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CHAPTER 12 CHEMICAL OXIDATION1 Philip C Singer, Ph.D Professor, Department of Environmental Sciences and Engineering University of North Carolina Chapel Hill, North Carolina David A Reckhow, Ph.D Professor of Civil and Environmental Engineering University of Massachusetts Amherst, Massachusetts Chemical oxidation processes play several important roles in the treatment of drinking water Chemical oxidants are used for the oxidation of reduced inorganic species, such as ferrous iron, Fe(II); manganous manganese, Mn(II); and sulfide, S(-II); and hazardous synthetic organic compounds such as trichloroethylene (TCE) and atrazine Oxidants can also be used to destroy taste- and odor-causing compounds and eliminate color In addition, in some cases, they may improve the performance of, or reduce the required amount of, coagulants Because many oxidants also have biocidal properties, they can be used to control nuisance aquatic growths, such as algae, in pretreatment basins, and may be used as primary disinfectants to meet CT (disinfectant concentration times contact time) requirements (see Chapter 14) These oxidants are often added at the head of the treatment plant, prior to or at the rapid mix basin, but they can also be employed after clarification, prior to filtration, after a substantial portion of the oxidant demand has been removed The most common chemical oxidants used in water treatment are chlorine, ozone, chlorine dioxide, and permanganate Ozone is sometimes used in conjunction with hydrogen peroxide or ultraviolet irradiation to produce radicals that have powerful oxidative properties Mixed oxidant technologies are also available Free chlorine has traditionally been the oxidant (and disinfectant) of choice in the United States, but concerns about the formation of potentially harmful halogenated disinfection by-products (DBPs) produced by reactions between free chlorine and natural organic material (NOM), exacerbated in some cases by the presence of bromide, have caused many water systems to adopt alternative chemical oxidants (and disinfectants) to lower halogenated DBP formation These other oxidants may also react with NOM and bromide to various degrees, depending upon Acknowledgment: We would like to thank Dr William H Glaze of the University of North Carolina at Chapel Hill who wrote the earlier version of this chapter, which provided a starting point for the current material 12.1 12.2 CHAPTER TWELVE the properties of the oxidant, to form oxidation by-products, some of which also have adverse public health effects or result in downstream operational problems in the treatment plant or distribution system This chapter reviews thermodynamic and kinetic principles associated with the use of chemical oxidants in general, the types and properties of the chemical oxidants used in water treatment, specific applications of oxidation processes for the treatment of drinking water, and the formation and control of oxidation and disinfection by-products Comparisons among the different oxidant choices are presented where information is available PRINCIPLES OF OXIDATION Thermodynamic Considerations Thermodynamics establishes the bounds or constraints for oxidation reactions Chemical kinetics fills in much of the detail In many cases there are simply no other available data than the thermodynamic enthalpies and entropies of reaction Despite its limitations, the domain of thermodynamics is where one must begin the task of characterizing and understanding oxidation reactions In this section, the most basic thermodynamic concepts relating to oxidation reactions will be presented For a more comprehensive treatment of the subject, there are many excellent textbooks that can be consulted (e.g., Stumm and Morgan, 1996; Pankow, 1991) Electrochemical Potentials Oxidation reactions are often viewed as reactions involving the exchange of electrons Since acids are frequently defined as proton donors and bases as proton acceptors, one can think of oxidants as electron acceptors and reductants as electron donors In fact, it’s not quite this simple Many oxidants actually donate an electron-poor element or chemical group, rather than simply accept a lone electron Nevertheless, it’s useful to treat all oxidation reactions as simple electron transfers for the purpose of balancing equations and performing thermodynamic calculations Thermodynamic principles can be used to determine if specific oxidation reactions are possible This generally involves the calculation of some form of reaction potential Although in most cases oxidation equilibria lie very far to one side or the other, it is sometimes instructive to calculate equilibrium concentrations of the reactants and products The first step is to identify the species being reduced and those being oxidized Appropriate half-cell reactions and their standard half-cell potentials (Eored and Eoox , respectively) are available in tables of thermodynamic constants (a few are listed in Tables 12.1 and 12.2) These may be combined to get the overall standard cell potential Eonet (Eq 12.1) Eonet = Eoox + Eored (12.1) Much as a pKa describes the tendency of an acid to give up a hydrogen ion, an electrochemical potential E describes the tendency of an oxidant to take up an electron, or a reductant to give one up The standard-state Gibbs Free Energy of reaction ∆Go is related to the standard electrochemical cell potential by Faraday’s constant F and the number of electrons transferred n: ∆Go = −nFEonet (12.2) For a one-electron transfer reaction, this becomes: ∆Go (K cal) = −23Eonet (volts) (12.3) 12.3 CHEMICAL OXIDATION TABLE 12.1 Standard Half-Cell Potentials for Chemical Oxidants Used in Water Treatment Oxidant Eored, volts Reduction half-reaction ⁄2O3(aq) + H + e → 1⁄2O2(aq) + 1⁄2H2O OH + H+ + e− → H2O ⁄2H2O2 + H+ + e− → H2O ⁄3MnO4− + 4⁄3H+ + e− → 1⁄3MnO2(s) + 2⁄3H2O ClO2 + e− → ClO2− ⁄2HOCl + 1⁄2H+ + e− → 1⁄2Cl− + 1⁄2H2O ⁄2OCl− + H+ + e− → 1⁄2Cl− 1⁄2H2O ⁄2HOBr + 1⁄2H+ + e− → 1⁄2Br− + 1⁄2H2O ⁄2NH2Cl + H+ + e− → 1⁄2Cl− + 1⁄2NH4+ ⁄4NHCl2 + 3⁄4 H+ + e− → 1⁄2Cl− + 1⁄4NH4+ ⁄4O2(aq) + H+ + e− → 1⁄2H2O + Ozone Hydroxyl radical Hydrogen peroxide Permanganate Chlorine dioxide Hypochlorous acid Hypochlorite ion Hypobromous acid Monochloramine Dichloramine Oxygen − 2.08 2.85 1.78 1.68 0.95 1.48 1.64 1.33 1.40 1.34 1.23 Sources: Lide (1995); American Water Works Assoc (1990); Stumm and Morgan (1996) Classical thermodynamics indicates that reactions with a negative Gibbs Free Energy (or a positive Eo) will spontaneously proceed in the direction as written (i.e., from left to right), and those with a positive value (or negative Eo) will proceed in the reverse direction Consider a generic oxidation reaction: aAox + bBred → aAred + bBox (12.4) where substance A picks up one electron from substance B In order to determine which substance is being reduced and which is being oxidized, one must calculate and compare oxidation states of the reactant atoms and product atoms The equilibrium constant K for this reaction defines the concentration quotient for the reactants and products at equilibrium: [Ared]a[Box]b K = ᎏᎏ [Aox]a[Bred]b (12.5) The overall standard cell potential is then directly related to this equilibrium constant by: RT Eonet = ᎏ ln K nF (12.6) TABLE 12.2 Standard Half-Cell Potentials for Some Oxidation Reactions That Can Occur During Drinking Water Treatment E oox, volts Oxidation half-reaction ⁄2Br + ⁄2H2O → ⁄2HOBr + ⁄2H + e ⁄2Mn+2 + H2O → 1⁄2MnO2(s) + 2H+ + e− Fe+2 + 3H2O → Fe(OH)3(s) + 3H+ + e− ⁄8NH4+ + 3⁄8H2O → 1⁄8NO3− + 11⁄4H+ + e− ⁄2NO2− + 1⁄2H2O → 1⁄2NO3− + H+ + e− ⁄8H2S + 1⁄2H2O → 1⁄8SO4−2 + 11⁄4H+ + e− ⁄2H2S → 1⁄2S(s) + H+ + e− ⁄2HCOO− → 1⁄2CO2(g) + 1⁄2H+ + e− 1 − 1 + − −1.33 −1.21 −1.01 −0.88 −0.84 −0.30 −0.14 +0.29 Sources: Lide (1995); American Water Works Assoc (1990); Stumm and Morgan (1996) 12.4 CHAPTER TWELVE and for a one-electron-transfer reaction at 25°C, this simplifies to: log K = ᎏ Eonet 0.059 (12.7) Oxidation-Reduction Reactions Oxidation State Oxidation state is characterized by an oxidation number, which is the charge one would expect for an atom if it were to dissociate from the surrounding molecule or ion (assigning any shared electrons to the more electronegative atom) Oxidation number may be either a positive or a negative number—usually an integer between −VII and +VII, although in their elemental forms, for example, S(s), O2(aq), atoms have an oxidation number of zero This concept is useful in balancing chemical equations and performing certain calculations The rules for calculating oxidation number are described in most textbooks on general chemistry Balancing Equations The first step in working with oxidation reactions is to identify the role of the reacting species At least one reactant must be the oxidizing agent (i.e., containing an atom or atoms that become reduced), and at least one must be a reducing agent (i.e., containing an atom or atoms that become oxidized) The second step is to balance the gain of electrons from the oxidizing agent with the loss of electrons from the reducing agent Next, oxygen atoms are balanced by adding water molecules to one side or another, and hydrogens are balanced with H+ ions For a more detailed treatment on calculations using oxidation reactions, the reader is referred to a general textbook on aquatic chemistry (e.g., Stumm and Morgan, 1996; Pankow, 1991) As an example consider the oxidation of manganese by ozone (Eq 12.8) The substance being oxidized is manganese (i.e., the reducing agent) and the one doing the oxidizing (i.e., being itself reduced) is ozone Mn + O3 → products (12.8) Next, the products formed need to be evaluated It might be known from experience that reduced soluble manganese (i.e., Mn+2) can be oxidized in water to the relatively insoluble manganese dioxide It might also be known that ozone ultimately forms hydroxide and oxygen after it becomes reduced Mn+2 + O3 → MnO2 + O2 + OH− (12.9) The next step is to determine the oxidation state of all atoms involved (Eq 12.10) −II +I } } +IV −II } } } } } +II Mn + O3 → MnO2 + O2 + OH− +2 (12.10) From this analysis, it is clear that manganese is oxidized from +II to +IV, which involves a loss of two electrons per atom On the other side of the ledger, the ozone undergoes a gain of two electrons per molecule, as one of the three oxygen atoms goes from an oxidation state of to −II The two half-reactions can be written as single electron transfers These half-reactions are balanced by adding water molecules and H+ ions to balance oxygen and hydrogen, respectively Mn+2 + H2O → ᎏ12ᎏMnO2 + 2H+ + e− ᎏᎏ (12.11) By convention, when hydroxide appears in a half-reaction, additional H+ ions are added until all of the hydroxide is converted to water This is done to the reduction half-reaction O3 + H+ + e− → + ᎏ12ᎏO2 + ᎏ12ᎏH2O ᎏᎏ (12.12) 12.5 CHEMICAL OXIDATION From this point, it is a simple matter of combining the equations and canceling out terms or portions of terms that appear on both sides At the same time, the standard electrode potentials can be combined to get the overall potential Mn+2 + H2O → ᎏ12ᎏMnO2(S) + 2H+ + e− −1.21 V (Eoox) O3(aq) + H+ + e− → ᎏ12ᎏO2(aq) + ᎏ12ᎏH2O +2.04 V (Eored) O3(aq) + ᎏ12ᎏMn+2 + ᎏ12ᎏH2O → ᎏ12ᎏO2(aq) + ᎏ12ᎏMnO2(S) + H+ +0.83 V (Eonet) ᎏᎏ ᎏᎏ ᎏᎏ (12.13) Immediately, it is seen that this reaction will proceed toward the right (the Eonet is positive) But how far to the right will it go? To answer this, Eq 12.7 is rearranged to get o K = e16.95E net (12.14) K = e16.95*0.83 = 1.29 × 106 (12.15) So for this reaction and using the concentration quotient from the reaction stoichiometry, [O2(aq)]0.5[MnO2(s)]0.5[H+] 1.29 × 106 = ᎏᎏᎏ [O3(aq)]0.5[Mn+2]0.5[H2O]0.5 (12.16) Because the activity of solvents (i.e., water) and solid phases are, by convention, equal to 1, [O2(aq)]0.5[H+] 1.29 × 106 = ᎏᎏᎏ [O3(aq)]0.5[Mn+2]0.5 (12.17) Furthermore, if the pH is 7.0 and a dissolved oxygen concentration of 10 mg/L and an ozone concentration of 0.5 mg/L is maintained in the contactor, an equilibrium Mn+2 concentration of 1.8 × 10−25 M or about 10−27 mg/L can be calculated.Thermodynamic principles therefore indicate that this reaction essentially goes to completion Now, knowing that the Mn+2 should react essentially completely to form manganese dioxide, it might be desirable to determine if ozone can possibly oxidize the manganese dioxide to a higher oxidation state, that is, to permanganate To examine this, the preceding ozone equation must first be combined with the reverse of the permanganate reduction equation (from Table 12.1) O3(aq) + H+ + e− → ᎏ12ᎏO2(aq) + ᎏ12ᎏH2O ᎏᎏ MnO2 + ᎏ23ᎏH2O → ᎏ13ᎏMnO4− + ᎏ43ᎏH+ + e− ᎏᎏ O3(aq) + ᎏ13ᎏMnO2 + ᎏ13ᎏH2O → ᎏ13ᎏMnO4− + ᎏ13ᎏH+ + ᎏ12ᎏO2(aq) ᎏᎏ (12.18) This allows the net potential to be calculated: Eonet = Eoox + Eored = (−1.68V) + (+2.04V) = +0.36V (12.19) Again, this is a favorable reaction The equilibrium constant is: 1 log K = ᎏ Eonet = ᎏ (+0.36V) = 6.1 0.059 0.059 K = 1.26 × 106 (12.20) 12.6 CHAPTER TWELVE The equilibrium quotient can now be formulated directly from the balanced equation Note that neither manganese dioxide (MnO2) nor water (H2O) appears in this quotient This is because both are presumed present at unit activity Manganese dioxide is a solid and as long as it remains in the system, it is considered to be in a pure, undiluted state The same may be said for water As long as the solutes remain dilute, the concentration of water is at its maximum and remains constant [MnO4−]0.33[H+]0.33[O2]0.5 K = ᎏᎏᎏ = 106.1 [O3]0.5 (12.21) So under typical conditions where the pH is near neutrality (i.e., [H+] = 10−7), dissolved oxygen is near saturation (i.e., [O2(aq)] = × 10−4 M), and the ozone residual is 0.25 mg/L (i.e., [O3(aq)] = × 10−6 M), the expected equilibrium permanganate concentration should be: [MnO4−]0.33[10−7]0.33[3 × 10−4]0.5 K = ᎏᎏᎏᎏ = 106.1 [5 × 10−6]0.5 (12.22) and solving for permanganate [MnO4−]0.33 = 3.5 × 107 [MnO4−] = 327 (12.23) Obviously, one cannot have 327 mol/L of permanganate Nevertheless, the system will be forced in this direction so that all of the manganese dioxide would be converted to permanganate Once the manganese dioxide is gone, the reaction must stop As already mentioned, the preceding thermodynamic analysis is quantitatively accurate when all reactions are at equilibrium However, this is rarely the case Many oxidation reactions are quite slow or, in some cases, kinetically unfavored, and the actual concentrations of reactants and products observed during water treatment are far from those predicted by classical thermodynamics For this reason, oxidation chemistry must rely heavily on kinetics Kinetics and Mechanism Reaction Kinetics Thermodynamics indicates whether a reaction will proceed as written However, it will not indicate whether this reaction will produce significant change within milliseconds or thousands of years For this, chemical kinetics must be considered As an example, consider the reaction between hypochlorous acid and bromide ion HOCl + Br− = HOBr + Cl− (12.24) In order for a molecule of hypochlorous acid and a molecule of bromide to combine to form products, the two molecules must come into contact with each other (contact meaning approach within a certain distance so that bonding forces can play a role) The probability that a single HOCl:Br− molecular encounter will occur within any fixed time period is directly proportional to the number of molecules of each type in the system It will also depend on the rate of movement of each of the reactant molecules As a consequence, the rate of formation of products—for example, CHEMICAL OXIDATION 12.7 HOBr—will be dependent on a number of factors, including the concentration of hypochlorous acid and the concentration of bromide in the reacting solution This is the kinetic law of mass action, which is expressed mathematically in Eq 12.25 d[HOBr] ᎏᎏ = kf [HOCl][Br−] dt (12.25) The reactants and products are expressed in molar units of concentration and kf is called the forward reaction rate constant The units for kf are liters/mole per unit time The reaction rate constant is going to be a function of such things as the rate of movement of the molecules and the probability of HOBr formation, given that a collision between hypochlorous acid and bromide has already occurred Because the concentrations of HOCl and Br− that appear in Eq 12.25 are raised to the first power, it is said that this rate law is first order in both reactants The overall order of the reaction is the sum of the individual orders (i.e., second order in this case) In a more general sense, Eq 12.26 is the rate law for any elementary reaction of the type described by Eq 12.27 d[A] − ᎏ = kfa[A]a[B]b dt (12.26) aA + bB → cC + dD (12.27) where the capital letters represent chemical species participating in the reaction and the small letters are the stoichiometric coefficients (i.e., the numbers of each molecule or ion required for the reaction) The overall order describes the extent of dependence of the reaction rate on reactant concentrations For the reaction in Eq 12.27, it is equal to (a + b) The order with respect to species A is a, and the order with respect to species B is b Thus, the reaction in Eq 12.27 is first order in both reactants and second order overall Chemical reactions may be either elementary or nonelementary Elementary reactions are those reactions that occur exactly as they are written, without any intermediate steps These reactions almost always involve just one or two reactants The number of molecules or ions involved in elementary reactions is called the molecularity of the reaction Thus, for all elementary reactions, the overall order equals the molecularity Nonelementary reactions involve a series of two or more elementary reactions Many complex environmental reactions are nonelementary In general, reactions with an overall reaction order greater than or reactions with some noninteger reaction order are nonelementary Reaction rate constants for the various oxidants with similar solutes are often positively correlated In other words, a compound favored for oxidation by one oxidant is generally favored by others as well Those that are relatively resistant to oxidation by one will likewise be unreactive with others A good case study is the extensive research done on the oxidation of phenolic compounds, as presented by Tratnyek and Hoigne (1994).These data highlight the similarities between the chemical structure of a reactant and its reactivity with various oxidants Chemists have used such relationships to develop quantitative structure-activity relationships (QSARs) The Hammett equations are one of the most widely used QSARs (see Brezonik [1994] for more detail on this subject) Temperature Dependence As mentioned previously, the reaction rate constant k is a function of temperature The Arrhenius equation (Eq 12.28) is the classic model that describes this relationship: k = koe−Ea /RT (12.28) 12.8 CHAPTER TWELVE where ko is called the frequency factor or the preexponential factor, Ea is the activation energy, R is the universal gas constant (199 cal/°K-mole), and T is the temperature in °K The values for ko and Ea may be either found in the literature or determined from experimental measurements Types of Reactions To this point, considerations have addressed whether or not a certain oxidation reaction can occur, and perhaps how fast it can occur However, it is sometimes quite useful to know how the reaction occurs on a molecular scale In other words, by what mechanism or pathway does it go from reactants to products? For example, the problem of disinfection by-products is one of chemical pathways There is no inherent problem with oxidizing natural organic matter using chlorine However, when that reaction occurs through addition and substitution reactions (see below) rather than simple electron transfer reactions, chlorinated organic byproducts such as the trihalomethanes (THMs) are obtained Oxidation reactions can generally be categorized as those involving electron transfer and those involving transfer of atoms and groups of atoms.They may also be characterized as reactions involving species with paired electrons (ionic) and those involving unpaired electrons (radical) Aqueous chlorine presents a wide array of ionic reactions (e.g., see Morris [1975]) that will serve as illustrative examples for this discussion Table 12.3 presents a summary of the major types of ionic reactions occurring in drinking water Hypochlorous and hypobromous acid can be added to olefinic bonds (i.e., carbon-carbon double bonds), forming halohydrins This is an electrophilic reaction where the initial attack is by the halogen atom (on the positive side of the HOX dipole) The most stable configuration places the halogen on the carbon with the most hydrogen atoms (producing the most stable carbonium ion: Markovnikov’s rule) The other carbon becomes a carbonium ion, which subsequently reacts with the HO portion of the HOX species or with water Activated ionic substitution can occur with both aromatic and aliphatic compounds As with the addition reactions, this type of reaction will also lead to the formation of organohalide compounds Aromatic substitution reactions occur readily when an electron-donating substituent is bound to the ring Functional groups on the aromatic ring, such as OH and NH2, can be thought of as creating a partial negative charge on the ortho and para positions (second closest and farthest carbon atoms from the functional group, respectively) The halogen end of the HOX molecule attacks one of these carbons Next, there is a loss of the OH end of the molecule, and displacement of the H atom from the carbon under attack Substitution on aliphatic species is also a multistep reaction, as exemplified by the haloform reaction (see next section on pathways) When substitution (or transfer) of a halogen occurs onto a nitrogen atom, a relatively reactive N-halo organic compound results These compounds retain some of the oxidizing capabilities of hypohalous acid and, consequently, the reactions are not considered to cause an oxidant demand The rates of substitution reactions with nitrogenous organic compounds generally increase as the basicity of the nitrogen atom increases Oxidation reactions with halogens are characterized by the formation of the inorganic halide ion, and an oxidized (nonhalogenated) form of the reacting compound With organic compounds, it is quite common to observe addition of an oxygen atom For example, oxidation reactions transform unsaturated hydrocarbons into alcohols, then to aldehydes and ketones, and finally to carboxylic acids Some oxidations not result in a net transfer of atoms For these electron transfer reactions, it is common to form free radical intermediates When this happens, chain reactions can occur, sometimes leading to the types of reactions listed in Table 12.4 12.9 CHEMICAL OXIDATION TABLE 12.3 Major Types of Ionic Reactions Reaction type Example Addition to an olefinic bond Activated aromatic substitution Substitution onto nitrogen Oxidation with oxygen transfer Oxidation with electron transfer In addition to these ionic reactions, there are several reactions involving free radical species that can occur following addition of drinking water oxidants These types of reactions are commonly encountered with ozone, chlorine dioxide, and especially the advanced oxidation processes (see below) For example, addition of ozone will always lead to some decomposition and subsequent formation of hydroxyl radicals (•OH) These reactive species engage in reactions that generally lead to the formation of new free radical species The most common types are addition reactions, hydrogen abstractions, and single electron transfers Catalysis Many types of oxidation reactions are strongly affected by the presence of catalysts These are compounds that alter reaction rates without being formed or consumed in the reaction They typically participate in some key, rate-limiting step and are regenerated during some later step Catalysts generally provide an alternative pathway to a reaction with a lower activation energy Probably the most important catalytic processes involve the participation of acids and bases Specific acid and specific base catalysis involve H+ and OH−, respectively General acid and base catalysis involves any electron acceptor (e.g., a proton donor) and electron donor (e.g., a proton acceptor), respectively.A good example of general base catalysis is the classic haloform reaction (Figure 12.1) Here the rate-limiting 12.10 CHAPTER TWELVE TABLE 12.4 Major Types of Radical Reactions Reaction type Example Radical addition reaction Hydrogen abstraction reaction Radical oxidation reaction with single electron transfer step is the loss of a proton giving the enol While it is shown in Figure 12.1 as occurring by reaction with hydroxide, any strong base will participate Also the base (or hydroxide) consumed in the first step is regenerated in the second Another type of catalysis that is important in oxidation processes comes from the initiation of a free radical chain reaction Examples include the decomposition of ozone by hydroxide and the decomposition of chlorine by iron (e.g., see Brezonik [1994]) In either case, the original oxidant will not react appreciably with recalcitrant compounds such as oxalate However, in the presence of sufficient catalyst, decomposition is initiated, leading to a series of chain propagation reactions whereby oxalate can be easily converted to carbon dioxide Reaction Pathways Oxidation reactions in drinking waters can be very complex They may begin with one of the mechanisms discussed here, but then may be followed by a wide range of nonoxidation processes, such as elimination reactions, hydrolysis reactions, radical chain reactions, and rearrangement reactions The formation of trihalomethanes may occur through many different reaction mechanisms One of the most widely discussed is the haloform reaction (Figure 12.1), which involves the stepwise chlorine substitution of the enolate form of a methyl ketone.This classic reaction begins with a base-catalyzed halogenation ultimately leading to a carboxylic acid and chloroform It is base-catalyzed because the species that reacts with hypochlorous acid is the enol form of the methyl ketone This undergoes electrophilic substitution, forming a monohalogenated intermediate The presence of halogens on this carbon speeds subsequent enolization, which leads to complete halogenation of the α-carbon The resulting intermediate (a trihalogenated acetyl compound) is subject to base-catalyzed hydrolysis,giving a trihalomethane and a carboxylic acid If hypochlorous acid is the only halogenating species, chloroform is the result Many early studies with acetone (propanone) indicated that the rate-limiting step was the initial enolization Once the enolate was formed, the molecule quickly CHEMICAL OXIDATION 12.37 identical to the results for chlorine at an equally high chlorine/carbon ratio Most of the chlorinated organic by-products remained tied up in high-molecular-weight compounds This contrasts with free chlorine, as monochloramine is a much weaker oxidant and is much less likely to create the smaller fragments that can be analyzed by gas chromatography Monochloramine is known to transfer active chlorine to the nitrogen of amines and amino acids forming organic chloramines (Scully, 1986) This reaction also occurs with free chlorine Model compound studies have shown that monochloramine can also add chlorine to activated aliphatic carbon-carbon double bonds (Johnson and Jensen, 1986) The adjacent carbon may become substituted with an amine group or with oxygen Other types of reactions involve the simple addition of amine or chloramine to unsaturated organic molecules Under certain conditions chlorine substitution onto activated aromatic rings has been observed For example, monochloramine will slowly form chlorophenols from phenol (Burttschell et al., 1959) Preformed chloramines will also add chlorine to phloroacetophenone, a very highly activated aromatic compound, to produce chloroform (Topudurti and Haas, 1991) However, the rate of reaction for this compound is low and the molar yield (i.e., moles DBP formed per mole precursor) is only percent as compared to 400 percent for free chlorine Chlorine Dioxide By-products Chlorine dioxide undergoes a wide variety of oxidation reactions with organic matter to form oxidized organics and chlorite (see Table 12.5 and Eq 12.50).The concentrations of chlorite account for 50 to 70 percent of the chlorine dioxide consumed (Rav Acha et al., 1984; Werdehoff and Singer, 1987) Chlorite may also be formed, along with chlorate (ClO3−), by the disproportionation of chlorine dioxide (see Eq 12.49) All three of the oxidized chlorine species (chlorine dioxide, chlorite, and chlorate) are considered to have adverse health effects, and their presence in finished water is a source of concern Chlorine dioxide can also undergo a limited number of chlorine substitution reactions Reactions with specific model compounds demonstrate the formation of chlorinated aromatic compounds and chlorinated aliphatics Trihalomethanes, however, have not been detected as reaction products As with the other chlorinecontaining oxidants, chlorine addition/substitution products are favored at low oxidant-to-carbon ratios and oxidation reactions are favored at high ratios Studies using drinking waters and NOM have shown that small amounts of TOX form upon treatment with typical levels of chlorine dioxide This may be due to the formation of HOCl when chlorine dioxide reacts with NOM, and subsequent reaction with other organics (Werdehoff and Singer, 1987) The relatively small amount of HOCl formed in this manner probably leads to sparsely halogenated macromolecular TOX, which would account for the lack of identifiable organo-halide by-products Ozonation By-products Ozone can lead to the formation of brominated byproducts when applied to waters with moderate to high bromide levels This is a direct result of ozone’s ability to oxidize bromide to hypobromous acid and related species (see Eq 12.66) Some of this oxidized bromine will continue to react to form bromate ion Much of the remaining hypobromous acid will react with NOM, forming brominated organic compounds These by-products encompass the same general classes reported for the halogenated by-products of chlorine—that is, THMs (bromoform), HAAs (dibromoacetic acid), HANs (haloacetonitriles, such as dibromoacetonitrile) One recent study showed that percent of the raw water bromide becomes incorporated as TOX (or total organic bromide, TOBr) following ozonation under conditions typical of drinking water treatment (Song, 1997) 12.38 CHAPTER TWELVE Most ozonation by-products are not halogenated The majority of ozone byproducts are similar to the general oxidation products reported for other disinfectants For example, parallel field studies with ozone and chlorine dioxide found both to produce about the same level of low-molecular-weight-carboxylic acids (Griffini and Iozzelli, 1996) Nevertheless, a number of studies suggest that ozone produces higher levels of simple aldehydes and keto-acids (or aldo-acids) than the other major disinfectants Figure 12.6 presents some dose-specific yield data for these major by-products Other oxidation by-products attributed to ozone include the hydroxy-acids, aromatic acids, and hydroxyaromatics (see Table 12.5) Organic peroxides and epoxides are also expected ozonation by-products, although their detection in treated drinking waters has proved to be a challenge Because of the strong link between ozonation by-products and biodegradable organic matter (BOM) in water, attempts have been made to attribute this BOM to specific compounds Krasner et al (1996) have shown that as much as 40 percent of the assimilable organic carbon and 20 percent of the biodegradable dissolved organic carbon can be assigned to the known major ozonation by-products However, this total still represents less than percent of the dissolved organic carbon (DOC) compound Factors Influencing By-product Formation Time Reaction time is among the most important factors determining DBP concentrations under conditions where a disinfectant residual persists The major halogenated DBPs (e.g., THMs and HAAs) are chemically stable They accumulate in disinfected waters and their concentrations will increase with reaction time for as FIGURE 12.6 Yields of major ozonation by-products normalized per mg of ozone applied, as reported in the literature (after Reckhow, 1999) CHEMICAL OXIDATION 12.39 long as a disinfectant residual exists (Figure 12.7) However, there are cases where HAA concentrations drop to near zero after long residence times in real drinking water distribution systems This phenomenon is generally attributed to biodegradation, and it does not appear to occur with the THMs (the latter appears to be biodegradable only under anaerobic conditions) Laboratory tests have shown that HAAs form more rapidly than THMs Studies have also shown the brominated analogs to form more rapidly than the purely chlorinated compounds This causes the HAA/THM or the TTHM/chloroform ratio to be high in the early stages of the reaction and to slowly drop with reaction time These observations are supported by data collected from full-scale treatment plants and distribution systems In contrast to the case for chlorine, ozonation by-products form quickly and show little long-term increase This is due to the rapid dissipation of ozone residuals Once dissolved ozone is gone, by-product formation can only continue by means of hydrolysis reactions, which represent a relatively minor contribution to the total byproduct concentration Many halogenated DBPs are chemically unstable and are subject to hydrolysis or further oxidation For these compounds, there is a reaction time during which the concentration reaches a maximum (see subsequent discussion on pH effects) Some DBPs, like dichloroacetonitrile, decompose slowly and reach a maximum concentration only after reaction times on the order of days Others are more reactive and decay to undetectable levels within minutes to hours (e.g., 1,1-dichloropropanone) FIGURE 12.7 Effect of reaction time on the major chlorination by-products (from Reckhow and Singer, 1984) 12.40 CHAPTER TWELVE Disinfectant Dose Disinfectant dose has a variable impact on DBP formation Small changes in the dose used for residual disinfection often have minor effects on DBP formation This is because these systems have an excess of disinfectant, and they are therefore precursor-limited.When the residual drops below about 0.3 mg/L, DBP formation becomes disinfectant-limited Under these circumstances, changes in disinfectant dose will have a large effect Figure 12.8 shows that when and mg/L of chlorine are added to Connecticut River water, it produces THMs in direct proportion to those doses However, when an excess is added (>5 mg/L), a residual persists and the THM formation levels off The other two waters in this figure show only the latter behavior, because they had lower TOC values and their chlorine demand was less than the minimum dose tested (3 mg/L) As a general rule, disinfectant dose plays a greater role in DBP formation during primary disinfection than during secondary disinfection This is because primary disinfectants are usually added in amounts well below the long-term demand Therefore, the disinfectant is the limiting reactant, not the organic precursors Figure 12.8 shows a near linear relationship between ozone dose and glyoxalic acid formation When the ozone is applied after coagulation and filtration, the start of a plateau appears This is a reflection of the removal of ozone-demanding organics and a closer approach to precursor-limiting conditions The relationship between disinfectant dose and DBP formation can be illustrated with a simple kinetic model Figure 12.9 shows the results of a kinetic simulation for the simple second-order reaction as follows: k C A + B→ where A represents the NOM precursor material, B is the disinfectant, C is the particular DBP, and k is the second-order rate constant If the initial precursor level (Ao) is held constant, and a rate constant and reaction time are arbitrarily chosen, FIGURE 12.8 Observations on effect of disinfectant dose on DBP formation Chlorination conditions: pH 7, 20°C, 3-day reaction time (Figures redrawn from Coombs, 1990; Reckhow et al., 1993.) CHEMICAL OXIDATION 12.41 FIGURE 12.9 Theoretical second-order kinetic plot showing effect of reactant dose on product formation the DBP formation can be calculated as a function of the disinfectant dose The three curves in Figure 12.9 are the result of this calculation for three different combinations of k and reaction time t As the product of k and t gets large, the DBP formation curve approaches two straight lines (one where DBP formation increases linearly with dose, and one where DBP formation is independent of dose) The upward sloping line corresponds to the disinfectant-limiting case.The horizontal line corresponds to the precursor-limiting case This model is consistent with the experimental observations in Figure 12.8 When using chloramines, the chlorine-to-nitrogen ratio must also be a consideration As the Cl2/N ratio increases from to (weight basis), the water’s exposure to a transient free chlorine residual increases substantially This results in a shift in the DBP character (yields and types) from that typical of pure chloramines to that of free chlorine Model compound studies (DeLaat et al., 1982) have shown that highly reactive precursors (e.g., resorcinol) will experience this shift at lower Cl2/N ratios than the more slowly reacting compounds (e.g., acetone) It’s likely that the same could be said for waters with highly reactive NOM versus less reactive waters pH The overall reaction between chlorine and NOM is relatively insensitive to pH However, the concentration of TOX and specific halogenated by-products is strongly influenced by pH (e.g., Fleischacker and Randtke, 1983; Reckhow and Singer, 1984) Nearly all by-products (THAA, TCP, DHAN, etc.) decrease in concentration with increasing pH One important exception is the THMs Although pH can influence chlorination reactions in many ways, it is probably base-catalyzed hydrolysis mechanisms which have the greatest overall effect Many DBPs (e.g., 1,1,1-trichloropropanone) are decomposed by base-catalyzed hydrolysis These compounds are less prevalent in waters with high pH, and they tend to drop in con- 12.42 CHAPTER TWELVE centration at long residence times (Type 1, Table 12.6) The THMs increase at high pH because many hydrolysis reactions actually lead to THM formation (Type 2) Other by-products are themselves unaffected by base hydrolysis (e.g., the HAAs), but their formation pathways may be altered at high pH, resulting in less compound produced at high pH (Type 3) None of the major ozonation by-products are subject to base hydrolysis Instead, pH plays a role by altering the rate of decomposition of ozone to hydroxyl radicals (see above) As pH increases, ozone decomposition accelerates, and this is thought to be responsible for a decrease in the classical by-products of ozonation (e.g., aldehydes; Schechter and Singer 1995) However, the chemistry of ozone and hydroxyl radical reactions with NOM are not well understood, and there may be cases where some carbonyl by-products increase with pH (e.g., Itoh and Matsuoka, 1996) Bromate is formed in ozonated waters from a series of reactions between ozone or hydroxyl radicals and naturally occurring bromide (see above) A key intermediate in this mechanism is hypobromite ion At lower pHs, more of the hypobromite becomes protonated, forming the less reactive hypobromous acid This causes the overall formation of bromate to decrease as pH is decreased On the other hand, lower pHs favor the formation of TOX (presumably TOBr) during ozonation (Song, 1997) This is probably due to suppressed decomposition of ozone, changes in the speciation of the oxidized bromine, and base-catalyzed hydrolysis of brominated by-products Temperature The rate of formation of DBPs generally increases with increasing temperature Laboratory and full-scale data suggest that chloroform formation is more sensitive to temperature than DCAA formation The relationship is not as clear for TCAA formation At high temperatures, accelerated depletion of residual chlorine slows DBP formation, even though the rate constants for DBP-producing reactions might increase This may be especially true for TCAA, as its formation is more sensitive to chlorine residual than is chloroform or DCAA formation High temperatures may also accelerate DBP degradation processes Biodegradation of HAAs would be expected to proceed more quickly at high temperature Reactive DBPs would undergo abiotic reactions with bases or active chlorine at a faster rate as temperature increases For this reason, increases in temperature may actually cause a decrease in the concentration of certain DBPs Precursor Material It is well established that halogenated DBP formation is, by a first approximation, proportional to the TOC concentration at the point of chlorinaTable 12.6 Impact of Base-Catalyzed Hydrolysis on DBP Concentration Profiles pH of max concentration Drop at high residence times* DBP susceptible to base hydrolysis Low Yes Haloketones, chloral hydrate DBP formed as a result of base hydrolysis High No THMs DBP not susceptible to hydrolysis, but precursors are Low No THAAs Type Characteristics * Not considering possible biodegradation Examples CHEMICAL OXIDATION 12.43 tion This is why, in waters with high DBP precursor concentrations, shifting the point of chlorination from before coagulation/settling to after yields substantial reductions in DBP formation Furthermore, a 40 percent reduction in TOC across conventional treatment often translates to a 50 percent or more reduction in some DBPs when this switch is made This is partly attributed to preferential removal of precursor organics by coagulation and the accompanying shift in the nature of the NOM at the point of chlorination Chlorination studies with NOM extracts and whole waters have suggested that the activated aromatic content is an important determinant in its tendency to form major chlorination by-products (Reckhow et al., 1993) This probably explains why UV absorbance is such a good predictor of a water’s tendency to form THMs and HAAs, as UV absorbance by NOM is generally attributed to activated aromatic chromophores Bromide Bromide is readily oxidized by aqueous chlorine and ozone to form hypobromous acid (Eqs 12.43 and 12.66, respectively) This reacts in concert with other oxidants to form brominated DBPs As bromide levels increase, formation of the brominated DBPs increases When chlorine is the disinfectant, families of mixed bromo-chloro DBPs will result, such as the THMs Figure 12.10 shows how increasing bromide concentration causes shifts in the THM speciation from chloroform to progressively more brominated members of the group A similar effect has been observed for the HAAs, halopicrins, halopropanones, haloacetaldehydes, and haloacetonitriles (e.g., Trehy and Bieber, 1981; Pourmoghaddas and Dressman, 1993; Xie and Reckhow, 1993; Cowman and Singer, 1996) Most laboratory studies have also shown that the molar formation of THMs increases slightly with bromide concentration Since the brominated forms are heavier than their chlorinated analogs, the mass-based THM level (e.g., µg/L) increases even more sharply with increasing bromide level FIGURE 12.10 The effect of bromide concentration on THM speciation (redrawn from Minear and Bird, 1980) 12.44 CHAPTER TWELVE The formation of nonhalogenated DBPs is probably insensitive to varying bromide levels One study of ozonation by-products showed only very minor effects of bromide on aldehyde levels (Schechter and Singer, 1995) Seasonal Effects DBP concentrations are strongly influenced by seasons The underlying factors are temperature and seasonally induced changes in water quality (NOM characteristics, bromide, pH) In general, DBP concentrations are highest in summer This is attributed to high temperatures, which accelerate DBP-forming reactions Furthermore, disinfectant demand reactions are faster, requiring higher disinfectant doses in order to maintain target residuals Competing with this is the fact that chemical degradation is faster at high temperatures and biodegradation may play a larger role With some sources, water quality changes may also be important For example, the San Francisco Bay delta, which provides raw drinking water to a large portion of the population in central and southern California, experiences significant seasonal variations in bromide concentrations During the spring season, high freshwater flows from the Sierras keep salt water, with its concomitant high bromide content, out of the delta so that chlorinated DBPs constitute most of the DBP content of the treated drinking water On the other hand, during the summer and early fall when freshwater flows are low, more saltwater intrusion into the delta occurs and bromine-containing species are the dominant DBPs formed Control of Oxidation/Disinfection By-products There are three general approaches to controlling DBP concentrations: (1) minimize DBP formation by reducing the organic precursor material at the point of disinfection; (2) minimize DBP formation by reducing the disinfectant dose, changing the nature of the disinfectant, or optimizing the conditions of disinfection; and (3) removal of DBPs after their formation Most efforts have focused on the first approach, as represented by the interest in “enhanced coagulation.” However, DBP control strategies should also consider process changes in accord with the second approach The third approach (DBP removal) is most appropriate for control of biodegradable ozonation by-products Minimizing Organic Precursors The objective of this particular strategy is to minimize the quantity and reactivity of the organic precursors at the point of disinfection This can be done by either reducing the precursor content of the raw water, improving precursor removal through the plant, shifting the point of disinfection to a later stage of treatment, or some combination of the three It has been proposed that watershed management practices that help reduce primary productivity in impoundments will also result in reduced THM precursor levels (Karimi and Singer, 1991; Chapra et al., 1997) Reductions in organic precursors and bromide levels may also be achieved through careful source water blending Once they have entered the plant, DBP precursors may be removed by coagulation, adsorption, oxidation, or membrane separation Coagulation with aluminum and ferric salts has proven to be a very effective means of removing hydrophobic organic precursors (see Chapter 6) Adsorption with granular activated carbon (GAC) can also be quite effective (see Chapter 9) Other approaches that have been shown to remove TOC are also likely to remove DBP precursors (e.g., membrane processes, Chapter 11; ion exchange, Chapter 13) Modifying Disinfection Various elements of the disinfection step can sometimes be changed to help control DBP formation One obvious approach is to minimize 12.45 CHEMICAL OXIDATION the disinfectant dose to the extent possible within the constraints of good disinfection practice For example, primary and secondary disinfectant doses may be reduced in some plants, while relying on booster disinfection out in the distribution system to take care of areas susceptible to low residuals pH adjustment can also be used to favorably affect DBP formation Base addition for corrosion control might Table 12.7 General Advantages and Disadvantages of Common Water Treatment Oxidants Oxidant Advantages Disadvantages Chlorine Strong oxidant Strong disinfectant Produces persistent residual Inexpensive Easy to use Long history of use Produces halogenated DBPs May contribute to taste and odor problems Chloramines Does not produce THMs Produces persistent residual Controls microbial growth Easy to use Long history of use Weak oxidant Weak disinfectant Produces unidentified TOX Can lead to nitrification problems in distribution system May contribute to taste and odor problems Ozone Very strong oxidant Very strong disinfectant Effective for taste and odor Does not produce halogenated DBPs except in bromide-rich waters May aid in coagulation and flocculation Does not produce a persistent disinfectant residual Relatively costly Produces bromate in bromide-rich waters Produces biodegradable organic material that must be controlled Advanced oxidation processes Very strong oxidant Most effective for taste and odor control Does not produce halogenated DBPs Limited disinfection properties Relatively costly Produces biodegradable organic material that must be controlled Chlorine dioxide Strong oxidant Strong disinfectant Effective for certain types of taste and odor Does not produce halogenated DBPs Does not react with ammonia Produces chlorite as an inorganic DBP May produce undesirable odors Difficult to maintain a persistent disinfection residual Permanganate Easy to feed Effective for Fe and Mn oxidation Does not produce halogenated DBPs Effective for certain types of taste and odor Produces manganese dioxide which must be removed Can lead to pink water if dosage not carefully controlled Limited disinfection capabilities Oxygen Easy to feed Does not produce halogenated by-products Relatively weak oxidant for most water treatment applications except Fe(II) and sulfide oxidation 12.46 CHAPTER TWELVE be delayed until after disinfection with chlorine to help depress formation of THMs A widely used approach is to substitute one disinfectant for another, “alternative” disinfectant Common examples include the use of chloramines as a secondary disinfectant in place of free chlorine, or the use of preozonation as a primary disinfectant in place of prechlorination Removing By-products In general, DBPs are hydrophilic and low in molecular weight, characteristics which make them difficult to remove by most physicochemical processes For example, the removal of THMs in disinfected water requires either air stripping or GAC adsorption with frequent regeneration On the other 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H2O → 12 MnO2(S) + 2H+ + e− −1.21 V (Eoox) O3(aq) + H+ + e− → 12 O2(aq) + 12 H2O +2.04 V (Eored) O3(aq) + 12 Mn+2 + 12 H2O → 12 O2(aq) + 12 MnO2(S) + H+ +0.83 V (Eonet) ᎏᎏ ᎏᎏ ᎏᎏ (12. 13)... is converted to water This is done to the reduction half-reaction O3 + H+ + e− → + 12 O2 + 12 H2O ᎏᎏ (12. 12) 12. 5 CHEMICAL OXIDATION From this point, it is a simple matter of combining the equations... equation (from Table 12. 1) O3(aq) + H+ + e− → 12 O2(aq) + 12 H2O ᎏᎏ MnO2 + ᎏ23ᎏH2O → ᎏ13ᎏMnO4− + ᎏ43ᎏH+ + e− ᎏᎏ O3(aq) + ᎏ13ᎏMnO2 + ᎏ13ᎏH2O → ᎏ13ᎏMnO4− + ᎏ13ᎏH+ + 12 O2(aq) ᎏᎏ (12. 18) This allows