Ozone Reaction Kinetics for Water and Wastewater Systems - Chapter 2 potx

©2004 CRC Press LLC 2 Reactions of Ozone in Water Due to its electronic configuration, ozone has different reactions in water. These fall into three categories: •Oxidation–reduction reactions •Dipolar cycloaddition reactions • Electrophilic substitution reactions A possible fourth way of reaction could be some sort of nucleophilic addition, although it has only been checked in nonaqueous systems. 1 In some cases, free radicals are formed from these reactions. These free radicals propagate themselves through mechanisms of elementary steps to yield hydroxyl radicals. These hydroxyl radicals are extremely reactive with any organic (and some inorganic) matter present in water. 2 For this reason, ozone reactions in water can be classified as direct and indirect reactions. The direct reactions are the true ozone reactions, that is, the reactions the molecule of ozone undergoes with any other type of chemical species (molecular products, free radicals, etc.). The indirect reactions are those between the hydroxyl radical, formed from the decomposition of ozone, or from other direct reactions of ozone, with compounds present in water. It can be said that direct ozone reactions are the initiation step leading to indirect reactions. 2.1 OXIDATION–REDUCTION REACTIONS Redox reactions are characterized through the transfer of electrons from one species (reductor) to another one (oxidant). 3 The oxidizing or reducing character of any chemical species is given by the standard redox potential. Ozone presents one of the highest standards of redox potentials, 4 only lower to those of the fluorine atom, oxygen atom, and hydroxyl radical (see Table 2.1). Because of its high standard redox potential, the ozone molecule presents a high capacity to react with numerous compounds by means of this reaction type. This reactivity is particularly important in the case of some inorganic species such as Fe 2+ or Γ . However, in most of these reactions there is no explicit transfer of electron, but rather an oxygen transfer from the ozone molecule to the other compound. Examples of explicit electron transfer reactions are scarce but the reactions between ozone and the hydroperoxide ion and the superoxide ion radical could be catalogued in this group 6 : (2.1)OHO O HO 32 3 2 +→•+• −− ©2004 CRC Press LLC (2.2) In most of the cases, however, one oxygen atom is transferred as, for example, in the reaction with Fe 2+ : (2.3) Nonetheless, in all these reactions, some atom of the inorganic species goes to a higher valence state, that is, it looses electrons, so that these reactions could theoretically be catalogued as oxidation–reduction reactions since, in an implicit way, there is an electron transfer. The reaction of ozone with nitrite is an example of this. The two half reactions are (2.4) (2.5) The standard redox potential allows checking the possibility that ozone reacts through redox reactions with a given compound. Main electron ion half reactions of ozone in water are reactions (2.5) and (2.6): (2.6) TABLE 2.1 Standard Redox Potential of Some Oxidant Species 5 Oxidant Species E o , Volts Relative Potential of Ozone Fluorine 3.06 1.48 Hydroxyl radical 2.80 1.35 Atomic oxygen 2.42 1.17 Ozone 2.07 1.00 Hydrogen peroxide 1.77 0.85 Hydroperoxide radical 1.70 0.82 Permanganate 1.67 0.81 Chlorine dioxide 1.50 0.72 Hypochlorous acid 1.49 0.72 Chlorine 1.36 0.66 Bromine 1.09 0.53 Hydrogen peroxide 0.87 0.42 Iodine 0.54 0.26 Oxygen 0.40 0.19 OO O O 32 3 2 +•→•+ −− OFe FeO O 3 2 22 +→+ ++ NO H O e NO H 22 3 22 −−−+ +−→+ OH e OHO 322 22++→+ +− OHOe O OH E v o 32 2 2124++→+ = −− . ©2004 CRC Press LLC From these data the importance of pH on ozone redox reactions can be deduced. Detailed information of standard redox potential of different substances can be obtained elsewhere. 3,4 2.2 CYCLOADDITION REACTIONS Addition reactions are those resulting from the combination of two molecules to yield another one. 7 One of the molecules usually presents atoms-sharing more than two electrons (i.e., unsaturated compounds such as olefinic compounds with a carbon double bond) and the other one presents an electrophilic character. These unsaturated compounds present π electrons that in a lesser extent keep bonded the carbon atoms of the double bond. These π electrons are quite available to electrophilic compounds. It can also be said that an addition reaction develops between one base compound (a compound with π electrons) and an acid compound (an electrophilic compound). As a general rule, the following scheme would correspond to an addition reaction: (2.7) In practice, there could be different types of addition reactions such as those between ozone and any olefinic compound. In this case, the reaction follows the mechanism of Criegge 8 that constitutes an example of cycloaddition reaction. The mechanism of Criegge develops through three steps as shown in Figure 2.1. In the first step, a very unstable five-member ring or primary ozonide is formed. 9 This breaks up, in a second step, to give a zwitterion. In the third step, this zwitterion reacts in a different way, depending on the solvent where the reaction develops, on experimental conditions, and on the nature of the olefinic compound. Thus, in a neutral solvent, it decomposes to yield another ozonide, some peroxide or ketone, and polymer substances as shown in Figure 2.2. When the reaction is in a partici- pating solvent (i.e., a protonic or nuclephilic solvent) some oxy-hydroperoxide species is generated (Figure 2.3). Finally, a third possibility is the so-called abnormal ozonolysis that could develop both in participating and nonparticipating solvents. In this way, some ketone, aldehyde, or carboxylic acids can be formed (Figure 2.4). The cycloaddition reaction, then, leads to the break up of both σ and π bonds of the olefinic compound while the basic addition reaction (2.7) leads only to the break up of the π bond. Compounds with different double bonds (C=N or C=O) do not react with ozone through this type of reaction. 10,11 This is not the case of aromatic compounds that could also reacts with ozone through 1,3-cycloaddition reactions leading to the break up of the aromatic ring. However, in these cases, the cycload- dition reaction also is less probable than the electrophilic attack of one terminal oxygen of the ozone molecule on any nucleophilic center of the aromatic compound. The reason of this is due to the stability of the aromatic ring because of the resonance. Notice that the cycloaddition reaction leads to the break up of the aromatic ring, then to the loss of aromaticity, while the electrophilic reaction (see later) retains the aromatic ring. −=−+ →− − −CC XY XC CY ©2004 CRC Press LLC FIGURE 2.1 Criegee mechanism. FIGURE 2.2 Decomposition ways of primary ozonide in an inert solvent. FIGURE 2.3 Decomposition ways of primary ozonide in a participating solvent. C C C C C C Different ways of reaction (see Figures 2.2 to 2.4) III I II C C C C 2 C + 2 CC x ( ) CC CC C -PH C OOH P -P = OH C OOH OH -P = -N H C OOH N H -P = -C O O C OOH O-C -P = -O OOH O O O C O ©2004 CRC Press LLC 2.3 ELECTROPHILIC SUBSTITUTION REACTIONS In these reactions, one electrophilic agent (such as ozone) attacks one nucleophilic position of the organic molecule (i.e., an aromatic compound), giving rise to the substitution of one part (i.e., atom, functional group, etc.) of the molecule. 7 This type of reaction is the base of the ozonation of aromatic compounds such as phenols as shown later. Aromatic compounds are prone to undergo electrophilic substitution reactions rather than cycloaddition reactions because of the stability of the aromatic ring. For example, the benzene molecule is strongly stabilized by the resonance phenomena. The benzene molecule can be represented by different electronic struc- tures that constitute the benzene hybrid. The difference of stability between individ- ual structures and the hybrid is the energy of resonance. In the case of benzene, the individual structure is the cyclohexatriene, and the resonance energy is 36 kcal, that is, the energy difference between those of the cyclohexatriene and the benzene hybrid. This is the reason of the aromatic properties: the higher the resonance energy, the stronger the aromatic properties. The reactions of aromatic compounds depend on these aromatic properties. Thus, after the electrophilic substitution, the aromatic properties are still valid, and the resulting molecules present the aromatic stability. This situation is lost when cycloaddition takes place. In a general way, an aromatic substitution reaction develops in two steps as shown in Figure 2.5 for the case of benzene and one electrophilic agent YZ. In the first step, a carbocation (C 6 H 5 + HY) is formed and, in the second, a proton is taken due to the action of a base compound. FIGURE 2.4 Examples of abnormal ozonolysis. FIGURE 2.5 Basic steps of the aromatic electrophilic substitution reaction. C C C C +C C C R R C Ketone H R C Aldehyde HO R C Acid + H + E Slow H E H E + :N Fast E + H:N ©2004 CRC Press LLC Another important fact to consider is the presence of substituting groups in the aromatic molecule (i.e., phenols, cresols, aromatic amines, etc.). These groups strongly affect the reactivity of the aromatic ring with electrophilic agents. Thus, groups such as HO–, NO 2 – , Cl – , etc., activate or deactivate the aromatic ring for the electrophilic substitution reaction. Also, depending on the nature of the substituting group, the substitution reaction can take place in different nucleophilic points of the aromatic ring. Thus, activating groups promote the substitution of hydrogen atoms from their ortho and para positions with respect to these groups, while the deacti- vating groups facilitate the substitution in the meta position. Table 2.2 shows the effect of different substituting groups on the electrophilic reaction of the benzene molecule. In fact, both the resulting products of the electrophilic substitution reaction and the relative importance of the reaction rate can be predicted after considering the nature of substituting groups. Differences in the rate of substitution reaction should theoretically be due to differences in the slow step of the process, that is, the formation of the carbocation: the higher the stability of the carbocation, the faster the electrophilic substitution reaction rate. The carbocation is a hybrid of different possible structures where the positive charge is distributed throughout the aromatic ring, although positions ortho and para, regarding the substituting group position, present the higher nucleophilic character. As a consequence, these positions have the highest probability to undergo the electrophilic substitution reaction (see Figure 2.6). Factors that affect the spreading of the positive charge are those that stabilize the carbocation or intermediate state. Also, the substituting group can increase or decrease the carbocation stability, depending on the capacity to release or take electrons. From Figure 2.6, it is evident that the stabilizing or destabilizing effect is especially important when the substituting group is bonded to the ortho or para carbon atom with respect to the attacked nucleophilic TABLE 2.2 Activating and Deactivating Groups of the Aromatic Electrophilic Substitution Reaction 7 Groups Action on Reaction Importance –OH – , –O – , –NH 2 , –NHR, –NR 2 Activation Strong –OR, –NHCOR Activation Intermediate –C 6 H 5 , –Alkyl Activation Weak –NO 2 , –NR 3 + Deactivation Strong –C ≡ N, –CHO, –COOH Deactivation Intermediate –F, –Cl, –Br, –I Deactivation Weak FIGURE 2.6 Resonance forms of the hybride carbocation. H E H E H E H E ©2004 CRC Press LLC position. Groups such as alkyl radicals or –OH activate the aromatic ring because they tend to release electrons while groups such as –NO 2 deactivate the aromatic ring since they attract electrons. In the first case, the carbocation is stabilized, while in the second case it is not. For example, in the case of the ozonation of phenols, this property is particularly important due to the strong electron donor character of the hydroxyl group. In addition, the carbocation formed in the case of phenol is a hybrid constituted not only by the contribution of structures I to III (see Figure 2.6) but also by a fourth structure (see Figure 2.7) where the positive charge is on the oxygen atom. Structure IV is especially stable since each atom (except the hydrogen atom) has completed the orbitals (eight electrons). This carbocation is more stable than those from the electrophilic sub- stitution in the benzene molecule (where there is no substituting group) or in the meta position with respect to the –OH group in the molecule of phenol (Figure 2.8). In these two cases, structure IV is not possible, then the ozonation of phenol is faster than that of benzene and goes mainly at ortho and para positions with respect to the –OH group. In fact, literature reports kinetic studies (see Chapters 3 and 5) of the ozonation of aromatic compounds where the rate constant of the direct reactions between ozone and phenol, and ozone and benzene have been found to be 2 × 10 6 and 3 M –1 sec –1 , respec- tively. 12–14 It should be noticed, however, that these values correspond to pH 7 and 20˚C. As shown later, rates of phenol ozonation are largely influenced by the pH of water because of the dissociating character of phenols. More information on the stability of carbocations in electrophilic substitution reactions in different aromatic structures can be obtained from organic chemistry books. 7 In the case of the ozonation of phenol, the mechanism goes through different electrophilic substitution and cycloaddition reactions as shown in Figure 2.9. 15–17 2.4 NUCLEOPHILIC REACTIONS According to the resonance structures of the ozone molecule (see Figure 1.2), there exists a negative charge on one of the terminal oxygen atoms. This fact confers, at FIGURE 2.7 Resonance forms of the carbocation formed during the ozonation of phenol (attack to ortho position). FIGURE 2.8 Resonance forms of the carbocation formed during the ozonation of phenol (attack to meta position). H E H E OH OH H E OH H E OH H E OH H E OH H E OH ©2004 CRC Press LLC least theoretically, a nucleophilic character to the ozone molecule. Thus, ozone could react with molecules containing electrophilic positions. These reactions belong to the nucleophilic addition type, and molecules with double (and triple) bonds between atoms of different electronegativity could theoretically be involved. In the case of ozonation, the nucleophilic activity can be shown in the presence of carbonyl or double and triple carbon nitrogen bonds. 1 Thus, the following example shows two possible ways (nucleophilic and electrophilic) of an ozone attack on a ketone. For example, the nucleophilic reaction of ozone on Schiff bases with carbon-nitrogen double bonds has been reported. Figure 2.10 shows this example. It should be noted, however, that most of the information related to the mechanism of the ozonation of organic compounds has been obtained in an organic medium, and that there is scarce information on this matter when water is the solvent. 2.5 INDIRECT REACTIONS OF OZONE These reactions are due to the action of free radical species coming from the decomposition of ozone in water. The free radical species are formed in the initiation or propagation reactions of the mechanisms of advanced oxidation processes involv- ing ozone and other agents, such as hydrogen peroxide or UV radiation, among others. 18 An advanced oxidation process (AOP) is defined as that producing hydroxyl radicals which are strong oxidant species. 2 In the ozone decomposition mechanisms, FIGURE 2.9 General mechanism of the ozonation of phenol (AO = Abnormal ozonolysis). HO OH O O O O OO HO OH OH OH HO OH OH O O CO 2 + H 2 O C C O HO O HO H 2 O 2 O HO C C H OH C O OH C O OH C O OH C O H H 2 O 2 H 2 O 2 HO C O OH O H C OH O C C O C H H O C C O C OH O C C O OH HO O C C O H HO O C C O H H AO AO AO ©2004 CRC Press LLC the hydroxyl radical is the main responsible species of the indirect reactions. Then, the reaction between the hydroxyl radical and compounds (that could be called pollutants) present in water constitute the indirect reactions of ozone. Numerous studies have been developed to clarify the mechanism of decompo- sition of ozone in water since Weiss in 1934 19 proposed the first model. Today, the mechanism of Staehelin, Hoigné, and Buhler (SHB) 20–23 is generally accepted that ozone follows its decomposition in water, although when pH is high, another mech- anism by Tomiyasu, Fukutomi, and Gordon (TFG) is also considered as the most representative. 24 In Tables 2.3 and 2.4 both mechanisms together to values of the rate constants of their reactions are shown. The reactions of ozone with the hydroxyl and hydroperoxide ions can be con- sidered as the main initiation reactions of the ozone decomposition mechanism in water. However, other initiation reactions develop when other agents, such as UV radiation or solid catalysts are also present. Thus, the direct photolysis of ozone that yields hydrogen peroxide and then free radicals, 26 or the ozone adsorption and decomposition on a catalyst surface to yield active species (in some cases hydroxyl radicals, 27 as will be discussed in other chapters) are also examples of initiation reactions. The reaction of ozone and the superoxide ion radical [reaction (2.2)] is one of the main propagating reactions of the ozone decomposition mechanism. There are also other reactions that lead to the decomposition or stabilization of ozone in water. Thus, substances of different nature can also contribute to the appear- ance or inhibition of free radicals. These substances are called initiators, inhibitors, and promoters of the decomposition of ozone. 21 The initiators are those substances, such as the hydroperoxide ion (the ionic form of hydrogen peroxide) mentioned above, that directly react with ozone to yield the superoxide ion radical [reaction (2.1)]. These reactions are initiation reactions. The superoxide ion radical is the key to propagating free radical species because it rapidly reacts with ozone to yield free FIGURE 2.10 An ozone nucleophilic substitution reaction. (From Riebel, A.H. et al., Ozo- nation of carbon-nitrogen bonds. I. Nucleophilic attack of ozone, J. Am. Chem. Soc ., 82, 1801–1807, 1960. With permission.) + OOO C 6 H 5 CH N R C 6 H 5 CH N R C 6 H 5 CH N R OO O OO O OO O C 6 H 5 CH N R HHH C 6 H 5 CH N R O C 6 H 5 CHO N RC 6 H 5 C NHR O + ©2004 CRC Press LLC radicals, such as the ozonide ion radical [reaction (2.2)] that eventually leads to the hydroxyl radical (see Table 2.3 or 2.4). Promoters are those species that, through their reaction with the hydroxyl radical, propagate the radical chain to yield the key free radical: the superoxide ion radical. Examples of these substances are methanol, formic acid, or some humic substances. 21 Of particular interest is the role of hydrogen peroxide in the mechanism of ozone decomposition. In fact, hydrogen peroxide is the initiating agent of ozone decomposition as proposed by Tomiyasu et al. 24 but it also acts as promoter of ozone decomposition according to the following reactions 28 : TABLE 2.3 Ozone Decomposition Mechanism in Pure Water According to Staehelin, Hoigné, and Bühler 22,23 Reaction Rate constant Reaction # Initiation Reaction a 70 M –1 sec –1 (2.8) Propagation Reactions 7.9 × 10 5 sec –1 25 (2.9) 5 × 10 10 M –1 sec –1 25 (2.10) 1.6 × 10 9 M –1 sec –1 (2.2) 5.2 × 10 10 M –1 sec –1 (2.11) 3.3 × 10 2 sec –1 (2.12) 1.1 × 10 5 sec –1 (2.13) 2 × 10 9 M –1 sec –1 (2.14) 2.8 × 10 4 sec –1 (2.15) Termination Reactions b 5 × 10 9 M –1 sec –1 25 (2.16) b 5 × 10 9 M –1 sec –1 25 (2.17) a Later, Hoigné 6 considered reaction (2.8) should be reaction (2.18)] of Tomiyashu et al. mechanism 24 (see Table 2.4) although the rate constant value kept the same (70 M–1sec–1). This reaction change implies that reaction (2.1), hydrogen peroxide equilibrium reactions (2.22) and (2.23) (see Table 2.4) and reactions between hydrogen peroxide and the hydroxyl radical (2.27) and (2.28) also take part of the mechanism. b Reaction products, H 2 O 2 and O 3 were tentatively proposed. OOH HO O k i 322 1 +→•+• −− HO O H k 22 1 • →•+ −+ OH HO k 22 1 −+ ′ •+  →• OO O O k 32 3 2 2 +•→•+ −− OH HO k 33 3 −+ •+  →• HO O H k 33 4 • →•+ −+ HO HO O k 32 5 • →•+ OHO HO k 34 6 +•→• HO HO O k 422 7 • →•+ HO HO H O O k T 44 223 1 2•+ • →•+ HO HO H O O O k T 43 2223 2 •+ • →•++ [...]... –1sec–1 − − O3 + O2 •  → O3 • +O2  1.6 × 109 M –1sec–1 (2. 2) − 10 O3 • + H2 O → HO • +O2 + OH − 20 –30 M –1sec–1 (2. 19) − − O3 • + HO • → HO2 • +O2 • 6 × 109 M –1sec–1 (2. 20) O3 + HO •  6 → HO2 • +O  3 × 109 M –1sec–1 (2. 21) − HO2 + H + → H2 O2 5 × 1010 M –1sec–1 k k k2 k k11 k k 12 k′ − 12 H2 O2 → HO2 + H + 0 .25 sec–1 25 (2. 9) 25 25 25 (2. 10) (2. 22) (2. 23) Termination Reactions T3 O3 + HO... 104 M –1sec–1 (2. 59) BrO − + H2 O2  → Br − + H2 O  2 × 105 M –1sec–1 (2. 60) BrO − + Br •  → Br − + BrO •  4.1 × 109 M –1sec–1 (2. 61) − 2 BrO • + H2 O  → BrO − + BrO2 + 2 H +  4.9 × 109 M –1sec–1 (2. 62) − BrO2 + HO •  → BrO2 • +OH −  2 × 109 M –1sec–1 (2. 63) 2 BrO2 •  → Br2 O4  1.4 × 109 M –1sec–1 (2. 64) Br2 O4  → 2 BrO2 •  7 × 109 sec–1 (2. 65) − − Br2 O4 + OH −  → BrO3 + BrO2 + H +  7 ×... –1sec–1 (2. 52) 2 Br2− •  → Br • + Br3− •  2 × 109 M –1sec–1 (2. 53) Br2 + H2 O  → HBrO + Br − + H +  8 .24 M –1sec–1 (2. 54) BrO − + Br2− •  → BrO • +2 Br −  8 × 107 M –1sec–1 (2. 55) BrO − + HO •  → BrO • +OH −  4.5 × 109 M –1sec–1 (2. 56) HBrO + HO •  → BrO • + H2 O  2 × 109 M –1sec–1 (2. 57) _ HBrO + O2 •  → Br • +OH − + O2  9.5 × 108 M –1sec–1 (2. 58) HBrO + H2 O2  → Br − + H2 O + H + + O2  7... Reaction products, O2, CO2 and O2–• were tentatively proposed but not confirmed Reactions (2. 27) and (2. 28) (see text) have to be added to this mechanism 7 −1 −1 9 −1 −1 H 1 = 2. 7 × 10 M s HO • + H2O2 k → HO2 • + H2O  − H 2 = 7.5 × 10 M s HO • + HO2 k → HO2 • +OH −  (2. 27) (2. 28) However, as shown in Chapter 8, hydrogen peroxide can also act as indirect inhibitor of the ozone decomposition,... experiments, 21 ºC, pH 6.3–7.9, phosphate buffers adjusted with NaOH and HClO4 2nd part of Reference 22 Batch reactors, conditions as in Reference 52 Spectrophotometric and colorimetric analysis Stopped flow cell Spectrophotometric analysys 20 ºC, pH 12 NaOH, Na2CO3 effect 20 04 CRC Press LLC 60 (19 82) 20 (19 82) 61 (19 82) 62 (1983) 22 (1984) 23 (1984) (21 ) 1985 24 (1985) TABLE 2. 7 (continued) Works on Aqueous Ozone. ..TABLE 2. 4 Ozone Decomposition Mechanism in Pure Water at Alkaline Conditions According to Tomiyasu, Fukutomi, and Gordon24 Reaction Rate Constant Reaction # Initiation Reaction − k8 − 2 − 2 ki 2 O3 + OH  → HO + O2 *  40 M –1sec–1 O3 + HO → HO2 • +O • (2. 18) 2. 2 × 106 M –1sec–1 − 3 (2. 1) Propagation Reactions − HO2 •  9 → O2 • + H +  7.9 × 105 sec–1 ′ − O2 • + H +  9 → HO2 •  5 × 1010... 2. 5 × 109 M –1sec–1 (2. 24) = − C2 HO • +CO3  → OH − + CO3 • * 4 .2 × 108 M –1sec–1 (2. 25) k k − − CO3 • +O3  →(O2 + CO2 + O2 •) * kT 4 No data was given (2. 26) * Carbonates were assumed to be present because of alkaline conditions In fact, reactions (2. 25) and (2. 26) are not true termination reactions since the superoxide ion radical, O2–•, would propagate the radical chain Reaction products, O2,... Technol., 19, 120 6– 121 2, 1985 22 Buhler, R E., Staehelin, J., and Hoigné, J., Ozone Decomposition in Water Studied by Pulse Radiolusis 1 HO2/O 2- and HO3/O 3- as Intermediates, J Phys Chem., 88, 25 60 25 64, 1984 23 Staehelin, J., Buhler, R E., and Hoigné, J., Ozone Decomposition in Water Studied by Pulse Radiolysis 2 OH and HO4 as Chain Intermediates, J Phys Chem., 88, 5999–6004, 1984 24 Tomiyasu, H.,... Akata, A., Kinetics of ozone photolysis in aqueous solution, AIChE J., 42, 328 3– 329 2, 1996 72 Nemes, A., Fábián, I., and Gordon, G., Experimental aspects of mechanistic studies on aqueous ozone decomposition in alkaline solution, Ozone Sci Eng., 22 , 28 7–304, 20 00 73 Hsu, Y et al., Ozone transfer into water in a gas-inducing reactor, Ind Eng Chem Res., 41, 120 – 127 , 20 02 74 von Gunten, U and Laplanche,... –1sec–1 (2. 29) = − CO3 + HO •  → CO3 • +OH − 4 .2 × 108 M –1sec–1 (2. 30) − = C3 HCO3  → CO3 + H + 2. 2 sec–1 * (2. 31) 5 × 1010 * (2. 32) 2. 25 × 104 * (2. 33) − C4 HCO3 + H +  → H2 CO3 5 – 1010 * (2. 34) − HCO3 •  → CO3 • + H + 500 sec–1 * (2. 35) − C5 CO3 • + H +  → HCO3 • 5 × 1010 M –1sec–1 * (2. 36) − − C6 CO3 • + H2 O2  → HCO3 + HO2 • 4.3 × 105 M –1sec–1 (2. 37) CO • + HO  → CO + HO2 • 5.6 . O+→+++ −+ 22 2 2 BrO H O Br H O −− +→+ 22 2 BrO Br Br BrO −− +•→+• 22 22 BrO H O BrO BrO H•+ →++ −−+ BrO HO BrO OH 22 −− +•→•+ 2 224 BrO Br O•→ Br O BrO 24 2 2→• Br O OH BrO BrO H 24 3 2 +→++ −−−+ CO. 0. 42 Iodine 0.54 0 .26 Oxygen 0.40 0.19 OO O O 32 3 2 +•→•+ −− OFe FeO O 3 2 22 +→+ ++ NO H O e NO H 22 3 22 −−−+ +−→+ OH e OHO 322 22 ++→+ +− OHOe O OH E v o 32 2 21 24++→+ = −− . 20 04. 25 (2. 10) 1.6 × 10 9 M –1 sec –1 (2. 2) 20 –30 M –1 sec –1 (2. 19) 6 × 10 9 M –1 sec –1 (2. 20) 3 × 10 9 M –1 sec –1 (2. 21) 5 × 10 10 M –1 sec –1 25 (2. 22) 0 .25 sec –1 25 (2. 23) Termination