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CHAPTER 14 DISINFECTION Charles N Haas, Ph.D LD Betz Professor of Environmental Engineering Drexel University Philadelphia, Pennsylvania Disinfection is a process designed for the deliberate reduction of the number of pathogenic microorganisms While other water treatment processes, such as filtration or coagulation-flocculation-sedimentation, may achieve pathogen reduction, this is not generally their primary goal A variety of chemical or physical agents may be used to carry out disinfection The concept of disinfection preceded the recognition of bacteria as the causative agent of disease Averill (1832), for example, proposed chlorine disinfection of human wastes as a prophylaxis against epidemics Chemical addition during water treatment for disinfection became accepted only after litigation on its efficacy (Race, 1918) The prophylactic benefits of water disinfection soon became apparent, particularly with respect to the reduction of typhoid and cholera While significant progress is being made in controlling the classic waterborne diseases, newly recognized agents have added to the challenge These include viruses (Melnick et al., 1978; Mosley, 1966), certain bacteria (Campylobacter, Palmer et al., 1983; Yersinia, Brennhovd et al., 1992; Reasoner, 1991; or Mycobacteria, Geldreich, 1971; Iivanainen et al., 1993; Reasoner, 1991; for example), and protozoans (Giardia, Brown et al., 1992; Le Chevallier et al., 1991; Miller et al., 1978; Reasoner, 1991; Renton et al., 1996; Rose et al., 1991; Cryptosporidium, Bridgman et al., 1995; Centers for Disease Control and Prevention, 1995; Gallaher et al., 1989; Goldstein et al., 1996; Hayes et al., 1989; Le Chevallier et al., 1991; Leland et al., 1993; Mac Kenzie et al., 1994; Miller, 1992; Reasoner, 1991; Richardson et al., 1991; Rose et al., 1991; Rush et al., 1990; Smith, 1992) Occasional outbreaks of drinking-water-associated hepatitis have also occurred (Nasser, 1994; Rosenberg et al., 1980) In addition, new viral agents are continually being found to be capable of waterborne transmission The state of disinfection practice in the United States in the late 1980s was summarized in a survey of the AWWA Disinfection Committee (Haas et al., 1992) Most water utilities continue to rely on chlorine or hypochlorite as their primary disinfection chemicals (Table 14.1), although increasing numbers are using ammonia (for pre- or postammoniation) or chlorine dioxide or ozone With the increasing concern for removing and inactivating some of the more resistant pathogens, such as Giardia and Cryptosporidium, while minimizing disinfection by-products, options other than 14.1 14.2 CHAPTER FOURTEEN TABLE 14.1 Water Utility Disinfection Practices According to 1989 AWWA Survey (N = 267) Chlorine alone Gas Hypochlorite Chlorine + ClO2 Ozone Other No ammonia Ammonia 67.42% 5.99% 3.37% 0.37% 0.75% 19.85% 0.75% 1.50% Source: Haas et al., 1992 traditional chlorination are gaining popularity This chapter will cover the use of chlorine, as well as the major alternative agents, for the purpose of disinfection HISTORY OF DISINFECTION Chlorine Chlorine gas was first prepared by Scheele in 1774, but chlorine was not regarded as a chemical element until 1808 (Belohlav and McBee, 1966) Early uses of chlorine included the use of Javelle water (chlorine gas dissolved in an alkaline potassium solution) in France for waste treatment in 1825 (Baker, 1926) and its use as a prophylactic agent during the European cholera epidemic of 1831 (Belohlav and McBee, 1966) Disinfection of water by chlorine first occurred in 1908 at Bubbly Creek (Chicago) and the Jersey City Water Company Within two years, chlorine was introduced as a disinfectant at New York City (Croton), Montreal, Milwaukee, Cleveland, Nashville, Baltimore, and Cincinnati, as well as other smaller treatment plants Frequently, dramatic reductions in typhoid accompanied the introduction of this process (Hooker, 1913) By 1918, over 1000 cities, treating more than billion gal/day (1.1 × 107 m3/day) of water, were employing chlorine as a disinfectant (Race, 1918) Chloramination, the addition of both chlorine and ammonia either sequentially or simultaneously, was first employed in Ottawa, Canada, and Denver, Colorado, in 1917 Both of these early applications employed prereaction of the two chemicals prior to their addition to the full flow of water Somewhat later, preammoniation (the addition of ammonia prior to chlorine) was developed In both cases, the process was advocated for its ability to prolong the stability of residual disinfectant during distribution and for its diminished propensity to produce chlorophenolic taste and odor substances Shortages of ammonia during World War II, and recognition of the superiority of free chlorine as a disinfectant, reduced the popularity of the chloramination process Recent concerns about organic by-products of chlorination, however, have increased the popularity of chloramination (Wolfe et al., 1984) Chlorine Dioxide Chlorine dioxide was first produced from the reaction of potassium chlorate and hydrochloric acid by Davy in 1811 (Miller et al., 1978) However, not until the industrial-scale preparation of sodium chlorite, from which chlorine dioxide may more readily be generated, did its widespread use occur (Rapson, 1966) DISINFECTION 14.3 Chlorine dioxide has been used widely as a bleaching agent in pulp and paper manufacture (Rapson, 1966) Despite early investigations on the use of chlorine dioxide as an oxidant and disinfectant (Aston and Synan, 1948), however, its ascendancy in both water and wastewater treatment has been slow As recently as 1971 (Morris, 1971), it was stated that “ ClO2 has never been used extensively for water disinfection.” By 1977, 84 potable water treatment plants in the United States were identified as using chlorine dioxide treatment, although only one of these relied upon it as a primary disinfectant (Miller et al., 1978) In Europe, chlorine dioxide was being used as either an oxidant or disinfectant in almost 500 potable water treatment plants (Miller et al., 1978) Ozone Ozone was discovered in 1783 by Van Marum, and named by Schonbein in 1840 In 1857, the first electric discharge ozone generation device was constructed by Siemens, with the first commercial application of this device occurring in 1893 (Water Pollution Control Federation, 1984) Ozone was first applied as a potable water disinfectant in 1893 at Oudshoorn, Netherlands In 1906, Nice, France, installed ozone as a treatment process, and this plant represents the oldest ozonation installation in continuous operation (Rice et al., 1981) In the United States, ozone was first employed for taste and odor control at New York City’s Jerome Park Reservoir in 1906 In 1987, five water treatment facilities in the United States were using ozone oxidation primarily for taste and odor control or trihalomethane precursor removal (Glaze, 1987) Since the 1993 Milwaukee Cryptosporidium outbreak, there has been an upsurge in interest in ozone as a disinfectant UV Radiation The biocidal effects of ultraviolet radiation (UV) have been known since it was established that short-wavelength UV was responsible for microbial decay often associated with sunlight (Downes and Blount, 1877) By the early 1940s, design guidelines for UV disinfection were proposed (Huff et al., 1965) UV has been accepted for treating potable water on passenger ships (Huff et al., 1965) Historically, however, it has met with little enthusiasm in public water supply applications because of the lack of a residual following application In wastewater treatment, in contrast, over 600 plants in the United States are either using, currently designing, or constructing UV disinfection facilities (Scheible et al., 1992) Other Agents A variety of other agents may be used to effect inactivation of microorganisms These include heat, extremes in pH, metals (silver, copper), surfactants, permanganate, and electron beam irradiation Heat is useful only in emergencies as in “boil water” orders, and is uneconomical An alkaline pH (during high lime softening) may provide some microbial inactivation, but is not usually sufficient as a sole disinfectant Potassium permanganate has been reported to achieve some disinfecting effects; however, the magnitudes have not been well characterized High-energy electrons for disinfection of wastewaters and sludges have also been studied (Farooq et al., 1993); however, their 14.4 CHAPTER FOURTEEN feasibility in drinking water is uncertain In this chapter, therefore, primary consideration will be given to chlorine compounds, ozone, chlorine dioxide, and UV Regulatory Issues for Disinfection Processes SWTR and GDR Requirements Amendments to the Safe Drinking Water Act require that all surface water suppliers in the United States filter and/or disinfect to protect the health of their customers The filtration and disinfection treatment requirements for public water systems using surface water sources or groundwater under the direct influence of surface water are included in what is called the Surface Water Treatment Rule (SWTR, June 1989) The SWTR requires that all surface water treatment facilities provide filtration and disinfection that achieves at least (1) a 99.9 percent (3-log) removal-inactivation of Giardia lamblia cysts and (2) a 99.99 percent (4-log) removal-inactivation of enteric viruses The SWTR assumes that for effective filtration, a conventional treatment plant achieves 2.5-log removal of Giardia and a 2-log removal of viruses Disinfection is required for the remainder of the removal-inactivation The amount of disinfection credit to be awarded is determined with the CT concept, CT being defined as the residual disinfectant concentration (C, mg/L) multiplied by the contact time (T, min) between the point of disinfectant application and the point of residual measurement The SWTR Guidance Manual provides tables of CT values for several disinfectants, which indicate the specific disinfection or CT credit awarded for a calculated value of CT A large safety factor is incorporated into the CT values included in the Guidance Manual tables In addition to relying on the CT tables to calculate disinfection credit, the SWTR allows utilities to demonstrate the effectiveness of their disinfection systems through pilot-scale studies, which may be prohibitively expensive for smaller operations The SWTR is being revised to take into account knowledge developed since the mid-1980s, and the anticipated formal promulgation of the Enhanced Surface Water Treatment Rule (ESWTR) will further affect the level of required disinfection A more complete discussion of the SWTR is included in Chapter Furthermore, under the Safe Drinking Water Act, EPA is required to promulgate rules for the disinfection of groundwaters While the regulatory development of the anticipated Groundwater Disinfection Rule is currently pending, this is expected to require a level of disinfection either by chemical agents or by virtue of aquifer passage of all groundwaters being used in community water supply systems Disinfection By-product Requirements Along with disinfection requirements, since 1974 there have been explicit regulations on disinfection by-products—first with respect to trihalomethanes, and more recently with respect to haloacetic acids, bromate, and other possible by-products The combination of the requirement to achieve disinfection along with the requirement to minimize disinfection byproducts has led to an increasing spectrum of options being considered DISINFECTANTS AND THEORY OF DISINFECTION Basic Chemistry Chlorine and Chlorine Compounds Chlorine may be used as a disinfectant in the form of compressed gas under pressure that is dissolved in water at the point of DISINFECTION 14.5 application, solutions of sodium hypochlorite, or solid calcium hypochlorite The three forms are chemically equivalent because of the rapid equilibrium that exists between dissolved molecular gas and the dissociation products of hypochlorite compounds Elemental chlorine (Cl2) is a dense gas that, when subject to pressures in excess of its vapor pressure, condenses into a liquid with the release of heat and with a reduction in specific volume of approximately 450-fold Hence, commercial shipments of chlorine are made in pressurized tanks to reduce shipment volume When chlorine is to be dispensed as a gas, supplying thermal energy to vaporize the compressed liquid chlorine is necessary The relative amount of chlorine present in chlorine gas, or hypochlorite salts, is expressed in terms of available chlorine The concentration of hypochlorite (or any other oxidizing disinfectant) may be expressed as available chlorine by determining the electrochemical equivalent amount of Cl2 to that compound Equation 14.1 shows that mole of elemental chlorine is capable of reacting with two electrons to form inert chloride: Cl2 + e− = Cl− (14.1) Equation 14.2 shows that mole of hypochlorite (OCl−) may react with two electrons to form chloride: OCl− + e− + 2H+ = Cl− + H2O (14.2) Hence, mole of hypochlorite is electrochemically equivalent to mole of elemental chlorine, and may be said to contain 70.91 g of available chlorine (identical to the molecular weight of Cl2) Calcium hypochlorite (Ca(OCl)2) and sodium hypochlorite (NaOCl) contain and moles of hypochlorite per mole of chemical, respectively, and, as a result, 141.8 and 70.91 g available chlorine per mole, respectively The molecular weights of Ca(OCl)2 and NaOCl are, 143 and 74.5, respectively, so that pure preparations of the two compounds contain 99.2 and 95.8 weight percent available chlorine; hence, they are effective means of supplying chlorine for disinfection purposes Calcium hypochlorite is available commercially as a dry solid In this form, it is subject to a loss in strength of approximately 0.013 percent per day (Laubusch, 1963) Calcium hypochlorite is also available in a tablet form for use in automatic feed equipment at low-flow treatment plants Sodium hypochlorite is available in to 16 weight percent solutions Higherconcentration solutions are not practical because chemical stability rapidly diminishes with increasing strength At ambient temperatures, the half-life of sodium hypochlorite solutions varies between 60 and 1700 days, respectively, for solutions of 18 and percent available chlorine (Baker, 1969; Laubusch, 1963) It should be noted that the loss of strength in sodium hypochlorite solutions may also result in the formation of by-products that may be undesirable Thermodynamically, the autodecomposition of hypochlorite to chlorate is highly favored by the following overall process (Bolyard et al., 1992): ClO− → Cl− + ClO3− (14.3) Measurements of sodium hypochlorite disinfectant solutions at water utilities have revealed that the mass concentration of chlorate is from 1.7 to 220 percent of the mass concentration of free available chlorine (Bolyard et al., 1992, 1993) The concentration of chlorate present in these stock solutions is kinetically controlled 14.6 CHAPTER FOURTEEN and may be related to the solution strength, age, temperature, pH, and presence of metal catalysts (Gordon et al., 1993, 1995) When a chlorine-containing compound is added to a water containing insignificant quantities of kjeldahl nitrogen, organic material, and other chlorine-demanding substances, a rapid equilibrium is established among the various chemical species in solution The term free available chlorine is used to refer to the sum of the concentrations of molecular chlorine (Cl2), hypochlorous acid (HOCl), and hypochlorite ion (OCl−), each expressed as available chlorine The dissolution of gaseous chlorine to form dissolved molecular chlorine is expressible as a phase equilibrium, and may be described by Henry’s law: Cl2(g) = Cl2(aq) H(mol/L-atm) = [Cl2(aq)]/PCl2 (14.4) where quantities within square brackets represent molar concentrations, PCl2 is the gas phase partial pressure of chlorine in atmospheres, and H is the Henry’s law constant, estimated from the following equation (Downs and Adams, 1973): H = 4.805 × 10−6 exp (2818.48/T) (mol/L-atm) (14.5) Dissolved aqueous chlorine reacts with water to form hypochlorous acid, chloride ions, and protons as indicated by Equation 14.6 Cl2(aq) + H2O = H+ + HOCl + Cl− [H+][HOCl][Cl−] KH = ᎏᎏ [Cl2(aq)] ΂ ΃ 2581.93 = 2.581 exp − ᎏ (mol2/L2) T (14.6) This reaction typically reaches completion in 100 ms (Aieta and Roberts, 1985; Morris, 1946) and involves elementary reactions between dissolved molecular chlorine and hydroxyl ions The extent of chlorine hydrolysis, or disproportionation (because the valence of chlorine changes from on the left to +1 and −1 on the right), as described by Equation 14.6, decreases with decreasing pH and increasing salinity; hence, the solubility of gaseous chlorine may be increased by the addition of alkali or by the use of fresh, rather than brackish, water Hypochlorous acid is a weak acid and may dissociate according to Equation 14.7: HOCl = OCl− + H+ Ka = [OCl−][H+]/[HOCl] (14.7) The pKa of hypochlorous acid at room temperature is approximately 7.6 (Brigano et al., 1978) Morris (1966) has provided a correlating equation for Ka as a function of temperature: ln(Ka) = 23.184 − 0.0583 T − 6908/T (14.8) where T is specified in degrees Kelvin (K = °C + 273) Figure 14.1 illustrates the effect of pH on the distribution of free chlorine between OCl− and HOCl One practical consequence of the reactions described by Equations 14.4 through 14.8 is that the chlorine vapor pressure over a solution depends on solution pH, decreasing as pH increases (because of the increased formation of nonvolatile hypochlorite acid) Therefore, the addition of an alkaline material such as lime or DISINFECTION 14.7 FIGURE 14.1 Effect of pH on relative amount of hypochlorous acid and hypochlorite ion at 20°C sodium bicarbonate will reduce the volatility of chlorine from accidental spills or leaks and thus minimize danger to exposed personnel The acid-base properties of gaseous chlorine, or the hypochlorite salts, will also result in a loss or gain, respectively, of alkalinity, and a reduction or increase, respectively, in pH For each mole of free chlorine (i.e., mole of Cl2, or of NaOCl or 0.5 mole of Ca(OCl)2), there will be a change of one equivalent of alkalinity (increase for sodium and calcium hypochlorite, and decrease for chlorine gas) The solution produced by a gas chlorinator contains 3500 mg/L available chlorine at a pH of What is the equilibrium vapor pressure of this solution at 20°C (given that the value of the hydrolysis constant KH is 4.5 × 104 at this temperature)? EXAMPLE 14.1 The pH is sufficiently low that the dissociation of hypochlorous acid to form hypochlorite can be ignored Therefore, a balance over chlorine species yields: [Cl2] + [HOCl] = (3500 × 10−3)/71 The factor of 71 reflects the fact that mole of either dissolved chlorine or hypochlorous acid contains 71 g of available chlorine The hydrolysis equilibrium constant can be used to develop an additional equation: 4.5 × 104 = [H+][Cl−][HOCl]/[Cl2] or, because the pH is given, 4.5 × 107 = [Cl−][HOCl]/[Cl2] 14.8 CHAPTER FOURTEEN Because chlorine gas was used to generate the dissolved free chlorine, the disproportionation reaction requires that for each mole of HOCl produced, mole of Cl− must have been produced If the initial concentration of chloride (in the feedwater to the chlorinator) was minimal, then a third equation results: [Cl−] = [HOCl] These three equations can be manipulated to produce a quadratic equation in the unknown [Cl2]1/2: [Cl2] + 6708[Cl2]1/2 − 0.05 = The single positive root is the only physically meaningful one, hence: [Cl2]1/2 = 7.45 × 10−6 or [Cl2] = 5.55 × 10−11 The Henry’s law constant can be computed from Equation 14.5 as 0.072 moles/ L-atm, and therefore the partial pressure of chlorine gas is found: PCl2 = 5.55 × 10−11/0.072 = 7.7 × 10−10 atm = (0.77 ppb) The OSHA permissible exposure limit (PEL) is reported as ppm (ACGIH, 1994) Therefore, this level is of no apparent health concern to the workers Chlorine Dioxide Chlorine dioxide (ClO2) is a neutral compound of chlorine in the +IV oxidation state It has a boiling point of 11°C at atmospheric pressure The liquid is denser than water and the gas is denser than air (Noack and Doeff, 1979) Chemically, chlorine dioxide is a stable free radical that, at high concentrations, reacts violently with reducing agents It is explosive, with the lower explosive limit in air variously reported as 10 percent (Downs and Adams, 1973; Masschelein, 1979b) or 39 percent (Noack and Doeff, 1979) As a result, virtually all applications of chlorine dioxide require the synthesis of the gaseous compound in a dilute stream (either gaseous or liquid) on location as needed The solubility of gaseous chlorine dioxide in water may be described by Henry’s law, and a fit of the available solubility data (Battino, 1984) results in the following relationship for the Henry’s law constant (in units of atm−1): ln(H) = mole fraction dissolved ClO2(aq)/PClO2 = 58.84621 + (47.9133/T) − 11.0593 ln(T) (14.9) Under alkaline conditions, the following disproportionation into chlorite (ClO2−) and chlorate (ClO3−) occurs (Gordon et al., 1972): ClO2 + OH− = H2O + ClO3− + ClO2− (14.10) In the absence of catalysis by carbonate, the reaction (Equation 14.10) is governed by parallel first- and second-order kinetics (Gordon et al., 1972; Granstrom and Lee, 1957) The half-life of aqueous chlorine dioxide solutions decreases substantially with increasing concentration and with pH values above Even at neutral pH values, however, in the absence of carbonate at room temperature, the half-life DISINFECTION 14.9 of chlorine dioxide solutions of 0.01, 0.001, and 0.0001 mol/L is 0.5, 4, and 14 h, respectively Hence, the storage of stock solutions of chlorine dioxide for even a few hours is impractical The simple disproportionation reaction to chlorate and chlorite is insufficient to explain the decay of chlorine dioxide in water free of extraneous reductants Equation 14.10 predicts that the molar ratio of chlorate to chlorite formed should be 1:1 Medir and Giralt (1982), however, found that the molar ratio of chlorate to chlorite to chloride to oxygen produced was 5:3:1:0.75, and that the addition of chloride enhanced the rate of decomposition and resulted in the predicted 1:1 molar ratio of chlorite to chlorate Thus, the oxidation of chloride by chlorate, and the possible formation of intermediate free chlorine, may be of significance in the decay of chlorine dioxide in demand-free systems (Gordon et al., 1972) The concentration of chlorine dioxide in solution is generally expressed in terms of g/L as chlorine by multiplying the molarity of chlorine dioxide by the number of electrons transferred per mole of chlorine dioxide reacted and then multiplying this by 35.5 g Cl2 per electron mole Conventionally, the five-electron reduction (Equation 14.11) is used to carry out this conversion ClO2 + 5e− + 4H+ = Cl− + H2O (14.11) Note, however, that the typical reaction of chlorine dioxide in water, being reduced to chlorite, is a one-electron reduction as follows: ClO2 + e− = ClO2− (14.11a) Hence, according to Equation 14.11, mole of chlorine dioxide contains 67.5 g of mass, and is equivalent to 177.5 (=5 × 35.5) g Cl2 Therefore, g of chlorine dioxide contains 2.63 g as chlorine In examining any study on chlorine dioxide, due care with regard to units of expression of disinfectant concentration is warranted Ozone Ozone is a colorless gas produced from the action of electric fields on oxygen It is highly unstable in the gas phase; in clean vessels at room temperature the half-life in air is 20 to 100 h (Manley and Niegowski, 1967) The solubility of ozone in water can be described by a temperature- and pHdependent Henry’s law constant The following provisional relationship (H in atm−1) has been suggested (Roy, 1979): H = 3.84 × 107 [OH−] exp (−2428/T) (14.12) Practical ozone generation systems have maximum gaseous ozone concentrations of about 50 g/m3; thus, the maximum practical solubility of ozone in water is about 40 mg/L (Stover et al., 1986) Upon dissolution in water, ozone can react with water itself, with hydroxyl ions, or with dissolved chemical constituents, as well as serving as a disinfecting agent Details of these reactions will be discussed later in this chapter and in Chapter 12 DISINFECTANT DEMAND REACTIONS Chlorine Reactions with Ammonia In the presence of certain dissolved constituents in water, each of the disinfectants may react and transform to less biocidal chemical 14.10 CHAPTER FOURTEEN forms In the case of chlorine, these principally involve reactions with ammonia and amino nitrogen compounds In the presence of ammonium ion, free chlorine reacts in a stepwise manner to form chloramines This process is depicted in Equations 14.13 through 14.15: NH4+ + HOCl = NH2Cl + H2O + H+ (14.13) NH2Cl + HOCl = NHCl2 + H2O (14.14) NHCl2 + HOCl = NCl3 + H2O (14.15) These compounds, monochloramine (NH2Cl), dichloramine (NHCl2 ), and trichloramine (NCl3), each contribute to the total (or combined) chlorine residual in a water.The terms total available chlorine and total oxidants refer, respectively, to the sum of free chlorine compounds and reactive chloramines, or total oxidating agents Under normal conditions of water treatment, if any excess ammonia is present, at equilibrium the amount of free chlorine will be much less than percent of total residual chlorine Each chlorine atom associated with a chloramine molecule is capable of undergoing a two-electron reduction to chloride; hence, each mole of monochloramine contains 71 g available chlorine; each mole of dichloramine contains × 71 or 142 g; and each mole of trichloramine contains × 71 or 223 g of available chlorine Inasmuch as the molecular weights of mono-, di-, and trichloramine are 51.6, 86, and 110.5, respectively, the chloramines contain, respectively, 1.38, 1.65, and 2.02 g available chlorine per gram The efficiency of the various combined chlorine forms as disinfectants differs, however, and thus the concentration of available chlorine does not completely characterize process performance On an approximate basis, for example, for coliforms, the biocidal potency of HOCl:OCl−:NH2Cl:NHCl2 is approximately 1:0.0125:0.005:0.0166; and for viruses and cysts, the combined chlorine forms are considerably less effective (Chang, 1971) As Equation 14.12 indicates, the formation of monochloramine is accompanied by the loss of a proton, because chlorination reduces the affinity of the nitrogen moiety for protons (Weil and Morris, 1949a) The significance of chlorine speciation on disinfection efficiency was graphically demonstrated by Weber et al (1940) as shown in Figure 14.2 As the dose of chlorine is increased, the total chlorine residual (i.e., remaining in the system after 30 min) increases until a dose of approximately 50 mg/L, whereupon residual chlorine decreases to a very low value, and subsequently increases linearly with dose indefinitely The “hump and dip” behavior is paralleled by the sensitivity of microorganisms to the available chlorine residual indicated by the time required for 99 percent inactivation of Bacillus metiens spores At the three points indicated, the total available chlorine is approximately identical at 22 to 24 mg/L, yet a 32-fold difference in microbial sensitivity occurred The explanation for this behavior is the “breakpoint” reaction between free chlorine and ammonia (Figure 14.3) At doses below the hump in the chlorine residual curve (zone 1), only combined chlorine is detectable At doses between the hump and the dip in the curve, an oxidative destruction of combined residual chlorine accompanied by the loss of nitrogen occurs (zone 2) (Taras, 1950) One possible reaction during breakpoint is: NH3 + HOCl = N2 + H+ + Cl− + H2O (14.16) This reaction also may be used as a means to remove ammonia nitrogen from water or wastewaters (Pressley et al., 1972) Finally, after the ammonia nitrogen has 14.46 CHAPTER FOURTEEN ozone inactivation of microorganisms are poor predictors of field performance, unless detailed aspects of field-scale hydraulic features are considered In one pilot study of surface water treated with granular activated carbon (GAC) with no other pretreatment (Morin et al., 1975), it was found that a contactor with 57 s of contact and an ozone dosage of 1.45 mg/l (with a transferred dose of 1.13 mg/l) resulted in complete viral inactivation (>7 logs) Satisfactory coliform results (