International Journal of Photoenergy

International Journal of Photoenergy / 2014 / Article

Review Article | Open Access

Volume 2014 |Article ID 821674 | 21 pages | https://doi.org/10.1155/2014/821674

Recent Developments in Homogeneous Advanced Oxidation Processes for Water and Wastewater Treatment

Academic Editor: Vincenzo Augugliaro
Received31 May 2013
Accepted21 Oct 2013
Published26 Feb 2014

Abstract

This paper reports on recent developments in homogeneous Advanced Oxidation Processes (AOPs) for the treatment of water and wastewater. It has already been established that AOPs are very efficient compared to conventional treatment methods for degradation and mineralization of recalcitrant pollutants present in water and wastewater. AOPs generate a powerful oxidizing agent, hydroxyl radical, which can react with most of the pollutants present in wastewater. Therefore, it is important to discuss recent developments in AOPs. The homogeneous AOPs such as O3, UV/O3, UV/O3/H2O2, and UV/H2O2, Fe2+/H2O2, UV/Fe2+/H2O2 on the degradation of pollutants are discussed in this paper. The influence on the process efficiency of various experimental parameters such as solution pH, temperature, oxidant concentration, and the dosage of the light source is discussed. A list of contaminants used for degradation by various AOPs and the experimental conditions used for the treatment are discussed in detail.

1. Introduction

Wastewater is water that contains various pollutants, which means it cannot be used like pure water and should not be disposed of in a manner dangerous to humans, living organisms, and the environment. Water pollution has a serious impact on all living creatures, adversely affecting water use for drinking, household needs, recreation, fishing, transportation, and commerce. It has been estimated that the total global volume of wastewater produced in 1995 was in excess of 1,500 km3 [1]. On July 28, 2010, the United Nations General Assembly declared safe and clean drinking water and sanitation a human right essential to the full enjoyment of life and all other human rights [2]. It is a concern that nearly 900 million people in the world do not have access to safe drinking water. Approximately 1.5 million children under five die every year as a result of diseases linked to a lack of access to water and sanitation as indicated by World Health Organization (WHO) [3]. It was estimated that about 1.8 million deaths annually are due to lack of access to safe drinking water and poor sanitation.

In the past, economically viable chlorination has been used for water treatment. Yet the potentially adverse health effects of the by-products formed, together with raised drinking water standards, have led researchers to search for effective and economical alternatives to chlorinating drinking water [4, 5]. Various wastewater treatment processes have been tried using physical, chemical, and biological methods [612]. Some of these methods have disadvantages, however, and cannot be applied for large scale treatment. For example, one drawback of precipitation methods is sludge formation. Chemical coagulation and flocculation use a large amount of chemicals and the generated sludge may contain hazardous materials, so sludge disposal remains a problem. Adsorption techniques have been used widely for the removal of various water and wastewater pollutants. Their disadvantage is that the pollutants may only transfer to the adsorbent, which needs to be regenerated regularly, resulting in additional costs. Membrane technologies such as ultrafiltration, nanofiltration, and reverse osmosis have been used for the full scale treatment and reuse of water and chemicals. Yet these methods have several operational difficulties in addition to high capital costs. Thus physical methods may not be suitable for the complete removal of pollutants from the environment. Similarly, two different basic biological wastewater treatment methods have been employed: aerobic and anaerobic treatments. These methods also do not completely remove the high concentration of pollutants present in wastewater. Other biological methods involve cost-effectiveness or operational difficulties, making biological means unsuitable for wastewater treatment.

Among the chemical methods, oxidation is efficient and applicable to large scale wastewater treatment. Generally air, oxygen, ozone, and oxidants such as NaOCl and H2O2 are used for chemical treatment. The oxidation potential of some of the oxidants is listed in Table 1. The basic chemical oxidation process with air and oxygen also occurs in nature, but it is no longer sufficient for highly polluted wastewater. Therefore there is a significant need to develop a wastewater treatment process which can remove the pollutants effectively by a simple method.


SubstancePotential (V)

Hydroxyl radical (OH)2.86
Oxygen (O)2.42
Ozone molecule (O3)2.07
Hydrogen peroxide (H2O2)1.78
Chlorine (Cl2)1.36
Chlorine dioxide (ClO2)1.27
Oxygen molecule (O2)1.23

Advanced oxidation processes (AOPs) for wastewater treatment have received a great deal of attention in recent years. AOPs generate the highly reactive hydroxyl radical (OH) to degrade the recalcitrant chemicals present in wastewater [1315]. These OH radicals attack the most organic molecules rapidly and nonselectively. The versatility of AOPs is also enhanced by the fact that they offer various alternative methods of hydroxyl radical production, thus allowing a better compliance with specific treatment requirements. The eco-friendly end product is the special feature of these AOPs, which are more efficient as they are capable of mineralizing a wide range of organic pollutants. Interestingly, AOPs can make use of solar energy rather than artificial light sources. The latter rely on high electrical power, which is costly and hazardous.

AOPs such as ozonation (O3), ozone combined with hydrogen peroxide (O3/H2O2) and UV irradiation (O3/UV) or both (O3/H2O2/UV), ozone combined with catalysts (O3/catalysts), UV/H2O2, Fenton and photo-Fenton processes (Fe2+/H2O2 and Fe2+/H2O2/UV), and the ultrasonic process and photocatalysis have been successfully used for wastewater treatment [1623]. This review reports on recent advances in the aforementioned AOPs for water and wastewater treatment. The authors discuss the principle of hydroxyl radical generation from each AOP, the influence of various experimental parameters, and their consequences for the treatment process.

1.1. Ozone Based Advanced Oxidation Processes

(i) Ozonation. Ozone is an environmentally friendly oxidant since it decomposes into oxygen without producing self-derived by-products in the oxidation reaction. It is widely used in the purification of drinking water, the treatment of wastewater and process water, the sterilization of water in artificial pools, and so forth. In an ozonation process, two possible oxidizing actions may be considered. The first or direct method involves the reaction between ozone dissolved compounds. The second is known as the radical method because of the reactions between the hydroxyl radicals generated in ozone decomposition and the dissolved compounds [24]. Some oxidation products are refractory to further oxidative conversion by means of ozone, thus preventing a complete abatement of TOC. Yet the high energy cost of direct ozonation limits many practical applications. To increase the efficiency of the ozonation process, the ozone is combined with H2O2 and UV light, which is expected to increase the removal rate substantially by producing more hydroxyl radicals in the treatment system.

(ii) O3/H2O2 Process. Hydroxyl radicals are generated by a radical chain mechanism through the interaction between ozone and H2O2 as shown in (1). Degradation is facilitated by both ozone and hydroxyl radical:

(iii) O3/UV Process. Hydroxyl radicals are generated in the O3/UV process by photolysis of ozone in the presence of water as shown in the following:

(iv) O3/UV/H2O2. This combined process may generate hydroxyl radicals in different ways as mentioned in (1)-(2). It is considered to be the most effective treatment process for highly polluted effluents.

Wastewater was treated using the Fenton process or homogeneous AOP employing iron salt with hydrogen peroxide. The combination of Fenton’s reagent with UV light is called a photo-Fenton reaction. UV light irradiation enhances the efficiency of the Fenton process. The hydroxyl radical generated in the Fenton process is due to the iron catalysed decomposition of H2O2 as shown in the following: In addition to the above reaction the formation of hydroxyl radical also occurs by the following reactions in the photo-Fenton process shown in the following: This is also attributed to the decomposition of the photoactive Fe(OH)2+ which leads to the addition of the HO radicals: A considerable increase in oxidation power is observed mainly due to the photoreduction of Fe(III) to Fe(II), which can react with H2O2, establishing a cycle: Among various AOPs, Fenton’s reagent (H2O2/Fe2+) is one of the most effective methods of organic pollutant oxidation. The advantage of Fenton’s reagent is that no energy input is necessary to activate hydrogen peroxide. These processes are economic and can be operated and maintained easily.

1.2. UV/H2O2 Process

The UV/H2O2 process is a homogeneous advanced oxidation process employing hydrogen peroxide with UV light. Hydrogen peroxide requires activation by an external source such as UV light and the photolysis of hydrogen peroxide generates the effective oxidizing species hydroxyl radical (OH). The rate of photolysis of H2O2 depends directly on the incident power or intensity. The hydrogen peroxide decomposition quantum yield is 0.5 at UV (254 nm) irradiation. Solar light could also be used as a radiation source but the rate of photolysis may be low compared to UV light. In this process the dosage needs to be optimized, however, since excess H2O2 may scavenge hydroxyl radical.

1.3. Heterogeneous AOPs
1.3.1. Catalytic Ozonation Process

Heterogeneous catalytic ozonation is a novel type of AOP that combines ozone with the adsorptive and oxidative properties of solid phase catalysts to decompose pollutants at room temperature. Catalytic ozone decomposition at room temperature is advantageous compared to thermal decomposition in terms of energy conservation since it does not require large volumes of air to be heated. It is therefore a promising advanced oxidation technology for water treatment.

1.3.2. Photocatalysis

Heterogeneous photocatalysis through illumination by UV or visible light on a semiconductor surface generates hydroxyl radicals. The photocatalyst can be used successfully for the effective treatment of pollutants in water and wastewater.

2. Ozone Based AOPs

As noted above, ozone reacts with various organic and inorganic compounds in an aqueous solution, either by direct reaction of molecular ozone or through a radical mechanism involving hydroxyl radical induced by the ozone decomposition. Figure 1 shows the experimental setup of the ozonation process. This process is strongly influenced by a number of experimental parameters such as solution pH, influent ozone dosage rate, and temperature. The primary reactions initiated by ozone in water are strongly pH dependent. Ozone reacts with organic substrate at low pH as a molecular form, but at high pH it decomposes before reacting with the substrate. Ozone decomposition is catalyzed by hydroxide ions and proceeds more rapidly with increasing pH, eventually to produce hydroxyl radicals. The influence of solution pH on ozonation process efficiency has been observed in a number of studies. For example, Jung et al. investigated the effect of pH on the ozonation of ampicillin from pH 5 to 9, concluding that higher pH conditions are necessary for effective removal [66]. They also discussed how changing pH influences the charge of some specific functional groups on the ozonation process. Can and Gurol investigated the effect of solution pH on the ozonation of humic substances. They found that rapid ozone decomposition was caused by the interaction of ozone with the humic substance, which eventually yielded hydroxyl radical [67]. They further noted that increasing humic substance concentration facilitates fast ozone decomposition into hydroxyl radical. Similarly, the influence of solution pH and temperature on the ozonation of six dichlorophenols was investigated by Qiu et al. [68]. They revealed that the changing solution pH was strongly influenced the decomposition and the rate was increased by raising the hydroxyl radical concentration from acidic to alkaline pH [68].

Although hydroxyl radical formation is highly favourable to produce more OH radicals by ozone self-decomposition at pH 10, a portion of carbonate or bicarbonate ion formation could play a key scavenging role in trapping OH radicals, appreciably decreasing the degradation rate. Wu et al. found that 2-propanol degradation decreases at pH 10 and suggested bicarbonate formation as the possible reason for the decreasing degradation rate at this pH [19]. Other studies reached quite different results. Moussavi and Mahmoudi noted a higher removal rate of Reactive Red 198 azo dye in an ozonation process at pH 10 [69]. Interestingly, Begum and Gautam noted that as the pH increased from 9 to 12 in the ozonation process the endosulfan and lindane removal rate also increased [70]. In contrast to the above results other authors noted that the oxidation rate is relatively independent of solution pH values [71]. Hong and Zeng found that the rates of pentachlorophenol decomposition were very similar between pH 7 and 12, indicating then negligible influence of pH values [72]. These results clearly showed that the nature of pollutants being used for the ozonation process played an important role besides the favourable hydroxyl radical formation at higher pH. Based on the above discussion it is concluded that the influence of pH on the ozonation process needs to be optimized.

Several investigations were conducted into the effect of temperature on the ozonation process. Changing the temperature generally influences the ozonation process in two ways. Firstly, when the temperature increases, the solubility of ozone may decrease, since Henry’s law coefficient of ozone increases with rising temperatures. Secondly, raising the temperature increases the activation energy which may positively assist the ozonation process. Muruganandham et al. noted that N-methyl pyrrolidone (NMP) mineralization was substantially increased when the ozonation temperature rose from 5 to 50°C [13]. They also concluded that the increasing removal rate due to the higher reaction temperature is not balanced by the lower solubility of ozone. Similar results were noted in other ozonation studies [68, 7376]. Some researchers found, however, that increasing temperature in the ozonation process decreases the removal rate by decreasing the ozone solubility [77, 78]. Interestingly, Ku et al. found that the reaction rates of phorate decomposition were relatively independent of solution temperatures and pH values [71]. Yet some mineralization formation of products such as phosphate and carbonate was increased significantly with raised solution temperature.

Another important experimental parameter influencing ozonation process efficiency is influent ozone dosage. Treatment cost increases with a higher applied ozone dose, so it is necessary to optimize this dosage. For semibatch experiments, increasing the ozone dosage will enhance the mass transfer rate of ozone from the gas phase to the liquid phase, which is expected to enhance the degradation rate appreciably. As the ozone concentration in the liquid phase is saturated, however, ozone mass transfer is limited at a very high ozone dosage [79]. Many authors investigated how the influent ozone dosage affects the degradation rate in the ozonation process within different experimental parameters. Muruganandham et al. reported that the optimal ozone dosage for NMP mineralization is 18.4 mg min−1 [13]. Moreover, an ozone dosage of 27.6 mg min−1 was noted as optimum for the degradation of dimethyl sulphoxide (DMSO) [80]. Begum and Gautam reported an optimum ozone dosage of 57 mg min−1 for endosulfan and lindane degradation although a higher ozone dosage slightly increased the endosulfan decomposition [70]. Yet other studies reported a linear increase in removal efficiencies with ozone dosage [81]. The above discussion clearly indicates that ozone dosage needs to be optimized in an ozonation process and that a number of experimental factors could influence the removal rate.

Though the ozonation process is effective for treating some organic compounds, a key problem is the accumulation of refractory compounds which interfere with the mineralization of the organic matter present in water. Some compounds were even found to be refractory to the ozonation process [15, 82, 83]. To improve its efficiency, ozonation was therefore combined with other oxidants. The combination of single oxidants can offer very effective treatment by producing more hydroxyl radicals.

It was reported that ozone in the presence of UV light enhances the decomposition rate of pollutants present in wastewater. The hydroxyl radicals generated in the UV/O3 process are shown in (2). Decomposition may proceed in three different ways: (i) by ozonation, (ii) by direct UV photolysis, and (iii) by photolysis of ozone which generates hydroxyl radicals. The detail of the ozone photolysis mechanism is shown in (6)–(10) [84]. These combined processes should synergistically accelerate the degradation rate in the UV/O3 process compared to the ozonation process. Many authors found that the UV/O3 process is more efficient than ozonation in organic compounds degradation and mineralization [8588]:

Recent studies also combined H2O2, and TiO2 with the UV/O3 process [8991]. These processes are more efficient than O3/UV alone due to their synergistic effect. The dosage of H2O2 in the O3/H2O2 process needs to be optimized, however. For example, Medellin-Castillo et al. studied diethyl phthalate degradation using the O3/H2O2 process with 0.45 to 1.80 mM of H2O2, noting a linear relation with the degradation rate [25]. Kwon et al. studied 1,4-dioxane degradation with H2O2/O3 (w/w) ratios of 0, 0.25, 0.5, 0.75, 1, 1.25, and 1.5. They found the optimum dosage ratio to be 0.5 and noted a strong retardation effect at a ratio of 1.5 [26]. Similar results were also reported by Kusic et al. for the mineralization of phenol [27]. This could be because an excess of hydrogen peroxide could react with the hydroxyl radical produced during the decomposition process as shown in (11) and (12). So the H2O2 dosage needs to be optimized in the degradation process. Many authors have used various combined AOPs for pollutant degradation and their conclusions are summarized in Table 2. Consider


ReferenceAOPs applied Pollutant(s)Conclusions

[25]O3, UV/H2O2, O3/H2O2, O3/ACDiethyl phthalate (DEP) in ultrapure water, surface water, and wastewaterThe O3/AC process was the most efficient for the removal of DEP in all three types of water. The O3/H2O2 and O3/AC processes are more efficient than ozonation alone.

[26]O3, UV photolysis, O3/UV, H2O2/O3, O3/H2O2/UV1,4-DioxaneThe O3/H2O2/UV process was most efficient for 1,4-dioxane removal at pH 10, with H2O2 : O3 ratio of 0.5.

[27]O3, O3/H2O2, UV/H2O2, UV/O3 UV/H2O2/O3PhenolThe UV/H2O2/O3 process at pH 7 with H2O2 = 10 mM was most ecoeffective with 100% of phenol removal within 30 min and 58.0% TOC removal after 1 h. UV/H2O2/O3 was the most effective process for phenol wastewater mineralization.

[28]O3, O3/H2O2, UV/H2O2Twenty-four micropollutants including endocrine disrupting compounds, pharmaceuticals, and personal care productsThe general trend of ozone and hydroxyl radical reactivity with the selected micropollutants was explained. Suitable technology for the removal of these micropollutants was suggested based on the micropollutant reactivity with ozone and hydroxyl radical.

[29]O3, UV photolysis, O3/UV, O3/catalyst, UV/catalyst, O3/UV/catalyst, H2O2/UV, H2O2/UV/catalystPyruvic acidThe UV/H2O2 process with or without perovskite catalysts facilitates pyruvic acid removal fastest. The O3/UV/perovskite process was efficient for mineralization.

[30]UV, UV/H2O2, UV/O3, UV/H2O2/O3,
UV/Fenton
p-ChlorophenolOperating conditions such as initial pH, concentration of H2O2, and ferrous salt were optimized for each process. The UV/Fenton and UV/H2O2/O3 processes were found to be the most effective for the degradation and mineralization of p-CP.

[31]O3, H2O2/O3, UV/ /H2O2Nitroaromatics such as
nitrotoluene, dinitrotoluene, and trinitrotoluene
Ozonation and/or Fenton’s reagent were found to be efficient for TNT degradation. The O3/H2O2 process at pH > 7 was most efficient for 2-MNT and 2.4-DNT removal.

[32] O3/UV, H2O2/UV, O3/H2O2, O3/H2O2/UVHaloacetic acids (HAAs), dichloroacetic acid (DCAA), and trichloroacetic acid
(TCAA)
The O3/UV process was the most efficient of the six degradation methods for DCAA and TCAA in water. Decomposition by AOPs was easier for DCAA than for TCAA.

[33]O3, O3/UV, O3/H2O2, UV/H2O2, O3/UV/H2O2O-NitrotolueneThe optimum H2O2 dosage and solution pH were studied. Adding H2O2 to the ozonation process accelerated the oxidation of O-nitrotoluene by a factor of 8. The O3/UV and UV/H2O2 processes are 20 and 10 times more efficient than the ozonation process, respectively.

[34]O3, UV-vis,
O3 + UV-vis, TiO2 + UV-vis, O3 + UV-vis + TiO2, and O3 + TiO2
Phenol,
4-nitrophenol, and
4-chlorophenol
AOP efficiencies are in the following order: adsorption < TiO2 + UV-vis < UV-vis < O3 + TiO2 O3 < O3 + UV-vis O3 + UV-vis + TiO2. The O3 + UV-vis and O3 + UV-vis + TiO2 methods are the most economically attractive.

[35]O3, O3/H2O2, O3/activated carbonAcid Blue 92 (AB92)Ozone treatment was a very effective method for complete removal of colour but in COD removal it was not efficient. The removal of COD in ozonation, O3/H2O2, and O3/AC processes, 30%, 80% and 100%, respectively.

[36]O3 or O3/H2O2, O3/powdered activated carbon (PAC)Sodium Dodecylbenzenesulfonate(SDBS)Comparison of the O3/PAC system with the O3 and O3/H2O2 processes showed that the O3/PAC system was more effective in the removal of SDBS.

[37]O3, O3/UV, UV/H2O2 Dye house effluentThe AOP efficiency is dependent on the pH and dosage of H2O2. The UV/H2O2 process is 50 times more efficient than the O3/H2O2 process.

[38]Ozonation, sonication, UV photolysis, O3/ultrasound, UV/ultrasound, O3/UV/ultrasoundPhenolThe efficiency of the various AOPs at two different pH was in the following order. At pH 2, US/UV/O3 O3/UV > US/O3 > US/UV > O3 > US > UV.
At pH 10.0, US/UV/O3 > O3/UV > O3 > US/O3 > US/UV > UV > US. The maximum rate of phenol degradation was observed in the combined application of O3/UV/US at basic pH.

[39]US
O3
O3/US
O3/UV
US/UV
O3/US/UV
Acid Orange 7The UV/O3 process was more effective at all times than the US and/or O3 process. The O3/US/UV process was the most efficient for colour and aromatic removal and AO7 dye mineralization.

The presence of transition metal ions such as Mn2+, Co2+, Ag+, and Fe2+ in the ozonation process has significant catalytic effects in producing hydroxyl radical [92, 93]. Abd El-Raady and Nakajima studied the degradation of formic, oxalic, and maleic acids in the presence of first row transition metal ions such as Co2+, Ni2+, Mn2+, Cu2+, Zn2+, Cr3+, and Fe2+ and compared the process efficiency with the O3 and O3/H2O2 processes [93]. They concluded that the presence of Co2+ and Mn2+ ions has the highest catalytic activity for the decomposition of oxalic acid and that O3/Co2+ and O3/Mn2+ are more efficient than the O3/H2O2 process. Similarly, Cortes et al. reported that the O3/Mn2+ and O3/Fe2+ processes were more effective in the removal of organochloride compounds than the O3/Fe3+ and O3/high pH systems [92]. Beltrán et al. found that the presence of Co2+ in water significantly enhances the ozonation rate of oxalic acid at acidic pH and that catalytic ozonation proceeds through the formation of a Co(HC2O4)2 complex [94]. Heterogeneous catalytic ozonation has received increasing attention due to its potentially higher effectiveness in the degradation of recalcitrant pollutants [95101].

3. Fenton and Photo-Fenton Based AOPs

3.1. Fenton Reaction

The Fenton process has its root in the finding reported in 1894 that ferrous ion strongly elevated the oxidation of tartaric acid by hydrogen peroxide [102]. In the Fenton process, hydrogen peroxide is added to wastewater in the presence of ferrous salts, generating species that are strongly oxidative with respect to organic compounds. OH is traditionally regarded as the key oxidizing species in Fenton processes. The Fenton process mechanism is quite complex and is described in detail with equations in the literature [103, 104]. In summary, the classic Fenton free radical mechanism in the absence of organic compounds mainly involves the following sequence of reactions [102, 105]:

OH radicals are rapidly generated through (14). In the above reactions, iron cycles between Fe2+ and Fe3+ and plays the role of catalyst. The net reaction of (14)–(20) is the decomposition of H2O2 into water and O2 catalyzed by iron as follows (21):

As iron(II) acts as a catalyst, it has to be regenerated, which seems to occur through the following scheme [102]

Generally speaking, Fenton’s oxidation process is composed of four stages including pH adjustment, oxidation reaction, neutralization and coagulation, and precipitation. The organic substances are removed at two stages of oxidation and coagulation [106, 107]. OH radicals are responsible for oxidation, and coagulation is ascribed to the formation of ferric hydroxo complexes [107, 108]. The relative importance of oxidation and coagulation depends primarily on the H2O2/Fe2+ ratio. Chemical coagulation predominates at a lower H2O2/Fe2+ ratio, whereas chemical oxidation is dominant at higher H2O2/Fe2+ ratios [107, 109]. Wang et al. [110] and Lau et al. [111] reported that, in Fenton treatment of biologically stabilized leachate, oxidation and coagulation were responsible for approximately 20% and 80% of overall COD, removal respectively. Fenton oxidation has been tested with a variety of synthetic wastewaters containing a diversity of target compounds, such as phenols [112114], chlorophenols [115], formaldehyde [116], 2,4-dinitrophenol [116], 2,4,6-trinitrotoluene [117], 2,4-dinitrotoluene, chlorobenzene, tetrachloroethylene [118], halomethanes, amines, and hexahydro-1,3,5-trinitro-1,3,5-triazine (RDX) [119]. Many chemicals are refractory to Fenton oxidation, however, such as acetic acid, acetone, carbon tetrachloride, methylene chloride, oxalic acid, maleic acid, malonic acid, n-paraffins, and trichloroethane [116]. It has been demonstrated that these compounds are resistant under the usual mild operating conditions of Fenton oxidation [114, 120, 121]. In addition to these basic studies, the process has been applied to industrial wastewaters (such as chemical, pharmaceutical, textile, paper pulp, cosmetic, and cork processing wastewaters), sludge, and contaminated soils [122] resulting in significant reductions of toxicity, improvement of biodegradability, and colour and odour removal [116].

The oxidation rate was influenced by many factors such as pH value, Fe2+ : H2O2 ratio, and the amount of iron salt. Some of these parameters are discussed in detail in the following sections. The Fenton process seems to be the best compromise because it is technologically simple, there is no mass transfer limitation (homogeneous nature), and both iron and hydrogen peroxide are cheap and nontoxic. From the economic point of view, using the Fenton process as a pretreatment can lower the cost and improve biological treatment efficiency [107].

A batch Fenton reactor essentially consists of a pressurized stirred reactor with metering pumps for the addition of acid, a base, a ferrous sulphate catalyst solution and industrial strength (35–50%) hydrogen peroxide. It is recommended that the reactor vessel be coated with an acid resistant material, because Fenton’s reagent is very aggressive and corrosion can be a serious problem. The pH of the solution must be adjusted to maintain the stability of the catalyst, as at pH 6 iron hydroxide is usually formed. For many chemicals the ideal pH for the Fenton reaction is between 3 and 4, and the optimum catalyst to peroxide ratio is usually 1 : 5 wt/wt. Reactants are added in the following sequence: wastewater followed by dilute sulphuric acid catalyst in acidic solutions, base or acid for the adjustment of pH at a constant value, and lastly hydrogen peroxide (which must be added slowly, maintaining a steady temperature). Since wastewater compositions are highly changeable, there are some design considerations to enable the Fenton reactor to operate within flexible parameters. The discharge from the Fenton reactor is fed into a neutralizing tank to adjust the pH of the stream, followed by a flocculation tank and a solid-liquid separation tank for adjusting the TDS (total dissolved solids) content of the effluent stream. A schematic representation of the Fenton oxidation treatment is shown in Figure 2 [123].

As mentioned above, Fenton oxidation was applied to wastewater treatment based on the following observed optimum pH conditions, since this has been shown to affect the degradation of pollutants significantly [106, 124, 125]. The best value pH has been observed to be 2.8–3 in the majority of cases; [116, 126, 127], hence this is the recommended operating pH. At lower pH (pH = 2.5), the formation of (Fe(II) (H2O))2+ occurs, which reacts more slowly with hydrogen peroxide, producing a smaller amount of reactive hydroxyl radicals by reducing the degradation efficiency [123].

Furthermore, the scavenging effect of hydroxyl radicals by hydrogen ions becomes important at a very low pH, at which the reaction of Fe3+   with hydrogen peroxide is also inhibited. At an operating pH of >3, the decomposition rate decreases because of the decreased free iron species in the solution, probably due to the formation of Fe(II) complexes with the buffer inhibiting the formation of free radicals. At a pH higher than 3, Fe3+ starts precipitating as ferric oxyhydroxides and breaks down the H2O2 into O2 and H2O [124, 128], inhibiting the generation of ferrous ions. Additionally, the oxidation potential of OH radical is known to decrease with an increase in pH [123].

Usually the rate of degradation increases with an increased concentration of ferrous ions [125], though the increase is sometimes observed to be marginal above a certain concentration [106, 129]. Additionally, an enormous increase in ferrous ions will lead to an increased unutilized quantity of iron salts, contributing to increased TDS content in the effluent treatment, which is not permitted. Thus laboratory scale studies are required to establish the optimum loading of ferrous ions under similar conditions, unless data are available in the open access literature [123].

The concentration of hydrogen peroxide plays a more crucial role in the overall efficacy of the degradation process. Usually it has been observed that the percentage degradation of the pollutant increases with an increased dosage of hydrogen peroxide [106, 129]. Care should be taken however in selecting the operating oxidant dosage. The residual hydrogen peroxide contributes to COD, so an excess amount is not recommended. The presence of hydrogen peroxide is also harmful to many microorganisms and affects the overall degradation efficiency significantly where Fenton oxidation is used as a pretreatment to biological oxidation. One more negative effect of hydrogen peroxide, if present in large quantities, is that it acts as a scavenger for the generated hydroxyl radicals. Thus hydrogen peroxide loading should be adjusted so that the entire amount is utilized. This can be decided based on laboratory scale studies with the effluent in question [123].

It should be noted that the dose of H2O2 and the concentration of Fe2+ are two relevant and closely related factors affecting the Fenton process. The H2O2 dose has to be fixed according to the initial pollutant concentration. An amount of H2O2 corresponding to the theoretical stoichiometric H2O2 to chemical oxygen demand (COD) ratio is frequently used [116], although it depends on the response of the specific contaminants to oxidation and on the objective pursued in terms of reducing the contaminant load. Usually a lower initial pollutant concentration is favoured [125], but the negative effects of treating a large quantity of effluent need to be analyzed before the dilution ratio can be set. For real industrial wastes, some dilution is often essential before any degradation is observed using Fenton oxidation [123].

As noted above, as the maximum degradation rates are observed at a pH of approximately 3, the operating pH should be maintained constant around this optimum value. The type of buffer solution used also affects the degradation process [125]. Acetic acid/acetate buffer provides maximum oxidation efficiency, at least as observed for phosphate and sulphate buffers. This can be attributed to the formation of stable Fe3+ complexes under these conditions [123].

Not many studies are available depicting the effect of temperature on degradation rates and ambient conditions can safely be used with good efficiency [123]. Besides, reaction temperature is another crucial parameter in the Fenton process. In principle, increasing the temperature should enhance the kinetics of the process, but it also favours the decomposition of H2O2 towards O2 and H2O. This increases at a rate of around 2.2 times each 10°C in the range of 20–100°C [130]. Oxidation with Fenton’s reagent has already been proved effective and promising for the destruction of several compounds and consequently for the treatment of a wide range of wastewaters, as described in several reviews (e.g. [102, 104, 116, 123, 131, 132]). Table 3 summarizes recent Fenton processes for some wastewater treatments.


ReferenceProcess conditionsPollutant(s)Conclusions

[40]A temperature controllable magnetic stirrer ensures perfect mixing at a constant rate of 300 rpm during all experiments. The effect of Fe2+ concentration on COD removal varied in the range of 0.5–10 mM (these factors were kept constant: H2O2 = 30 mM; pH = 3;  min; COD = 2741 mg/L). The selected H2O2 concentration was in the range of 10–100 mM while pH = 3 and Fe2+  = 10 mM at 30 min. The tested pH values ranged between 2 and 5.Synthetic acid dye baths (SADB) consist of three different acid dyestuffs (C.I. Acid Yellow 242, C.I. Acid Red 360, and C.I. Acid Blue 264) and two dye auxiliaries (a levelling agent and an acid donor)Optimum experimental conditions for the simulated acid dye bath effluent were established as follows: Fe2+ = 10 mM, H2O2 = 30 mM, and pH = 3 at room temperature ( °C), which yielded an overall COD removal efficiency of 23%. The corresponding colour removal efficiency was 92% and the first-order COD abatement rate constant increased from 0.02 min−1 to 0.03 min−1 by increasing the temperature from 20 to 50°C. The first-order reaction rate constant for H2O2 consumption increased from 0.15 min−1 to 0.34 min−1 by increasing the temperature from 20 to 50°C. Further increases in temperature did not improve oxidation and oxidant consumption rates. H2O2 consumption ran parallel to COD removal at a rate approximately 10 times faster than COD abatement.

[41]The Fenton reactor was stirred at room temperature in an open-batch system with a magnetic stirring bar and was treated for 2 h. The Fe+2 : H2O2 ratio was varied in the range of 1 : 5, 1 : 10, 1 : 20, 1 : 30, 1 : 40, and 1 : 50, pH in the range of 2–4, and Fe2+ in the range 0.5 and 1 mM.RB49 Reactive Blue 49
RB137 Reactive Blue 137
The Fenton process was decolourized more than 90% in all cases. The best mineralization extent, that is, maximal TOC removal, 72.1%, was obtained for degradation of RB49 by Fenton process, Fe2+ : H2O2 = 1 : 20, Fe2+ = 0.5 mM at pH = 3. The molecular structure of the dyes studied plays a significant role in oxidation by Fenton type processes.

[42]The oxidation studies were conducted in brown 500 mL glass bottles. The pH of wastewater and bleach was first adjusted to 3 with H2SO4. Degradation of EDTA in distilled water was conducted by Fenton’s reagent with Fe concentrations 0–0.9 mM and a maximum reaction time of 15 min. The temperature reaction and pH were fixed at 60°C and 3, respectively.Ethylenediamine tetra acetic acid (EDTA), novel complexing agents, namely, BCA5 and BCA6Fenton’s process proved highly effective in the degradation of EDTA in spiked integrated wastewater. With an initial molar ratio of 70 : 1 (H2O2 and EDTA) or higher, EDTA degradation was nearly complete within 3 min of reaction time. Lower EDTA degradation levels at pH 4 and low temperature in bleaching effluent are a major drawback in this study.

[43]The initial concentrations of Fe(II) used in this study were 8.37, 13.95, 19.53, 25.11, and 33.40 mg/L, the Fe2+ : H2O2 ratios were set at 0.016, 0.028, 0.039, 0.05, and 0.067, and the concentration of H2O2 was kept constant at 500 mg/L. The initial concentrations of H2O2 used in this study were 50, 100, 200, 500, and 700 mg/L, the Fe2+ : H2O2 ratios were set at 0.0199, 0.0279, 0.06975, 0.1395, and 0.279, and the concentration of Fe(II) was fixed at 13.95 mg/L.Azo dye C.I. Acid Yellow 23 (AY 23)The decolourization rate is strongly dependent on the initial concentrations of Fe2+ and H2O2. The optimum operational conditions were obtained at pH 3. The results show that as much as 98% of AY 23 can be decolourized by 13.95 mg/L ferrous ions and 500 mg/L H2O2.

[44]All tests were conducted in a 200 mL double glass cylindrical jacketed reactor, which allows cycle water to maintain the reaction mixture at a constant temperature. Temperature control was realized through a thermostat and a magnetic stirrer was used to stir reaction solutions. Operating pH was in the range of 2.5–6.0 and decolouration time was 60 min. Hydrogen peroxide in the range of to  M and the Fe2+ dosage on the decolourization of OG with different initial concentrations from to  M. Reaction temperature was varied in the range of 20–50°C. The effect of the presence of chloride ion ( to  M) on the decolourization of OG was investigated. The decolourization of different concentrations of OG was studied in the range of to  M.Azo dye Orange G (OG)The results showed a suitable decolourization condition of initial pH 4.0, H2O2 dosage  M, and molar ratio of H2O2 / Fe2+ 286 : 1. The decolourization efficiencies within 60 min were more than 94.6%. It was found that the decolourization efficiency of OG enhanced with increased reaction temperature but the presence of chloride ion had a negative impact on the decolourization of OG. The decolourization kinetics of OG by Fenton oxidation process followed the second-order reaction kinetics, and the apparent activation energy was detected to be 34.84 kJ/mol.

[45]Chemical oxidation of the red dye solutions with Fenton’s reagent was carried out in a closed jacketed batch reactor (1 L capacity). The reactor was provided with constant stirring, accomplished through a magnetic bar and a Falc magnetic stirrer. The temperature of the reaction mixture was kept constant by coupling the reactor to a Huber thermostatic bath. Operating pH and H2O2 concentration were varied in the range of 2–5 and 5.9–8.8 mM, respectively. The effect of the Fe2+ concentration and reaction temperature was investigated in the range of 0.13–1.1 mM and 20–70°C, respectively.Azo dye (Procion Deep Red H-EXL gran)Total organic carbon (TOC) reduction occurred after 120 min of reaction; however, the reaction time required to achieve colour removal levels above 95% is around 15 min. Four operating variables must be considered, namely, the pH, the concentration of hydrogen peroxide, the temperature, and the concentration of ferrous ion, between 3-4, 5.9 mM, 20 min, and 0.27 mM, respectively. It was concluded that temperature and ferrous ion concentration are the only-variables that affect TOC removal, and, due to cross interactions, the effect of each variable depends on the value of the other one, thus affecting the process response positively or negatively.

[46]Fenton’s reagent experiments were carried out at room temperature ( °C) using different H2O2 and Fe(II) doses at pH 3.5. The percentage variation of simazine removal was investigated with H2O2 concentration at different simazine doses between 0.5 and 5.0 mg/L and at different Fe(II) doses between 5 and 30 mg/L at the end of a 6 min reaction time.SimazineAt a constant simazine concentration, the percentage of TOC removal increased with increasing H2O2 and Fe(II) concentrations up to 15 mg/L Fe(II) and 50 mg/L peroxide above which mineralization decreased due to the scavenging effects of H2O2 on hydroxyl radicals. Maximum pesticide (100%) and TOC removals (32%) were obtained with H2O2/Fe(II)/simazine ratio of 55 : 15 : 3 (mg/L). Simazine degradation was incomplete, yielding the formation of intermediates which were not completely mineralized to CO2 and H2O.

[47]
The experiments were performed in an insulated vessel with a capacity of 1 L mounted on a steel frame and stirred at 130 rpm. The pH of initial solutions was set at 3. Gradation efficiencies were compared by varying Fenton’s reagent concentration and ratios. The parallel monitoring of Fenton’s reagent concentrations allowed the evidencing hydrogen peroxide or ferrous ion contents as limiting factors for TNT removal. The H2O2 0/ Fe(II) 0 ratio was varied in the range of 0.1–2 mM.TNTFenton oxidation is an effective method to transform TNT totally in contaminated aqueous solution. This is feasible by the efficient generation of hydroxyl radicals during H2O2 catalytic decomposition with Fe(II) ions. TNT degradation kinetics and efficiency are largely influenced by H2O2 and Fe2+ concentrations. Using [H2O2]0 : [Fe(II)]0 molar ratios equal to or lower than 0.5 leads to the formation of the maximum number of intermediates. The absolute rate constant of the reaction between hydroxyl radicals and TNT is 9.6– M−1 s−1.

[48]The Fenton reactor was a 0.5 L beaker placed in a thermostat water bath with constant temperature and stirred by a magnetic stirrer, with operating pH values of 2.50, 3.00, 3.50, 4.00, and 5.00, initial H2O2 concentration in the range of 0.10 mM to 4.00 mM, initial concentration of Fe2+ from 0.01 mM to 0.10 mM, and initial Amido Black 10B concentration on its degradation in the range of 10–100 mg/L. A series of experiments were conducted by varying the temperature from 15°C to 45°C.Azo dye Amido Black 10BThe optimal operation parameters for the Fenton oxidation of Amido Black 10B were 0.50 mM H2O2 0 and 0.025 mM Fe2+ 0 for 50 mg/L [dye]0 at an initial pH of 3.50 at a temperature of 25°C. Under these conditions, 99.25% dye degradation efficiency in aqueous solution was achieved after 60 min of reaction. The Fenton treatment process showed that it was easier to destruct the –N = N-group than to destruct the aromatic rings of Amido Black 10B.

[49]Fenton oxidation was performed in a batch reactor under initially anaerobic conditions to determine the effect of [MTBE]0 on the degradation of MTBE with FR: MTBE degradation at different [MTBE]0 in the range of 1, 2, and 5 mg/L when treated with the same amount of FR. This study was performed using solutions containing [MTBE]0 of 11.4 and 22.7 mM, each one in individual experiments at pH values of 3.0, 3.6, 5.0, 6.3, and 7.0. The FR to MTBE molar ratio varied in the range of 0.5 : 1 and 200 : 1. The initial concentration of pollutant was 22.7  M and FR was used in a 1 : 1 molar ratio of ferrous iron (Fe2+) and hydrogen peroxide (H2O2) at pH = 3. Methyl tert-butyl ether (MTBE)FR partially degraded low MTBE 0 in water (11.4 and 22.7  M). Experiments at acidic pH yielded the best results of MTBE degradation (>90%), and small differences were observed between the results at pH 3.0 and 5.0. The majority of MTBE degradation and generation of intermediates occurred during the initial phase and followed pseudo first-order kinetics.

[50]The experiments were conducted in batch mode. 4L borosilicate reactors were filled with 3.6 L of deionized (DI) water at pH = 3.0 and purged with high-purity nitrogen until the dissolved oxygen (DO) reading was below 0.01 mg/L and the oxygen concentration in the head space was negligible ( 0.01%).Methyl tert-butyl ether (MTBE)The added amount of FR proved to be an important controlling parameter for the overall MTBE degradation mineralization efficiency. An FR to MTBE molar ratio of 20 : 1 was the minimum required to achieve complete MTBE degradation. Kinetic analysis is reported to be pseudo first-order given the good linear correlation found between and FMMR. Other intermediates not identified in this study are generated in significant concentrations at these conditions.

[51]A series of experiments were conducted at pH 3 for 5, 15, or 60 min of mixing followed by 30 min clarification. The studied H2O2/Fe2+ stoichiometric molar ratios were 1, 2, 3, 4, 5, and 10 with H2O2 dose of 1000 mgL−1, and the H2O2/Fe2+ stoichiometric molar ratios were 0.5, 2, 3, 5, and 10 with H2O2 dose of 500 mgL−1. A further series of experiments were conducted at an initial pH of 3, 4, 5, 6, or 7 with 5 min mixing followed by 30 min clarification. Comparisons between the Fenton process and Fe3+ coagulation were carried out at an initial pH of 3 and 7.Nuclear laundry waterThe experimental data generally indicated decreased removal efficiencies of organic compounds with an increasing H2O2/Fe2+ ratio. Yet taking into account all factors, thermostat cost-effective degradation conditions were at H2O2/Fe2+ stoichiometric molar ratio of 2 with 5 min mixing and an H2O2 dose of 1000 mgL−1. The initial pH of the laundry water can be as high as 7. Fe3+  coagulation experiments were conducted in order to interpret the nature of the Fenton process. Since the removal efficiency of organic compounds in the Fenton process was slightly higher than in coagulation, the treatment of the nuclear laundry water can be called Fenton-based Fe3+ coagulation.

3.2. Photo-Fenton Processes

The photo-Fenton process, as its name suggests, is rather similar to the Fenton one but also employs radiation [102, 104, 123, 133]. The photo-Fenton reaction is also well known in the literature [104, 134] as an efficient and inexpensive method of wastewater and soil treatment [104, 135]. Photo-Fenton process is known to be capable of improving the efficiency of dark Fenton or Fenton-like reagents by means of the interaction of radiation (UV or Vis) with Fenton’s reagent [136]. This technique has been suggested as feasible and promising for removing pollutants from natural and industrial waters and increasing the biodegradability of chlorophenols when used as a pretreatment method to decrease water toxicity [104]. Some of its most innovative applications include oxalate as a ligand of iron ions [104, 137].

The effectiveness of photo-Fenton processes is attributed to the photolysis of Fe(III) cations in acidic media yielding Fe(II) cations (24), in conjunction with reaction between Fe(II) and OH to yield hydroxyl radicals (Fenton’s reaction, step 24):

In this process, the photolytic decomposition of Fe(OH)2+ (24) is accelerated, providing an additional source of highly oxidative hydroxyl radicals compared to the “simple” Fenton process [102]. The photo-Fenton process produces more hydroxyl radicals than the conventional Fenton method (Fe(II) with hydrogen peroxide) or photolysis, thus promoting organic pollutant degradation rates. This process consists of a combination of Fenton reagents (Fe2+/H2O2) and light energy [138, 139] and thus of two reactions [136]:

The first reaction is a reaction of Fe2+ with H2O2, which generates the powerful reactive species OH radicals and oxidizes Fe2+ to Fe3+. In other words, the hydroxyl radical generation in Fenton processes is due to the iron catalyzed decomposition of H2O2. The first photo-Fenton reaction causes the formation of hydroxyl radicals. The second reaction of the photo-Fenton process is a reaction of Fe3+ with water, which occurs when light is used at a wavelength from 300 nm to 650 nm. This generates OH radicals and reduces Fe3+ to Fe2+. These two oxidation-reduction reactions occur repeatedly and completely mineralize organic pollutants to CO2 and H2O [136]:

The oxidation power of the photo-Fenton process is attributed to the generation of OH radicals. Without irradiation, a Fenton-like reaction occurred instead of a photo-Fenton reaction. The Fenton-like reaction is a reaction of Fe3+ with H2O2, which causes the reduction of Fe3+ to Fe2+:

Since Reaction (30) occurs instead of Reaction (27), organic pollutants are mineralized even without irradiation. It should be noted, however, that Reaction (30) is rather slower than Reaction (27). Thus the degradation rate under dark conditions is rather lower than that of the photo-Fenton reaction [136]. Figure 3 shows the reaction pathways for the process starting with the primary photoreduction of the dissolved Fe(III) complexes to Fe(II) ions followed by Fenton’s reaction and the subsequent oxidation of organic compounds. Additional hydroxyl radicals generated in the first step also take part in the oxidation reaction [140].

Appropriate implementation of the photo-Fenton treatment depends mainly on the operating variables—H2O2/COD molar ratio, H2O2/Fe2+ molar ratio, and irradiation time. The conventional method is to optimize the operating variables by changing one factor at a time; that is, a single factor is varied while all other factors are kept unchanged for a particular set of experiments. Likewise, other variables are individually optimized through single-dimensional searches, which are time consuming and incapable of reaching the actual optimum as interaction among variables is not taken into consideration [141]. Some illustrative works from recent years are discussed in detail in Table 4.


ReferenceProcess conditionsPollutant(s)Conclusions

[52]Natural pH conditions with phenol concentrations in the range of 180–733 mg/L. The photochemical treatment was mediated with ferrioxalate and peroxide in two photoreactors of different volumes and operation conditions (batch and with closed flow).WastewaterPhenol transformation efficiencies of 100% and total COD reduction percentages of 85% were reached within the first hour of phototreatment, with an aromatic free effluent as the final product in both types of reactor. The ferrioxalate type complexes using mass ratios of oxalate/phenol = 1.5, oxalate/Fe3+ = 15, and H2O2/phenol > 5.0 were shown to be very effective in the treatment of these effluents, even at pH conditions close to neutral, the pH region in which Fenton type processes begin to lose efficiency due to the precipitation of iron as a hydroxide.

[53]Photo-Fenton process in a CPC solar photoreactor. The effect of solar activated photo-Fenton reagent at pH 5.0 before and after a slow sand filtration (SSF) process in waters containing natural iron species was investigated.Natural organic matter (NOM) model compounds (dihydroxy-benzene)The results showed that the total transformation of dihydroxybenzene compounds was obtained with a mineralization of over 80%. The mineralization of organic compounds dissolved in natural water was higher than in Milli-Q water, suggesting that the aqueous organic and inorganic components (metals, humic acids, and photoactive species) positively affect the photocatalytic process. When 1.0 mg/L of Fe3+ was added to the system, photo-Fenton degradation improved.

[54]Two laboratory scale photo-Fenton experiments were performed with the solar simulator and SMX dissolved in diluted water (DW) and in seawater (SW) at the same concentration (50 mg/L; DOC = 23.75 mg/L) as in the pilot plant experiments for their comparison with natural solar radiation. The initial DOC of SW was 2.6 mg C/L. The experiments were performed at three different initial concentrations of FeSO4·7H2O (2.6, 5.2, and 10.4 mg/L). Initial H2O2 concentrations ranged from 30 to 210 mg/L. The solar pilot plant reactor consisted of a compound parabolic collector (CPC) with a 3.0 m2 irradiated surface and total volume of 39 L.Antibiotic sulfamethoxazole (SMX)The photo-Fenton degradation of SMX was strongly influenced by the seawater matrix when compared to distilled water. Indeed, in seawater it is proposed that degradation occurs mainly through and Cl1− (or ) and not through HO. The increased iron concentration showed a slight improvement on the pollutant degradation and mineralization rate. The increase of H2O2 concentration up to 120 mg/L in distilled water reduced the sample toxicity during the photo-Fenton process, which demonstrates that this is a feasible technology for the treatment of wastewater containing this compound.

[55]
Photo-Fenton oxidation was carried out using a cylindrical Pyrex thermostatic cell with a 300 mL capacity ( °C), equipped with a magnetic stirrer. The dye solution volume was 250 mL. A 6 W Philips black light fluorescent lamp which basically emits at 350 nm was used as an artificial light source. The incident light intensity, measured with a uranyl actinometer, was Einstein s−1. A few Fenton reagent doses were tested in the present work (a series of three experiments): 5 mg/L Fe(II) and 125 mg/L H2O2, 10 mg/L Fe(II) and 125 mg/L H2O2 10 mg/L Fe(II) and 250 mg/L H2O2. Contaminants with a ratio of BOD5/COD ≥ 0.4 are generally accepted as biodegradable, while those with ratios between 0.2 and 0.3 units were partially biodegradable. Homo-bireactive dye (Procion Red H-E7 B)The results demonstrated that a photo-Fenton reaction can be used successfully as a pretreatment process to biocompatibilize Procion Red H-E7B reactive dye solutions. The best pretreatment results were obtained with 60 min of photo-Fenton irradiation time and 10 mg/L Fe(II) and 125 mg/L H2O2 of initial reagent concentration. Under these conditions, the BOD5/COD index increased from 0.10 to 0.35 units with 39% mineralization and 16.5 mg/L of residual H2O2. The use of photo-Fenton type reactions as a pretreatment allows the SBR system to remove Procion Red H-E7B Reactive Dye from aqueous solution, which improves the low success rate of aerobic biological removal of dye colour.

[56]This study explored the application of the solar photoFenton process to the degradation of PNA in water. The operating pH value was varied in the range of 3–6. The effect, of H2O2 and Fe2+ dosage on the degradation of PNA by solar photo-Fenton process were investigated between 2.5–40 and 0.025–0.1 mM, respectively. Also the effect, of temperature and initial pollutant concentration were investigated in the range of 20–50°C and  mM.P-Nitroaniline (PNA)The optimum conditions for the degradation of PNA in water were considered to be pH 3.0, 10 mmol/L H2O2, 0.05 mmol/L Fe2+, 0.072–0.217 mmol/L PNA, and temperature 20°C. Under optimum conditions, the degradation efficiencies of PNA were more than 98% within a 30 min reaction time. The degradation characteristic of PNA showed that the conjugated systems of the aromatic ring in the PNA molecules were effectively destroyed. The experimental results indicated that the solar photo-Fenton process has advantages over the classic Fenton process, such as higher oxidation power, a wider working pH range, and a lower ferrous ion usage.

[57]During the experiment, H2O2 was added continuously to the reactor at a flow rate of 1 mL/min with a syringe pump. Two 8 W monochromatic UV lamps of 312 nm (with an emission range between 280 and 360 nm) were placed axially in the reactor and kept in place with a quartz sleeve. The UV intensity of one 8 W UV lamp is 60  W/cm. The reaction temperature was maintained at °C using a water bath. A two factor CCD was carried out using H2O2 dosage rate ranging from 1 to 10 mg/L min and Fe3+ dosage from 1 to 100 mg/L to investigate their influence on carbofuran degradation under the photo-Fenton process.CarbofuranUnder these conditions, the toxicity unit measured by Microtox test with 5 min exposure was decreased from 47 to 6 and the biodegradability evaluated by BOD 5/COD ratio was increased from 0 to 0.76 after a 60 min reaction. The results obtained in this study demonstrate that the photo-Fenton process is a promising pretreatment to biological treatment for carbofuran removal from contaminated water or wastewater.

[58]Experiments were carried out in a Pyrex glass cylindrical reactor of 0.10 m diameter and 0.20 m height. The working volume was 1 L and all experiments were conducted in batch mode. The initial solution pH was adjusted to 3 which is the optimal value for the Fenton and photo-Fenton reactions using sulphuric acid. All experiments except those in the dark and at night were carried out between 10 am and 4 pm. The mean solar radiation during the experiments from October to January was in the range of 2.55–3.01 kWh (m2 day)−1. The effects of solar light, initial Fe concentration, and initial H2O2 concentration were investigated.Acid Orange 7With increasing Fe dosage the decolourization rate increased, but the enhancement was not pronounced beyond 10 mg/L. Although the addition of H2O2 increased the decolourization rate up to around 1000 mg/L of H2O2, further additions of H2O2 did not enhance colour removal. At excess dosages of Fenton reagents, colour removal was not improved, due to their scavenging of hydroxyl radicals. It was found that the pseudo first-order decolourization kinetic constant based on the accumulated solar energy is the sole parameter unifying solar photo-Fenton decolourization processes under different weather conditions.

4. UV/H2O2 Process

Like other AOPs, the oxidizing ability of UV/H2O2 may be attributed to the formation of OH, , and as reflected by their mechanistic pathways (Reactions (1), (19)–(23)). In fact, the AOP occurs via a reaction with OH radicals, produced by UV irradiation of H2O2. The molar absorptivity of hydrogen peroxide is low at 253.7 nm, about 20 M−1 cm−1, and OH radicals are formed per incident photon absorbed [104]. At this wavelength, the photolysis rate of aqueous hydrogen peroxide is about 50 times slower than that of ozone [104]. This technique requires a relatively high dose of H2O2 and/or a much longer UV exposure time than, for example, the UV/O3 process. In contrast, the rate of photolysis of hydrogen peroxide has been found to be pH dependent and increases when more alkaline conditions are used, because at 253.7 nm peroxide anions may be formed, which display a higher molar absorptivity than hydrogen peroxide, namely, 240 M−1 cm−1 [104, 142]. In this AOP, the formation of OH radicals is directly facilitated by the photolysis of H2O2 [143]. The radicals, which are formed by the homolytic splitting of the oxidant’s O–O bonds, transform the chemical structures of target chelating agents [143, 144]. Consider the following:Initiation (Rate Constant) Propagation [144] (Rate Constant) Termination [144] (Rate Constant) It is important to note that the effectiveness of UV/H2O2 systems depends on various conditions that affect their ability to degrade chelating agents. The variables include the type and concentration of contaminants or dissolved inorganic substances (such as carbonates and iron cations), organic substances present in surface water, light transmittance in solutions (as indicated by turbidity or colour), pH, temperature, and the optimum oxidant dose [145]. An excessive concentration of H2O2 would act as a radical scavenger, slowing down the rate of oxidation [146], while a low concentration of H2O2 insufficiently forms OH radicals in aqueous solutions, leading to a slower oxidation rate [143, 147]. The UV/H2O2 process is sensitive to the scavenging effects of carbonate ions at a pH ranging from 8 to 9. Furthermore, the UV/H2O2 process requires a long UV exposure time with a powerful output at a wide range of wavelengths. Nevertheless this treatment is more economically attractive than the UV/O3 process, due to its lower energy consumption [143, 148].

Tubular reactor configurations are usually employed for direct photolysis and photo-Fenton processes or processes based on H2O2/UV reagent, in order to achieve a good interaction between CPs, other intermediates, and radiation [104, 149]. Also various lamps are employed to generate the radiation supplied to CP samples for direct UV photolysis and for techniques based on UV/H2O2, UV/O3, photo-Fenton processes, and photocatalysis. The various commercial radiation sources employed include high, medium, and low pressure mercury vapour lamps for the generation of UV radiation [149151] and solar-simulated xenon lamps as a source of visible radiation [152]. The lamp can be located either in an axial position housed by a sleeve [150] or vertically, in its centre [104]. The typical findings observed in the UV/H2O2 process are listed in Table 5.


ReferenceProcess conditionsPollutant(s)Conclusions

[59]For photolytic experiments, the samples were irradiated with a UV lamp with an output of 254 nm operating at 50–60 Hz with a current intensity of 0.12 A at ambient temperature. The photolytic decolouration of carmine via UV radiation in the presence of H2O2 was optimized using response surface methodology (RSM) utilizing Design-Expert 7.1.Carmine
(C.I. Natural Red 4)
Under the optimized conditions of 62  M dye, 5.5 mM H2O2, and pH 4, the experimental values were as predicted, indicating the suitability of the model and the success of RSM in optimizing photooxidation conditions for carmine dye. In the optimization, and correlation coefficients for the quadratic model were evaluated quite satisfactorily at 0.998 and 0.997, respectively.

[60]UV/TiO2/H2O2, UV/TiO2, and UV/H2O2 were compared as pretreatment processes to detoxification and treatment. The tubes were then irradiated for 40 h (initial concentrations of 50 mg/L) or 56 h (initial concentrations of 100 mg/L) at 300  W cm−2 with two 18 W UV bluelamps and an initial chlorophenol concentration of 50 mg/L.4-Chlorophenol (4CP),
2,4-dichlorophenol (DCP), 2,4,6-trichlorophenol (TCP), and pentachlorophenol (PCP)
Chlorophenol photodegradation was well described by a first-order model kinetic ( > 0.94) and the shortest 4CP, DCP, TCP, and PCP half-lives were achieved during UV/TiO2/H2O2 treatment at 8.7, 7.1, 4.5, and 3.3 h, respectively.

[61]A 60 W mercury vapour lamp (UV C, 253.7 nm) with a frequency of 50 Hz and a voltage of 240 V was used. The initial concentrations of H2O2 and melanoidin were manipulated while pH, flow rate, irradiated surface area, volume, lamp intensity, and temperature were kept constant. The relative change of each constituent was identified at various initial concentrations of H2O2 (up to 12000 mg/L) and melanoidin (263–5314 mg-Pt Co/L).MelanoidinUV/H2O2 was shown to remove the colour associated with melanoidin effectively. The process was less effective in removing the DON and DOC present in the melanoidin solution. At the optimum H2O2 dose (3300 mg/L), with an initial melanoidin concentration of 2000 mg/L, the removal of colour, DOC and DON was 99%, 50%, and 25%, respectively.

[62]This study compared the efficacy of UV photodegradation with that of different advanced oxidation processes (O3, UV/H2O2, O3/activated carbon). Photo-irradiations were carried out using a merry-go-round photoreactor (MGRR), DEMA equipped with a 500 WTQ 718 Heraeus medium-pressure mercury lamp (239–334 nm) or a TNN 15/32 Heraeus low-pressure mercury lamp (254 nm). The temperature in the MGRR was kept at °C during all irradiations. The concentration of H2O2 used was 3 mM.Naphthalene sulphonic acidsThese results demonstrated that the treatment of naphthalene sulphonic acids with UV radiation is not effective in their removal from aqueous solutions. The presence of duroquinone and 4-carboxybenzophenone during the irradiation of naphthalene sulphonic acids increases their elimination rate. O3/activated carbon and UV/H2O2 based systems were found to be more efficient than the irradiation process in the removal of naphthalene sulphonic acids from aqueous solutions.

[63]The reactor had a 1 L capacity and was equipped with a mercury medium-pressure steam UV lamp which was 110 mm in length and used 1000 W, 145 V, and 7.5 A. In the UV light/H2O2 flow reactor system, the initial concentration of sulphide was 6.34 mg L−1. The initial concentrations of sulphurous water were 6.34 mg L−1 of HS, 1000 mg L−1 of , and 1.5 mg L−1 of . The amount of hydrogen peroxide added was of  mL L−1.Sulphurous waterIn a batch reactor it was possible to demonstrate that the sulphur compounds of the sulphurous waters could be oxidized to sulphate in a UV light/H2O2 air system with very small concentrations of hydrogen peroxide ( mL L−1). In a flow reactor it was possible to obtain the same results by adding only  mL L−1 of hydrogen peroxide.

[64]Radiation energy was supplied by two lamps. Two different types of lamp were used: (1) two Philips TUV lamps with an input power of 15 W each and (2) two Heraeus UV-C lamps operated with an input power of 40 W each. Both types of lamp are low pressure mercury vapour lamps with one single significant emission wavelength at 253.7 nm. DCA concentration and radiation absorbing species concentration (H2O2) were 60 ppm, 145 ppm and pH and temperature were kept at 3.4 and 20°C, respectively.Dichloroacetic acid (DCA)The fastest degradation rate was obtained with the H2O2/UV40W system, followed by H2O2/UV15W. Although the photocatalytic process was effective in degrading DCA, the reaction rate was much slower when compared with the homogeneous processes. For the H2O2/UV40W reaction, the DCA conversion at  s (ca. 4 h of reaction) is more than 80%, whereas the H2O2/UV15W system reaches half of this value. The DCA and TOC conversion values are similar in each process. This is in agreement with the fact that there are no stable reaction intermediates and DCA is rapidly converted into HCl and CO2.

[65]Low pressure mercury vapour lamps with a maximum emission primarily at 253.7 nm were used as the light source. The changes in the pH of dye solutions as a function of the irradiation time for different initial pH values are carried out. The effect of the initial H2O2 concentration in a range of 10–100 mM on the rate of RO16 decolourization was investigated. The effect of the initial RO16 concentration in a range from 20 to 80 mg dm−3 on the efficiency of dye degradation was also investigated. The influence of UV light intensity on the decolourization of RO16 azo dye was monitored by varying the light intensity from 730 up to 1950  W cm−2.Azo dye Reactive Orange 16The UV/H2O2 process could be used efficiently for the decolourization of aqueous solutions of the azo dye Reactive Orange 16. It was found that the rate of decolourization is significantly affected by the initial pH, the initial hydrogen peroxide concentration, the initial dye concentration, and the UV light intensity. The decolourization follows pseudo first-order reaction kinetics. Peroxide concentrations in the range from 20 to 40 mM appear to be optimal. Colour removal was observed to be faster in neutral pH solutions than in acidic and basic ones. The hydroxyl radical scavenging effect of the examined inorganic anions increased in the order phosphate < sulphate < nitrate < chloride.

Ultimate oxidation of CPs to carbon dioxide and water has rarely been obtained under typical test conditions. As summarized in Table 4, typical half-life times are between 0.3 and 20.1 minutes for CP degradation, depending on the initial concentration of CP and hydrogen peroxide, the intensity of radiation, and the degree of chlorination. It is observed that the degradation rates increase when the number of chlorine substituents decreases [104].

5. Conclusion

Recent developments in various homogeneous AOPs have been analysed comprehensively. The principle of individual and combined AOPs and their efficiency on the degradation of various pollutants was discussed. The influence of various experimental parameters such as oxidant dosage, solution pH, flow rates, substrate concentrations, water matrix, and light intensity on the AOPs was explored. This review also listed various AOPs applied for the degradation of contaminants under different experimental conditions. Combined AOPs substantially enhanced the degradation rate by generating more reactive radicals under suitable conditions. The optimum oxidant dosage and solution for efficient removal were reported.

Conflict of Interests

The authors declare that there is no conflict of interests regarding the publication of this paper.

Acknowledgments

The authors wish to thank the National Science Council (NSC) in Taiwan for their financial support under the Contract no. NSC-101-2221-035-031-MY3. The Laboratory of Green Chemistry, Mikkeli, Finland, and Water and Environmental Technology (WET) Center, Temple University, are also gratefully acknowledged for their partial financial support of this study.

References

  1. http://www.unesco.org/water/news/newsletter/184.shtml.
  2. http://www.un.org/News/Press/docs/2010/ga10967.doc.htm.
  3. http://www.who.int/water_sanitation_health/recognition_safe_clean_water/en/index.html.
  4. H. Kim, J. Shim, and S. Lee, “Formation of disinfection by-products in chlorinated swimming pool water,” Chemosphere, vol. 46, no. 1, pp. 123–130, 2002. View at: Publisher Site | Google Scholar
  5. G. C. White, Handbook of Chlorination and Alternative Disinfectants, Van Nostrand Reinhold, New York, NY, USA, 3rd edition, 1992.
  6. Z. Aksu and J. Yener, “A comparative adsorption/biosorption study of mono-chlorinated phenols onto various sorbents,” Waste Management, vol. 21, no. 8, pp. 695–702, 2001. View at: Publisher Site | Google Scholar
  7. S. C. J. M. Van Hoof, A. Hashim, and A. J. Kordes, “The effect of ultrafiltration as pretreatment to reverse osmosis in wastewater reuse and seawater desalination applications,” Desalination, vol. 124, no. 1–3, pp. 231–242, 1999. View at: Publisher Site | Google Scholar
  8. D. Keenan and A. Sabelnikov, “Biological augmentation eliminates grease and oil in bakery wastewater,” Water Environment Research, vol. 72, no. 2, pp. 141–146, 2000. View at: Google Scholar
  9. G. Chen, X. Chen, and P. L. Yue, “Electrocoagulation and electroflotation of restaurant wastewater,” Journal of Environmental Engineering, vol. 126, no. 9, pp. 858–863, 2000. View at: Publisher Site | Google Scholar
  10. M. J. Dietrich, T. L. Randall, and P. J. Canney, “Wet air oxidation of hazardous organics in wastewater,” Environmental Progress, vol. 4, no. 3, pp. 171–177, 1985. View at: Google Scholar
  11. M. Muruganandham and M. Swaminathan, “Advanced oxidative decolourisation of reactive yellow 14 azo dye by UV/TiO2, UV/H2O2, UV/H2O2/Fe2+ processes—a comparative study,” Separation & Purification Technology, vol. 48, no. 3, pp. 297–303, 2006. View at: Publisher Site | Google Scholar
  12. M. Paleia, G. Ifrim, M. Barbu, G. Bahrim, and S. Caraman, “Strategies for the aerobic biological treatment of the dairy wastewaters in controlled conditions,” Environmental Engineering & Management Journal, vol. 9, no. 3, pp. 399–405, 2010. View at: Google Scholar
  13. M. Muruganandham, S. H. Chen, and J. J. Wu, “Mineralization of N-methyl-2-pyrolidone by advanced oxidation processes,” Separation & Purification Technology, vol. 55, no. 3, pp. 360–367, 2007. View at: Publisher Site | Google Scholar
  14. J. J. Wu, M. Muruganandham, L. T. Chang, G. J. Lee, V. N. Batalova, and G. M. Mokrousov, “Catalytic ozonation of oxalic acid using SrTiO3 catalyst,” Ozone: Science & Engineering, vol. 33, no. 1, pp. 74–79, 2011. View at: Publisher Site | Google Scholar
  15. J. J. Wu, S. H. Chen, and M. Muruganandham, “Catalytic ozonation of oxalic acid using carbon-free rice husk ash catalysts,” Industrial & Engineering Chemistry Research, vol. 47, no. 9, pp. 2919–2925, 2008. View at: Publisher Site | Google Scholar
  16. S. Vilhunen, M. Vilve, M. Vepsäläinen, and M. Sillanpää, “Removal of organic matter from a variety of water matrices by UV photolysis and UV/H2O2 method,” Journal of Hazardous Materials, vol. 179, no. 1–3, pp. 776–782, 2010. View at: Publisher Site | Google Scholar
  17. M. Muruganandham and M. Swaminathan, “Decolourisation of reactive orange 4 by Fenton and photo-Fenton oxidation technology,” Dyes & Pigments, vol. 63, no. 3, pp. 315–321, 2004. View at: Publisher Site | Google Scholar
  18. A. Matilainena and M. Sillanpää, “Removal of natural organic matter from drinking water by advanced oxidation processes,” Chemosphere, vol. 80, no. 4, pp. 351–365, 2011. View at: Google Scholar
  19. J. J. Wu, J. S. Yang, M. Muruganandham, and C. C. Wu, “The oxidation study of 2-propanol using ozone-based advanced oxidation processes,” Separation & Purification Technology, vol. 62, no. 1, pp. 39–46, 2008. View at: Publisher Site | Google Scholar
  20. N. N. Mahamuni and Y. G. Adewuyi, “Advanced oxidation processes (AOPs) involving ultrasound for wastewater treatment: a review with emphasis on cost estimation,” Ultrasonics Sonochemistry, vol. 17, no. 6, pp. 990–1003, 2010. View at: Publisher Site | Google Scholar
  21. M. Muruganandham, J.-S. Yang, and J. J. Wu, “Effect of ultrasonic irradiation on the catalytic activity and stability of goethite catalyst in the presence of H2O2 at acidic medium,” Industrial & Engineering Chemistry Research, vol. 46, no. 3, pp. 691–698, 2007. View at: Publisher Site | Google Scholar
  22. M. Muruganandham, R. Amutha, G.-J. Lee, S.-H. Hsieh, J. J. Wu, and M. Sillanpaa, “Facile fabrication of tunable Bi2O3 self-assembly and its visible light photocatalytic activity,” Journal of Physical Chemistry C, vol. 116, no. 23, pp. 12906–12915, 2012. View at: Google Scholar
  23. M. Muruganandham, R. Amutha, E. Repo, M. Sillanpää, Y. Kusumoto, and M. Abdulla-Al-Mamun, “Controlled mesoporous self-assembly of ZnS microsphere for photocatalytic degradation of Methyl Orange dye,” Journal of Photochemistry and Photobiology A, vol. 216, no. 2-4, pp. 133–141, 2010. View at: Publisher Site | Google Scholar
  24. T. Garoma and M. D. Gurol, “Degradation of tert-butyl alcohol in dilute aqueous solution by an O3/UV process,” Environmental Science & Technology, vol. 38, no. 19, pp. 5246–5252, 2004. View at: Publisher Site | Google Scholar
  25. N. A. Medellin-Castillo, R. Ocampo-Perez, R. Leyva-Ramos, M. Sanchez-Polo, J. Rivera-Utrilla, and J. D. Mendez-Diaz, “Removal of diethyl phthalate from water solution by adsorption, photo-oxidation, ozonation and advanced oxidation process (UV/H2O2, O3/H2O2 and O3/activated carbon),” Science of the Total Environment, vol. 442, pp. 26–35, 2013. View at: Google Scholar
  26. S. C. Kwon, J. Y. Kim, S. M. Yoon, W. Bae, K. S. Kang, and Y. W. Rhee, “Treatment characteristic of 1,4-dioxane by ozone-based advanced oxidation processes,” Journal of Industrial & Engineering Chemistry, vol. 18, no. 6, pp. 1951–1955. View at: Google Scholar
  27. H. Kusic, N. Koprivanac, and A. L. Bozic, “Minimization of organic pollutant content in aqueous solution by means of AOPs: UV- and ozone-based technologies,” Chemical Engineering Journal, vol. 123, no. 3, pp. 127–137, 2006. View at: Publisher Site | Google Scholar
  28. X. Jin, S. Peldszus, and P. M. Huck, “Reaction kinetics of selected micropollutants in ozonation and advanced oxidation processes,” Water Research, vol. 46, no. 19, pp. 6519–6530, 2012. View at: Google Scholar
  29. F. J. Rivas, M. Carbajo, F. Beltrán, O. Gimeno, and J. Frades, “Comparison of different advanced oxidation processes (AOPs) in the presence of perovskites,” Journal of Hazardous Materials, vol. 155, no. 3, pp. 407–414, 2008. View at: Publisher Site | Google Scholar
  30. I. Peternel, N. Koprivanac, and I. Grcic, “Mineralization of p-chlorophenol in water solution by AOPs based on UV irradiation,” Environmental Technology, vol. 33, no. 1, pp. 27–36, 2012. View at: Publisher Site | Google Scholar
  31. A. K. Biń, P. Machniewski, R. Sakowicz, J. Ostrowska, and J. Zieliński, “Degradation of nitroaromatics (MNT, DNT and TNT) by AOPs,” Ozone: Science & Engineering, vol. 23, no. 5, pp. 343–349, 2001. View at: Google Scholar
  32. K. Wang, J. Guo, M. Yang, H. Junji, and R. Deng, “Decomposition of two haloacetic acids in water using UV radiation, ozone and advanced oxidation processes,” Journal of Hazardous Materials, vol. 162, no. 2-3, pp. 1243–1248, 2009. View at: Publisher Site | Google Scholar
  33. M. Y. Ghaly, G. Härtel, R. Mayer, and R. Haseneder, “Aromatic compounds degradation in water by using ozone and AOPS. A comparative study. O-Nitrotoluene as a model substrate,” Ozone: Science & Engineering, vol. 23, no. 2, pp. 127–138, 2001. View at: Google Scholar
  34. O. Gimeno, M. Carbajo, F. J. Beltrán, and F. J. Rivas, “Phenol and substituted phenols AOPs remediation,” Journal of Hazardous Materials, vol. 119, no. 1–3, pp. 99–108, 2005. View at: Publisher Site | Google Scholar
  35. J. B. Parsa and S. H. Negahdar, “Treatment of wastewater containing acid blue 92 dye by advanced ozone-based oxidation methods,” Separation & Purification Technology, vol. 98, no. 19, pp. 315–320, 2012. View at: Google Scholar
  36. J. Rivera-Utrilla, J. Méndez-Díaz, M. Sánchez-Polo, M. A. Ferro-García, and I. Bautista-Toledo, “Removal of the surfactant sodium Dodecylbenzenesulfonate from water by simultaneous use of ozone and powdered activated carbon: comparison with systems based on O3 and O3/H2O2,” Water Research, vol. 40, no. 8, pp. 1717–1725, 2006. View at: Publisher Site | Google Scholar
  37. I. Arslan, I. A. Balcioglu, and T. Tuhkanen, “Advanced oxidation of synthetic dyehouse effluent by O3, H2O2/O3 and H2O2/UV processes,” Environmental Technology, vol. 20, no. 9, pp. 921–931, 1999. View at: Google Scholar
  38. R. Kidak and N. H. Ince, “Catalysis of advanced oxidation reactions by ultrasound: a case study with phenol,” Journal of Hazardous Materials, vol. 146, no. 3, pp. 630–635, 2007. View at: Publisher Site | Google Scholar
  39. G. Tezcanli-Güyer and N. H. Ince, “Individual and combined effects of ultrasound, ozone and UV irradiation: a case study with textile dyes,” Ultrasonics, vol. 42, no. 1–9, pp. 603–609, 2004. View at: Publisher Site | Google Scholar
  40. I. A. Alaton and S. Teksoy, “Acid dyebath effluent pretreatment using Fenton's reagent: process optimization, reaction kinetics and effects on acute toxicity,” Dyes & Pigments, vol. 73, no. 1, pp. 31–39, 2007. View at: Publisher Site | Google Scholar
  41. H. Kušić, A. Lončarić Božić, and N. Koprivanac, “Fenton type processes for minimization of organic content in coloured wastewaters: part I: processes optimization,” Dyes & Pigments, vol. 74, no. 2, pp. 380–387, 2007. View at: Publisher Site | Google Scholar
  42. K. Pirkanniemi, S. Metsärinne, and M. Sillanpää, “Degradation of EDTA and novel complexing agents in pulp and paper mill process and wastewaters by Fenton's reagent,” Journal of Hazardous Materials, vol. 147, no. 1-2, pp. 556–561, 2007. View at: Publisher Site | Google Scholar
  43. M. A. Behnajady, N. Modirshahla, and F. Ghanbary, “A kinetic model for the decolorization of C.I. Acid Yellow 23 by Fenton process,” Journal of Hazardous Materials, vol. 148, no. 1-2, pp. 98–102, 2007. View at: Publisher Site | Google Scholar
  44. S.-P. Sun, C.-J. Li, J.-H. Sun, S.-H. Shi, M.-H. Fan, and Q. Zhou, “Decolorization of an azo dye Orange G in aqueous solution by Fenton oxidation process: effect of system parameters and kinetic study,” Journal of Hazardous Materials, vol. 161, no. 2-3, pp. 1052–1057, 2009. View at: Publisher Site | Google Scholar
  45. C. S. D. Rodrigues, L. M. Madeira, and R. A. R. Boaventura, “Optimization of the azo dye Procion Red H-EXL degradation by Fenton's reagent using experimental design,” Journal of Hazardous Materials, vol. 164, no. 2-3, pp. 987–994, 2009. View at: Publisher Site | Google Scholar
  46. E. C. Catalkaya and F. Kargi, “Advanced oxidation and mineralization of simazine using Fenton's reagent,” Journal of Hazardous Materials, vol. 168, no. 2-3, pp. 688–694, 2009. View at: Publisher Site | Google Scholar
  47. K. Ayoub, S. Nélieu, E. D. Van Hullebusch, A. Maia-Grondard, M. Cassir, and A. Bermond, “TNT oxidation by Fenton reaction: reagent ratio effect on kinetics and early stage degradation pathways,” Chemical Engineering Journal, vol. 173, no. 2, pp. 309–317, 2011. View at: Publisher Site | Google Scholar
  48. J.-H. Sun, S.-P. Sun, G.-L. Wang, and L.-P. Qiao, “Degradation of azo dye Amido black 10B in aqueous solution by Fenton oxidation process,” Dyes & Pigments, vol. 74, no. 3, pp. 647–652, 2007. View at: Publisher Site | Google Scholar
  49. A. A. Burbano, D. D. Dionysiou, M. T. Suidan, and T. L. Richardson, “Oxidation kinetics and effect of pH on the degradation of MTBE with Fenton reagent,” Water Research, vol. 39, no. 1, pp. 107–118, 2005. View at: Publisher Site | Google Scholar
  50. A. A. Burbano, D. D. Dionysiou, and M. T. Suidan, “Effect of oxidant-to-substrate ratios on the degradation of MTBE with Fenton reagent,” Water Research, vol. 42, no. 12, pp. 3225–3239, 2008. View at: Publisher Site | Google Scholar
  51. M. Vilve, A. Hirvonen, and M. Sillanpää, “Effects of reaction conditions on nuclear laundry water treatment in Fenton process,” Journal of Hazardous Materials, vol. 164, no. 2-3, pp. 1468–1473, 2009. View at: Publisher Site | Google Scholar
  52. D. Prato-Garcia, R. Vasquez-Medrano, and M. Hernandez-Esparza, “Solar photoassisted advanced oxidation of synthetic phenolic wastewaters using ferrioxalate complexes,” Solar Energy, vol. 83, no. 3, pp. 306–315, 2009. View at: Publisher Site | Google Scholar
  53. A. Moncayo-Lasso, C. Pulgarin, and N. Benítez, “Degradation of DBPs' precursors in river water before and after slow sand filtration by photo-Fenton process at pH 5 in a solar CPC reactor,” Water Research, vol. 42, no. 15, pp. 4125–4132, 2008. View at: Publisher Site | Google Scholar
  54. A. G. Trovó, R. F. P. Nogueira, A. Agüera, A. R. Fernandez-Alba, C. Sirtori, and S. Malato, “Degradation of sulfamethoxazole in water by solar photo-Fenton. Chemical and toxicological evaluation,” Water Research, vol. 43, no. 16, pp. 3922–3931, 2009. View at: Publisher Site | Google Scholar
  55. J. García-Montaño, F. Torrades, J. A. García-Hortal, X. Domènech, and J. Peral, “Degradation of Procion Red H-E7B reactive dye by coupling a photo-Fenton system with a sequencing batch reactor,” Journal of Hazardous Materials, vol. 134, no. 1–3, pp. 220–229, 2006. View at: Publisher Site | Google Scholar
  56. J.-H. Sun, S.-P. Sun, M.-H. Fan, H.-Q. Guo, Y.-F. Lee, and R.-X. Sun, “Oxidative decomposition of p-nitroaniline in water by solar photo-Fenton advanced oxidation process,” Journal of Hazardous Materials, vol. 153, no. 1-2, pp. 187–193, 2008. View at: Publisher Site | Google Scholar
  57. L.-A. Lu, Y.-S. Ma, M. Kumar, and J.-G. Lin, “Photo-Fenton pretreatment of carbofuran—analyses via experimental design, detoxification and biodegradability enhancement,” Separation & Purification Technology, vol. 81, pp. 325–331, 2011. View at: Publisher Site | Google Scholar
  58. A. Özcan, M. A. Oturan, N. Oturan, and Y. Şahin, “Removal of Acid Orange 7 from water by electrochemically generated Fenton's reagent,” Journal of Hazardous Materials, vol. 163, no. 2-3, pp. 1213–1220, 2009. View at: Publisher Site | Google Scholar
  59. B. K. Körbahti and M. A. Rauf, “Determination of optimum operating conditions of carmine decoloration by UV/H2O2 using response surface methodology,” Journal of Hazardous Materials, vol. 161, no. 1, pp. 281–286, 2009. View at: Publisher Site | Google Scholar
  60. T. Essam, M. A. Amin, O. El Tayeb, B. Mattiasson, and B. Guieysse, “Sequential photochemical-biological degradation of chlorophenols,” Chemosphere, vol. 66, no. 11, pp. 2201–2209, 2007. View at: Publisher Site | Google Scholar
  61. J. Dwyer, L. Kavanagh, and P. Lant, “The degradation of dissolved organic nitrogen associated with melanoidin using a UV/H2O2 AOP,” Chemosphere, vol. 71, no. 9, pp. 1745–1753, 2008. View at: Publisher Site | Google Scholar
  62. M. Sánchez-Polo, J. Rivera-Utrilla, J. D. Méndez-Díaz, S. Canonica, and U. von Gunten, “Photooxidation of naphthalenesulfonic acids: comparison between processes based on O3, O3/activated carbon and UV/H2O2,” Chemosphere, vol. 68, no. 10, pp. 1814–1820, 2007. View at: Publisher Site | Google Scholar
  63. F. Hernández and G. Geissler, “Photooxidative treatment of sulfurous water for its potabilization,” Photochemistry & Photobiology, vol. 81, no. 3, pp. 636–640, 2005. View at: Publisher Site | Google Scholar
  64. C. S. Zalazar, M. L. Satuf, O. M. Alfano, and A. E. Cassano, “Comparison of H2O2/UV and heterogeneous photocatalytic processes for the degradation of dichloroacetic acid in water,” Environmental Science & Technology, vol. 42, no. 16, pp. 6198–6204, 2008. View at: Publisher Site | Google Scholar
  65. J. Mitrović, M. Radović, D. Bojić, T. Anbelković, M. Purenović, and A. Bojić, “Decolorization of the textile azo dye reactive orange 16 by the UV/H2O2 process,” Journal of the Serbian Chemical Society, vol. 77, no. 4, pp. 465–481, 2012. View at: Publisher Site | Google Scholar
  66. Y. J. Jung, W. G. Kim, Y. Yoon, T.-M. Hwang, and J.-W. Kang, “pH effect on ozonation of ampicillin: kinetic study and toxicity assessment,” Ozone: Science & Engineering, vol. 34, no. 3, pp. 156–162, 2012. View at: Google Scholar
  67. Z. S. Can and M. Gurol, “Formaldehyde formation during ozonation of drinking water,” Ozone: Science & Engineering, vol. 25, no. 1, pp. 41–51, 2003. View at: Google Scholar
  68. Y. Qiu, C.-H. Kuo, and M. E. Zappi, “Ozonation kinetics of six dichlorophenol isomers,” Ozone: Science & Engineering, vol. 24, no. 2, pp. 123–131, 2002. View at: Google Scholar
  69. G. Moussavi and M. Mahmoudi, “Degradation and biodegradability improvement of the reactive red 198 azo dye using catalytic ozonation with MgO nanocrystals,” Chemical Engineering Journal, vol. 152, no. 1, pp. 1–7, 2009. View at: Publisher Site | Google Scholar
  70. A. Begum and S. K. Gautam, “Endosulfan and lindane degradation using ozonation,” Environmental Technology, vol. 33, no. 8, pp. 943–949, 2012. View at: Publisher Site | Google Scholar
  71. Y. Ku, H. S. Lin, W. Wang, and C. M. Ma, “Decomposition of phorate in aqueous solution by ozonation,” Journal of Environmental Science & Health Part B, vol. 42, no. 2, pp. 143–149, 2007. View at: Publisher Site | Google Scholar
  72. P. K. A. Hong and Y. Zeng, “Degradation of pentachlorophenol by ozonation and biodegradability of intermediates,” Water Research, vol. 36, no. 17, pp. 4243–4254, 2002. View at: Publisher Site | Google Scholar
  73. H. Hoigne and H. Bader, “Rate constants of reactions of ozone with organic and inorganic compounds in water. II. Dissociating organic compounds,” Water Research, vol. 17, no. 2, pp. 185–194, 1983. View at: Publisher Site | Google Scholar
  74. K. Y. Li, C. H. Kuo, and J. L. Weeks Jr., “A kinetic study of ozone-phenol reaction in aqueous solutions,” AIChE Journal, vol. 25, no. 4, pp. 583–591, 1979. View at: Google Scholar
  75. T. Otake, S. Tone, K. Kono, and K. Nakao, “Photo-oxidation of phenols with ozone,” Journal of Chemical Engineering of Japan, vol. 12, no. 4, pp. 289–293, 1979. View at: Google Scholar
  76. M. S. Yazgan, C. Kinaci, and I. Toroz, “Ozonation of aqueous solution of alpha endosulfan,” Journal of Environmental Science & Health Part B, vol. 38, no. 6, pp. 843–853, 2003. View at: Publisher Site | Google Scholar
  77. J. A. Roth and D. E. Sullivan, “Solubility of ozone in water,” Industrial & Engineering Chemistry Fundamentals, vol. 20, no. 2, pp. 137–140, 1981. View at: Google Scholar
  78. J. L. Sotelo, F. J. Beltrán, F. J. Benítez, and J. Beltrán-Heredia, “Ozone decomposition in water: kinetic study,” Industrial & Engineering Chemistry Research, vol. 26, no. 1, pp. 39–43, 1987. View at: Google Scholar
  79. J. Wu, M. A. Eiteman, and S. E. Law, “Evaluation of membrane filtration and ozonation processes for treatment of reactive-dye wastewater,” Journal of Environmental Engineering, vol. 124, no. 3, pp. 272–277, 1998. View at: Google Scholar
  80. J. J. Wu, M. Muruganandham, and S. H. Chen, “Degradation of DMSO by ozone-based advanced oxidation processes,” Journal of Hazardous Materials, vol. 149, no. 1, pp. 218–225, 2007. View at: Publisher Site | Google Scholar
  81. J.-M. Fanchiang and D.-H. Tseng, “Decolorization and transformation of anthraquinone dye Reactive Blue 19 by ozonation,” Environmental Technology, vol. 30, no. 2, pp. 161–172, 2009. View at: Publisher Site | Google Scholar
  82. J. J. Wu, M. Muruganandham, L. T. Chang, and S. H. Chen, “Oxidation of propylene glycol methyl ether acetate using ozone-based advanced oxidation processes,” Ozone: Science & Engineering, vol. 30, no. 5, pp. 332–338, 2008. View at: Publisher Site | Google Scholar
  83. F. J. Beltrán, J. F. García-Araya, and I. Giráldez, “Gallic acid water ozonation using activated carbon,” Applied Catalysis B, vol. 63, no. 3-4, pp. 249–259, 2006. View at: Publisher Site | Google Scholar
  84. E. Reisz, W. Schmidt, H.-P. Schuchmann, and C. Von Sonntag, “Photolysis of ozone in aqueous solutions in the presence of tertiary butanol,” Environmental Science & Technology, vol. 37, no. 9, pp. 1941–1948, 2003. View at: Publisher Site | Google Scholar
  85. B. Xu, Z. Chen, F. Qi, J. Ma, and F. Wu, “Comparison of N-nitrosodiethylamine degradation in water by UV irradiation and UV/O3: efficiency, product and mechanism,” Journal of Hazardous Materials, vol. 179, no. 1–3, pp. 976–982, 2010. View at: Publisher Site | Google Scholar
  86. Y.-S. Shen and Y. Ku, “Decomposition of gas-phase chloroethenes by UV/O3 process,” Water Research, vol. 32, no. 9, pp. 2669–2679, 1998. View at: Publisher Site | Google Scholar
  87. M.-S. Chou and K.-L. Chang, “Oxidation of aqueous dimethyl sulfoxide (DMSO) using UV, O3, and UV/O3,” Ozone: Science & Engineering, vol. 29, no. 5, pp. 391–397, 2007. View at: Publisher Site | Google Scholar
  88. M. Trapido, A. Hirvonen, Y. Veressinina, J. Hentunen, and R. Munter, “Ozonation, ozone/UV and UV/H2O2 degradation of chlorophenols,” Ozone: Science & Engineering, vol. 19, no. 1, pp. 75–96, 1997. View at: Google Scholar
  89. M. S. Lucas, J. A. Peres, and G. Li Puma, “Treatment of winery wastewater by ozone-based advanced oxidation processes (O3, O3/UV and O3/UV/H2O2) in a pilot-scale bubble column reactor and process economics,” Separation & Purification Technology, vol. 72, no. 3, pp. 235–241, 2010. View at: Publisher Site | Google Scholar
  90. H. Qi, D.-Z. Sun, and G.-Q. Chi, “Formaldehyde degradation by UV/TiO2/O3 process using continuous flow mode,” Journal of Environmental Sciences, vol. 19, no. 9, pp. 1136–1140, 2007. View at: Publisher Site | Google Scholar
  91. S.-J. Kim, S.-C. Kim, S.-G. Seo et al., “Photocatalyzed destruction of organic dyes using microwave/UV/O3/H2O2/TiO2 oxidation system,” Catalysis Today, vol. 164, no. 1, pp. 384–390, 2011. View at: Publisher Site | Google Scholar
  92. S. Cortes, J. Sarasa, P. Ormad, R. Gracia, and J. L. Ovelleiro, “Comparative efficiency of the systems O3/high pH and O3/catalyst for the oxidation of chlorobenzenes in water,” Ozone: Science & Engineering, vol. 22, no. 4, pp. 415–426, 2000. View at: Google Scholar
  93. A. A. Abd El-Raady and T. Nakajima, “Decomposition of carboxylic acids in water by O3, O3/H2O2, and O3/catalyst,” Ozone: Science & Engineering, vol. 27, no. 1, pp. 11–18, 2005. View at: Publisher Site | Google Scholar
  94. F. J. Beltrán, F. J. Rivas, and R. Montero-de-Espinosa, “Ozone-enhanced oxidation of oxalic acid in water with cobalt catalysts. 1. Homogeneous catalytic ozonation,” Industrial & Engineering Chemistry Research, vol. 42, no. 14, pp. 3210–3217, 2003. View at: Google Scholar
  95. T. Zhang and J. Ma, “Catalytic ozonation of trace nitrobenzene in water with synthetic goethite,” Journal of Molecular Catalysis A, vol. 279, no. 1, pp. 82–89, 2008. View at: Publisher Site | Google Scholar
  96. B. Kasprzyk-Hordern, M. Ziółek, and J. Nawrocki, “Catalytic ozonation and methods of enhancing molecular ozone reactions in water treatment,” Applied Catalysis B, vol. 46, no. 4, pp. 639–669, 2003. View at: Publisher Site | Google Scholar
  97. T. Oyama, T. Otsu, Y. Hidano, T. Koike, N. Serpone, and H. Hidaka, “Enhanced remediation of simulated wastewaters contaminated with 2-chlorophenol and other aquatic pollutants by TiO2-photoassisted ozonation in a sunlight-driven pilot-plant scale photoreactor,” Solar Energy, vol. 85, no. 5, pp. 938–944, 2011. View at: Publisher Site | Google Scholar
  98. E. M. Rodríguez, G. Fernández, P. M. Alvarez, and F. J. Beltrán, “TiO2 and Fe (III) photocatalytic ozonation processes of a mixture of emergent contaminants of water,” Water Research, vol. 46, no. 1, pp. 152–166, 2012. View at: Publisher Site | Google Scholar
  99. F. J. Beltrán, F. J. Rivas, and O. Gimeno, “Comparison between photocatalytic ozonation and other oxidation processes for the removal of phenols from water,” Journal of Chemical Technology and Biotechnology, vol. 80, no. 9, pp. 973–984, 2005. View at: Publisher Site | Google Scholar
  100. S. Song, M. Xia, Z. He, H. Ying, B. Lü, and J. Chen, “Degradation of p-nitrotoluene in aqueous solution by ozonation combined with sonolysis,” Journal of Hazardous Materials, vol. 144, no. 1-2, pp. 532–537, 2007. View at: Publisher Site | Google Scholar
  101. G. Xu, S. Chen, J. Shi, S. Wang, and G. Zhu, “Combination treatment of ultrasound and ozone for improving solubilization and anaerobic biodegradability of waste activated sludge,” Journal of Hazardous Materials, vol. 180, no. 1–3, pp. 340–346, 2010. View at: Publisher Site | Google Scholar
  102. J. Herney-Ramirez, M. A. Vicente, and L. M. Madeira, “Heterogeneous photo-Fenton oxidation with pillared clay-based catalysts for wastewater treatment: a review,” Applied Catalysis B, vol. 98, no. 1-2, pp. 10–26, 2010. View at: Publisher Site | Google Scholar
  103. H. Gallard and J. De Laat, “Kinetic modelling of Fe(III)/H2O2 oxidation reactions in dilute aqueous solution using atrazine as a model organic compound,” Water Research, vol. 34, no. 12, pp. 3107–3116, 2000. View at: Publisher Site | Google Scholar
  104. M. Pera-Titus, V. García-Molina, M. A. Baños, J. Giménez, and S. Esplugas, “Degradation of chlorophenols by means of advanced oxidation processes: a general review,” Applied Catalysis B, vol. 47, no. 4, pp. 219–256, 2004. View at: Publisher Site | Google Scholar
  105. Y. Deng and J. D. Englehardt, “Treatment of landfill leachate by the Fenton process,” Water Research, vol. 40, no. 20, pp. 3683–3694, 2006. View at: Publisher Site | Google Scholar
  106. Y. W. Kang and K.-Y. Hwang, “Effects of reaction conditions on the oxidation efficiency in the Fenton process,” Water Research, vol. 34, no. 10, pp. 2786–2790, 2000. View at: Publisher Site | Google Scholar
  107. Q. Zhou, W. Li, and T. Hua, “Removal of organic matter from landfill leachate by advanced oxidation processes: a review,” International Journal of Chemical Engineering, Article ID 270532, 2010. View at: Publisher Site | Google Scholar
  108. W. Z. Tang, Physicochemical Treatment of Hazardous Wastes, Lewis, Boca Raton, Fla, USA, 2003.
  109. E. Neyens and J. Baeyens, “A review of classic Fenton's peroxidation as an advanced oxidation technique,” Journal of Hazardous Materials, vol. 98, no. 1–3, pp. 33–50, 2003. View at: Publisher Site | Google Scholar
  110. P. Wang, I. W. C. Lau, H. H. P. Fang, and D. Zhou, “Landfill leachate treatment with combined UASB and Fenton coagulation,” Journal of Environmental Science and Health Part A, vol. 35, no. 10, pp. 1981–1988, 2000. View at: Google Scholar
  111. I. W. C. Lau, P. Wang, and H. H. P. Fang, “Organic removal of anaerobically treated leachate by Fenton Coagulation,” Journal of Environmental Engineering, vol. 127, no. 7, pp. 666–669, 2001. View at: Publisher Site | Google Scholar
  112. J. Beltran De Heredia, J. Torregrosa, J. R. Dominguez, and J. A. Peres, “Kinetic model for phenolic compound oxidation by Fenton's reagent,” Chemosphere, vol. 45, no. 1, pp. 85–90, 2001. View at: Publisher Site | Google Scholar
  113. N. Kang, D. S. Lee, and J. Yoon, “Kinetic modeling of Fenton oxidation of phenol and monochlorophenols,” Chemosphere, vol. 47, no. 9, pp. 915–924, 2002. View at: Publisher Site | Google Scholar
  114. J. A. Zazo, J. A. Casas, A. F. Mohedano, M. A. Gilarranz, and J. J. Rodríguez, “Chemical pathway and kinetics of phenol oxidation by Fenton's reagent,” Environmental Science & Technology, vol. 39, no. 23, pp. 9295–9302, 2005. View at: Publisher Site | Google Scholar
  115. S. Basu and I. W. Wei, “Mechanism and kinetics of oxidation of 2,4,6-trichlorophenol by Fenton's reagent,” Environmental Engineering Science, vol. 17, no. 5, pp. 279–290, 2000. View at: Google Scholar
  116. P. Bautista, A. F. Mohedano, J. A. Casas, J. A. Zazo, and J. J. Rodriguez, “An overview of the application of Fenton oxidation to industrial wastewaters treatment,” Journal of Chemical Technology and Biotechnology, vol. 83, no. 10, pp. 1323–1338, 2008. View at: Publisher Site | Google Scholar
  117. Z. M. Li, P. J. Shea, and S. D. Comfort, “Fenton oxidation of 2,4,6-trinitrotoluene in contaminated soil slurries,” Environmental Engineering Science, vol. 14, no. 1, pp. 55–66, 1997. View at: Google Scholar
  118. M. Yoshida, B.-D. Lee, and M. Hosomi, “Decomposition of aqueous tetrachloroethylene by Fenton oxidation treatment,” Water Science & Technology, vol. 42, no. 1-2, pp. 203–208, 2000. View at: Google Scholar
  119. K.-D. Zoh and M. K. Stenstrom, “Fenton oxidation of hexahydro-1,3,5-trinitro-1,3,5-triazine (RDX) and octahydro-1,3,5,7-tetranitro-1,3,5,7-tetrazocine (HMX),” Water Research, vol. 36, no. 5, pp. 1331–1341, 2002. View at: Publisher Site | Google Scholar
  120. J. A. Zazo, J. A. Casas, A. F. Mohedano, and J. J. Rodríguez, “Catalytic wet peroxide oxidation of phenol with a Fe/active carbon catalyst,” Applied Catalysis B, vol. 65, no. 3-4, pp. 261–268, 2006. View at: Publisher Site | Google Scholar
  121. C. B. Molina, J. A. Casas, J. A. Zazo, and J. J. Rodríguez, “A comparison of Al-Fe and Zr-Fe pillared clays for catalytic wet peroxide oxidation,” Chemical Engineering Journal, vol. 118, no. 1-2, pp. 29–35, 2006. View at: Publisher Site | Google Scholar
  122. S. Kanmani and S. Muthuselvi, “Soil remediation by Fenton/photo Fenton oxidation processes,” Journal of the Institute of Public Health Engineers, vol. 2, pp. 5–9, 2005. View at: Google Scholar
  123. P. R. Gogate and A. B. Pandit, “A review of imperative technologies for wastewater treatment I: oxidation technologies at ambient conditions,” Advances in Environmental Research, vol. 8, no. 3-4, pp. 501–551, 2004. View at: Publisher Site | Google Scholar
  124. F. K. Nesheiwat and A. G. Swanson, “Clean contaminated sites using Fenton's reagent,” Chemical Engineering Progress, vol. 96, no. 4, pp. 61–66, 2000. View at: Google Scholar
  125. F. J. Benitez, J. L. Acero, F. J. Real, F. J. Rubio, and A. I. Leal, “The role of hydroxyl radicals for the decomposition of p-hydroxy phenylacetic acid in aqueous solutions,” Water Research, vol. 35, no. 5, pp. 1338–1343, 2001. View at: Publisher Site | Google Scholar
  126. G. Lofrano, S. Meriç, V. Belgiorno, and R. M. A. Napoli, “Fenton's oxidation of various-based tanning materials,” Desalination, vol. 211, no. 1–3, pp. 10–21, 2007. View at: Publisher Site | Google Scholar
  127. M. I. Badawy, M. Y. Ghaly, and T. A. Gad-Allah, “Advanced oxidation processes for the removal of organophosphorus pesticides from wastewater,” Desalination, vol. 194, no. 1–3, pp. 166–175, 2006. View at: Publisher Site | Google Scholar
  128. L. Szpyrkowicz, C. Juzzolino, and S. N. Kaul, “A comparative study on oxidation of disperse dyes by electrochemical process, ozone, hypochlorite and Fenton reagent,” Water Research, vol. 35, no. 9, pp. 2129–2136, 2001. View at: Publisher Site | Google Scholar
  129. F. J. Rivas, F. J. Beltran, J. Frades, and P. Buxeda, “Oxidation of p-hydroxybenzoic acid by Fenton's reagent,” Water Research, vol. 35, no. 2, pp. 387–396, 2001. View at: Publisher Site | Google Scholar
  130. C. W. Jones, “Activation of hydrogen peroxide using inorganic and organic species,” in Applications of Hydrogen Peroxide and Derivatives, J. H. Clark, Ed., chapter 2, pp. 37–78, Royal Society of Chemistry, Cambridge, UK, 1999. View at: Publisher Site | Google Scholar
  131. R. Andreozzi, V. Caprio, A. Insola, and R. Marotta, “Advanced oxidation processes (AOP) for water purification and recovery,” Catalysis Today, vol. 53, no. 1, pp. 51–59, 1999. View at: Google Scholar
  132. E. G. Garrido-Ramírez, B. K. G. Theng, and M. L. Mora, “Clays and oxide minerals as catalysts and nanocatalysts in Fenton-like reactions—a review,” Applied Clay Science, vol. 47, no. 3-4, pp. 182–192, 2010. View at: Publisher Site | Google Scholar
  133. S. Malato, J. Blanco, A. Vidal, and C. Richter, “Photocatalysis with solar energy at a pilot-plant scale: an overview,” Applied Catalysis B, vol. 37, no. 1, pp. 1–15, 2002. View at: Publisher Site | Google Scholar
  134. J. Kiwi, C. Pulgarin, P. Peringer, and M. Grätzel, “Beneficial effects of homogeneous photo-Fenton pretreatment upon the biodegradation of anthraquinone sulfonate in wastewater treatment,” Applied Catalysis B, vol. 3, no. 1, pp. 85–99, 1993. View at: Google Scholar
  135. R. Sauleda and E. Brillas, “Mineralization of aniline and 4-chlorophenol in acidic solution by ozonation catalyzed with Fe2+ and UVA light,” Applied Catalysis B, vol. 29, no. 2, pp. 135–145, 2001. View at: Publisher Site | Google Scholar
  136. M. Tokumura, H. T. Znad, and Y. Kawase, “Decolorization of dark brown colored coffee effluent by solar photo-Fenton reaction: effect of solar light dose on decolorization kinetics,” Water Research, vol. 42, no. 18, pp. 4665–4673, 2008. View at: Publisher Site | Google Scholar
  137. R. Aplin, A. J. Feitz, and T. D. Waite, “Effect of Fe(III)-ligand properties on effectiveness of modified photo-Fenton processes,” Water Science & Technology, vol. 44, no. 5, pp. 23–30, 2001. View at: Google Scholar
  138. I. Muñoz, J. Rieradevall, F. Torrades, J. Peral, and X. Domènech, “Environmental assessment of different solar driven advanced oxidation processes,” Solar Energy, vol. 79, no. 4, pp. 369–375, 2005. View at: Publisher Site | Google Scholar
  139. M. Rodríguez, S. Malato, C. Pulgarin et al., “Optimizing the solar photo-Fenton process in the treatment of contaminated water. Determination of intrinsic kinetic constants for scale-up,” Solar Energy, vol. 79, no. 4, pp. 360–368, 2005. View at: Publisher Site | Google Scholar
  140. P. R. Gogate and A. B. Pandit, “A review of imperative technologies for wastewater treatment II: hybrid methods,” Advances in Environmental Research, vol. 8, no. 3-4, pp. 553–597, 2004. View at: Publisher Site | Google Scholar
  141. G. C. Heng, E. S. Elmolla, and M. Chaudhuri, “Optimization of photo-Fenton treatment of mature landfill leachate,” Nature Environment and Pollution Technology, vol. 11, no. 1, pp. 65–72, 2012. View at: Google Scholar
  142. W. H. Glaze, J. W. Kang, and D. H. Chapin, “The chemistry of water treatment processes involving ozone, hydrogen peroxide and ultraviolet radiation,” Ozone: Science & Engineering, vol. 9, pp. 335–352, 1987. View at: Google Scholar
  143. M. E. T. Sillanpää, T. A. Kurniawan, and W.-H. Lo, “Degradation of chelating agents in aqueous solution using advanced oxidation process (AOP),” Chemosphere, vol. 83, no. 11, pp. 1443–1460, 2011. View at: Publisher Site | Google Scholar
  144. W. Z. Tang, Physico-Chemical Treatment of Hazardous Wastes, CRC Press, London, UK, 2004.
  145. K. B. Ko, J. Y. Lee, Y. H. Yoon et al., “Effects of nitrate on the UV photolysis of H2O2 for 2,4-dichlorophenol degradation in treated effluents,” Desalination and Water Treatment, vol. 2, no. 1–3, pp. 6–11, 2009. View at: Publisher Site | Google Scholar
  146. P. Lei, C. Chen, J. Yang, W. Ma, J. Zhao, and L. Zang, “Degradation of dye pollutants by immobilized polyoxometalate with H2O2 under visible-light irradiation,” Environmental Science & Technology, vol. 39, no. 21, pp. 8466–8474, 2005. View at: Publisher Site | Google Scholar
  147. D. Jiang, S. Zhang, and H. Zhao, “Photocatalytic degradation characteristics of different organic compounds at TiO2 nanoporous film electrodes with mixed anatase/rutile phases,” Environmental Science & Technology, vol. 41, no. 1, pp. 303–308, 2007. View at: Publisher Site | Google Scholar
  148. M. A. Brown and J. E. Casida, “Daminozide: oxidation by photochemically generated singlet oxygen to dimethylnitrosamine and succinic anhydride,” Journal of Agricultural and Food Chemistry, vol. 36, no. 5, pp. 1064–1066, 1988. View at: Google Scholar
  149. R. Bauer, G. Waldner, H. Fallmann et al., “The photo-Fenton reaction and the TiO2/UV process for wastewater treatment—novel developments,” Catalysis Today, vol. 53, no. 1, pp. 131–144, 1999. View at: Google Scholar
  150. F. J. Benitez, J. Beltran-Heredia, J. L. Acero, and F. J. Rubio, “Oxidation of several chlorophenolic derivatives by UV irradiation and hydroxyl radicals,” Journal of Chemical Technology and Biotechnology, vol. 76, no. 3, pp. 312–320, 2001. View at: Publisher Site | Google Scholar
  151. J. Yoon, S. Kim, D. S. Lee, and J. Huh, “Characteristics of p-chlorophenol degradation by photo Fenton oxidation,” Water Science & Technology, vol. 42, no. 3-4, pp. 219–224, 2000. View at: Google Scholar
  152. S. Sabhi and J. Kiwi, “Degradation of 2,4-dichlorophenol by immobilized iron catalysts,” Water Research, vol. 35, no. 8, pp. 1994–2002, 2001. View at: Publisher Site | Google Scholar

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