Sangchul Hwang
Department of Civil Engineering and Surveying, University of Puerto Rico, Mayagu¨ez, Puerto Rico
INTRODUCTION
Chemical oxidation technologies are defined as the processes that use oxidizing agents to degrade or transform complex hazardous chemicals to simpler nontoxic ones. Advanced oxidation processes (AOPs) constitute, in general, the generation and the use of hydroxyl radicals (OH) to oxidize hazardous chemi-cals, which are otherwise very recalcitrant to conven-tional oxidation processes.
Advanced oxidation processes have been used for the treatment of drinking water, wastewater, and soil=
groundwater contaminated with unwanted and hazardous substances. The processes are, in general, based on the generation and the use of highly reactive hydroxyl radicals (OH) that react indiscriminately with many organic and inorganic substances. This entry provides the readers with an overview of such AOPs in terms of fundamentals of the reaction mechanisms and their application to drinking water, wastewater, and soil=groundwater treatment processes.
REDUCTION AND OXIDATION
This section includes a brief technical discussion on the fundamentals related to the basic reduction and oxidation reactions.
Redox Reaction
A reduction and oxidation (redox, hereafter) reaction is an electron transfer reaction between an oxidizing agent and a reducing agent. Oxidizing agents (oxi-dants) are substances that cause oxidation, whereas reducing agents (reductants) are those that cause reduction. Losing an electron is oxidation and gaining an electron is reduction. Hence, oxidizing agents gain electrons or are reduced and reducing agents lose electrons or are oxidized. Oxidation state is a measure of the charge on an atom in any chemical substance.
More information on the oxidation state of the substance is addressed in the following section.
A redox reaction can be separated into a reduction half-reaction and an oxidation half-reaction. In each of these reactions, the number of electrons lost or gained is equal to the change in oxidation state of the oxidized
or reduced substances. Also, both reactions are needed to be properly balanced.[1,2]
Electrode Potential
The power of an oxidant or a reductant is measured by the electrode potential of the substance.[3] Under standard conditions, the electrode potential is defined as the standard electrode potential, E. Standard conditions are the cases at 25C, 1 atm pressure, and unit activity of all species. Table 1 lists the values of E for some chemical oxidants used in water and wastewater treatment processes.
The standard free energy DGis defined as follows:
DG ¼ nFE ¼ RT ln K ð1Þ where n is the number of electrons transferred in the redox reaction, F is Faraday’s constant (23,062.4 0.3 cal=V=eq), R is the gas constant (1.9872 cal=
deg=mol), and T is the absolute temperature (K). For example, when water containing manganese ions is ozonated under acidic conditions at 25C, a reduction half-reaction and an oxidation half-reaction can be written, respectively, as follows:
5 O½ 3 þ 2Hþ þ 2e ! O2 þ H2O ð2Þ 2 Mn 2þ þ 4H2O ! MnO4 þ 8Hþ þ 5e
ð3Þ Consequently, the net redox reaction for such condi-tions is as follows:
5O3 þ 2Mn2þ þ 3H2O ! 2MnO4 þ 6Hþ5O2 ð4Þ For the first reduction half-reaction [Eq. (2)] E is found to be 2.07 V, and for the second oxidation half-reaction [Eq. (3)] the value is obtained to be ()1.49 V by reversing the sign of the E value for the reduction half-reaction of permanganate under acidic conditions.[1,2] Therefore, the value of E for the net redox reaction is the sum of two values [2.07 þ ()1.49 ¼ 0.58 V]. The values of DG and K of the net redox reaction are calculated from Eq. (1) to be ()133 kcal=mol and 5 1097, respectively, indicating that the reaction is very favored thermodynamically.
Encyclopedia of Chemical ProcessingDOI: 10.1081/E-ECHP-120008088
Copyright # 2006 by Taylor & Francis. All rights reserved. 41
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At conditions where the thermodynamic activities of all substances are not unity (i.e., nonstandard condi-tions), such concentration effects on redox potential are expressed by the Nernst equation as follows:
E ¼ E RT nF ln ared
aox ð5Þ
where ared and aoxrepresent the chemical activities of all of the species that appear on the reduced side and the oxidized side, respectively, in the redox reaction.
Analogous to pH, it is convenient to express the activity of electron as pE because electron activities may vary many orders of magnitude. In this sense, the value of pE (pE if at standard conditions) is defined as follows:
pE ¼ logðaeÞ ¼ E=2:303RT
F ¼ E
0:0591 ð6Þ where ae is the activity of electron in solution.
Accordingly, if the electron activity were increased by a factor of 10, the pE value would be changed by ()1.0. For example, if E (or pE) for Eqs. (2), (3), or (4) is positive, the reaction will proceed to the right direction as written because DG becomes negative.
The greater the positive value of Eor pE, the greater the tendency of the reaction to proceed.
FUNDAMENTALS OF AOP
The use of chemical reagents with high oxidizing potential is the most effective way of oxidizing sub-stances. The most highly oxidizing reagent available is the hydroxyl radical (OH). In this regard, most AOPs are performed in conjunction with the genera-tion ofOH to initiate oxidations.
A radical is a compound containing an atom with a single unpaired electron.[4] Structurally, OH is a highly reactive radical because of its orbital
characteristics. The four outermost orbitals of OH have only seven electrons so that it has a great tendency to gain the eighth one to form the stable state. In the presence of an organic contaminant(s),
OH can abstract a hydrogen atom, thereby provoking contaminant oxidation. Such abstraction reaction is thermodynamically favored strongly, releasing about 119 kcal=mol of energy:
OH þ H ! H2O DH ¼ ðÞ119 kcal=mol ð7Þ
where DH is the enthalpy of the reaction. If the reac-tion has a negative DH, the reacreac-tion releases the heat of the reaction (i.e., exothermic reaction). Conversely, if the reaction has a positive DH, it is called an endothermic reaction.
The required dissociation energy (i.e., DH for breaking bond) for a C–H bond ranges from 80 to 104 kcal=mol, depending on the C–H location from which H is abstracted. In this regard, DH for the chemical oxidation via H abstraction byOH is nega-tive so that the reaction is energetically permitted.[4]
In this section, brief fundamental reaction mecha-nisms for each AOP are addressed. Included as AOPs are individual and combinational processes in the use of ultraviolet (UV) irradiation, catalyzed titanium dioxide oxidation, Fenton’s reagent oxidation, ozona-tion, peroxone oxidaozona-tion, and permanganate oxidation.
Photochemical Oxidation
Ultraviolet and visible lights have sufficient energy to alter the electronic configuration of a molecule that is known to be at its ground state. When a molecule absorbs light of energy hn, one of the two electrons with opposite spin occupying the same orbital can be promoted to a vacant orbital of higher energy. This phenomenon is called electron transition, allowing the molecule to be in an excited state. In some cases, the promoted electron maintains a spin opposite to that of its former partner, resulting in an excited singlet state. At other times, a spin is reversed, leading to an excited triplet state.[4,5]
The relative energies of the ground and excited states are depicted in Fig. 1. The easiest electronic tran-sition is the excitation of a nonbonded electron (n) into a p antibonding orbital (i.e., n ! p transition). The other one is excitation from a p bonding orbital to a pantibonding orbital (i.e., p ! ptransition). Excita-tion of an electron from a s bonding molecular orbital to a s antibonding orbital (i.e., s ! s transition) is impractical. This is, for example, because in dilute aqu-eous solutions the solvent is present at much greater concentrations and would absorb most of the light.[4,5]
Table 1 Standard electrode potentials for selected chemical oxidants
In general, a molecule in the ground state X absorbs hn to produce an excited state Y or Z as follows:
X !hn Y ! products ð8Þ
Y ! Z ð9Þ
Z ! products or X ð10Þ
Typical reactions of excited states include rearrange-ment, abstraction, addition, and cleavage reactions.
Photocatalyzed Titanium Dioxide Oxidation Titanium dioxide (TiO2) has been the most commonly applied photocatalyst because of its chemical stability and low toxicity. The decomposition of organic contaminants in TiO2 suspensions is initiated by photogenerated electron=hole pairs as follows:[6,7]
TiO2 !hn TiO2ðe þ hþÞ ð11Þ The holes (hþ) react with electron donors (H2O and OH) or are adsorbed to TiO2 to produce OH [Eq. (12)]. At the same time, dioxygen molecules react with the electrons to yield superoxide radical anions (O2
) [Eq. (13)], which are in turn protonated to generate hydroperoxy radicals (HO2
) [Eq. (14)].
hþ þ H2Oðor OHsurfaceÞ ! OH þ Hþ ð12Þ
O2 þ e ! O2 ð13Þ
O2 þ Hþ ! HO2 ð14Þ
The surface of TiO2 exhibits the positive charge because of an excess of protons from H2O or in an
acidic solution. In this regard, OH photogenerated on the TiO2surface acts as the major oxidant for the negatively charged contaminants that are attracted to the TiO2 surface via coulombic forces. Contaminants can also be oxidized at the solution bulk phase as follows:
OH þ compounds ! products ð15Þ
Oxidation in Use of Hydrogen Peroxide
The simplest way for the OH generation is the direct photochemical cleavage of H2O2as follows:[8]
H2O2 !hn 2OH ð16Þ
However, the above theoretical amount ofOH quan-tum yield can be reduced to about 0.5 because of the radical–radical recombination [Eq. (17)] and other scavenging compounds, such as H2O2itself [Eq. (18)]
and bicarbonate ions (HCO3) [Eqs. (19–21)]
OH þ OH ! H2O2 ð17Þ
OH þ H2O2 ! H2O þ HO2 ð18Þ
OH þ HCO3 ! CO3 þ H2O ð19Þ
OH þ CO32 ! CO3 þ OH ð20Þ CO3 þ H2O2 ! HCO3 þ O2 þ Hþ ð21Þ Another common method of generating OH for the oxidation of organic compounds is via Fenton’s reagent. Fenton’s reagent oxidation involves the decomposition of hydrogen peroxide in the presence Fig. 1 The relative energies of the ground and excited states.
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of reduced iron salts into OH.[9] Hydrogen peroxide (H2O2) reacts with ferrous iron (Fe2þ) to yield OH and ferric iron (Fe3þ) [Eq. (22)]; Fe3þis reduced back to Fe2þ via reaction with H2O2 [Eq. (23)], the super-oxide radical (O2) [Eq. (24)], or the perhydroxyl radical (HO2) [Eq. (25)], which is the protonated form of O2 (pKa ¼ 4.8). The primary reactions are as follows:
H2O2 þ Fe2þ ! Fe3þ þ OH þ OH ð22Þ H2O2 þ Fe3þ ! Fe2þ þ O2 þ 2Hþ ð23Þ
O2 þ Fe3þ ! Fe2þ þ O2 ð24Þ
HO2 þ Fe3þ ! Fe2þ þ HO2 ð25Þ Hydroxyl radicals are highly reactive and are involved in nonspecific reactions with a wide range of compounds.[10] These reactions involve moderate to moderately high second-order reaction rate constants (107–109M1sec1) between the radical and com-pounds as shown in Table 2 and represent a promising option for the remediation of hazardous chemicals.[11]
Fenton’s reagent remediation is discussed in detail later in this entry.
Ozonation
Ozone (O3) can directly oxidize organic compounds.
Typical direct ozonation involves the insertion of the O3 molecule into unsaturated carbon–carbon bond, resulting in the formation of an ozonide. This direct mechanism is very selective. The other mode of oxidation is through the reaction with OH, which is generated via several O3-based reactions.[12]
Ozone is very unstable in water with the half-life being in the range of seconds to hours. This is because
the bonds that hold the O atoms together in O3 molecule are very weak. The stability of ozone is largely dependent on the pH of the water. In fact, hydroxide ions play an important role and accelerate the O3decomposition as follows:
O3 þ OH ! HO2 þ O2 ð26Þ
O3 þ HO2 ! O2 þ OH þ O2 ð27Þ
O3 þ O2 ! O3 þ O2 ð28Þ
Peroxone Oxidation
Ozone decomposition can be accelerated by the addition of H2O2. The reaction between H2O2 and O3 is known to produce the OH. This reaction is called peroxone.[13] The formation of theOH during peroxone oxidation is as follows:
H2O2 þ H2O ! HO2 þ H3Oþ ð29Þ
O3 þ HO2 ! HO2 þ O3 ð30Þ
HO2 ! Hþ þ O2 ð31Þ
O2 þ O3 ! O3 þ O2 ð32Þ
O3 þ Hþ ! HO3 ð33Þ
HO3 ! O2 þ OH ð34Þ
As such, the peroxone process results in the forma-tion ofOH through the reaction of O3with H2O2. A difference between the ozonation and peroxone process is that the former relies mainly on the direct oxidation by O3, whereas the latter depends primarily on the oxidation with OH. The O3 residual in the peroxone process is short-lived because the O3 decom-position is accelerated by the addition of H2O2, leading to a more reactive and faster oxidation in the peroxone process compared to the ozonation.
Permanganate Oxidation
Oxidation of organic chemicals using permanganate (MnO4, either as potassium permanganate, KMnO4
or sodium permanganate, NaMnO4) involves, in general, direct electron transfer rather than free radical processes. The primary redox reactions for Table 2 Rate constants forOH reactions with selected
environmental contaminants
Reactants Rate constants (M1sec1)
Benzene 7.8 109
2-Chlorophenol 1.2 1010
Naphthalene 5 109
Nitrobenzene 3.9 109
Phenol 6.6 109
Tetrachloroethylene 2.6 109
Toluene 3.0 109
Trichloroethylene 4.2 109
Vinyl chloride 1.2 1010
(From Ref.[10].)
permanganate are pH-dependent as follows:[14] The basic reactions involved in KMnO4oxidation in the presence of water are as follows:
RH2 þ 2KMnO4 þ 4H2O
! ROH þ 2MnO2 þ 2KOH ð38Þ
For example, the stoichiometric reaction for the complete destruction of trichloroethylene (C2HCl3) is as follows:
2KMnO4 þ C2HCl3
! 2CO2 þ 2MnO2 þ 2Kþ þ Hþ þ 3Cl ð39Þ Similar to the scavenging reactions in Fenton systems, in permanganate systems, there is a background oxidant demand that imposes a demand on the permanganate ion, which in turn reduces process efficiency. This demand, resulting from reaction with a wide range of reduced, naturally occurring chemical species, can often be greater than the demand imposed by the target contaminant to be oxidized. Permanganate is more expensive (KMnO4—$1.40=lb, $162=1000 eq;
NaMnO4—$5.95=lb, $620=1000 eq) than H2O2
($0.26=lb; $39=1000 eq). For further information on principles and practices of permanganate oxidation, the readers are highly recommended to refer to Ref.[14].
APPLICATION OF AOP
This section provides brief examples of the application of AOPs to the treatment of drinking water, waste-water, and soil=groundwater contaminated with unwanted and=or hazardous materials.
Drinking Water Treatment
The application of O3 in drinking water treatment is prevalent because of its capability of disinfection and oxidation.[15] Ozone, as a disinfectant, is unstable in water and undergoes reactions with water components, whereas O3 decomposes to OH so that advanced oxidation occurs. Unfortunately, undesired oxidation=
disinfection by-products (DBPs) can be formed from
the reaction of O3andOH with water matrix compo-nents. Such organic and inorganic DBPs can be treated with a subsequent treatment process, such as biological filtration. Care should be taken during ozonation of bromide-containing waters, which forms bromate that is not degraded by biological means. In addition, micro-organisms such as Cryptosporidium parvum oocysts are so resistant against disinfection that higher O3exposures are required and in turn more DBPs are formed.[11]
In an attempt to remove arsenic from potable muni-cipal water and groundwater, Krishna et al. employed the treatment with Fenton’s reagent followed by passing through iron scrap.[16] Their results indicated that the two-stage approach was capable of removing 2.5 mg=L arsenic to lower than the United States Environmental Protection Agency’s (USEPA) guide-line value of 10 mg=L.
Wastewater Treatment
Arslan et al. have investigated the AOPs for the treat-ment of reactive dye wastewater.[7] The investigators found that all investigated AOPs (e.g., UV=H2O2, photo-Fenton, UV=TiO2) were capable of completely decolorizing and partially mineralizing the dye waste-water within 1 hr. Among those processes, photo-Fenton oxidation achieved the highest removal efficiency.
Wanpeng et al. have also applied a Fenton reagent oxidation for the treatment of a dye H-acid containing wastewater.[17] These investigators found that the process not only removed chemical oxidant demand effectively, but also improved the biodegradability by combining with the coagulation process.
Using a batch recycle mode, oily wastewater was treated in the photoassisted advanced oxidation systems.[18] Their results indicated that UV=H2O2
oxidized oily compounds into organic acids with the efficiency being greater at acidic pH. The oxidation rate was enhanced in the presence of Fe3þ.
Ozone, H2O2, and UV irradiation were applied sepa-rately and in all possible combinations for the treatment of textile wastewater.[19]Among the investigated AOPs, the most effective system was the simultaneous use of all three agents (i.e., O3=UV=H2O2). Effective perfor-mance was also achieved in combined treatment O3=H2O2. The authors pointed out that it was advisable to assess the costs of different AOPs to make practical comparison among the treatment schemes.[19]
Soil and Groundwater Treatment
One of the commonly used methods of treating soil and groundwater contaminated with hazardous
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materials is the use of catalyzed hydrogen peroxide in the presence of iron species (i.e., Fenton’s reagent oxida-tion). Watts et al. applied a higher H2O2concentration than would be stoichiometrically required to oxidize pentachlorophenol (PCP) in soil systems to enhance mineralization.[20]With a pH of 2–3 being an optimum parameter, they achieved a rapid decrease in both PCP and total organic carbon contents within 24 hr.
Photochemical oxidation processes have also been used for the remediation of contaminated soil and groundwater environments. For example, Lewis et al.
documented a field evaluation of the UV=oxidation technology (UV=O3=H2O2) to treat volatile organic contaminants in groundwater.[21] Greater than 90%
removal efficiency was reported for most contaminants.
Another AOP commonly used in in situ oxidation of contaminants in soil and groundwater environments is the permanganate oxidation. Although the process is active at most environmental pHs [Eq. (36)], Eqs. (36), (38), and (39) also indicate that manganese dioxide (MnO2) is formed as a reaction by-product. Under most environmental conditions (pH 3.5–12), MnO2is present as a precipitate, i.e., MnO2(s). Under some conditions where background oxidant demand is high and large volumes and concentrations of MnO4
are needed, accumulation of large quantities of MnO2(s) in the sub-surface can have negative consequences. For example, MnO2(s) encrustment on dense nonaqueous phase liquids (DNAPLs) may interfere with mass transfer of DNAPL components to the aqueous phase. Permeabil-ity reductions in the porous media may result from the formation of MnO2(s) on the well screen, sand=gravel pack, or aquifer material. Permeability loss can also result from the ion exchange of Naþor Kþfor divalent cations in the aquifer matrix. Undesirable reaction by-products may result from the oxidation and mobilization of metals or from the heavy metal content of the MnO4
product itself. Finally, there is a secondary drinking water standard established by the USEPA for manga-nese (0.05 mg=L) based on color, staining, and taste.
Manganese imparts a black to brown color to water, can cause black staining on almost everything but glass, and imparts a bitter metallic taste.
Advanced Oxidation Processes as Pretreatment for Biological Treatment
Because of the presence of biorefractory materials, the efficiency of the biological treatments is often hampered.
In general, AOP could reduce the toxicity that would otherwise detrimentally affect the biological processes.
Additionally, AOP could transform biorefractory materials to more biodegradable compounds.[22,23]
As such, the use of AOP as pretreatment for the biological processes is promising. However, it should
be noted that there is a great necessity to improve our understanding of such combined chemical and biological systems. This is because, for example, excess hydrogen peroxide concentration may show toxicity to the micro-organisms, and the by-products produced in the advanced oxidation process may sometimes be toxic to the micro-organisms as well.
FENTON’S REAGENT OXIDATION
This section introduces unconventional reaction mech-anisms and the results of Fenton’s reagent oxidation.
Hydroxyl-Radical-Independent Fenton Remediation
As mentioned earlier, the classical Fenton reagent oxi-dation constitutes the generation of OH as follows:[9]
H2O2 þ Fe2þ ! Fe3þ þ OH þ OH ð22Þ However, because of the small rate constant of the above Fenton reaction (<102M1sec1), the oxidation of Fe2þbyO2[Eq. (40)] may completely overshadow the Fenton reaction and may have a significant role because of its much higher rate constant (107M1sec1).
O2 þ Fe2þ ! Fe3þ þ H2O2 ð40Þ In Fe2þ-aerobic condition, an iron–oxygen complex, perferryl ion can be produced as an intermediate as follows:
Fe2þ þ O2 $ Fe2þO2
$ Fe3þO2
$ Fe3þ þ O2 ð41Þ A common representation of perferryl ion is Fe5þ¼O, with the formal charge of iron as (þ)5. Its high electron affinity may replaceOH as the oxidant.
Another iron–oxygen complex can be ferryl ion as follows: A common representation of ferryl ion is Fe4þ¼O or Fe2þO, with the formal charge of iron (þ)4. Like perferryl ion, the high electron affinity of ferryl ion
Another iron–oxygen complex can be ferryl ion as follows: A common representation of ferryl ion is Fe4þ¼O or Fe2þO, with the formal charge of iron (þ)4. Like perferryl ion, the high electron affinity of ferryl ion