Phsicochemical Treatment of Hazardous Wastes - Chapter 4 docx

4 Advanced Oxidation Processes 4.1 Introduction Any oxidation process in which hydroxyl radical is the dominant species is defined as an advanced oxidation process (AOP). For any oxidation reaction, two factors determine the rate of reaction. First, if a reaction has a high free energy or high electrical potential, the reaction is very likely to occur and it is considered to be thermodynamically favorable. The oxidation potentials for common oxidants suitable for environmental applications are listed in Table 4.1. As can be seen in the table, the hydroxyl radical has an oxidation poten- tial of 2.80 V. The hydroxyl radical is a short-lived and extremely potent oxidizing agent, according to its potential as shown in the table. Because they are extremely potent oxidizing agents, hydroxyl radicals react with organic compounds by three mechanisms: hydrogen abstraction, electron transfer, and hydroxylation (Huang et al., 1993). From a thermodynamic point of view, the higher the oxidation potential is, the stronger the oxidant species will be. Another factor is how fast the reaction is. The fundamental theory under- lining the mechanisms involved in AOPs is the transition state theory (TST), which provides theoretical guidance for the search of the most efficient AOP. According to the TST, hydroxyl radicals may accelerate the oxidation rates of an organic compound by several orders of magnitude compared with oxidation rates for common oxidants. This is because the radical reaction will have a much lower activation energy barrier than regular reactions do; therefore, oxidants such as oxygen, hydrogen peroxide, and ozone are com- bined with catalysts such as transition metals, metal oxides, photons, and ultrasound to generate hydroxyl radicals. For each AOP, the degradation rate is investigated to search for the most efficient process. We begin this chapter with basic chemical kinetics followed by discussion on the TST, oxidants, and catalysts used in AOPs. TX69272_C04.fm Page 85 Wednesday, December 3, 2003 9:41 AM © 2004 by CRC Press LLC 86 Physicochemical Treatment of Hazardous Wastes 4.2 Chemical Kinetics Chemical kinetics focuses on the rate of a reaction through studying the concentration profile with time. Based on the number of reactants involved in the chemical reaction, the reaction can be classified as zero, first, or second order. Third-order reactions are rare because the probability of three reactants colliding and reacting is low. The following are simplified mathematic descriptions of the chemical kinetics of the various orders. 4.2.1 Zero-Order Reactions The rate law for a reaction that is zero order can be expressed as: TABLE 4.1 Oxidation–Reduction Potentials of Chemical Reagents for Water and Wastewater Treatment Reactions Potential in Volts (E°) at 25°C F 2 + 2e = 2F – 2.87 OH • + H + + e – = H 2 O 2.33 O 3 + 2H + + 2e = O 2 + H 2 O 2.07 H 2 O 2 + 2H + + 2e = 2H 2 O 2 (acid) 1.76 MnO 4 – + 4H + + 3e = MnO 2 + 2H 2 O 1.68 HClO 2 + 3H + + 4e = Cl – + 2H 2 O 1.57 MnO 4 – + 8H + + 5e = Mn 2+ + 4H 2 O 1.49 HOCl + H + + 2e = Cl – + H 2 O 1.49 Cl 2 + 2e = 2 Cl – 1.36 HOBr + H + + 2e = Br – + H 2 O 1.33 O 3 + H 2 O + 2e = O 2 + 2 OH – 1.24 ClO 2 (gas) + e = ClO 2 – 1.15 Br 2 + 2e = 2Br – 1.07 HOI + H + + 2e = I – + H 2 O 0.99 ClO 2 (aq.) + e = ClO 2 – 0.95 ClO – + 2H 2 O + 2e = Cl – + 2OH – 0.9 H 2 O 2 + 2H 3 O + 2e = 4H 2 O (basic) 0.87 ClO 2 – + 2H 2 O + 4e = Cl – + 4OH – 0.78 OBr – + H 2 O + 2e = Br – + 4OH – 0.70 I 2 + 2e = 2 I – 0.54 I 3 + 2e = 3 I – 0.53 OI – + H 2 O + 2e = I – + 2OH – 0.49 O 2 + 2H 2 O + 4e = 4OH – 0.40 Source: Lide, D.R. et al., CRC Handbook of Chemistry and Physics , 73rd ed., CRC Press, Boca Raton, FL, 1992. With permission. TX69272_C04.fm Page 86 Wednesday, December 3, 2003 9:41 AM © 2004 by CRC Press LLC Advanced Oxidation Processes 87 (4.1) The above equation can be integrated as: (4.2) Therefore, the time required to reduce the concentration of reactant A to half is: (4.3) As a result, the rate constant can be found as: (4.4) Because the general form of the units of rate constants is (time) –1 (concentra- tion) 1– n , the unit of the rate constant of a zero-order reaction is (time) –1 (con- centration) 1 . The rate of zero-order reaction is independent of the concentration of the reactant, which is often encountered in heterogeneous reactions on the surface such as activated carbon adsorption. 4.2.2 First-Order Reactions The following is a typical first-order reaction: (4.5) The rate law for the first-order reaction is: (4.6) Integrating this equation, the time dependence of concentration A becomes: (4.7) The concentration profile of reactant A is: −= ′ [] = d d A t kA k 0 0 dAkdt t A A =− ∫∫ 0 0 t A k A k 12 00 12 2 / / == k A t = 0 12 2 / AB k → d d A t kA=− [] −=−+ln[ ]Aktconstant TX69272_C04.fm Page 87 Wednesday, December 3, 2003 9:41 AM © 2004 by CRC Press LLC 88 Physicochemical Treatment of Hazardous Wastes (4.8) Therefore, the t 1/2 can be found as follows: (4.9) where the rate constant k is (time) –1 . 4.2.3 Second-Order Reactions The second-order reaction has the following general form: (4.10) The rate law for a second-order reaction can be expressed as: (4.11) Assume the x moles of reactants A and B have been reacted and x moles of C have been produced, then the production rate of C should be: (4.12) (4.13) (4.14) (4.15) (4.16) [] [] A A kt t 0 =− k t = ln / 2 12 AB k +→ C d d A t kA B=− [][] d d x t kA x B x=− −()() 0 1 0 1 dx kA x B x dt ()() 00 −− = dx kA x B x dt xt ()() 00 00 −− = ∫∫ a Ax b Bx AxBx 00 00 1 − + − = −−()() k tB A Bx Ax I=⋅ − − − + 11 00 0 0 ln TX69272_C04.fm Page 88 Wednesday, December 3, 2003 9:41 AM © 2004 by CRC Press LLC Advanced Oxidation Processes 89 (4.17) (4.18) (4.19) where the unit of rate constant k should be (time) –1 (concentration) –1 . 4.2.4 n th Order Reactions For the n th order reaction with respect to one reactant, the general solution for rate expression is: (4.20) A simple integration of this equation results in: (4.21) which can be rewritten as: (4.22) As mentioned before, the unit of the rate constant is (time) –1 (concentra- tion) 1–n . 4.3 Transition State Theory Chemical reactions are studied in terms of elementary reactions involving only one step for bond breaking, bond formation, or electron transfer. A characteristic of elementary reactions is the molecularity. In other words, if kt BA AB x BA x = − − () − () 1 00 00 00 ln A A kA t = + 0 0 1 k tA = 1 12 0 / r A t kA n ==− [] d d 1 1 11 1 0 1 nA A kt t nn − −       = −− [] [] 11 1 1 0 1 [] [] () AA nkt t nn−− −       =− TX69272_C04.fm Page 89 Wednesday, December 3, 2003 9:41 AM © 2004 by CRC Press LLC 90 Physicochemical Treatment of Hazardous Wastes a reaction takes place in a single irreducible act at a molecular level without any detectable intermediates, the reaction is called as an elementary reaction: (4.23) However, when reactants A and B are approaching each other, most of the motions in a reacting molecular system are ordinary vibrations, rotations, and translations. Only one normal mode corresponding to the reaction coor- dinates is involved in breaking or forming a chemical bond to form a new molecule. The new chemical bond results in the rearrangement of atoms. The collection of these atoms is defined as the reactive center. As the distance between the two atoms becomes shorter and shorter, the electron clouds interact with each other due to the rapid motion of electrons. As a result, a multidimensional and continuous potential surface is devel- oped. The potential field becomes stronger and stronger as the two molecules approach each other. On this multidimensional surface exists a most eco- nomic energy path for reactants A and B to interact with each other. The reaction path is referred as the reaction coordinate. Along this coordinate, the highest energy along the most economic reaction path defines an activated complex (AC), which is expressed as [AB] ‡ . The contour plot of such potential energy can be seen in Figure 4.1. The energy between reactants A and B and [AB] ‡ is defined as the activation energy barrier, as shown in Figure 4.2. According to the TST, the following reaction will represent the formation of the activated complex and subsequent production of C: (4.24) The AC usually has several characteristics. The molecules going over the barrier are in equilibrium with all the other reactant molecules, and the AC can be treated as a normal molecular species except that one of its vibrational modes is missing and must be replaced by translation along the reaction coordinate. Also, the rate of formation of product C through the AC is the universal frequency ν =; therefore, the reaction rate k can be expressed as follows: (4.25) because: (4.26) AB C kp +→ AB AB K kp + === [] → ‡ ‡ C kT h B k kT h AB B =⋅ [] ‡ K AB AB ‡ ‡ = [] [][] TX69272_C04.fm Page 90 Wednesday, December 3, 2003 9:41 AM © 2004 by CRC Press LLC Advanced Oxidation Processes 91 Therefore, (4.27) Concentrations A and B in solution should be replaced by the activities of reactants A and B. Then, the rate expression becomes: (4.28) On the other hand, thermodynamics offers the following relationship: (4.29) Therefore, Equation (4.29) becomes: (4.30) FIGURE 4.1 Potential energy surface with a late transition state. (From Boudart, M., Kinetics of Chemical Processes, Prentice Hall, Englewood Cliffs, NJ, 1968. With permission.) r BC r AB Transition State Complex A + BC AB + C k kT h KAB B = [][] ‡ k kT h vv v KAB BAB =⋅ ⋅ [][] ‡ ‡ Ke GRT‡ ‡ = −∆ / k kT h vv v e BAB GRT =⋅ ⋅ − ‡ ‡ ∆ / TX69272_C04.fm Page 91 Wednesday, December 3, 2003 9:41 AM © 2004 by CRC Press LLC 92 Physicochemical Treatment of Hazardous Wastes Furthermore, (4.31) The equation takes the final form as follows: (4.32) Experimentally, the Arrhenius equation describes the relationship between the rate constant k and the temperature as follows: (4.33) By taking the natural logarithms of both sides, the equation becomes: (4.34) FIGURE 4.2 Potential energy profile for elementary reaction. (From Boudart, M., Kinetics of Chemical Processes, Prentice Hall, Englewood Cliffs, NJ, 1968. With permission.) E O,2 =E O,C +E O,D Eo Activation Barrier Transition state: Z Products: C+D Reactants: A+B Reaction coordinate Potential energy Heat of reaction Eo,z E O,1 =E O,A + E O,B ∆∆ ∆GHTS ‡‡ ‡ =− k kT h vv v ee BAB SR HRT =⋅ ⋅ ⋅ − ‡ ‡‡ ∆∆// kAe ERT a =⋅ − / ln ln k E RT A a =− ⋅ + 1 TX69272_C04.fm Page 92 Wednesday, December 3, 2003 9:41 AM © 2004 by CRC Press LLC Advanced Oxidation Processes 93 The above equation suggests that the activation energy barrier can be exper- imentally determined. In addition, if the transition state equation is com- pared with the Arrhenius equation, the following will also be true: (4.35) Because E a is the potential energy of activation, the difference between ∆H ‡ and E a is the kinetic energy of activation, which is usually small compared to E a . For a bimolecular reaction, this difference is RT. As a result, the acti- vation energy will be the same as the change of activation enthalpy: (4.36) It is important to point out that for exothermic reaction, the activation energy ∆E ‡ equals the activation energy: (4.37) However, for endothermic reaction, the activation energy equals the activa- tion barrier plus the heat of reaction ∆H: (4.38) The transition state theory indicates that the rate of a reaction is not a matter of energy alone, but also requires a favorable configuration by a change of entropy. In addition, the rate of a reaction can be speeded up through the following methods. These methods are the guiding principles in the search for the most efficient AOPs: • By increasing the temperature to increase the universal collision frequency ν =, where is the Boltzmann constant (1.38*10 –23 J•K –1 ) and h is the Planck constant (6.63*10 –24 J•s); supercritical water oxidation uses this method. • By increasing the ground energy of reactants using ultraviolet pho- tons and ultrasound to reduce the activation energy barrier. • By increasing pressure to increase the positive entropy of activa- tion, . In supercritical water oxidation, either air or oxygen is used as the oxidant. Although the activation energy barrier is extremely high, ∆S is large due to extremely high pressure. As a result, the reaction is still fast enough to oxidize concentrated organic waste. • By decreasing the enthalpy of the reaction, ∆H ‡ . A kT h vv v e BAB SR =⋅ ⋅ ‡ ‡ ∆ / EH a =∆ ‡ ∆EE a ‡ = ∆∆EE H a ‡ =+ kT h B k B e SR∆ ‡ / TX69272_C04.fm Page 93 Wednesday, December 3, 2003 9:41 AM © 2004 by CRC Press LLC 94 Physicochemical Treatment of Hazardous Wastes The last approach is the fundamental approach used in AOPs. In terms of thermodynamics, ordinary oxidants such as oxygen, ozone, and hydrogen peroxide will form activated complexes with organic pollutants with large enthalpy, ∆H ‡ , and the reactions will be thermodynamically less favorable than reactions between hydroxyl radicals and organic compounds; therefore, one way to increase reaction rates is to convert these oxidants to hydroxyl radicals first. As a result, the enthalpy change when the hydroxyl radical attacks an organic molecule is several orders smaller than when a common oxidant attacks an organic molecule; thus AOPs can be found by any com- bination of oxidants such as oxygen, ozone, and hydrogen peroxide with catalysts such as UV photons, transition metals, and ultrasound. Based on these guiding principles in searching for AOPs, Table 4.2 provides possible AOPs with different combinations of oxidants and catalysts. Because AOPs take advantage of the high reactivity of hydroxyl radicals, initial, propagation, promotion, recombination, and reversible reactions are commonly involved in the degradation of organic pollutants. Table 4.3 lists these major elementary reactions. 4.4 Oxidants 4.4.1 Oxygen Because 20% of air is oxygen, it is not surprising that oxygen is the most common oxidant. The redox reaction of oxygen is: O 2 + 2H 2 O + 4e = 4OH – (4.39) TABLE 4.2 Possible AOPs with Different Combinations of Oxidants and Catalysts Catalyst Metals and Ions Metal Oxides Oxidants Photon Ultra- sound Electron Oxidant Fe 2+ Fe Pt TiO 2 Fe 2 O 3 OH – O 3 H 2 O 2 UV US e – (reductant) O 3 XXX X X X X X H 2 O 2 XXXXX X X X O 2 XXXX X H 2 OX XX TiO 2 X Note: X represents a combination that can generate hydroxyl radicals. TX69272_C04.fm Page 94 Wednesday, December 3, 2003 9:41 AM © 2004 by CRC Press LLC [...]... reductive product of the above reaction The pH effect on the thermodynamics of the reduction of oxygen is summarized in Figure 4. 3 © 20 04 by CRC Press LLC TX69272_C 04. fm Page 96 Wednesday, December 3, 2003 9 :41 AM 96 Physicochemical Treatment of Hazardous Wastes +1.66 +1.08 O2 -0 .05 HO2 +1 .44 · H2O2 +0.7 14 H2O + OH · +2.813 2 H2O pH 0 (1M H+) +2 .40 2 H2O pH 7 +1.985 4 OH- pH 14 (1M OH-) +1.763 +0.695... heat of fusion of H2O2 is 87. 84 cal/g For the liquid/vapor phase relationship for aqueous hydrogen peroxide, partial pressures of the vapors over the liquid are lower than the calculated value for ideal solutions Table 4. 5 shows TABLE 4. 4 Freezing Point of Hydrogen Peroxide H2O2 Concentration (wt%) Freezing Point (°C) H2O2 Concentration (wt%) Freezing Point (°C) 0 10 20 30 35 40 45 45 .2 48 .6 0 -6 .4 -1 4. 6... Oxidation Processes 99 TABLE 4. 6 Total Heat of Vaporization of Aqueous H2O2 Heat of Vaporization (cal/g solution) H2O2 Concentration (wt%) 25°C 60°C 0 20 40 60 80 100 582.1 5 34. 5 503.1 46 0 .4 4 14. 1 362.7 563.2 526 .4 487 .4 446 .0 40 1.3 351.3 Source: Othmer, D.F., Hydrogen peroxide, in Encyclopedia of Chemical Technology, 4th ed., Howe-Grant, M Ed., John Wiley & Sons, New York, 1991 With permission resulting... 107 106 k obs 105 1 04 103 102 10 1.0 0.1 1 2 3 4 5 6 7 pH 8 9 10 11 12 13 FIGURE 4. 4 Effect of pH on the second-order rate constant for the dismutation of superoxide (From Bielski, B.H.J., Reevaluation of the spectral and kinetic properties of •HO2 and of •O 2- free radicals, Photobiol., 28, 645 – 649 , 1978 With permission.) The heat for Reaction (4. 49) is 88 kcal/mol, and for Reaction (4. 50) it is about... -1 4. 6 -2 5.7 33 41 .4 51.7 52.2a 52b 50 60 61 65 70 80 90 100 — –52.2 –55.5 –56.1a 49 40 .3 – 24. 8 –11.5 –0 .43 — a Eutectic Compound, H2O2•2H2O45 (1966) 42 56 Source: Othmer, D.F., Hydrogen peroxide, in Encyclopedia of Chemical Technology, 4th ed., Howe-Grant, M Ed., John Wiley & Sons, New York, 1991 With permission b © 20 04 by CRC Press LLC TX69272_C 04. fm Page 98 Wednesday, December 3, 2003 9 :41 AM 98... five general types, as follows: 1) Decomposition 2H2O2 → 2H2O + O2 (4. 41) 2) Molecular addition H2O2 + Y → Y•H2O2 (4. 42) 3) Substitution H2O2 + RX → ROOH + HX (4. 43) H2O2 + 2RX → ROOR + 2HX (4. 44) 4) H2O2 as a reducing agent H2O2 + Oxidant → Oxidant–H2 + O2 (4. 45) 5) H2O2 as an oxidizing agent H2O2 + Reductant → Reductant–O + H2O (4. 46) While undergoing these reactions, hydrogen peroxide may react as... by 2 .4 kcal/mol and less stable than the H2O•H2O complex by approximately 2.8 kcal/mol 4. 4.2.13 Frequencies Table 4. 9 illustrates the harmonic vibrational frequencies and the corresponding infrared intensities for the OH•H2O2 complex In comparison with the © 20 04 by CRC Press LLC TX69272_C 04. fm Page 1 04 Wednesday, December 3, 2003 9 :41 AM 1 04 Physicochemical Treatment of Hazardous Wastes FIGURE 4. 5... 50.8 (3.8) 3781.6 (–69.5) 3755.5 (–87.2) 149 1 .4 (+32.6) 1319.7 (+17.9) 922.5 (+0.1) 517.0 40 7.5 (+12.7) 260.8 237.1 165.3 129.5 3650.1 (–125.9) 3601 .4 (–105.5) 149 5.7 ( +42 .0) 205.6 (2.9) 239.1 (3.5) 13.5 (0.7) 21.2 (1.6) 8.6 13.2 1316.9 (+20 .4) 109.8 (1.1) 98.6 (1.1) 931.5 (-3 .1) 559.5 41 5.2 ( +48 .3) 281.7 225.1 182.3 143 .5 1.5 (1 .4) 267.3 253 .4 (1.1) 31.5 85.6 3.7 45 .2 1.9 (1.7) 250.1 276.6 (1.2) 51.6... R3 R4 HO O R3 C (II) R4 H O H HO O R3 C O R4 + H2O2 R3 (III) CH HO R4 FIGURE 4. 9 Generation of zwitterion and hydroxy–hydroperoxide (From Langlais, B et al., Ozone in Water Treatment: Application and Engineering, Lewis Publishers, Boca Raton, FL, 1991 With permission.) © 20 04 by CRC Press LLC TX69272_C 04. fm Page 110 Wednesday, December 3, 2003 9 :41 AM 110 Physicochemical Treatment of Hazardous Wastes. .. heat of pure liquid H2O2 to water and oxygen at 25°C is –23 .44 kcal/mol 4. 4.2 .4 Reaction Mechanism Depending on its usage, hydrogen peroxide is a versatile and effective oxidizing agent as a source of active oxygen, compared with molecular oxygen © 20 04 by CRC Press LLC TX69272_C 04. fm Page 99 Wednesday, December 3, 2003 9 :41 AM Advanced Oxidation Processes 99 TABLE 4. 6 Total Heat of Vaporization of Aqueous . 60 –55.5 20 -1 4. 6 61 –56.1 a 30 -2 5.7 65 49 35 33 70 40 .3 40 41 .4 80 – 24. 8 45 51.7 90 –11.5 45 .2 52.2 a 100 –0 .43 48 .6 52 b —— a Eutectic. b Compound, H 2 O 2 •2H 2 O45 (1966) 42 56. Source:. 60°C 0 20 40 60 80 100 582.1 5 34. 5 503.1 46 0 .4 4 14. 1 362.7 563.2 526 .4 487 .4 446 .0 40 1.3 351.3 Source: Othmer, D.F., Hydrogen peroxide, in Encyclopedia of Chemical Technology, 4th ed., Howe-Grant, M. Ed., John. Decomposition 2H 2 O 2 → 2H 2 O + O 2 (4. 41) 2) Molecular addition H 2 O 2 + Y → Y•H 2 O 2 (4. 42) 3) Substitution H 2 O 2 + RX → ROOH + HX (4. 43) H 2 O 2 + 2RX → ROOR + 2HX (4. 44) 4) H 2 O 2 as a reducing