6 Fenton’s Reagent 6.1 Introduction In 1881, Fenton published a brief description of the powerful oxidizing properties of a mixture of hydrogen peroxide and ferrous salts. This mixture became known as Fenton’s reagent , and the reaction has become known as the Fenton’s reaction . Initially, Fenton applied this reaction to oxidize organic acids such as formic, glycolic, lactic, tartronic, malic, saccharic, mucic, glyc- eric, benzoic, picric, dihydroxytartaric, dihydroxymaleic, and acetylenedi- carboxylic (Fenton, 1900). In the absence of ferrous salt, the degradation of hydrogen peroxide proceeds at very slow rates, with little or no oxidation of the organic acids (Fenton, 1899, 1900). In addition, Cross et al. (1900) further confirmed that ferrous salts significantly enhance the kinetics of hydrogen peroxide decomposition. Goldhammer (1927) investigated the effect of Fenton’s reagent on phenols and found that for each equivalent of Fe 2+ three equivalents of H 2 O 2 were decomposed. They also noted that in concentrated hydrogen peroxide solutions each mole of Fe 2+ decomposed 24 equivalents of hydrogen peroxide. Haber and Weiss (1934) were the first to propose that free radicals existed as intermediates during the chemical reactions in solution. The next year, Haber and Weiss further investigated the Fenton chemistry and concluded that Fenton’s reaction can be expressed as a series of chain reactions with reaction pathways dependent on the concentration of the reactants. The study disproved the original theory of Fenton’s reaction, which suggested that the interaction between an intermediate, six-valent, iron–oxygen com- plex and hydrogen peroxide was the most significant reaction step. In 1934, Haber and Weiss proposed that breaking rate of chain length was increased at lower pH so the propagation cycle was extended before termi- nation. The concentration of free hydroxyl radicals was determined to be directly proportional to the concentration of hydrogen peroxide. Baxendale and Wilson (1957) reported that in an oxygen-free environment Fenton’s reagent initiates very rapid polymerization of methyl acrylate, methacrylic acid, methyl methacrylate, acrylonitrile, and styrene, and the TX69272_C06.fm Page 165 Friday, November 14, 2003 2:09 PM © 2004 by CRC Press LLC 166 Physicochemical Treatment of Hazardous Wastes reaction is a function of the concentration of hydroxyl radicals. In the pres- ence of oxygen, no polymerization occurs. Barb et al. (1951) conducted an extensive investigation of Fenton’s reagent chemistry. When [H 2 O 2 ]/[Fe 2+ ] ratios are low, the reaction rate is second order and stoichiometry is 2[Fe 2+ ] ≅ [H 2 O 2 ]; however, in the presence of polymerizable vinyl compound the reaction remains second order but the stoichiometry changes to [Fe 2+ ] ≅ [H 2 O 2 ]. Thus, they concluded that polymerization of vinyl compounds occurs and results in a polymer with terminal hydroxyl groups. An inhibition effect of hydroxyl radicals due to the higher concentration of hydrogen peroxide was also suggested. To explain this mechanism, it was proposed that hydroxyl radicals react with hydrogen peroxide to form hydrogen diox- ide. This process decreases the hydroxyl radicals generated by the reaction between ferrous iron and hydrogen peroxide. In addition, Barb et al. (1951) suggested that hydrogen dioxide is not a strong oxidizing agent capable of breaking the bonds of vinyl compound or oxidizing other organics. Merz and Waters (1949) showed that oxidation of organic compounds by Fenton’s reagent could proceed by chain as well as non-chain mechanisms, which was later confirmed by Ingles (1972). Kremer (1962) studied the effect of ferric ions on hydrogen peroxide decomposition for Fenton’s reagent. It was confirmed that once ferric ions are produced the ferric–ferric system is catalytic in nature, which accounts for relatively constant concentration of ferrous ion in solutions. In the late 1970s, two major theories were considered: the free radical mechanism by Walling and Cleary (1977) and complex formation by Kremer and Stein (1977). Walling proposed that Fenton’s oxidation predominantly takes place by the free-radical mechanism. On the other hand, Kremer pro- posed that complexation between the iron and the organic molecules has a significant role and thus concluded that both mechanisms occur simulta- neously. In the late 1980s a simultaneous effort was made to apply Fenton’s reagent to the field of environmental science. Various contaminants were studied in the laboratory to determine the optimum conditions. Practical applications of Fenton’s reagent to treat contaminants have also been exam- ined by pilot-plant and continuous treatment systems in textile wastewater, etc.; for example, Bigda (1996) applied Fenton’s reaction to the design of a reactor for treatment of organic contaminants. 6.2 Kinetic Models Although Fenton (1894) studied the violet color in caustic alkali during oxidation of tartaric and racemic acids by ferrous salt and hydrogen peroxide, no reaction kinetic model was offered. Fenton reported that the color disappeared when acid was added. Also, it has been observed that fresh external air is more active than room air. Fenton performed different experiments using various amounts of ferrous and hydrogen peroxides and TX69272_C06.fm Page 166 Friday, November 14, 2003 2:09 PM © 2004 by CRC Press LLC Fenton’s Reagent 167 proposed that iron catalyzed this reaction. For example, a small amount of iron is sufficient to determine oxidation of an unlimited amount of tartaric acid. In tartaric acid, two atoms of hydrogen are removed from a molecule of acid, resulting in the production of dihydroxymaleic acid. Among com- mon oxidants such as chlorine, potassium permanganate, atmospheric oxy- gen, and electrolysis, the most effective oxidizing agent is hydrogen peroxide. Fenton’s work was extended to alcohols (Fenton, 1899) and other organic acids (1900). Attempts to identify the intermediates and products of several organic acids and alcohols were made without success. 6.2.1 Chain Reaction Mechanism by Merz and Waters Goldschmidt and Pauncz (1933) suggested that Fenton’s reaction is a chain reaction involving the same reactive intermediates occurring during catalytic decomposition of H 2 O 2 rather than via formation of peroxides of iron: 2H 2 O 2 = 2H 2 O + O 2 (6.1) It was also shown that the ratio of oxidized alcohol to oxidized Fe 2+ could be greater then one. Baxendale and Wilson (1957) showed that hydroxyl radical initiating the chain polymerization of olefins by hydrogen peroxide was the same process as the rapid oxidation of glycolic acid. Merz and Waters (1947) confirmed that simple water-soluble alcohols are oxidized rapidly by Fenton’s reagent. The primary alcohols are oxidized to aldehydes, which are further oxidized at comparable rates by exactly the same mechanism. Merz and Waters proposed a mechanism of chain oxidation of alcohols and alde- hydes by sodium persulfate, hydrogen peroxide, and an excess of ferrous salt as follows: 1. Chain initiation: Fe 2+ + H 2 O 2 = Fe 3+ + OH – + OH • (6.2) 2. Chain propagation: RCH 2 OH + OH • = R–CHOH • + H–OH (reversible) (6.3) R–CHOH • + HO–OH = R–CHO + OH • + H 2 O (6.4) 3. Chain ending at low substrate concentration: Fe 2+ + OH • = Fe 3+ + OH – (6.5) 4. Chain ending at high substrate (alcohol) concentration: 2R–CHOH• = R–CHO + R–CH 2 OH (disproportionation) (6.6) TX69272_C06.fm Page 167 Friday, November 14, 2003 2:09 PM © 2004 by CRC Press LLC 168 Physicochemical Treatment of Hazardous Wastes In 1949, Merz and Waters determined the values for the ratio of rate constants k 2 / k 3 that indicated which particular radical reduced hydrogen peroxide. Based on the reaction pathways, they classified the reacting com- pounds into two groups. The first group of substrates reacts by chain process. Only a small amount of reducing agent is required. The second group is comprised of substrates that react by non-chain processes — in this case, the oxidation is caused by the hydroxyl radical, and considerable loss of hydroxyl radical occurs. For the first group, the reaction rate can be expressed by Equation (6.7): d[H 2 O 2 ]/d[RH] = 1 + k 2 [Fe 2+ ]/ k 3 [RH] (6.7) For non-chain reactions, the kinetic rates are described by Equation (6.8): d[H 2 O 2 ]/d[RH] = 2 + k 2 [Fe 2+ ]/ k 3 [RH] (6.8) The values for the ratio of rate constants k 2 / k 3 can be determined from the intercept of their graphs. The results will suggest which particular radical reduced hydrogen peroxide. 6.2.2 Redox Formulation by Barb et al. Barb et al. (1951) gave a redox formulation that involves the following reaction sequence: (6.9) (6.10) (6.11) (6.12) (6.13) (6.14) Fe H O Fe HO H 3 22 2 2 •+++ +=++ k 1 Fe H O Fe OH OH 2 22 3–•++ +=++ k 2 HO H O HO HO • 22 2 2 • +=+ k 3 HO Fe Fe H O 2 •3 2 2 +=++ +++ k 4 HO Fe Fe OH 2 •2 3 2 – +=+ ++ k 5 OH Fe Fe OH •2 3 – +=+ ++ k 6 TX69272_C06.fm Page 168 Friday, November 14, 2003 2:09 PM © 2004 by CRC Press LLC Fenton’s Reagent 169 where k 1 and k 2 showed inverse [H + ] dependence. 6.2.3 Complex Mechanism by Kremer and Stein The following scheme was presented by Kremer and Stein (1959), and further elaborated by Kremer (1963): (6.15) (6.16) FeO 3+ + H 2 O 2 = Fe 3+ + H 2 O + O 2 (6.17) Let C 1 = [H + ] and C 2 = [FeO 3+ ], k a and k d showed inverse [H + ] dependence and k b >> k a >> k c , C 1 could be taken as a low concentration intermediate to a good approximation [ C 1 ] = K [H 2 O 2 ][Fe 3+ ], K = k a / k b (6.18) [Fe 3+ ] t = [ C 2 ] + [Fe 2+ ] (6.19) –d[H 2 O 2 ]/dt = k c K[Fe 3+ ] t [H 2 O 2 ] + (k d – k c K)[C 2 ][H 2 O 2 ] (6.20) d[O 2 ]/dt = k d [C 2 ][H 2 O 2 ] (6.21) d[C 2 ]/dt = k c K[Fe 3+ ] t [H 2 O 2 ] – (k d + k c K)[C 2 ][H 2 O 2 ] (6.22) [C 2 ] rises continually during the reaction, approaching a saturation value of k c K[Fe 3+ ] t /(k c K + k d ), and –d[H 2 O 2 ]/dt is always greater than twice d[O 2 ]/ dt. At the end of the reaction, some hydrogen peroxide will be stored as C 2 , and less than 0.5 mol of O 2 will be liberated per mole of H 2 O 2 decom- posed. 6.2.4 Walling’s Modified Kinetic Model Walling and Kato (1971) modified the reaction mechanism proposed by Merz and Waters as follows: Fe 2+ + H 2 O 2 = Fe 3+ + OH – + OH • , k 1 = 76 (6.23) Fe H O FeOOH H 3 22 2+++ ++ ⇔ k k b a FeOOH HO FeO 2–3++ =+ k c TX69272_C06.fm Page 169 Friday, November 14, 2003 2:09 PM © 2004 by CRC Press LLC 170 Physicochemical Treatment of Hazardous Wastes OH • + Fe 2+ = Fe 3+ + OH – , k 2 = 3 × 10 8 (6.24) (6.25) (6.26) , k 3 = 10 7 –10 10 (6.27) (6.28) (6.29) (6.30) where k is in L/mol/s, taken from the literature. The reaction conditions were chosen to minimize the competing processes as follows: HO • + H 2 O 2 = H 2 O + HO 2 , k = (1.2–4.5) × 10 7 (6.31) 2HO • = H 2 O 2 , k = 5.3 × 10 9 (6.32) Thus, the stoichiometry is: R = 2ar (1 – R) + b (6.33) where R = ∆[Fe 2+ ]/2∆[H 2 O 2 ], a = k 2 /Σk 3 , r = [Fe 2+ ]/2[RH], and b = (k 3j + 2k 3k )/ 2Σk 3 . This mechanism is referred to as the free-radical mechanism. 6.2.5 Ingles’ Approach In 1972, Ingles reported his studies of Fenton’s reagent using redox titration. He found evidence in support of Kremer’s complex mechanism theory and concluded that, when suitable complexes are formed, substrates are not oxidized by free radical; rather, electron transfer processes might be OH R H H O R • 2 • += + i i i k 3 OH R H H O R • 2 • += + j j j k 3 OH R H H O R • 2 • +=+ k k k 3 RFeFeproduct 3 4 2 i k • += + ++ 2R product (dimer) • 5 j k = RFe H Fe R H 2 6 3 kk k • += + + + + TX69272_C06.fm Page 170 Friday, November 14, 2003 2:09 PM © 2004 by CRC Press LLC Fenton’s Reagent 171 involved. Fenton’s reaction scheme was modified by Ingles for the case when substrate is present in large amounts in the form of substrate/iron-peroxide complexes. Ingles suggested that electron transfer occurs within this com- plex. (6.34) All substrates were considered to compete as ligands in iron complexes and to modify the reaction characteristics of each other and of the complex. Reaction 6.34 yields hydroxyl radicals, so the free-radical mechanism pro- posed by Walling appeared to be possible; however, Equation (6.35) to Equa- tion (6.38) involve electron transfer and do not lead to formation of hydroxyl radicals. Equation (6.37) and Equation (6.38) involve ionic mechanisms: (6.35) (6.36) (6.37) (6.38) 6.2.6 Transition State Approach by Tang and Huang 6.2.6.1 Competitive Method When modeling oxidation kinetics of chlorophenols by Fenton’s reagent, elementary rate constants are critical to obtain quantitative stoichiometric data in terms of optimal dosages for H 2 O 2 and Fe 2+ to achieve a given removal efficiency. If an elementary reaction rate constant for a given compound is not available in the literature, another method can be used to determine it experimentally. For example, the rate constants of 2,4-dichlorophenol (DCP) and 2,4,6-trichlorophenol (TCP) were determined by an alternative method by Tang and Huang (1996a). The equation used to calculate the rate constants is as follows: R I Fe II OOH R I Fe III O + OH - + . R 2 R 2 R I Fe II OOH . R I Fe III O + + OH - R I Fe II OOH + R I Fe II O + OH - + R I Fe II OOH R I Fe III O + OH - + + R 2 R 2 TX69272_C06.fm Page 171 Friday, November 14, 2003 2:09 PM © 2004 by CRC Press LLC 172 Physicochemical Treatment of Hazardous Wastes (6.39) where: = rate constant between any organic compound and hydroxyl radical. = rate constant between reference compound and hydroxyl rad- ical. [S] = concentration of the substrate at any time. [S 0 ] = initial concentration of the substrate. [R] = concentration of the reference compound at any time. [R 0 ] = initial concentration of the reference compound (2-chlorophenol). In their work, the reference compound is 2-chlorophenol, with a rate constant of 8.2 × 10 9 M –1 s –1 . Either 2,4-DCP or 2,4,6-TCP was mixed with 2-chlorophe- nol in a reactor, separately. Then, H 2 O 2 was mixed with the organic com- pound and the pH was adjusted to 3.5. The organic concentrations were measured by gas chromatography (GC) before and after Fe 2+ was added. The results are shown in Figure 6.1. According to the slopes of the straight line, the rate constants between hydroxyl radicals and 2,4-DCP and 2,4,6- TCP can be determined as 7.22 × 10 9 M –1 s –1 and 6.27 × 10 9 M –1 s –1 , respectively. The hydroxylation rate constants for 2,4-DCP and 2,4,6-TCP are clearly smaller than that for 2-chlorophenol; therefore, increasing chlorine content on the aromatic ring decreases the reactivity of the chlorinated phenols toward hydroxyl radical attack. 6.2.6.2 Dechlorination Kinetic Model 6.2.6.2.1 Pseudo First-Order Kinetic Model When an excess of H 2 O 2 and Fe 2+ is added at constant concentrations to the system, a steady-state concentration of hydroxyl radical can be assumed. The concentration of both H 2 O 2 and Fe 2+ can be considered as constant; therefore, the pseudo first-order kinetic can be developed as follows: Chlorinated phenols + • OH = intermediates (chlorinated aliphatic compounds) (6.40) where: k 1 is the pseudo first-order rate constant of oxidation. k 2 Intermediates + • OH = chloride ion + CO 2 + other products (6.41) kk HO ,S 0 0 HO ,R S]/[S R]/[R •• = ln([ ]) ln([ ]) k HO ,S • k HO ,R • k 1 TX69272_C06.fm Page 172 Friday, November 14, 2003 2:09 PM © 2004 by CRC Press LLC Fenton’s Reagent 173 where k 2 is the pseudo first-order rate constant of dechlorination. The deg- radation kinetics can be modeled as the following: d(CP)/dt = –k 1 (CP) (6.42) d(I)/dt = k 1 (CP) – k 2 (I) (6.43) d(Cl – )/dt = k 2 (I) (6.44) where CP is the concentration of chlorinated phenols at any time t; I is the concentration of intermediates formed at any time t; and Cl – is the concen- tration of chloride ion. The integrated form of the above equation is: (CP)/(CP) 0 = exp(–k 1 t) (6.45) (Cl)/(CP) 0 = 1 + [k 1 exp(–k 2 t) – k 2 exp(–k 1 t)]/(k 2 – k 1 ) (6.46) where (CP) 0 is the initial concentration. Figure 6.2 shows both the experi- mental data and the concentration profile predicted by the kinetic model for the oxidation and dechlorination of 2,4,6-TCP. FIGURE 6.1 The competitive oxidation kinetics of different chlorinated phenols at optimal ratio of H 2 O 2 to Fe 2+ of 25 and optimal pH of 3.5. Experimental conditions: (2-CP) = (2,4-DCP) = (2,4,6-TCP) = 5 × 10 –4 M; H 2 O 2 = 5 × 10 –3 M; Fe(ClO4) 2 = 2 × 10 –4 M; pH = 3; ionic strength = 0.05 M as Na 2 SO 4 . In[(Reference)/(Reference) 0 ] -In[(Substrate)/Substrate) 0 ] TX69272_C06.fm Page 173 Friday, November 14, 2003 2:09 PM © 2004 by CRC Press LLC 174 Physicochemical Treatment of Hazardous Wastes It is important to note that both H 2 O 2 and Fe 2+ have to be overdosed to maintain a steady-state concentration of hydroxyl radical and to obtain a satisfactory approximation of the mathematical model with the experimen- tal data. When H 2 O 2 and Fe 2+ concentrations are 5 × 10 –3 M and 2 × 10 –4 M, respectively, the relative rate constants of 2-chlorophenol (2-CP) and 2,4,6- TCP with respect to 2,4-DCP can be calculated. The oxidation and dechlo- rination constants of 2,4-DCP were found to be 0.995 1/min (k 1 ) and 0.092 1/min (k 2 ), as reported in a previous study (Tang and Huang, 1996). For comparison, Table 6.1 summarizes all the kinetic constants as determined in this study and in the related literature. FIGURE 6.2 Kinetic modeling of 2,4,6-trichlorophenol oxidation and chloride ion dissociation (in the math- ematical model, k 1 = 0.150 1/min and k 2 = 0.032 1/min). Experimental conditions: 2,4,6-TCP = 5 × 10 –4 M; H 2 O 2 = 5 × 10 –3 M, Fe(ClO4) 2 = 2 × 10 –4 M; pH = 3.5; ionic strength = 0.05 M as N 2 SO 4 . TABLE 6.1 Kinetic Rate Constants of Chlorinated Phenols by Fenton’s Reagent 2-CP 2,4-DCP 2,4,6-TCP Elementary rate constants (M –1 s –1 ) 8.2 × 10 9 7.2 × 10 9 6.3 × 10 9 Measured oxidation constants (1/min) 1.666 0.995 0.15 Measured dechlorination constants (1/min) 1.386 0.092 0.032 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1 051015 20 25 30 Time (minutes) Normalized 2,4,6-TCP or Cl- concentration Cl- concentration detected, conc. calculated by model 2.4.6-TCP measured, conc.calculated according to model Cl- 2,4,6-TCP TX69272_C06.fm Page 174 Friday, November 14, 2003 2:09 PM © 2004 by CRC Press LLC [...]... 2004 by CRC Press LLC TX69272_C 06. fm Page 181 Friday, November 14, 2003 2:09 PM Rate Constants of Dechlorination K Cl- (1/s) Fenton’s Reagent 181 0.35 0.3 0.25 0.2 0.15 0.1 0.05 0 0 2 4 6 8 10 12 14 16 18 20 (H2O2)/(CP) 2, 4 6- TCP 2,4-DCP 2-MCP FIGURE 6. 6 Dechlorination rate constants for 2-CP, 2,4-DCP, and 2,4 , 6- TCP vs H2O2/organic concentrations for 2-CP, 2,4-DCP, and 2,4 , 6- TCP Experimental conditions:... TX69272_C 06. fm Page 198 Friday, November 14, 2003 2:09 PM 198 Physicochemical Treatment of Hazardous Wastes The overall oxidation of 2-nitrophenol by OH• radicals could be expressed as follows: −H O 2 N - C 6 H 4 - OH + OH• → O 2 N - C 6 H 3 - (OH)2 → intermediates → CO2 + H2O + H+ + NO2–/NO3– (6. 111) The overall reaction is: C6H5NO3 + 6. OH(3H2O2) + 1/2O2 → 6CO2 + 5H2O2 + HNO3 (6. 112) Reaction (6. 110)... November 14, 2003 2:09 PM 182 Physicochemical Treatment of Hazardous Wastes TABLE 6. 3 Dechlorination Rate Constants during Oxidation of Dichlorophenols Dichlorophenol Rate Constants (×107 1/s) 2,4-Dichlorophenol 2,3-Dichlorophenol 2,5-Dichlorophenol 2 , 6- Dichlorophenol 3,5-Dichlorophenol 3,4-Dichlorophenol 2.28 8 .61 11.90 13. 46 13.82 14.12 atoms on the aromatic ring of the chlorinated phenols have the same... Experimental conditions: 2-CP = 2,4-DCP = 2,4 , 6- TCP = 5 × 10–4 M; H2O2 = 5 × 10–3 M; Fe(ClO4)2 = 2 × 10–4 M; pH = 3.5; ionic strength = 0.05 M as N2SO4 © 2004 by CRC Press LLC TX69272_C 06. fm Page 1 76 Friday, November 14, 2003 2:09 PM 1 76 Physicochemical Treatment of Hazardous Wastes predicting the rate constants of tetra-chlorophenol and penta-chlorophenol due to steric hindrance Figure 6. 3 indicates that... 2,4-Dichlorophenol 2,5-Dichlorophenol Pentachlorophenol o-Cresol m-Cresol p-Cresol 2,4-Dimethyl phenol 2,5-Dimethyl phenol 2-Nitrophenol 4-Nitrophenol 2,4-Dinitrophenol 2,5-Dinitrophenol α-Naphthol β-Naphthol % Oxidized (1 hr, no Fe2+) % Oxidized (1 hr, with Fe2+) 0 9 50 20 0 2 6 0 0 6 10 6 34 0 2 10 8.5 9.5 Source: Data from Barbeni et al., Chemosphere, 16, 2225–2237, 1987 © 2004 by CRC Press LLC 100 100...TX69272_C 06. fm Page 175 Friday, November 14, 2003 2:09 PM Fenton’s Reagent 175 TABLE 6. 2 Relative Ratios of Kinetic Constants Using 2,4-DCP as the Reference Compound (k/k2,4-DCP) 2-CP Elementary rate constants Observed oxidation constants Observed dechlorination constants 2,4-DCP 2,4 , 6- TCP 1.14 1 .67 15.07 1 1 1 0.88 0.15 0.35 Relative Rate Constants to K 2,4-DCP To evaluate the effect of the number of. .. be seen that 3-CP has three ortho and para positions enhanced by OH and Cl directors For 4-CP and 2-CP, however, no position is enhanced by both OH and Cl directors Because of these directors, intermediates with higher degrees of oxidation are expected to be produced in oxidation of 3-CP compared to the oxidation of 4-CP and 2-CP Therefore, the dechlorination rate constants of 4-CP and 2-CP will be smaller... H (OH)ClC6H3H + OH• = (OH)ClC•6H3 (6. 99) OH (OH)ClC6H3H(OH) = (OH)ClC6H3OH + H+ (6. 100) Thus, the degradative oxidation of chlorophenols proceeds by a hydroxylated species (Equation 6. 99 and Equation 6. 100), followed by ring opening to yield aldehydes and ultimate degradation of CO2 and Cl– It was suggested that the first step is the formation of radical cation by acid-catalyzed dehydration of radicals... and 6 Barbeni et al (1987) also compared the half-lives of different chlorophenols They demonstrated that the appearance of chloride ion is independent of the disappearance of parent organic compounds during Fenton’s oxidation TABLE 6. 4 Oxidation of Phenolic Compound, 3:1 Hydrogen Peroxide:Phenol Mole Ratio Phenolic Compound Phenol 2-Chlorophenol 3-Chlorophenol 4-Chlorophenol 2,4-Dichlorophenol 2,5-Dichlorophenol... opt ( Fe ) (6. 82) Therefore, the final form of the optimal ratio between H2O2 and Fe2+ can be expressed as: (H 2O2 ) k = t2 2+ Fe opt kt 3 ( © 2004 by CRC Press LLC ) (6. 83) TX69272_C 06. fm Page 1 86 Friday, November 14, 2003 2:09 PM 1 86 Physicochemical Treatment of Hazardous Wastes Substituting the numerical values into the above equation, we obtain the optimal ratio of H2O2 to Fe2+ . Equa- tion (6. 38) involve electron transfer and do not lead to formation of hydroxyl radicals. Equation (6. 37) and Equation (6. 38) involve ionic mechanisms: (6. 35) (6. 36) (6. 37) (6. 38) 6. 2 .6 Transition. PM © 2004 by CRC Press LLC 166 Physicochemical Treatment of Hazardous Wastes reaction is a function of the concentration of hydroxyl radicals. In the pres- ence of oxygen, no polymerization. unoccupied by chlorine FIGURE 6. 6 Dechlorination rate constants for 2-CP, 2,4-DCP, and 2,4 , 6- TCP vs. H 2 O 2 /organic concentrations for 2-CP, 2,4-DCP, and 2,4 , 6- TCP. Experimental conditions: