC0**

0.0

(Haque & Muneer, 2003).

0.2

0.4

0.6

0.8

1.0

1.2

**Irradiation time (min)** 0 20 40 60 80

Fig. 5. Depletion in TOC and change in absorption intensity as a function of irradiation time for an aqueous solution of isoproturon in the presence and absence of the photocatalyst. Experimental conditions**:** 0.5 mM isoproturon, V=250 mL, photocatalyst: TiO2 (Degussa P25, 1 gL-1), immersion well photoreactor, 125 W medium pressure Hg lamp, absorbance was followed at 238 nm after 75% dilution, cont. O2 purging and stirring, irradiation time = 80

It could be seen that 94.8% degradation and 78.45 % mineralization of the isoproturon takes place after 80 min of illumination. On the other hand no observable loss of the compound was found when the irradiations were carried out in the absence of the photocatalysts

Both the mineralization (depletion of TOC Vs irradiation time) and decomposition (decrease in absorption intensity/concentration Vs irradiation time) curves can be fitted reasonably well by an exponential decay curve suggesting the first order kinetics. Figure 6 shows the linear regression curve fit for the natural logarithm of the degradation of isoproturon Vs irradiation time for first order reaction. For each experiment, the degradation rate constant was calculated from the plot of the natural logarithm of the TOC depletion or absorbance / concentration of the pesticide as a function of irradiation time (the concentration of the pollutant was calculated by taking the absorbance of the pollutant at its λmax of standard concentration and then by plotting the graph of absorbance Vs standard concentration of the pollutant). The degradation rate for the mineralization and for the decomposition of the

–d[c]/dt = kcn (7)

pollutants for first order reaction was calculated using formula given below;

TOC = Total Organic Carbon, k = rate constant, c = concentration of the pollutant, n = order of reaction.

Fig. 6. Plot showing the linear regression curve fit for the natural logarithm of absorbance of the pesticide concentration of isoproturon against irradiation time for the first order reaction.

#### **3.2 Comparison of different photocatalysts**

Titanium dioxide is known to be the semiconductor with the highest photocatalytic activity and stable in aqueous solution. Several reviews have been written, regarding the mechanistic and kinetic details as well as the influence of experimental parameters. It has been demonstrated that degradation by photocatalysis can be more efficient than by other wet-oxidation technique (Weichrebe and Vogelpohl, 1995).

We have tested the photocatalytic activity of four different commercially available TiO2 powders (namely Degussa P25, Sachtleben Hombikat UV100, Milenium Inorganic PC500 and Travancore Titanium Product, India) on the degradation kinetics of the pollutants.

In most of the cases it has been observed that the degradation of pollutants under investigation proceed much more rapidly in the presence of Degussa P25 as compared with other TiO2 samples (Muneer and Bahnemann, 2001; Haque & Muneer, 2003; Bahnemann et al., 2007) as shown in Figure 7 (isoproturon). The better photocatalytic activity of Degussa P25 has also been reported by Pizzaro et al., in 2005. While in some cases Hombikat UV 100 is found to be better than Degussa P25 for the degradation of benzidine, eosine yellowish, and Remazol brilliant blue R (Muneer et al., 2002; Saquib and Muneer, 2002, 2003) as shown in Figure 8 (eosine yellowish). In an earlier study Lindner et al. 1995, showed that Hombikat UV100 was almost four times more effective than P25 when dichloroacetic acid was used as the model pollutant.

Photocatalytic Degradation of Organic Pollutants: Mechanisms and Kinetics 303

The differences in the photocatalytic activity of TiO2 are likely to be due to differences in the BET-surface, impurities, lattice mismatches or density of hydroxyl groups present on the catalyst's surface. Since they will affect the adsorption behaviour of a pollutant or

The reason for the better photocatalytic activity of Degussa P25, could be attributed to the fact that P25 being composed of small nano-crystallites of rutile being dispersed within an anatase matrix. The smaller band gap of rutile "catches" the photons, generating electronhole pairs. The electron transfer, from the rutile conduction band to electron traps in anatase phase takes place. Recombination is thus inhibited allowing the hole to move to the surface of the particle and react (Hurum et al., 2003). The better efficiency of photocatalyst Degussa P25 may also due to 'quantum size effect' (Nozik et al., 1993; weller, 1993). When the particles become too small, there is a 'blue shift' with an increase of the band gap energy, detrimental to the near UV-photon absorption, and an increase of the electron-hole

An important parameter in the photocatalytic reactions taking place on the particulate surfaces is the pH of the solution, since it dictates the surface charge properties of the photocatalyst and size of aggregates it forms. Employing Degussa P25 as photocatalyst the decomposition and mineralization of pollutant in aqueous suspensions of TiO2 was studied

The degradation rate for the decomposition and mineralization of acid red 29 was found to increase with the increase in pH from 3 to 11 as shown in Figure 9 (Qamar et al., 2005) while the degradation rate for the decomposition and mineralization of the pesticide derivative propachlor was found to decrease with the increase in pH from 3 to 11 as shown in Figure 10 (Muneer et al, 2005). Similar results on pH effect have also been reported earlier by Vaz et al., 1998 (for the degradation of uracil and 5-halogenouracil) and Lu et al., 1995 (for the

The interpretation of pH effects on the photocatalytic process is very difficult task because of its multiple roles such as electrostatic interactions between the semiconductor surface, solvent molecules, substrate and charged radicals formed during the reaction process. The ionization state of the surface of the photocatalyst can be protonated and deprotonated under acidic and alkaline conditions, respectively, as shown in following

The point of zero charge (pzc) of the TiO2 (Degussa P25) is widely reported at pH ~ 6.25 (Augustynski, 1988). Thus, the TiO2 surface will remain positively charged in acidic medium (pH < 6.25) and negatively charged in alkaline medium (pH > 6.25). The functional group present on the pollutants can be protonated and deprotonated depending on the pH of the reaction mixture. The better degradation rate in acidic or basic pH may also be attributed on the basis of the fact that the structural orientation of the molecule is favoured for the attack

(8) (9)

intermediate molecule and the lifetime and recombination rate of electron-hole pairs.

recombination.

as a function of pH.

equations:

**3.3 Influence of pH on the degradation kinetics**

degradation of dichlorvos, propoxur and 2,4-D).

of the reactive species under that condition.

Fig. 7. Comparison of degradation rate for the mineralization and for the decomposition of isoproturon in the presence of different photocatalysts. Experimental conditions**:** 0.5 mM isoproturon, V=250 mL, photocatalysts: TiO2 Degussa P25 (1 gL-1), Sachtleben Hombikat UV100 (1 gL-1), PC500 (1 gL-1), TTP (1 gL-1).

Fig. 8. Comparison of degradation rate for the mineralization and for the decomposition of eosine yellowish under different photocatalysts. Experimental conditions: 0.25 mM dye concentration, V=250 mL, photocatalysts: TiO2 Degussa P25 (1 gL-1), Sachtleben Hombikat UV100 (1 gL-1) and PC500 (1 gL-1).

Organic Pollutants Ten Years After 302 the Stockholm Convention – Environmental and Analytical Update

Fig. 7. Comparison of degradation rate for the mineralization and for the decomposition of isoproturon in the presence of different photocatalysts. Experimental conditions**:** 0.5 mM isoproturon, V=250 mL, photocatalysts: TiO2 Degussa P25 (1 gL-1), Sachtleben Hombikat

> **Types of photocatalyst** 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0 4.5

Fig. 8. Comparison of degradation rate for the mineralization and for the decomposition of eosine yellowish under different photocatalysts. Experimental conditions: 0.25 mM dye concentration, V=250 mL, photocatalysts: TiO2 Degussa P25 (1 gL-1), Sachtleben Hombikat

UV100 (1 gL-1), PC500 (1 gL-1), TTP (1 gL-1).

**Degradation rate (mol L-1 min-1 \* 10-3 )**

3

UV100 (1 gL-1) and PC500 (1 gL-1).

4

5

6

7

8

9

10

11

The differences in the photocatalytic activity of TiO2 are likely to be due to differences in the BET-surface, impurities, lattice mismatches or density of hydroxyl groups present on the catalyst's surface. Since they will affect the adsorption behaviour of a pollutant or intermediate molecule and the lifetime and recombination rate of electron-hole pairs.

The reason for the better photocatalytic activity of Degussa P25, could be attributed to the fact that P25 being composed of small nano-crystallites of rutile being dispersed within an anatase matrix. The smaller band gap of rutile "catches" the photons, generating electronhole pairs. The electron transfer, from the rutile conduction band to electron traps in anatase phase takes place. Recombination is thus inhibited allowing the hole to move to the surface of the particle and react (Hurum et al., 2003). The better efficiency of photocatalyst Degussa P25 may also due to 'quantum size effect' (Nozik et al., 1993; weller, 1993). When the particles become too small, there is a 'blue shift' with an increase of the band gap energy, detrimental to the near UV-photon absorption, and an increase of the electron-hole recombination.

#### **3.3 Influence of pH on the degradation kinetics**

An important parameter in the photocatalytic reactions taking place on the particulate surfaces is the pH of the solution, since it dictates the surface charge properties of the photocatalyst and size of aggregates it forms. Employing Degussa P25 as photocatalyst the decomposition and mineralization of pollutant in aqueous suspensions of TiO2 was studied as a function of pH.

The degradation rate for the decomposition and mineralization of acid red 29 was found to increase with the increase in pH from 3 to 11 as shown in Figure 9 (Qamar et al., 2005) while the degradation rate for the decomposition and mineralization of the pesticide derivative propachlor was found to decrease with the increase in pH from 3 to 11 as shown in Figure 10 (Muneer et al, 2005). Similar results on pH effect have also been reported earlier by Vaz et al., 1998 (for the degradation of uracil and 5-halogenouracil) and Lu et al., 1995 (for the degradation of dichlorvos, propoxur and 2,4-D).

The interpretation of pH effects on the photocatalytic process is very difficult task because of its multiple roles such as electrostatic interactions between the semiconductor surface, solvent molecules, substrate and charged radicals formed during the reaction process. The ionization state of the surface of the photocatalyst can be protonated and deprotonated under acidic and alkaline conditions, respectively, as shown in following equations:

$$\mathsf{TOT} \star \mathsf{H}^{\star} \xleftarrow{\star} \mathsf{H} \xrightarrow{\quad} \mathsf{TOT}\_{2}^{\star} \tag{8}$$

(9)

The point of zero charge (pzc) of the TiO2 (Degussa P25) is widely reported at pH ~ 6.25 (Augustynski, 1988). Thus, the TiO2 surface will remain positively charged in acidic medium (pH < 6.25) and negatively charged in alkaline medium (pH > 6.25). The functional group present on the pollutants can be protonated and deprotonated depending on the pH of the reaction mixture. The better degradation rate in acidic or basic pH may also be attributed on the basis of the fact that the structural orientation of the molecule is favoured for the attack of the reactive species under that condition.

Photocatalytic Degradation of Organic Pollutants: Mechanisms and Kinetics 305

It is important both from mechanistic and from application point of view to study the dependence of substrate concentration in the photocatalytic reaction rate. Effect of substrate concentration on the degradation of the pollutants was studied at different concentrations. As a representative example Figure 11 shows the degradation rate for the TOC depletion and for the decomposition of pollutant as a function of substrate concentration employing Degussa P25 as photocatalyst (Muneer et al, 2005). It was found that the degradation rate for the decomposition and for the mineralization increases gradually with the increase in substrate concentrations. Similar trend was found in most of the colourless organic pollutants as reported earlier (Sabin et al., 1992; Krosley et al., 1993; O'Shea et al., 1997). In coloured compound it has been found that the degradation rate increase upto a certain limit and after that a further increase in substrate concentration lead to decrease in the degradation rate. This may be due to the fact that as the initial concentrations of the pollutant increases, the irradiating mixture becomes more and more intense which prevents the penetration of light to the surface of the catalyst. Hence, the generation of relative amount of OH and O2 on the surface of the catalyst do not increase as the intensity of light and irradiation times are constant. Conversely, their concentrations will decrease with increase in concentration of the pollutant as the light photons are largely absorbed and

> **Substrate concentration (mM)** 0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4

Fig. 11. Influence of substrate concentration on the degradation rate for the mineralization and for the decomposition of propachlor. Experimental conditions: Substrate concentrations (0.20, 0.40, 0.60, 0.80 and 1.35 mM), V=250 mL, photocatalyst: TiO2 (Degussa P25, 1 gL-1).

As oxidation proceeds, less and less of the surface of the TiO2 particle is covered as the pollutant is decomposed. Evidently, at total decomposition, the rate of degradation is zero and a decreased photocatalytic rate is to be expected with increasing irradiation time. It has been agreed, with minor variation that the expression for the rate of photodegradation of

prevented from reaching the catalyst surface by the substrate molecules.

TOC Propachlor

**3.4 Effect of substrate concentration**

**Degradation rate (mol L-1 min-1 \* 10-3 )**

0.008

0.012

0.016

0.020

0.024

0.028

0.032

Fig. 9. Influence of pH on the degradation rate for the mineralization and for the decomposition of acid red 29. Experimental conditions: Reaction pH (3, 5.2, 7, 9 and 10.5), 0.25 mM dye concentration, V=250 mL, photocatalyst: TiO2 (Degussa P25, 1 gL-1).

Fig. 10. Influence of pH on the degradation rate for the mineralization and for the decomposition of propachlor. Experimental conditions: Reaction pH (3, 5, 7, 9 and 11), 0.6 mM propachlor, V=250 mL, photocatalyst TiO2 (Degussa P25, 1 gL-1).

#### **3.4 Effect of substrate concentration**

Organic Pollutants Ten Years After 304 the Stockholm Convention – Environmental and Analytical Update

**TOC Propachlor**

> **TOC Propachlor**

**Degradation rate (mol L-1 min-1 \* 10-3 )**

**Degradation rate (mol L-1 min-1 \* 10-3 )**

0.008

0.012

0.016

0.020

0.024

0.028

0.008

0.012

0.016

0.020

0.024

0.028

**pH** 2 3 4 5 6 7 8 9 10 11 12

**pH** 2 3 4 5 6 7 8 9 10 11 12

Fig. 10. Influence of pH on the degradation rate for the mineralization and for the

mM propachlor, V=250 mL, photocatalyst TiO2 (Degussa P25, 1 gL-1).

decomposition of propachlor. Experimental conditions: Reaction pH (3, 5, 7, 9 and 11), 0.6

Fig. 9. Influence of pH on the degradation rate for the mineralization and for the

0.25 mM dye concentration, V=250 mL, photocatalyst: TiO2 (Degussa P25, 1 gL-1).

decomposition of acid red 29. Experimental conditions: Reaction pH (3, 5.2, 7, 9 and 10.5),

It is important both from mechanistic and from application point of view to study the dependence of substrate concentration in the photocatalytic reaction rate. Effect of substrate concentration on the degradation of the pollutants was studied at different concentrations. As a representative example Figure 11 shows the degradation rate for the TOC depletion and for the decomposition of pollutant as a function of substrate concentration employing Degussa P25 as photocatalyst (Muneer et al, 2005). It was found that the degradation rate for the decomposition and for the mineralization increases gradually with the increase in substrate concentrations. Similar trend was found in most of the colourless organic pollutants as reported earlier (Sabin et al., 1992; Krosley et al., 1993; O'Shea et al., 1997). In coloured compound it has been found that the degradation rate increase upto a certain limit and after that a further increase in substrate concentration lead to decrease in the degradation rate. This may be due to the fact that as the initial concentrations of the pollutant increases, the irradiating mixture becomes more and more intense which prevents the penetration of light to the surface of the catalyst. Hence, the generation of relative amount of OH and O2 on the surface of the catalyst do not increase as the intensity of light and irradiation times are constant. Conversely, their concentrations will decrease with increase in concentration of the pollutant as the light photons are largely absorbed and prevented from reaching the catalyst surface by the substrate molecules.

Fig. 11. Influence of substrate concentration on the degradation rate for the mineralization and for the decomposition of propachlor. Experimental conditions: Substrate concentrations (0.20, 0.40, 0.60, 0.80 and 1.35 mM), V=250 mL, photocatalyst: TiO2 (Degussa P25, 1 gL-1).

As oxidation proceeds, less and less of the surface of the TiO2 particle is covered as the pollutant is decomposed. Evidently, at total decomposition, the rate of degradation is zero and a decreased photocatalytic rate is to be expected with increasing irradiation time. It has been agreed, with minor variation that the expression for the rate of photodegradation of

Photocatalytic Degradation of Organic Pollutants: Mechanisms and Kinetics 307

particles, i.e., surface exposed, are totally illuminated. When the catalyst concentration is very high, after travelling a certain distance on an optical path, turbidity impedes further penetration of light in the reactor. In any given application, this optimum catalyst concentration [(TiO2)OPT] has to be found, in order to avoid excess catalyst and insure total absorption of efficient photons. Our results on the effect of catalyst concentration on the degradation rate for the TOC depletion and decomposition of pollutants under investigation are in agreement with numerous studies reported in the literature (Shifu and Yunzhang