14. disinfection

CHAPTER 14DISINFECTION

Charles N. Haas, Ph.D.LD Betz Professor of Environmental Engineering

Drexel UniversityPhiladelphia, Pennsylvania

Disinfection is a process designed for the deliberate reduction of the number ofpathogenic microorganisms. While other water treatment processes, such as filtra-tion or coagulation-flocculation-sedimentation, may achieve pathogen reduction,this is not generally their primary goal.A variety of chemical or physical agents maybe used to carry out disinfection. The concept of disinfection preceded the recogni-tion of bacteria as the causative agent of disease. Averill (1832), for example, pro-posed chlorine disinfection of human wastes as a prophylaxis against epidemics.Chemical addition during water treatment for disinfection became accepted onlyafter litigation on its efficacy (Race, 1918). The prophylactic benefits of water disin-fection soon became apparent, particularly with respect to the reduction of typhoidand cholera.

While significant progress is being made in controlling the classic waterborne dis-eases, 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; Ren-ton et al., 1996; Rose et al., 1991; Cryptosporidium, Bridgman et al., 1995; Centers forDisease 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 etal., 1990; Smith, 1992). Occasional outbreaks of drinking-water-associated hepatitishave also occurred (Nasser, 1994; Rosenberg et al., 1980). In addition, new viralagents are continually being found to be capable of waterborne transmission.

The state of disinfection practice in the United States in the late 1980s was sum-marized in a survey of the AWWA Disinfection Committee (Haas et al., 1992). Mostwater utilities continue to rely on chlorine or hypochlorite as their primary disinfec-tion chemicals (Table 14.1), although increasing numbers are using ammonia (forpre- or postammoniation) or chlorine dioxide or ozone.With the increasing concernfor removing and inactivating some of the more resistant pathogens, such as Giardiaand Cryptosporidium, while minimizing disinfection by-products, options other than

14.1

traditional chlorination are gaining popularity. This chapter will cover the use ofchlorine, 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 achemical element until 1808 (Belohlav and McBee, 1966). Early uses of chlorineincluded the use of Javelle water (chlorine gas dissolved in an alkaline potassium solu-tion) in France for waste treatment in 1825 (Baker, 1926) and its use as a prophylacticagent 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 adisinfectant at New York City (Croton), Montreal, Milwaukee, Cleveland, Nashville,Baltimore, and Cincinnati, as well as other smaller treatment plants. Frequently, dra-matic reductions in typhoid accompanied the introduction of this process (Hooker,1913). By 1918, over 1000 cities, treating more than 3 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 sequentiallyor simultaneously, was first employed in Ottawa, Canada, and Denver, Colorado, in1917. Both of these early applications employed prereaction of the two chemicalsprior 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 pro-cess was advocated for its ability to prolong the stability of residual disinfectant dur-ing distribution and for its diminished propensity to produce chlorophenolic tasteand odor substances. Shortages of ammonia during World War II, and recognition ofthe superiority of free chlorine as a disinfectant, reduced the popularity of the chlo-ramination 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 andhydrochloric acid by Davy in 1811 (Miller et al., 1978). However, not until theindustrial-scale preparation of sodium chlorite, from which chlorine dioxide maymore readily be generated, did its widespread use occur (Rapson, 1966).

14.2 CHAPTER FOURTEEN

TABLE 14.1 Water Utility Disinfection PracticesAccording to 1989 AWWA Survey (N = 267)

No ammonia Ammonia

Chlorine aloneGas 67.42% 19.85%Hypochlorite 5.99% 0.75%

Chlorine + ClO2 3.37% 1.50%Ozone 0.37%Other 0.75%

Source: Haas et al., 1992.

Chlorine dioxide has been used widely as a bleaching agent in pulp and papermanufacture (Rapson, 1966). Despite early investigations on the use of chlorinedioxide as an oxidant and disinfectant (Aston and Synan, 1948), however, its ascen-dancy 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 waterdisinfection.”

By 1977, 84 potable water treatment plants in the United States were identifiedas using chlorine dioxide treatment, although only one of these relied upon it as aprimary disinfectant (Miller et al., 1978). In Europe, chlorine dioxide was being usedas 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. In1857, the first electric discharge ozone generation device was constructed bySiemens, 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 thisplant represents the oldest ozonation installation in continuous operation (Rice etal., 1981). In the United States, ozone was first employed for taste and odor controlat New York City’s Jerome Park Reservoir in 1906. In 1987, five water treatmentfacilities in the United States were using ozone oxidation primarily for taste andodor control or trihalomethane precursor removal (Glaze, 1987). Since the 1993 Mil-waukee Cryptosporidium outbreak, there has been an upsurge in interest in ozoneas a disinfectant.

UV Radiation

The biocidal effects of ultraviolet radiation (UV) have been known since it wasestablished that short-wavelength UV was responsible for microbial decay oftenassociated with sunlight (Downes and Blount, 1877). By the early 1940s, designguidelines for UV disinfection were proposed (Huff et al., 1965). UV has beenaccepted for treating potable water on passenger ships (Huff et al., 1965). Histori-cally, however, it has met with little enthusiasm in public water supply applicationsbecause of the lack of a residual following application. In wastewater treatment, incontrast, over 600 plants in the United States are either using, currently designing, orconstructing UV disinfection facilities (Scheible et al., 1992).

Other Agents

A variety of other agents may be used to effect inactivation of microorganisms.Theseinclude heat, extremes in pH, metals (silver, copper), surfactants, permanganate, andelectron 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 somemicrobial inactivation, but is not usually sufficient as a sole disinfectant. Potassiumpermanganate has been reported to achieve some disinfecting effects; however, themagnitudes have not been well characterized. High-energy electrons for disinfectionof wastewaters and sludges have also been studied (Farooq et al., 1993); however, their

DISINFECTION 14.3

feasibility in drinking water is uncertain. In this chapter, therefore, primary considera-tion 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 Actrequire that all surface water suppliers in the United States filter and/or disinfect toprotect the health of their customers. The filtration and disinfection treatmentrequirements for public water systems using surface water sources or groundwaterunder the direct influence of surface water are included in what is called the SurfaceWater Treatment Rule (SWTR, June 1989).

The SWTR requires that all surface water treatment facilities provide filtration anddisinfection that achieves at least (1) a 99.9 percent (3-log) removal-inactivation ofGiardia lamblia cysts and (2) a 99.99 percent (4-log) removal-inactivation of entericviruses. The SWTR assumes that for effective filtration, a conventional treatmentplant achieves 2.5-log removal of Giardia and a 2-log removal of viruses. Disinfectionis required for the remainder of the removal-inactivation. The amount of disinfectioncredit to be awarded is determined with the CT concept, CT being defined as theresidual 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 ofCT. A large safety factor is incorporated into the CT values included in the GuidanceManual tables. In addition to relying on the CT tables to calculate disinfection credit,the SWTR allows utilities to demonstrate the effectiveness of their disinfection sys-tems through pilot-scale studies, which may be prohibitively expensive for smalleroperations. The SWTR is being revised to take into account knowledge developedsince the mid-1980s, and the anticipated formal promulgation of the Enhanced Sur-face Water Treatment Rule (ESWTR) will further affect the level of required disin-fection.A more complete discussion of the SWTR is included in Chapter 1.

Furthermore, under the Safe Drinking Water Act, EPA is required to promulgaterules for the disinfection of groundwaters. While the regulatory development of theanticipated Groundwater Disinfection Rule is currently pending, this is expected torequire a level of disinfection either by chemical agents or by virtue of aquifer pas-sage 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—firstwith respect to trihalomethanes, and more recently with respect to haloacetic acids,bromate, and other possible by-products. The combination of the requirement toachieve disinfection along with the requirement to minimize disinfection by-products 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 theform of compressed gas under pressure that is dissolved in water at the point of


application, solutions of sodium hypochlorite, or solid calcium hypochlorite. Thethree forms are chemically equivalent because of the rapid equilibrium that existsbetween dissolved molecular gas and the dissociation products of hypochlorite com-pounds.

Elemental chlorine (Cl2) is a dense gas that, when subject to pressures in excessof its vapor pressure, condenses into a liquid with the release of heat and with areduction in specific volume of approximately 450-fold. Hence, commercial ship-ments of chlorine are made in pressurized tanks to reduce shipment volume. Whenchlorine is to be dispensed as a gas, supplying thermal energy to vaporize the com-pressed liquid chlorine is necessary.

The relative amount of chlorine present in chlorine gas, or hypochlorite salts, isexpressed in terms of available chlorine. The concentration of hypochlorite (or anyother oxidizing disinfectant) may be expressed as available chlorine by determiningthe electrochemical equivalent amount of Cl2 to that compound. Equation 14.1shows that 1 mole of elemental chlorine is capable of reacting with two electrons toform inert chloride:

Cl2 + 2 e− = 2 Cl− (14.1)

Equation 14.2 shows that 1 mole of hypochlorite (OCl−) may react with two elec-trons to form chloride:

OCl− + 2 e− + 2H+ = Cl− + H2O (14.2)

Hence, 1 mole of hypochlorite is electrochemically equivalent to 1 mole of ele-mental chlorine, and may be said to contain 70.91 g of available chlorine (identicalto the molecular weight of Cl2).

Calcium hypochlorite (Ca(OCl)2) and sodium hypochlorite (NaOCl) contain 2and 1 moles of hypochlorite per mole of chemical, respectively, and, as a result, 141.8and 70.91 g available chlorine per mole, respectively. The molecular weights ofCa(OCl)2 and NaOCl are, 143 and 74.5, respectively, so that pure preparations of thetwo compounds contain 99.2 and 95.8 weight percent available chlorine; hence, theyare effective means of supplying chlorine for disinfection purposes.

Calcium hypochlorite is available commercially as a dry solid. In this form, it issubject 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 automaticfeed equipment at low-flow treatment plants.

Sodium hypochlorite is available in 1 to 16 weight percent solutions. Higher-concentration solutions are not practical because chemical stability rapidly dimin-ishes with increasing strength. At ambient temperatures, the half-life of sodiumhypochlorite solutions varies between 60 and 1700 days, respectively, for solutions of18 and 3 percent available chlorine (Baker, 1969; Laubusch, 1963).

It should be noted that the loss of strength in sodium hypochlorite solutions mayalso result in the formation of by-products that may be undesirable. Thermodynam-ically, the autodecomposition of hypochlorite to chlorate is highly favored by the fol-lowing overall process (Bolyard et al., 1992):

3 ClO− → 2 Cl− + ClO3− (14.3)

Measurements of sodium hypochlorite disinfectant solutions at water utilitieshave revealed that the mass concentration of chlorate is from 1.7 to 220 percent ofthe mass concentration of free available chlorine (Bolyard et al., 1992, 1993). Theconcentration of chlorate present in these stock solutions is kinetically controlled

DISINFECTION 14.5

and may be related to the solution strength, age, temperature, pH, and presence ofmetal catalysts (Gordon et al., 1993, 1995).

When a chlorine-containing compound is added to a water containing insignifi-cant quantities of kjeldahl nitrogen, organic material, and other chlorine-demandingsubstances, a rapid equilibrium is established among the various chemical species insolution. The term free available chlorine is used to refer to the sum of the concen-trations of molecular chlorine (Cl2), hypochlorous acid (HOCl), and hypochloriteion (OCl−), each expressed as available chlorine.

The dissolution of gaseous chlorine to form dissolved molecular chlorine isexpressible 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 thegas phase partial pressure of chlorine in atmospheres, and H is the Henry’s law con-stant, 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, chlo-ride ions, and protons as indicated by Equation 14.6.

Cl2(aq) + H2O = H+ + HOCl + Cl−

KH =

= 2.581 exp �− � (mol2/L2) (14.6)

This reaction typically reaches completion in 100 ms (Aieta and Roberts, 1985;Morris, 1946) and involves elementary reactions between dissolved molecular chlo-rine and hydroxyl ions. The extent of chlorine hydrolysis, or disproportionation(because the valence of chlorine changes from 0 on the left to +1 and −1 on theright), as described by Equation 14.6, decreases with decreasing pH and increasingsalinity; hence, the solubility of gaseous chlorine may be increased by the addition ofalkali 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 (Bri-gano et al., 1978). Morris (1966) has provided a correlating equation for Ka as a func-tion 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 theeffect of pH on the distribution of free chlorine between OCl− and HOCl.

One practical consequence of the reactions described by Equations 14.4 through14.8 is that the chlorine vapor pressure over a solution depends on solution pH,decreasing as pH increases (because of the increased formation of nonvolatilehypochlorite acid). Therefore, the addition of an alkaline material such as lime or

2581.93�

T

[H+][HOCl][Cl−]��

[Cl2(aq)]


sodium bicarbonate will reduce the volatility of chlorine from accidental spills orleaks and thus minimize danger to exposed personnel.

The acid-base properties of gaseous chlorine, or the hypochlorite salts, will alsoresult in a loss or gain, respectively, of alkalinity, and a reduction or increase, respec-tively, in pH. For each mole of free chlorine (i.e., 1 mole of Cl2, or of NaOCl or 0.5mole of Ca(OCl)2), there will be a change of one equivalent of alkalinity (increasefor sodium and calcium hypochlorite, and decrease for chlorine gas).

EXAMPLE 14.1 The solution produced by a gas chlorinator contains 3500 mg/Lavailable chlorine at a pH of 3. What is the equilibrium vapor pressure of this solu-tion at 20°C (given that the value of the hydrolysis constant KH is 4.5 × 104 at thistemperature)?

1. The pH is sufficiently low that the dissociation of hypochlorous acid to formhypochlorite can be ignored. Therefore, a balance over chlorine species yields:

[Cl2] + [HOCl] = (3500 × 10−3)/71

2. The factor of 71 reflects the fact that 1 mole of either dissolved chlorine orhypochlorous acid contains 71 g of available chlorine.

3. The hydrolysis equilibrium constant can be used to develop an additional equa-tion:

4.5 × 104 = [H+][Cl−][HOCl]/[Cl2]

or, because the pH is given,

4.5 × 107 = [Cl−][HOCl]/[Cl2]

DISINFECTION 14.7

FIGURE 14.1 Effect of pH on relative amount of hypochlorous acid and hypochlo-rite ion at 20°C.

4. Because chlorine gas was used to generate the dissolved free chlorine, the dis-proportionation reaction requires that for each mole of HOCl produced, 1 moleof Cl− must have been produced. If the initial concentration of chloride (in thefeedwater to the chlorinator) was minimal, then a third equation results:

[Cl−] = [HOCl]

5. These three equations can be manipulated to produce a quadratic equation in theunknown [Cl2]1/2:

[Cl2] + 6708[Cl2]1/2 − 0.05 = 0

6. The single positive root is the only physically meaningful one, hence:

[Cl2]1/2 = 7.45 × 10−6

or

[Cl2] = 5.55 × 10−11

7. 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)

8. The OSHA permissible exposure limit (PEL) is reported as 1 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 inthe +IV oxidation state. It has a boiling point of 11°C at atmospheric pressure. Theliquid 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 inair 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 chlo-rine dioxide require the synthesis of the gaseous compound in a dilute stream (eithergaseous or liquid) on location as needed.

The solubility of gaseous chlorine dioxide in water may be described by Henry’slaw, and a fit of the available solubility data (Battino, 1984) results in the followingrelationship 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):

2 ClO2 + 2 OH− = H2O + ClO3− + ClO2

− (14.10)

In the absence of catalysis by carbonate, the reaction (Equation 14.10) is gov-erned by parallel first- and second-order kinetics (Gordon et al., 1972; Granstromand Lee, 1957). The half-life of aqueous chlorine dioxide solutions decreases sub-stantially with increasing concentration and with pH values above 9. Even at neutralpH values, however, in the absence of carbonate at room temperature, the half-life


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 fewhours is impractical.

The simple disproportionation reaction to chlorate and chlorite is insufficient toexplain the decay of chlorine dioxide in water free of extraneous reductants. Equa-tion 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 chloriteto chloride to oxygen produced was 5:3:1:0.75, and that the addition of chlorideenhanced the rate of decomposition and resulted in the predicted 1:1 molar ratio ofchlorite to chlorate.Thus, the oxidation of chloride by chlorate, and the possible for-mation of intermediate free chlorine, may be of significance in the decay of chlorinedioxide in demand-free systems (Gordon et al., 1972).

The concentration of chlorine dioxide in solution is generally expressed in termsof g/L as chlorine by multiplying the molarity of chlorine dioxide by the number ofelectrons transferred per mole of chlorine dioxide reacted and then multiplying thisby 35.5 g Cl2 per electron mole. Conventionally, the five-electron reduction (Equa-tion 14.11) is used to carry out this conversion.

ClO2 + 5e− + 4H+ = Cl− + 2 H2O (14.11)

Note, however, that the typical reaction of chlorine dioxide in water, beingreduced to chlorite, is a one-electron reduction as follows:

ClO2 + e− = ClO2− (14.11a)

Hence, according to Equation 14.11, 1 mole of chlorine dioxide contains 67.5 g ofmass, and is equivalent to 177.5 (=5 × 35.5) g Cl2. Therefore, 1 g of chlorine dioxidecontains 2.63 g as chlorine. In examining any study on chlorine dioxide, due care withregard to units of expression of disinfectant concentration is warranted.

Ozone. Ozone is a colorless gas produced from the action of electric fields on oxy-gen. It is highly unstable in the gas phase; in clean vessels at room temperature thehalf-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 pH-dependent 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 concentra-tions of about 50 g/m3; thus, the maximum practical solubility of ozone in water isabout 40 mg/L (Stover et al., 1986). Upon dissolution in water, ozone can react withwater itself, with hydroxyl ions, or with dissolved chemical constituents, as well asserving as a disinfecting agent. Details of these reactions will be discussed later inthis chapter and in Chapter 12.

DISINFECTANT DEMAND REACTIONS

Chlorine

Reactions with Ammonia. In the presence of certain dissolved constituents inwater, each of the disinfectants may react and transform to less biocidal chemical

DISINFECTION 14.9

forms. In the case of chlorine, these principally involve reactions with ammonia andamino nitrogen compounds. In the presence of ammonium ion, free chlorine reactsin a stepwise manner to form chloramines. This process is depicted in Equations14.13 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), andtrichloramine (NCl3), each contribute to the total (or combined) chlorine residual ina water.The terms total available chlorine and total oxidants refer, respectively, to thesum of free chlorine compounds and reactive chloramines, or total oxidating agents.Under normal conditions of water treatment, if any excess ammonia is present, atequilibrium the amount of free chlorine will be much less than 1 percent of totalresidual chlorine. Each chlorine atom associated with a chloramine molecule iscapable of undergoing a two-electron reduction to chloride; hence, each mole ofmonochloramine contains 71 g available chlorine; each mole of dichloramine con-tains 2 × 71 or 142 g; and each mole of trichloramine contains 3 × 71 or 223 g of avail-able chlorine. Inasmuch as the molecular weights of mono-, di-, and trichloramineare 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 chlo-rine forms as disinfectants differs, however, and thus the concentration of availablechlorine does not completely characterize process performance. On an approximatebasis, 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 chlo-rine forms are considerably less effective (Chang, 1971). As Equation 14.12 indi-cates, the formation of monochloramine is accompanied by the loss of a proton,because chlorination reduces the affinity of the nitrogen moiety for protons (Weiland Morris, 1949a).

The significance of chlorine speciation on disinfection efficiency was graphicallydemonstrated by Weber et al. (1940) as shown in Figure 14.2.As the dose of chlorineis 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 chlorinedecreases to a very low value, and subsequently increases linearly with dose indefi-nitely. The “hump and dip” behavior is paralleled by the sensitivity of microorgan-isms to the available chlorine residual indicated by the time required for 99 percentinactivation of Bacillus metiens spores. At the three points indicated, the total avail-able chlorine is approximately identical at 22 to 24 mg/L, yet a 32-fold difference inmicrobial sensitivity occurred.

The explanation for this behavior is the “breakpoint” reaction between free chlo-rine and ammonia (Figure 14.3). At doses below the hump in the chlorine residualcurve (zone 1), only combined chlorine is detectable. At doses between the humpand the dip in the curve, an oxidative destruction of combined residual chlorineaccompanied by the loss of nitrogen occurs (zone 2) (Taras, 1950). One possiblereaction during breakpoint is:

2 NH3 + 3 HOCl = N2 + 3 H+ + 3 Cl− + 3 H2O (14.16)

This reaction also may be used as a means to remove ammonia nitrogen fromwater or wastewaters (Pressley et al., 1972). Finally, after the ammonia nitrogen has


been completely oxidized, the residual remaining consists almost exclusively of freechlorine (zone 3). The minimum in the chlorine residual–versus–dose curve (in thiscase Cl2:NH4

+ − N weight ratio of 7.6/1) is called the breakpoint and denotes theamount of chlorine that must be added to a water before a stable free residual canbe obtained.

DISINFECTION 14.11

FIGURE 14.2 Effect of increased chloride dosage on residual chlorine and germicidalefficiency; pH 7.0, 20°C, NH3 10 mg/L. (Source: Adapted from Weber et al., 1940.)

FIGURE 14.3 Schematic idealization of breakpoint curve. (Source:Adapted from G. C.White, Disinfection of Wastewater and Water for Reuse,Van Nostrand Reinhold, New York. Copyright 1978.)

In their investigations of the chlorination of drinking water, Griffin and Cham-berlin (1941a,b) observed that:

1. The classical hump and dip curve is only seen at water pHs between 6.5 and 8.5.2. The molar ratio between chlorine and ammonia nitrogen dose at the breakpoint

under ideal conditions is 2:1, corresponding to a mass dose ratio (Cl2:NH4+ − N)

of 10:1.3. In practice, mass dose ratios of 15:1 may be needed to reach breakpoint.

The breakpoint reaction may also affect the pH of a water. If sodium hypochlo-rite is used as the source of active chlorine, as breakpoint occurs, the pH decreasesdue to an apparent release of protons during the breakpoint process (Equation14.16). If gaseous chlorine is used, this effect is reinforced by the release of protonsby hydrolysis of gaseous chlorine according to Equations 14.6 and 14.7 (McKee,1960).

The oxidation of ammonia nitrogen by chlorine to gaseous nitrogen at the break-point would theoretically require 1.5 mol of chlorine (Cl2) per mole of nitrogen oxi-dized according to Equation 14.16.The observed stoichiometric molar ratio betweenchlorine added and ammonia nitrogen consumed at breakpoint is typically about2:1, suggesting that more oxidized nitrogen compounds are produced at breakpointrather than N2 gas. Experimental evidence (Saunier and Selleck, 1979) indicates thatthe principal additional oxidized product may be nitrate formed via Equation 14.17:

NH4+ + 4 HOCl = NO3

− + 4 Cl− + 6 H+ + H2O (14.17)

Depending upon the relative amount of nitrate formed in comparison to nitrogenat breakpoint, between 1.5 and 4.0 mol of available chlorine may be required, whichis consistent with the available data.

Below the breakpoint, inorganic chloramines decompose by direct reactions withseveral compounds. For example, monochloramine may react with bromide ions toform monobromamine (Trofe, 1980). If trichloramine is formed, as would be the casefor applied chlorine doses in excess of that required for breakpoint, it may decom-pose either directly to form nitrogen gas and hypochlorous acid or by reaction withammonia to form monochloramine and dichloramine (Saguinsin and Morris, 1975).In distilled water, the half-life of monochloramine is approximately 100 h (Kinmanand Layton, 1976). Even in this simple circumstance, however, the decompositionproducts have not been completely characterized. Valentine (1986) found that thedecomposition of pure solutions of monochloramine produces an unidentified prod-uct that absorbs UV light at 243 nm and is capable of being oxidized or reduced.

Where the pH is below 9.0 (so that the dissociation of ammonium ion is negligi-ble), the amount of combined chlorine in dichloramine relative to monochloramineafter the reactions in Equations 14.13 and 14.14 have attained equilibrium is givenby the following relationship (McKee, 1960):

A = − 1 (14.18)

In Equation 14.18, A is the ratio of available chlorine in the form of dichloramineto available chlorine in the form of monochloramine, Z is the ratio of moles of chlo-rine (as Cl2) added per mole of ammonia nitrogen present, and B is defined by Equa-tion 14.19:

B = 1 − 4 Keq[H+] (14.19)

BZ��1 − �1� −� B�Z�(2� −� Z�)�


The equilibrium constant in Equation 14.19 refers to the direct interconversionbetween dichloramine and monochloramine as follows:

H+ + 2 NH2Cl = NH4+ + NHCl2

Keq = [NH4+][NHCl2]/[H+][NH2Cl]2 (14.20)

At 25°C, Keq has a value of 6.7 × 105 L/mol (Gray et al., 1978). From these rela-tionships, determination of the equilibrium ratio of dichloramine to monochlor-amine as a function of pH and applied chlorine dose ratio is possible (assuming nodissipative reactions other than those involving the inorganic chloramines). As pHdecreases and the Cl:N dose ratio increases, the relative amount of dichloraminealso increases (Figure 14.4). As the Cl:N molar dose ratio increases, the relativeamount of dichloramine also increases. As the Cl:N molar dose ratio increasesbeyond unity, the amount of dichloramine relative to monochloramine rapidlyincreases as well. For the conversion from dichloramine to trichloramine, the equi-librium constant given at 0.5 M ionic strength and 25°C indicates that the amount oftrichloramine to be found in equilibrium with di- and monochloramine at molardose ratios of up to 2.0 is negligible (Gray et al., 1978). This agrees with experimen-tal measurement of the individual combined chlorine species as a function ofapproach to breakpoint (White, 1972).

These findings, coupled with the routine observation of the breakpoint at molardoses at or below 2:1 (Cl2-to-N weight ratios below 10:1), indicate that trichloramineis not an important species in the breakpoint reaction. Rather, the breakpoint reac-tion leading to oxidation of ammonia nitrogen and reduction of combined chlorineis initiated with the formation of dichloramine.

DISINFECTION 14.13

FIGURE 14.4 Effect of pH and Cl2:NH4+ molar ratio on dichloramine-to-

monochloramine ratio (25°C).

The kinetics of formation of chloramine species have been investigated by vari-ous researchers since initial attempts by Weil and Morris (1949b). The formation ofmonochloramine is a first-order process with respect to both hypochlorous acid andun-ionized ammonia. However, determining whether this, or a process involvinghypochlorite ions reacting with ammonium cations, is the actual mechanism of reac-tion is not possible solely through kinetic arguments. If the neutral species areselected as the reactants, then the rate of formation of monochloramine (r) may bedescribed by (Morris and Isaac, 1983):

r (mol/L-s) = 6.6 × 108 exp(−1510/T) [HOCl][NH3] (14.21)

Because hypochlorous acid dissociates into hypochlorite with a pKa of approxi-mately 7.4 and ammonia is able to associate with a proton to form the ammoniumcation, with the pKa for the latter of approximately 9.3, for a constant chlorine:nitro-gen dose ratio, the maximum rate of monochloramine formation occurs at a pHwhere the product HOCl × NH3 is maximized, which is at the midpoint of the two pKvalues or 8.4. At this optimum pH and the usual temperatures encountered in prac-tice, the formation of monochloramine attains equilibrium in seconds to 1 min; how-ever, at either a higher or lower pH, the speed of the reaction slows.

A number of the other reactions in the chlorine-ammonia system may be kineti-cally limited. These have recently been reviewed; Table 14.2 is a compilation of theknown reaction kinetics involving chlorine, ammonia, and intermediate species.

The reaction of NH2Cl with HOCl to form NHCl2 is catalyzed by a number ofacidic species that may be present in water (Valentine and Jafvert, 1988). Possibly, anumber of the other reactions in Table 14.2 can also be catalyzed in a similar man-ner; however, insufficient data are available to evaluate this possibility.

When free chlorine is contacted with a water containing ammonia, the initialvelocity of monochloramine formation is substantially greater than the velocity ofthe subsequent formation of dichloramine. Hence, relative to equilibrium levels, aninitial accumulation of monochloramine will occur if large dose ratios are employed,until the dichloramine formation process can be driven (Palin, 1983).


TABLE 14.2 Summary of Chlorine Reaction Kinetics

Reaction Forward rate expression Reverse rate expression

NH3 + HOCl ⇔ NH2Cl + H2O 6.6 × 108 exp �− � 1.38 × 108 exp �− �

NH2Cl + HOCl ⇔ NHCl2 + H2O 3 × 105 exp �− � 7.6 × 10−7 L/mol-s*

NHCl2 + HOCl ⇔ NCl3 + H2O 2 × 105 exp �− � 5.1 × 103 exp �− �

2NH2Cl ⇔ NHCl2 + NH3 80 exp �− � 24.0 L/mol-s*

Rates are in units of L/mol-s.Concentrations are in mol/L.Reactions are elementary and water is at unit activity.* Rate constant at 25°C.Source: Morris and Isaac, 1983.

2160�

T

5530�

T3420�

T

2010�

T

8800�

T1510�

T

The kinetic evolution of the chlorine-ammonia speciation process in batch sys-tems is described by a series of coupled ordinary differential equations. While theseare highly nonlinear, various authors have applied numerical integration techniquesfor their solutions and, below the breakpoint, have found reasonable concordancebetween model predictions and experimental measurements (Haag and Lietzke,1980; Isaac et al., 1985; Saunier and Selleck, 1979; Valentine and Jafvert, 1988).

The breakpoint process involves a complex series of elementary reactions, ofwhich Equations 14.16 and 14.17 are the net results. Saunier and Selleck (1979) pro-posed that hydroxylamine (NH2OH) and NOH may be intermediates in this reac-tion. However, sufficient evaluation of their proposed kinetic scheme for thebreakpoint process has not yet been achieved to justify its use for design applications.

EXAMPLE 14.2 A water supply is to be postammoniated. If the water has a pH of7.0, a free chlorine residual of 1.0 mg/L, and a temperature of 25°C, how muchammonia should be added such that the ratio of dichloramine to monochloramine is0.1 (assume that, upon the addition of ammonia, none of the residual dissipates)?

1. From Equations 14.17 and 14.18, the following is determined:

B = 1 − 4 Keq (10−7) = 1 − 4 (6.7 × 105)(10−7)

= 0.732

2. From Equation 14.16, noting that the problem condition specifies A = 0.1, the fol-lowing equation is to be solved:

0.1 = −1 +

This can be rearranged into a quadratic equation

−0.289 Z2 + 0.134 Z = 0

3. The single nonzero root gives Z = 0.463, which is molar ratio of chlorine (as Cl2)to ammonia nitrogen. Because chlorine has a molecular weight of 70, 1 mg/L offree chlorine has a molarity of 1.43 × 10−5. Therefore, 3.09 × 10−5 molarity ofammonia is required, or (multiplying by the atomic weight of nitrogen, 14) a con-centration of 0.43 mg/L as N of ammonia must be added.

Reactions with Organic Matter. Morris (1967) has determined that organicamines react with free chlorine to form organic monochloramines. The rate laws forthese reactions follow patterns similar to the inorganic monochloramine formationprocess, except that the rate constants are generally less. In addition, the rate con-stants for this process correlate with the relative basicity of the amine reactant.Organic chloramines may also be formed by the direct reaction between monochlor-amine and the organic amine, and this is apparently the most significant mechanismof organic N-chloramine formation at higher concentrations such as might exist atthe point of application of chlorine to a water (Isaac and Morris, 1980). Pure solu-tions of amino acids and some proteins yield breakpoint curves identical in shape tothose of ammonia solutions (Baker, 1947; Wright, 1936).

Free chlorine reacts with organic constituents to produce chlorinated organic by-products. Murphy (1975) noted that phenols, amines, aldehydes, ketones, and pyrrolegroups are readily susceptible to chlorination. Granstrom and Lee (1957) found thatphenol could be chlorinated by free chlorine to form chlorophenols of variousdegrees of substitution. The kinetics of this process depend upon both phenolate

0.732 Z��1 − �1� −� 0�.7�3�2�(2� −� Z�)Z�

DISINFECTION 14.15

ions and hypochlorous acid. If excess ammonia was present, however, the formationof chlorophenols was substantially inhibited.

More recently, DeLaat (1982) determined that polyhydric phenols are substan-tially more reactive than simple ketones in the production of chloroform, and thatthe rates of these processes are first order with respect to the phenol concentrationand the free chlorine concentration. More significantly, the reactivity of these com-pounds was observed to be greater than the reactivity of ammonia with hypo-chlorous acid. Therefore, even if subbreakpoint chlorination is practiced, somechloroform may be formed rapidly prior to the conversion of free to combined chlo-rine. Chapter 12 presents additional discussion of the formation of trihalomethanesand other disinfection by-products that can arise from reactions with naturallyoccurring dissolved humic substances.

The reactivity of the chlorine species with compounds responsible for taste andodor depends on the predominant form of chlorine present. In field tests, Krasner(1986) determined that free chlorine, but not combined chlorine, could removetastes and odors associated with organic sulfur compounds.

Reactions with Other Inorganic Compounds. The rates of reaction between freechlorine residuals and other inorganic compounds likely to be present in water aresummarized in Table 14.3 (Wojtowicz, 1979). These reactions are generally firstorder in both the oxidizing agent (hypochlorous acid or hypochlorite anion) and thereducing agent.

Nitrites present in partially nitrified waters react with free chlorine via a com-plex, pH-dependent mechanism (Cachaza, 1976).While combined chlorine residualswere generally thought to be unreactive with nitrite,Valentine (1985) has found thatthe rate of decay of monochloramine in the presence of nitrite was far greater thanwould be predicted based on reaction of the equilibrium free chlorine, implicating adirect reaction between NH2Cl and NO2�N.

Overall Chlorine Demand Kinetics. Chlorine demand is defined as the differencebetween the applied chlorine dose and the chlorine residual measured at a particu-lar time. The rate of exertion of chlorine demand in complex aqueous solutions hasbeen the subject of numerous studies. The most systematic work has been that of


TABLE 14.3 Summary of Kinetics of HOCl and OCl− Reduction by Miscellaneous Reduc-ing Agents after Wojtowicz (1979)

Oxidizing agent Reducing agent Oxidation product Log k, L/m-s, 25°C

OCl− IO− IO4− −5.04

OCl− OCl− ClO2− −7.63

OCl− ClO2− ClO3

− −5.48OCl− SO3 SO4

2− 3.93HOCl NO2

− NO3− 0.82

HOCl HCOO− H2CO −1.38HOCl Br− BrO− 3.47HOCl OCN− HCO3

−, N2 −0.55HOCl HC2O4

− CO2 1.20HOCl I− IO− 8.52

Source: Feng, 1966.

Taras (1950), who chlorinated pure solutions of various organic compounds andfound that chlorine demand kinetics could be described by Equation 14.22:

D = ktn (14.22)

where t is the time in hours, D is the chlorine demand, and k and n are empirical con-stants. In subsequent work, Feben and Taras (1950, 1951) found that chlorinedemand exertion of waters blended with wastewater could be correlated to Equa-tion 14.22, with the value of n correlated to the 1-h chlorine demand.

Haas and Karra (1984) developed Equation 14.23 to describe chlorine demandexertion kinetics.

D = Co{1 − [x exp (−k1t) + (1 − x) exp (−k2t)]} (14.23)

where x is an empirical parameter, typically 0.4 to 0.6, k1 and k2 are rate constants,typically 1.0 min−1 and 0.003 min−1, respectively, and Co is the chlorine dose in mg/L.

Dugan et al. (1995) developed a Monod (Langmuir Hinshelwood) model fordescribing free chlorine decay in drinking water in the absence of ammonia. Itdescribes chlorine decay as a reaction with total organic carbon (TOC) in wateraccording to the following differential equation:

= − (14.24)

where TOC (assumed constant) is in mg C/L and C is the free chlorine concentrationin mg Cl2/L. Equation 14.24 can be integrated to the following implicit equation forchlorine concentration at time t (Ct):

Ct = K (TOC) ln � � − k (TOC) t + C0 (14.25)

Furthermore, k and K were correlated with the initial chlorine dose (C0) andTOC concentration. In tests conducted on a variety of waters, the constants werefound to be given by the following equations (it should be noted that the pH andtemperature were fixed at 8 and 20°C, and the waters were of relatively low ionicstrength, so that the applicability of these relationships under other conditions isunclear):

K = −0.85 � �k = 0.030 − 0.0060 � � (14.26)

Dechlorination. When the chlorine residual in a treated water must be loweredprior to distribution, the chlorinated water can be dosed with a substance that reactswith or accelerates the rate of decomposition of the residual chlorine. Compoundsthat may perform this function include thiosulfate, hydrogen peroxide, ammonia,sulfite/bisulfite/sulfur dioxide, and activated carbon; however, only the latter twomaterials have been widely used for this purpose in water treatment (Snoeyink andSuidan, 1975).

C0�TOC

C0�TOC

C0�Ct

k(TOC)C��K(TOC) + C

dC�dt

DISINFECTION 14.17

Chlorine Dioxide

The reaction of chlorine dioxide with material present in waters containing chlorinedioxide demand appears to be less significant than in the case of chlorine. Rather,the dominant causes of loss of chlorine dioxide during disinfection may be the directreactions with water and interconversions to chlorite and chloride, as outlined inEquations 14.10 and 14.11. At mg/L concentrations, ammonia nitrogen, peptone,urea, and glucose have insignificant chlorine dioxide demand in 1 h (Ingolls andRidenour, 1948; Sikorowska, 1961). However, a variety of inorganic and biologicalmaterials will react (Werderhoff and Singer, 1987).

Masschelein (1979a) concluded that only the following organic-ClO2 reactionsare of significance to water applications:

1. Oxidation of tertiary amines to secondary amines and aldehydes2. Oxidation of ketones, aldehydes, and (to a lesser extent) alcohols to acids3. Oxidation of phenols4. Oxidation of sulfhydryl-containing amino acids

Wajon et al. (1982) found a reaction stoichiometry of 2 mol of chlorine dioxideconsumed per mole of phenol (or hydroquinone) consumed. Products formedincluded chlorophenols, aliphatic organic acids, benzoquinone, and (in the case ofphenol) hydroquinone. The mechanism appeared to include the possible formationof hypochlorous acid as an active intermediate, and the rate of this process wasfound to be base catalyzed and first order in each of the reactants.

In general, chlorine dioxide itself has been found to produce fewer organic by-products with naturally occurring dissolved organic material, although some non-purgeable organic halogenated compounds are formed (Rav-Acha, 1984). Inpractice, however, chlorine dioxide may be generated in a manner in which chlo-rine is present as an impurity. Therefore, the reactions of such a stream may alsoinclude those discussed earlier regarding chlorine reactions. The inorganic by-products consist of chloride, chlorate, and chlorite; specific ratios may depend on the precise application conditions (Noack and Doeff, 1981; Werderhoff andSinger, 1987).

Ozone

Upon addition to water, ozone reacts with hydroxide ions to form hydroxyl radicalsand organic radicals. These radicals cause increased decomposition of ozone, andalso are responsible for nonselective (compared to the direct ozone reaction) oxida-tion of a variety of organic materials. Carbonate, and possibly other ions, may act asradical scavengers and slow this process (Hoigne and Bader, 1975, 1976).

Gurol and Singer (1982) determined that ozone decomposition kinetics in vari-ous aqueous solutions are second order in ozone concentration and base promoted.Some systematic difference between various buffer systems employed does occur,with borate giving higher decomposition rates than phosphate, and phosphate athigher ionic strength giving lower decomposition rates than phosphate at lowerionic strength (1 versus 0.1 M). This effect was suggested as being caused by phos-phate being a radical scavenger (and by radical decomposition being important athigher pH values).

As a result of these decomposition processes, the half-life of ozone in water, evenin the absence of other reactive constituents, is quite short, on the order of seconds


DISINFECTION 14.19

FIGURE 14.5 Schematic of bromate formation pathways. Solid lines: direct ozonereactions. Dashed lines: radical reactions (Source: Reprinted with permission from von Gunten, U., and J. Hoigne, 1994. Bromate formation during ozonation of bromide-containing waters: interaction of ozone and hydroxyl radical reactions. EnvironmentalScience and Technology 28(7): 1234–1242. Copyright 1994 American Chemical Society.)

to minutes. Water chemistry may exert a strong influence on the rate and extent ofozone demand in a given application. Reactions of ozone in aqueous solution arediscussed further in Chapter 12.

Bromide reacts with ozone under aqueous conditions typical of drinking waterdisinfection. Products of the reaction may be hypobromous acid, hypobromite,and/or bromate. Higher concentrations of bromide can reduce the rate of ozonedecomposition. Under alkaline conditions, this may be influenced by trace metalcatalysts and organic sinks for radicals and oxidized bromide species (Cooper etal., 1985).

Ozone will react with cyanides at a very fast rate. The mechanism involves reac-tion of the cyanide ion (to form unknown products), and the process is inhibited byiron complexes but catalyzed by copper complexes of cyanide (Gurol et al., 1985).

The reaction of ozone with bromide may proceed to the further product of bro-mate (BrO3

−) by a complex process that involves direct reaction as well as hydroxylradical mediation (von Gunten and Hoigne, 1994). The overall process is summa-rized in Figure 14.5. The formation of bromate by ozonation is highly important inview of the potential carcinogenicity of bromate in disinfected waters (Bull andKopfler, 1991).

Demand for UV

For UV disinfection, the “dose” may be described in terms of the emitted lamppower in the germicidal range per unit volume of fluid under irradiation, for exam-ple,W/m3.This can also be expressed as an integral over the disinfection reactor vol-ume of the surface intensity (in W/m2, for example) (Severin et al., 1983a, 1984b).

With ultraviolet light disinfection systems, the equivalent of demand results fromdissolved and suspended materials, such as proteins, humic material, and iron com-pounds, that absorb radiation and thus shield microorganisms. Huff et al. (1965)found that intensity monitoring within the reactor itself could be used to correct forsuch effects.

One particular problem unique to physical systems such as UV is the need toassure complete mixing in the transverse direction so that all microorganisms maycome equally close to the UV source. Cortelyou (1954) analyzed this effect for batchUV reactors, and the analysis was extended to flow-through reactors by Haas andSakellaropoulous (1979).This phenomenon results in the desirability to achieve tur-bulent flow conditions in a UV reactor.

ASSESSMENT OF MICROBIAL QUALITY(INDICATORS)

The microbial quality of a source water, or the efficacy of a treatment system forremoving microorganisms, can be assessed either by direct monitoring of pathogensor by the use of an indicator system. Because pathogens are a highly diverse group,generally requiring a highly specialized (and often insensitive and expensive) ana-lytical technique for each pathogen, the use of indicator organisms is a more popu-lar technique.

An indicator group of organisms can be used either to assess source water con-tamination or degree of treatment; however, the same indicator group is often usedto assess both properties. This places severe constraints on the group of indicatororganisms chosen. Bonde (1966) has proposed that an ideal indicator must:

1. Be present whenever the pathogens concerned are present2. Be present only when the presence of pathogens is an imminent danger, that is,

be unable to proliferate to any greater extent in the aqueous environment3. Occur in much greater numbers than pathogens4. Be more resistant to disinfectants and to the aqueous environment than pathogens5. Grow readily on relatively simple media6. Yield characteristic and simple reactions enabling, as far as possible, an unam-

biguous identification of the group7. Be randomly distributed in the sample to be examined, or be able to be uniformly

distributed by simple homogenization procedures8. Grow widely independent of other organisms present when inoculated in artifi-

cial media, that is, not be seriously inhibited in growth by the presence of otherbacteria

The use of coliforms as indicator organisms stems from the pioneering work ofPhelps (1909).The basic rationale was that coliforms and enteric bacterial pathogensoriginate from a common source—namely human fecal contamination. Subsequentwork by Butterfield et al. (1943, 1946), Kabler (1951), and Wattie and Butterfield(1944) confirmed that these organisms were at least as resistant to free or combinedchlorine as enteric bacterial pathogens.

The coliform group is a heterogeneous conglomerate of microorganisms, includ-ing forms native to mammalian gastrointestinal tracts as well as a number of exclu-sively soil forms. The common fermentation tube (FT) and membrane filter (MF)procedures are subtly different in the organisms they enumerate. Classically, coli-forms have been defined as “Gram-negative, non-sporeforming bacteria which [sic]ferment lactose at 35–37°C, with the production of acid and gas” (APHA, AWWA,and WPCF, 1989). The FT procedure, however, ignores anaerogenic and lactose-negative coliforms, and the MF procedure ignores non-lactose-fermenting strains(Clark and Pagel, 1977).

Furthermore, interferences can selectively reduce coliforms as measured by oneor the other method. Allen (1977), for example, found that high concentrations(>500 to 1000/mL) of standard plate count (SPC) organisms appeared to reduce therecovery of coliforms by the MF technique when compared to the FT technique.

The fecal coliform group of organisms is that subset of coliforms that are capableof growing at elevated temperature (44.5°C).The original rationale for developmentof this test was to provide a more selective indicator group, excluding mesophilic


coliforms primarily indigenous to soils. Total coliforms, however, continue to be thebasic U.S. microbiological standard for drinking water because the absence of col-iforms ensures the absence of fecal coliforms, which is a conservative standard.

While coliforms, either fecal or total, may be reasonably good indicators of fecalcontamination of a water supply, reservations were expressed as early as 1922(Anonymous, 1922) about the relative resistance of coliforms to chlorine vis-à-vispathogenic bacteria and the resulting adequacy of the coliform test as an indicatorof disinfection efficiency. In more recent work, coliforms have been found to bemore sensitive to disinfection by one or more forms of chlorine than various humanenteric viruses (Grabow et al., 1983; Kelly and Sanderson, 1958) and the protozoanpathogens Naegleria (Rubin et al., 1983), Giardia (Jarroll, 1981; Korich et al., 1990;Leahy, 1985; Rice et al., 1982), and Cryptosporidium (Kovich et al., 1990). In addi-tion, viruses (Scarpino et al., 1977) and protozoan cysts (Leahy, 1985) have beenfound to be more resistant to ClO2 inactivation than coliforms. Farooq (1976) hasdetermined that coliforms are more resistant to ozone than viruses. Rice and Hoff(1981) found that Giardia lamblia cysts survived exposure to UV doses sufficient toeffect over 99.99 percent inactivation of E. coli. Human enteric viruses have beenisolated in full-scale water treatment plants practicing conventional treatment, andmeeting turbidity and coliform standards in the presence of free residual chlorine(Payment et al., 1985; Rose et al., 1986).

As a result of the problems with the coliform group of organisms, a number ofworkers have investigated alternative indicator systems with greater resistance todisinfectants than coliforms. Among the most successful of these are the acid-fastbacteria and yeasts studied by Engelbrecht et al. (1977, 1979) and Haas et al. (1983a,b; 1985a, b). In addition, work using endotoxins (Haas and Morrison, 1981),Clostridia (Cabelli, 1977; Payment and Franco, 1993), and bacteriophage (Abad etal., 1994; Grabow, 1968; Grabow et al., 1983; Payment and Franco, 1993) has beencarried out. In addition, to some degree, heterotrophic plate count (HPC) organismsmay provide a conservative indicator of treatment efficiency. Despite these studies,however, in U.S. practice, no alternative to the total coliform group of organisms hasyet found widespread application.

PATHOGENS OF CONCERN

A variety of pathogenic organisms capable of transmission by the fecal-oral routemay be found in raw wastewaters.Waterborne outbreaks of shigellosis, salmonellosis,and various viral agents have been reported, in many cases associated with sewage-contaminated water supplies (Blostein, 1991; Drenchen and Bert, 1994; Haas, 1986;Herwaldt et al., 1991, 1992; Levine et al., 1990; Reeve et al., 1989; Rosenberg et al.,1976, 1980).Among the bacteria, Salmonella, Shigella, and Vibrio cholerae organismsare the classical agents of concern (Mosley, 1966). In more recent times, concern hasexpanded to other agents that have been found in wastewater—viruses and protozoa.

Among the viruses, enteroviruses (ECHO virus, Coxsackievirus), rotavirus,reovirus, adenovirus, and parvovirus have been isolated from wastewater (Melnicket al., 1978). New viruses that are suspected of waterborne transmission have beenidentified at the rate of about one organism per year (Gerba, personal communica-tion). Among the more important of these newly identified agents may be Norwalkvirus and calicivirus.

Over the past 15 years, significant concerns have increased over the risk frompathogenic protozoa in drinking water, particularly Giardia and Cryptosporidium

DISINFECTION 14.21

(Gallaher et al., 1989; Goldstein et al., 1996; LeChevallier et al., 1991; Leland et al.,1993; Richardson et al., 1991; Rose et al., 1991; Smith, 1992). The SWTR arose, to asignificant extent, from concerns over Giardia (Regli et al., 1988). Revisions todrinking water regulations presently under discussion are concerned with assuringan adequate degree of protection from Cryptosporidium.

DISINFECTION KINETICS

The information needed for the design of a disinfection system includes knowledgeof the rate of inactivation of the target, or indicator, organism(s) by the disinfectant.In particular, the effect of disinfectant concentration on the rate of this process willdetermine the most efficient combination of contact time (i.e., basin volume at agiven design flow rate) and the dose to employ.

Chick’s Law and Elaborations

The major precepts of disinfection kinetics were enunciated by Chick (1908), whorecognized the close similarity between microbial inactivation by chemical disinfec-tants and chemical reactions. A good overview of the principles of kinetic modelingof disinfection has been presented by Gyurek and Finch (1998). Disinfection is anal-ogous to a bimolecular chemical reaction, with the reactants being the microorgan-ism and the disinfectant, and can be characterized by a rate law as are chemicalreactions:

r = −kN (14.27)

where r is the inactivation rate (organisms killed/volume-time) and N is the concen-tration of viable organisms. In a batch system, this results in an exponential decay inorganisms, because the rate of inactivation equals dN/dt, assuming that the rate con-stant k is actually constant (e.g., the disinfectant concentration is constant).

Watson (1908) proposed Equation 14.28 to relate the rate constant of inactiva-tion k to the disinfectant concentration C:

k = k′Cn (14.28)

where n is termed the coefficient of dilution and k′ is presumed independent of disin-fectant concentration, and, by virtue of Equation 14.27, microorganism concentration.

From the Chick-Watson law, when C, n, and k′ are constant (i.e., no demand, con-stant concentration), the preceding rate law may be integrated so that in a thor-oughly mixed batch system,

ln(N/N0) = −k′Cnt (14.29)

where N and N0 are, respectively, the concentrations of viable microorganisms attime t and time 0. When disinfectant composition changes with time, or when a con-figuration other than a batch (or plug flow) system is used, the appropriate rate lawscharacterizing disinfectant transformation (Haas and Karra, 1984b) along with theapplicable mass balances must be used to obtain the relationship between microbialinactivation and concentration and time.


DISINFECTION 14.23

FIGURE 14.6 Chick’s law and its deviations.

Inactivation of microorganisms in batch experiments, even when disinfectantconcentration is kept constant, does not always follow the exponential decay pat-tern predicted by Equation 14.29. Indeed, two common types of deviations arenoted (Figure 14.6). In addition to the linear Chick’s law decay, the presence of“shoulders” or time lags until the onset of disinfection is often observed. Also,some microorganisms and disinfectants exhibit a “tailing” in which the rate ofinactivation progressively decreases. In some cases, a combination of both of thesebehaviors is seen.

Even if deviations from Chick-Watson behavior are observed, plotting combina-tions of disinfectant concentration and time to produce a fixed percent inactivationis generally possible. Such plots tend to follow the relationship Cnt = constant, wherethe constant is a function of the type of organism, pH, temperature, form of disin-fectant, and extent of inactivation. Such plots are linear on a log-log scale (Figure14.7). If the value of n is greater than 1, a proportionate change in disinfectant con-centration produces a greater effect than a proportionate change in time. In manycases (Hoff, 1986), the Chick-Watson law n value is close to 1.0, and hence a fixedvalue of the product of concentration and time (CT product) results in a fixeddegree of inactivation (at a given temperature, pH, etc.).

In the chemical disinfection of a water, the concentration of disinfectant maychange with time, and particularly during the initial moments of contact the chemi-cal form(s) of halogens such as chlorine undergo rapid transformations from thefree to the combined forms. Because C would thus not be a constant, typically disin-fection results obtained in batch systems exhibit tailing, the degree of which maydepend on the demand and the concentration of reactive constituents (such asammonia) in the system (Olivieri et al., 1971). Determination of the disinfectantresidual (and its chemical forms) is more critical than the disinfectant dose in thesesystems.

In the chlorine system, for example, knowing the rate laws for inactivation byindividual separate species and the dynamics of chlorine species interconversions asdescribed previously enables an overall model for chlorine inactivation to be for-mulated (Haas and Karra, 1984b). In doing this computation, the individual rates areusually assumed to be additive (Fair et al., 1948), although this assumption has notyet been experimentally verified.

The presence of shoulders in inactivation curves is often seen in organisms thatform clumps.This means that more than one cell must be inactivated to achieve inac-tivation of a colony or plaque-forming unit. For example, Rubin et al. (1983) foundthat cysts of Naegleria gruberii in demand-free water showed shoulder-type inacti-vation to free chlorine. Similarly, when cells of E. coli were agglutinated, they dis-played shoulder-type inactivation, which was absent in unagglutinated cultures(Carlson et al., 1975). Severin et al. (1984a) found shoulder-type inactivation curvesin the case of E. coli with preformed chloramines (i.e., solutions of ammonia andchlorine prereacted to form combined chlorine prior to addition of microorgan-isms), Candida parapsilosis (a yeast organism proposed as a possible disinfection-resistant indicator) with both preformed chloramines, and free chlorine andpoliovirus with iodine.

Shoulder inactivation curves may be explained by a multitarget model (Hiatt,1964), by a series event model (Severin et al., 1984a), or by a diffusional model


FIGURE 14.7 Concentration-time relationships for 99 percent inactivation of vari-ous microorganisms by various disinfectants. (1) Giardia lamblia; free chlorine, 5°C(Source: Hoff and Akin, 1986). (2) E. coli; free chlorine, 2 to 5°C, pH 8.5 (Source: Haasand Karra, 1984a). (3) E. coli; free chlorine, 20 to 25°C, pH 8.5 (Source: Haas andKarra, 1984a). (4) Poliovirus 1 (Mahoney); free chlorine, 2°C, pH 6 (Source: Haas andKarra, 1984a). (5) E. coli; combined chlorine, 3 to 5°C, pH 7 (Source: Haas and Karra,1984a). (6) Poliovirus 1 (Mahoney); ozone, 20°C, pH 7.2 (Source: Roy et al., 1981a).(7) Giardia muris; ozone, 5°C, pH 7 (Source: Wickramanayake et al., 1985).

(Haas, 1980). Tailing inactivation curves may be explained either by a vitalistichypothesis in which individuals in a population are nonidentical, and their inherentresistance is distributed in a permanent (time-independent) manner, or by a mecha-nistic concept (Cerf, 1977). In the latter case, four particular mechanisms have beenadvanced leading to tailing:

1. Conversion to resistant form during inactivation (hardening)

2. Existence of genetic variants of differing sensitivity

3. Protection of a subpopulation, or variations in received dose of disinfectant

4. Clumping of a subpopulation

The hardening process and resultant tailing have received wide attention, follow-ing discoveries of apparent hardening in the formaldehyde inactivation of poliovirusprepared for the Salk vaccine (Nathanson and Langmuir, 1963). Gard (1960) hasproposed an empirical rate law for this behavior, which has been used by Selleck etal. (1978) in the analysis of wastewater chlorination kinetics. Tailing behavior hasbeen found for viral and coliform inactivation by ozone (Katzenelson et al., 1974)and for coliform inactivation by free chlorine (Haas and Morrison, 1981; Olivieri etal., 1971).

Hom (1972) developed a flexible but highly empirical kinetic formulation for theinactivation rate based on modifying Equations 14.27 and 14.28 to the followingform:

r = −k′mN tm − 1Cn (14.30)

This equation is difficult to use as a rate model since it contains time as an explicitvariable. A formulation leading to the classical Hom integrated relationship can bewritten as (Haas and Joffe, 1994):

r = −mN(kCn)1/m �−ln � ��(1 − 1/m)

(14.31)

Upon integration, if C is constant, this results in the following relationship:

ln � � = −k′Cntm (14.32)

Depending upon the value of m, both shoulders and tailing may be depicted by Equation 14.32. In early work, Fair et al. (1948) used a model of the form ofEquation (14.32) with m = 2 to analyze E. coli inactivation by free and combinedchlorine.

EXAMPLE 14.3 A certain water supply has operational problems due to high levelsof HPC organisms. To maintain adequate system water quality, a decision has beenmade to keep the concentration of HPC organisms below 10/mL at the entry pointto the distribution system (i.e., following disinfection). Disinfection using free resid-ual chlorine is practiced. As part of the laboratory investigation to develop designcriteria for this system, the inactivation of the HPC organisms is determined in batchreactors (beakers). The pH and temperature are held constant at the expected finalwater conditions. Using water with an initial HPC of 1000/mL, the following data aretaken:

N�N0

N�N0

DISINFECTION 14.25

Cl2 residual, Contact time, HPC remaining,mg/L min number/mL

0.5 10 400.5 20 40.5 30 11 5 351 10 41 15 11.5 2 981.5 5 101.5 10 1

From this information, determine the best fit using the Hom inactivation model,and compute the necessary chlorine residual that will achieve HPC < 10/mL (froman initial concentration of 1000/mL) at a contact time of 10 min.

1. This problem can be solved using a maximum likelihood technique to fit theHom model to the data. Regression methods, however, can be used in two differ-ent ways—multiple linear regression and nonlinear regression (Haas and Heller,1989; Haas et al., 1988). In a batch system, without chlorine demand, the Hommodel becomes

ln � � = −kCntm

This can be rearranged as

ln �−ln � �� = ln (k) + n ln (C) + m ln (t)

2. A multiple linear regression using ln [− ln (N/N0)] as the dependent variable andln(C) and ln(t) as the independent variables produces an intercept (equal toln(k)) and slopes equal to n and m. This computation can be handled by commonspreadsheet programs as well as statistical packages. Transformation of the datagiven produces the following values of the dependent and independent variables:

ln (−ln (N/N0)) ln(t) ln(C)

1.169 2.303 −0.6931.709 2.996 −0.6931.933 3.401 −0.6931.210 1.609 01.709 2.303 01.933 2.708 00.843 0.693 0.4051.527 1.609 0.4051.933 2.303 0.405

The result is

k = 1.11

n = 0.68

m = 0.70

Correlation coefficient = 0.997

N�N0

N�N0


3. The final estimation equation then becomes:

ln � � = −1.11 C 0.68t 0.70

4. Inserting the design specifications gives the following results:

ln � � = −1.11 C 0.6810 0.70

C 0.6833 = 0.83

C = 0.77

Hence, if a 10-min contact time is accepted as a worst-case condition, and assum-ing good contactor hydraulics, the maximum chlorine residual required to achievethe design inactivation is 0.77 mg/L. From this information, and the chlorine demandof the water, the capacity of a chlorine feed system can be computed.

Another class of models can be obtained by assuming that inactivation is otherthan first order in surviving microbial concentrations. Depending upon the orderchosen, either tailing or shoulders can be produced. For example, Roy and cowork-ers (Roy, 1979; Roy et al., 1981a, b), using continuously stirred tank reactor studieson inactivation of Poliovirus 1 with ozone in demand-free systems, developed thefollowing rate law:

r = −kCN 0.69 (14.33)

A similar model was used by Benarde et al. (1967) to analyze E. coli inactivationby chlorine dioxide, and for analysis of various organisms in ozone contacting (Zhouand Smith, 1994, 1995).

Disinfection, like all other rate processes, is temperature dependent. This depen-dency may be quantified by the Arrhenius relationship:

k = ko exp(−E/RT) (14.34)

where k is a rate constant characterizing the reaction (such as the Chick-Watson k′value), T is the absolute temperature, R is the ideal gas constant, ko is called the fre-quency factor, and E, with units of energy/mole, is called the activation energy. E isalways positive, and as it increases, the effect of temperature becomes more pro-nounced. The values of E and ko may be determined from rates of inactivationobtained as a function of temperature. As E increases, the effect of temperature onthe rate increases. For example, an E of 10 kcal/mol doubles the rate between 10 and20°C. In contrast, activation energies for breaking hydrogen bonds are 3 to 7kcal/mol (Bailey and Ollis, 1986). Activation energies less than this range suggestphysical (e.g., transport) limitations rather than chemical reactions.

If the disinfectant concentration is not constant, then the rate laws for inactiva-tion must be combined with those for disinfectant demand. For example, in the caseof the Hom model with first-order decay in a batch system, the following approachhas been developed (Haas and Joffe, 1994). The time course of disinfectant concen-tration is given by:

C = C0 exp (−k*t) (14.35)

where C0 is the initial concentration and k* is the demand constant. This equation issubstituted into the Hom rate expression and then into a batch mass balance to yield:

= −mN[k(C0exp(−k*t))n]1/m �−ln � ��(1 − 1/m)

(14.36)N�N0

dN�dt

mg�L

10�1000

N�N0

DISINFECTION 14.27

This has an analytical solution in terms of an incomplete gamma function, whichunder conditions generally present in disinfection can be approximated as:

ln � � = −kC0ntm� �

m

(14.37)

For other disinfection decay rate laws, the solution may be obtained by numeri-cal integration (Finch et al., 1993).

In general, microbial inactivation kinetics have been determined in batch sys-tems. Real contactors, however, are continuous and may have nonideal flow patternswith backmixing or short-circuiting. If the residence time distribution of a continu-ous reactor is known, for example from tracer experiments, then it is possible todevelop reasonable estimates of inactivation efficiency in these systems (Haas, 1988;Haas et al., 1995; Lawler and Singer, 1993; Stover et al., 1986; Trussell and Chao,1977).

The residence time distribution is determined from a pulse or step tracer exper-iment. In a pulse experiment, a virtually instantaneous “slug” of tracer is intro-duced, while in a step experiment a virtually instantaneous step change inconcentration is effected.Aspects of these experiments in disinfection systems havebeen described by several recent authors (Bishop et al., 1993; Boulos et al., 1996;Teefy and Singer, 1990). From this experiment, descriptive characteristics of the res-idence time distribution, such as the mean residence time and dispersion, may beobtained; in addition, the cumulative residence time distribution (F curve) and itsdensity function (E curve) are also directly obtained.The E curve, written as a func-tion of time, E(t), is the probability density function that gives the fraction of thefluid elements leaving the reactor that were in the reactor a period of time betweent and t + dt.

Using the E curve and the batch kinetic rate expression, the composition of theeffluent may be obtained. The approach assumes a completely segregated system(Dankwerts, 1958; Levenspiel, 1972), and estimates the survival ratio in a disinfec-tion reactor by the following equation:

ln � �continuous

= �∞

0

S(t)E(t)dt (14.38)

where ln (N/N0)continuous = predicted continuous survival ratioS(t) = predicted batch survival at time t, given influent concen-

tration as initial condition (N/N0)E(t) = normalized density function for the residence time distri-

butionThe complete segregation model is one extreme of micromixedness behavior.

Prior work has suggested that, while there is a difference between the predictionsfrom the perfect segregation assumption and the perfect micromixed assumption,the numerical differences are less than those associated with variations in the resi-dence time distribution itself (Haas, 1988).The integration in Equation 14.38 may bemost effectively performed if the E curve is fitted to a parametric model, such as thetanks-in-series model or axial dispersion model (Haas et al., 1997; Nauman andBuffham, 1983) given, respectively, by:

E(t) = � �w

exp �− � (14.39)wt�θ

wt�θ

1�tΓ(w)

N�N0

1 − exp �− �n

mk*t��

��

��n

mk*t��

N�N0


and

E(t) = exp �− � (14.40)

where w is the number of equal-volume tanks in series and ν is the dimensionlessvariance, which is related to the Peclet number—the reciprocal of the dimensionlessdispersion number—by the following relationship:

ν = − (1 − e−Pe) (14.41)

The axial dispersion model also has the integrated solution that gives F, thecumulative fraction of fluid that spends a residence time ≤ t in the system as

F(t) = Φ �� − 1� � + exp � � Φ �−� + 1� � (14.42)

where Φ (z) is the standard normal-probability integral (i.e., area under the normaldistribution from negative infinity to z).

Application of this method is illustrated by the following example.

EXAMPLE 14.4 Batch studies of inactivation of Giardia in water by chlorine haveshown the applicability of the Hom model with first-order decay. The following arethe kinetic parameters:

k* = 0.03

k = 0.1

m = 1.2

n = 0.9

A contactor is to be designed with a mean residence time of 30 min and a Pecletnumber of 50. If a chlorine dose of 2 mg/L is to be used, what is the anticipated sur-vival ratio?

1. First, from Equation 14.41, the dimensionless variance is computed as:

ν = − (1 − e−50) = 0.0392

2. Now, the trapezoidal rule (Chapra and Canale, 1988) can be used to evaluate theintegral as shown in the following table.

Time, min S(t) E(t) S(t) × E(t) Trapezoid area

0 1 0 05 0.3000 8.189E-24 2.457E-24 6.141E-24

10 0.0750 1.435E-08 1.076E-09 2.690E-0915 0.0192 3.228E-04 6.206E-06 1.552E-0520 0.0053 0.0147 7.808E-05 2.107E-0425 0.0016 0.0577 9.234E-05 4.261E-0430 5.308E-04 0.0672 3.565E-05 3.200E-0435 1.934E-04 0.0393 7.609E-06 1.081E-0440 7.713E-05 0.0151 1.162E-06 2.193E-0545 3.349E-05 0.0044 1.461E-07 3.271E-0650 1.574E-05 0.0010 1.637E-08 4.062E-0755 7.954E-06 2.158E-04 1.716E-09 4.522E-0860 4.299E-06 4.035E-05 1.735E-10 4.724E-09

2�502

2�50

θ�tν

t�θ

2�ν

θ�tν

t�θ

2�Pe2

2�Pe

(t − θ)2

�2θtν

θ�2πt3ν

DISINFECTION 14.29

Time intervals are selected—the finer the time spacing, the more precise the eval-uation of the integral. More sophisticated integration approaches, such as Simpson’srule, or higher-order methods may also be chosen to improve precision (Chapra andCanale, 1988).The second and third columns are the computed values of the survival(S) and residence time density function (E) at the particular time (t).The fourth col-umn is the product of S and E. The final column represents the contribution of therectangle to the area. If t0 and t1 represent the prior row and the current row, then thevalues in the final column are given by:

� � (t1 − t0)

The sum of the final column is 0.0011061—indicating that there is just under 3 logs removal, that is, 0.11 percent survival.

Note that times beyond 60 min were not examined, since at these higher times, asshown in the right column, the contribution to the integral is very small.

THE CT APPROACH IN REGULATION

For regulatory purposes, under the SWTR, the adequacy of disinfection is judgedusing the product of the final residual concentration of disinfectant (in mg/L) andthe contact time (in minutes). The contact time is evaluated as that which isexceeded by 90 percent of the fluid (this is designated as the t10—denoting that 10percent of the fluid has a smaller residence time in the system). The t10 for disinfec-tion systems may be evaluated from tracer studies, or by use of default multipliers ofthe theoretical hydraulic residence time (V/Q). Critical values of the CT to achievevarying levels of disinfection, incorporating a margin of safety, have been publishedunder EPA guidance documents (Malcolm Pirnie and HDR Engineering, 1991).These are indicated in Table 14.4 for ozone, chlorine dioxide, and chloramines. In thecase of free chlorine, the CT values are functions of temperature, pH, and also con-centration (under the regulations, measured at the end of the contact chamber).Space does not permit a recapitulation of the full CT tables for free chlorine—thevalues at a few conditions are given in Table 14.5.

The use of the CT approach is illustrated in the following example.

EXAMPLE 14.5 The disinfection contactor described in Example 14.4 is to be usedto achieve a 3-log inactivation of Giardia. Compute the required residual necessaryto achieve this result from the CT tables if the pH is 7 and the temperature is 20°C.

S(t1)E(t1) + S(t0)E(t0)��

2


TABLE 14.4 CT Values for 99.9 Percent Reduction of Giardia lamblia with Ozone, ChlorineDioxide, and Chloramines

Temperature, °C

Disinfectant pH 1 5 10 15 20 25

Ozone 6–9 2.9 1.9 1.4 0.95 0.72 0.48Chlorine dioxide 6–9 63 26 23 19 15 11Chloramines 6–9 3800 2200 1850 1500 1100 750

Source: Malcolm Pirnie and HDR Engineering, 1991.

1. First, the t10 must be computed. Since the contactor is an axial dispersion contac-tor, equation 14.42 can be used. We therefore need to solve the following equa-tion for the time at which 10 percent of the fluid has exited:

0.1 = Φ �� − 1� � + exp � � Φ �− � + 1� �This equation must be solved by trial and error—however, this can readily bedone on a spreadsheet (in which the cumulative normal distribution function isavailable). Doing this, it is found that t10 = 25.3 min.

2. We now need to examine the last column of Table 14.5 to determine what chlo-rine residual would satisfy the requirement at this t10. The following table illus-trates the computation:

C CT t = (CT/C)

2.4 65 27.082.6 66 25.382.8 67 23.93

3. Finally, an inverse interpolation is needed to obtain the C that yields the desiredt, which is performed as follows:

C = 2.6 + (2.8 − 2.6) � �= 2.61 mg/L

Therefore, the final residual chlorine should exceed 2.61 mg/L.

25.3 − 25.38��23.93 − 25.38

30��t10(0.0392)

t10�30

2�(0.0392)

30��t10(0.0392)

t10�30

DISINFECTION 14.31

TABLE 14.5 CT Values for 99.9 Percent Reduc-tion of Giardia lamblia with Free Chlorine UnderSelected Conditions

Temperature, °C 5 5 20pH 7 9 7

C (mg/L)0.4 139 279 520.6 143 291 540.8 146 301 551 149 312 561.2 152 320 571.4 155 329 581.6 158 337 591.8 162 345 612 165 353 622.2 169 361 632.4 172 368 652.6 175 375 662.8 178 382 673 182 389 68

Source: Malcolm Pirnie and HDR Engineering, 1991.

UV PROCESSES

In the application of kinetic models to the analysis of UV disinfection processes, theconcentration of disinfectant is replaced by the incident light intensity (in units ofenergy per unit area). This may be determined either by direct measurement in theactual disinfection reactor or by modeling of physical aspects of light transmission(Stover et al., 1986; Water Pollution Control Federation, 1984). The potentialabsorbance of effective light by dissolved components and scattering by suspendedsolids may both contribute to demand for disinfectant.

MODE OF ACTION OF DISINFECTANTS

Chlorine

Since Nissen (1890), free chlorine at low pH has been known to be more biocidal thanfree chlorine at high pH. Holwerda (1928) proposed that hypochlorous acid was thespecific agent responsible for inactivation, and thus the pH effect. Fair et al. (1948)determined that the pH dependency of free chlorine potency correlated quantita-tively with the dissociation constant of hypochlorous acid. Chang (1944) determinedthat the association of chlorine with cysts of Entamoeba hystolytica was greater at lowpH than at high pH. Friberg (1957), Friberg and Hammarstrom (1956), and Haas andEngelbrecht (1980a), using radioactive free chlorine, found similar results appliedwith respect to bacteria, and also found that the microbial binding of chlorine couldbe described by typical chemical isotherms.With respect to viruses, this association ofchlorine parallels the biocidal efficacy of hypochlorous acid, hypochlorite, andmonochloramine (Dennis et al., 1979a, b; Olivieri et al., 1980).

Once taken into the environment of the living organism, chlorine may enter intoa number of reactions with critical components causing inactivation. In bacteria, res-piratory, transport, and nucleic acid activity are all adversely affected (Haas andEngelbrecht, 1980a, b; Venkobachar et al., 1975, 1977). In bacteriophage f2, themode of inactivation appears to be disruption of the viral nucleic acid (Dennis et al.,1979b). With poliovirus, however, the protein coat, and not the nucleic acid, appearsto be the critical site for inactivation by free chlorine (Fujioka et al., 1985; Tenno etal., 1980).

The rate of inactivation of bacteria by monochloramine is greater than could beattributed to the equilibrium-free chlorine present in solution. This argues stronglyfor a direct inactivation reaction of combined chlorine (Haas and Karra, 1984).Although organic chloramines are generally measured as combined or total chlorineby conventional methods, they are of substantially lower effectiveness as disinfec-tants than inorganic chloramines (Feng, 1966; Wolfe and Olson, 1985).

In general, the rate of inactivation of microorganisms by various disinfectantsincreases with increasing temperature. This may be characterized by an activationenergy, a Q10 (factor of increase for every 10°C temperature increase), or a temper-ature multiplier.

Surprisingly, Scarpino et al. (1972) reported that viruses were more sensitive tofree chlorine at high pH than at low pH. A variety of subsequent authors confirmedthese findings with viruses and with bacteria (Berg et al., 1989; Haas, 1981; Haas etal., 1986, 1990). Hypochlorite can form neutral ion pairs with sodium, potassium, andlithium, and (particularly at ionic strengths approaching 0.1 M) these can increasedisinfection efficiency by free chlorine at high pH (Haas et al., 1986). More recently,


calcium enhancement of chlorine inactivation of coliforms has been reported (Haasand Anotai, 1996).

Chlorine Dioxide

The dependence of inactivation efficiency on pH is weaker for chlorine dioxide thanfor chlorine, and more inconsistent. Benarde et al. (1965), working with E. coli, andScarpino et al. (1979), working with Poliovirus 1, found that the degree of inactiva-tion by chlorine dioxide increases as pH increases. However, for amoebic cysts, aspH increases, the efficiency of inactivation by chlorine dioxide decreases (Chen etal., 1985).The physiological mode of inactivation of bacteria by chlorine dioxide hasbeen attributed to a disruption of protein synthesis (Benarde et al., 1967). In the caseof viruses, chlorine dioxide preferentially inactivated capsid functions, rather thannucleic acids (Noss et al., 1985; Olivieri et al., 1985).

Benarde et al. (1967) computed the activation energy for the inactivation of E. coli by chlorine dioxide at pH 6.5 as 12 kcal/M.An identical number was computedfor the disinfection of Poliovirus 1 by chlorine dioxide at pH 7 (Scarpino et al., 1979).

Ozone

Understanding of the mode of inactivation of microorganisms by ozone remainshindered by difficulties in measuring low concentrations of dissolved ozone. Theeffect of pH on ozone inactivation of microorganisms appears to be predominantlyassociated with changing the stability of residual ozone, although additional work isneeded. Farooq (1976) found little effect of pH on the ability of dissolved ozoneresiduals to inactivate acid-fast bacteria. Roy (1979) found a slight diminution of thevirucidal efficacy of ozone residuals as pH decreased; however, Vaughn, as cited inHoff (1986) and Hoff and Akin (1986), noted the opposite effect. The principalaction of ozone as a disinfectant occurs via the dissolved ozone, rather than physicalcontact with ozone gas bubbles (Dahi and Lund, 1980; Farooq, 1976).

Bacterial cells lacking certain DNA polymerase gene activity were found to bemore sensitive to inactivation by ozone than wild-type strains, strongly implicatingphysicochemical damage to DNA as a mechanism of inactivation by ozone(Hamelin and Chung, 1978). For poliovirus, the primary mode of inactivation byozone also appears to be nucleic acid damage (Roy et al., 1981b).

Activation energies for ozone inactivation of Giardia and Naegleria cysts werereported by Wickramanayake et al. (1985).At pH 7, the activation energies were 9.7and 16.7 kcal/M. For Poliovirus 1, Roy et al. (1981a) estimated an activation energyof 3.6 kcal/M at pH 7.2. If the latter is correct, its low value suggests that ozone inac-tivates virus by a diffusional rather than a reaction-controlled process.

UV Light

The mode of inactivation of microorganisms by ultraviolet radiation is quite wellcharacterized. Specific deleterious changes in nucleic acid arise upon exposure toUV radiation (Jagger, 1967).These may be repaired by light-activated as well as darkrepair enzymes in vegetative microorganisms. The phenomenon of photoreactiva-tion of UV-disinfected microorganisms has been demonstrated in municipal efflu-ents (Scheible and Bassell, 1981). However, the operation of these repair processes

DISINFECTION 14.33

in microorganisms discharged to actual distribution systems or receiving waters isnot clear.

Severin et al. (1983b) have shown that the series event model for inactivationdescribes the kinetics of UV disinfection quite well. Kinetic parameters for inactiva-tion are shown in Table 14.6.

Activation energies for UV inactivation using this model were lower than forchemical disinfection (indicating the relative insensitivity to temperature) (Severinet al., 1983b).These are considerably lower than the activation energies for chemicaldisinfectants in accord with the apparent (purely physical) mechanism of UV inacti-vation.This also indicates, as a practical matter, that the effect of temperature on UVperformance is much less than for chemical agents such as ozone or chlorine.

The pH dependency of UV inactivation has not been characterized in controlledsystems. However, since the mechanism of UV inactivation appears to be purelyphysical, it is not anticipated that pH would dramatically alter the efficiency of UVdisinfection. Insofar as pH may affect the light absorption characteristics of humicmaterials, an indirect effect of pH (by changing the extent of demand) on inactiva-tion efficiency may exist.

Influence of Physical Factors on Disinfection Efficiency

The apparent increase in microbial resistance by clumping has already been discussed.

Solids Association. Microorganisms can also be partially protected against theaction of disinfectants by adsorption to or enmeshment in nonviable solid particlespresent in a water. Stagg et al. (1978) and Hejkal et al. (1979) found that fecal mate-rial protected poliovirus against inactivation by combined chlorine. Boardman andSproul (1977) found that kaolinite, alum flocs, and lime sludge increased the resis-tance of Poliovirus 1 to free chlorine in demand free systems.

For chlorine dioxide, bentonite turbidity protects poliovirus against the action ofchlorine dioxide (Brigano et al., 1978). Although there was some protection fromozone inactivation afforded to coliform bacteria and viruses by fecal matter and bycell debris, at ozone doses usually employed in disinfection, it was still possible toachieve more than 99.9 percent inactivation within 30 s (Sproul et al., 1978).

Influence of Physiological Factors on Disinfection Efficiency

The physiological state of microorganisms, especially vegetative bacteria, may influ-ence their susceptibility to disinfectants. Milbauer and Grossowicz (1959b) foundthat coliforms grown under minimal conditions were more resistant than cells grownunder enriched conditions. Similarly, Berg et al. (1985) found that chemostat-grown


TABLE 14.6 Kinetics of UV Inactivation

Activation energy k (cm2/(mW-s)) Number of events (kcal/mole)

E. coli 1.538 9 0.554Candida parapsilosis 0.891 15 0.562f2 virus 0.0724 1 1.023

Source: Severin et al., 1983.

cells produced at high growth rates were more sensitive to disinfectants than cellsharvested from low growth rates. A simple subculturing of aquatic strains ofFlavobacterium has been found to increase sensitivity to chlorine disinfection(Wolfe and Olson, 1985).

Postexposure conditions can also influence apparent microbial response to disin-fectants (Milbauer and Grossowicz, 1959a). In three New England water treatmentplants and distribution systems sampled for total coliforms using both standard MFtechniques and media to recover sublethally injured organisms (m-T7 agar medium),from 8 to 38 times as many coliforms were recovered on the latter medium than onthe former medium (McFeters et al., 1986). This sublethal injury does not adverselyaffect pathogenicity to mice (Singh et al., 1986).

With viruses, the phenomenon of multiplicity reactivation can occur when indi-vidual viruses inactivated by different specific events are combined in a single hostcell to produce a competent and infectious unit. This has been demonstrated tooccur in enteric viruses inactivated by chlorine (Young and Sharp, 1979).

The survivors of disinfection can exhibit inheritable increased resistance to sub-sequent exposure. This was first demonstrated for poliovirus exposed to chlorine(Bates et al., 1978). However, demonstration of this phenomenon in bacteria has notbeen consistent (Haas and Morrison, 1981; Leyval, 1984).

DISINFECTANT RESIDUALS FORPOSTTREATMENT PROTECTION

One factor that may be important in evaluating the relative merits of alternative dis-infectants is their ability to maintain microbial quality in a water distribution system.With respect to chlorine, it has been suggested that free chlorine residuals may serveto protect the distribution system against regrowth, or at least as a sentinel for thepresence of contamination (Snead et al., 1980). However, other studies have notedthe lack of correlation between distribution system water quality and the form orconcentration of chlorine residual (Haas et al., 1983b). Similarly, LeChevallier et al.(1990) have reported that microbial slimes grown in tap water may be more sensitiveto inactivation by combined chlorine than to free chlorine transported by the over-lying water. It must be recognized that, regardless of the disinfectant chosen, thewater distribution system can never be regarded as biologically sterile. As shown byMeans et al. (1986a), shifts in the dominant form of disinfectant (e.g., from free chlo-rine to monochloramine) can result in shifts in the taxonomic distribution ofmicroorganisms that inhabit the distribution system.

It must be noted that there is a difference in practice between U.S. systems andmany systems in Europe (Haas, 1999; Hydes, 1999; Trussell, 1999). In the UnitedStates, most utilities strive to maintain minimum chlorine residuals in the distribu-tion system (and in fact there is a strong incentive to do so under regulatory require-ments), with minimum residuals generally exceeding 0.2 mg/L. In a number ofcountries in Europe, the philosophical approach to maintaining distribution systemwater quality relies on nutrient control (principally degradable organic matter)rather than on disinfectant residuals.

The ability of chlorine dioxide residuals to maintain distribution system micro-bial water quality has not been well studied.With respect to both ozone and UV, theabsence of a residual may necessitate the addition of a second disinfectant if a resid-ual in the distribution system is desired. For further discussion on this point, seeChapter 18.

DISINFECTION 14.35

APPLICATION OF TECHNOLOGIES

With the increasing and conflicting objectives that must be met by disinfection pro-cesses, water utilities are moving toward the use of multiple disinfectants. In general,disinfectants may be applied at three locations within treatment. Application of adisinfectant prior to coagulation is generally termed preoxidation or predisinfection.While some inactivation may occur due to predisinfection, generally this is minordue to substantial disinfectant demand. Application of a disinfectant subsequent tosedimentation, either before or after filtration, but prior to a contactor, clear well, orhydraulic device with substantial contact, is termed primary disinfection. The major-ity of inactivation by disinfection processes is expected to occur due to this applica-tion. If another application occurs following primary disinfection (or if ammonia isadded to convert a free residual to a combined residual), this is termed secondarydisinfection. Generally the objective of secondary disinfection is to allow the pene-tration of an active disinfectant residual into the distribution system.

Chlorination

Chlorine may be obtained for disinfection in three forms, as well as generated on-site. For very small water treatment plants, solid calcium hypochlorite (Ca(OCl)2)can be used.This can be applied as a dry powder, or in proprietary tablet dispensers.Calcium hypochlorite is more expensive than the other chemical forms, and particu-larly in hard waters, its use can lead to scale formation.

Generally, on a per unit mass basis of active chlorine, the least expensive form atlarge usage rates is liquified chlorine gas. The use of liquified chlorine gas carrieswith it certain risks associated with accidental leakage of the gas. As a result, a num-ber of utilities have elected to use the somewhat more expensive sodium hypochlo-rite (NaOCl) as a source of disinfectant.

Upon addition to a water, chlorine gas will reduce the pH and alkalinity, whilesodium hypochlorite will raise the pH and alkalinity. In a poorly buffered water, theaddition of a pH control agent may thus be necessary to control the distribution sys-tem water aggressiveness.

Chlorine and hypochlorites have been produced from the electrolysis of brinesand saline solutions since the early 20th century (Rideal, 1908). This remains anattractive option for remote treatment plants near a cheap source of brine.The basicprinciple is the use of a direct current electrical field to effect the oxidation of chlo-ride ion with the simultaneous and physically separated reduction usually of waterto gaseous hydrogen.

In actual practice, it is necessary to operate electrolytic chlorine-generating unitsat voltages as high as 3.85 volts in order to provide reasonable rates of generation.At these overvoltages, however, additional oxidations such as chlorate formation,ohmic heating, and incomplete separation of hydrogen from oxidized products withsubsequent dissipative reaction combine to produce system inefficiencies. For typi-cal electrolytic generating units, current efficiencies of 97 percent may be obtainedalong with energy efficiencies of 58 percent (Downs and Adams, 1973). These effi-ciencies are related to the physical configuration of the electrolysis cells, brine con-centration, and desired degree of conversion to available chlorine (Bennett, 1978;Michalek and Leitz, 1972).

Chlorine can be applied at a variety of points within treatment. Table 14.7 fromthe 1989 AWWA Committee Survey illustrates the frequency with which chlorine isapplied at various locations.


Source Water Chlorination/Preoxidation. Prechlorination, the addition of chlo-rine at an early point within treatment, is designed to minimize operational prob-lems associated with biological slime formation on filters, pipes, and tanks, and alsorelease of potential taste and odor problems from such slimes. In addition, prechlo-rination can be used for the oxidation of hydrogen sulfide or reduced iron and man-ganese. Probably the most common point of addition of chlorine for prechlorinationis the rapid mix basin (where flocculant is added).

However, due to present concerns for minimizing the formation of chlorine by-products, the use of prechlorination is being supplanted by the use of other chemicaloxidants (e.g., ozone, permanganate) for the control of biological fouling, odor, orreduced iron or manganese.

Postchlorination. Postchlorination, or terminal disinfection, is the primary appli-cation for microbial reduction. It has been most common to add chlorine for thesepurposes either immediately before the clear well or immediately before the sandfilter. In the latter case, the filter itself serves, in effect, as a contact chamber for dis-infection.

In general, the use of specific contact chambers subsequent to the addition of chlo-rine to a water has been uncommon. Instead, the clear well, or finished water reservoir,serves the dual function of providing contact to ensure adequate time for microbialinactivation prior to distribution. The distribution system itself, from the entry pointuntil the first consumer’s tap, provides additional contact time (see Table 14.8).

The hydraulic characteristics of most finished water reservoirs, however, are notcompatible with the ideal characteristics of chlorine contact chambers. The latterare most desirably plug flow, while the former most usually have a large degree ofdispersion.

DISINFECTION 14.37

TABLE 14.7 Points of Application of Chlorine or Hypochlorite

Point of application Frequency of plants (N = 268)

Before coagulation 18.66%After coagulation 5.97%After sedimentation and before filtration 31.34%After filtration 47.01%Within the distribution system 15.67%

Percentages sum to more than 100 percent due to multiple points of application.Source: Haas et al., 1992.

TABLE 14.8 Residual and Contact Time to First Customer: 1989 AWWA Disinfection Committee Survey of Utilities (N = 178)

Contact time to first customer, min

Residual, mg/L 0 1–9 10–29 30–60 75–240 >240 Total

0–0.35 4.49% 0.56% 0.56% 0.56% 1.12% 0.56% 7.87%0.4–0.95 10.11% 3.93% 3.93% 3.93% 3.37% 2.81% 28.09%1–1.5 14.61% 5.62% 2.25% 4.49% 1.69% 5.62% 34.27%1.6–2 3.93% 1.69% 1.69% 3.93% 3.37% 1.69% 16.29%>2 6.74% 0.56% 0.56% 0.56% 2.81% 2.25% 13.48%Total 39.89% 12.36% 8.99% 13.48% 12.36% 12.92% 100.00%

Superchlorination/Dechlorination. In the process of superchlorination/dechlori-nation, which has generally been employed for treatment of a poor-quality water(with high ammonia nitrogen concentrations, or perhaps severe taste and odor prob-lems), chlorine is added beyond the breakpoint.This oxidizes the ammonia nitrogenpresent. Generally, the residual chlorine obtained at this point is higher than may bedesired for distribution. The chlorine residual may be decreased by the applicationof a dechlorinating agent (sulfur compounds or activated carbon).

A modern application of chlorination-dechlorination in water treatment may bejudicious where both high degrees of microbial inactivation and low levels of by-product formation are desired. It may be possible, for example, to hold a water withfree chlorine for a period (sufficient to ensure disinfection, but not so long as to pro-duce substantial by-products), then to partially (or completely) dechlorinate thewater to minimize the production of organic by-products.

Chloramination. Chloramination, the simultaneous application of chlorine andammonia or the application of ammonia prior to the application of chlorine, result-ing in a stable combined residual, has been a long-standing practice at many utilities(Table 14.9). As noted in Table 14.1, approximately 20 percent of U.S. utilities useammonia addition in conjunction with chlorine or hypochlorite.

Jefferson Parish, Louisiana, has been using simultaneous addition of ammoniaand chlorine in its disinfection process for over 35 years (Brodtmann and Russo,1979). To further reduce total trihalomethane formation, the point of disinfectantaddition was changed to immediately upstream of the filters (rather than upstreamof clarifiers). This change maintained satisfactory distribution system water quality,using a chloramine residual of 1.6 mg/L exiting the filters.

From a survey of utilities (Trussell and Kreft, 1984), it was found that:

1. Seventy percent use anhydrous ammonia, 20 percent aquo ammonia, 10 percentammonium sulfate.


TABLE 14.9 Utilities with Long Experienceof Chloramine Use

Approximate start of City chloramination

Denver 1914Portland (OR) 1924St. Louis 1934Boston 1944Indianapolis 1954Minneapolis 1954Dallas 1959Kansas City (MO) 1964Milwaukee 1964Jefferson Parish (LA) 1964Philadelphia 1969Houston 1982Miami (FL) 1982Orleans Parish (LA) 1982San Diego 1982

Source: Trussell and Kreft, 1984.

2. Most utilities use 3:1 to 4:1 chlorine-to-ammonia feed ratios. Excess ammonia isgenerally used to make monochloramine predominant; however, some utilitiesuse higher ratios to form more effective dichloramine.

3. There is no clear consensus as to the point of ammonia application (i.e., pre- orpostammoniation).

In current practice, chloramination has been regaining popularity as a means tominimize organic by-product formation.

A major concern that has been identified in chloramination arises during the tran-sition from free chlorination to chloramination (Means et al., 1986a). Data on distri-bution system water quality were collected at the Metropolitan Water District ofSouthern California before and after the distribution system was changed from a freechlorine residual to a combined (monochloramine) chlorine residual. There was noeffect on coliform counts; however, the plate counts on m-R2A medium increased dra-matically (with some increase also observed using pour plates/TGEA medium). Inone of the reservoirs, a precipitous drop in chlorine residual, associated with nitrifica-tion in the reservoir and growth of microorganisms, occurred following the switchover.This was postulated to occur as a reaction between nitrites and monochloramine.

It is important for water utilities to alert hospitals and kidney dialysis centers toswitchovers from free residual chlorination to chloramination. Birrell et al. (1978)and Eaton et al. (1973) reported cases of chloramine-induced hemolytic anemia insuch centers.

Influence of Relative Point ofAddition of Chlorine andAmmonia. Details of chlor-amination practice can dramatically influence process performance. Options as to pre-versus postammoniation or pH, in particular, must be considered.

Prereacted chloramine residuals are more effective bactericides (E. coli, demandfree batch tests) at pH 6 than at pH 8, and at high Cl2-to-N ratios (5:1) rather thanlow ratios (down to 2:1). Concurrent addition of ammonia and chlorine was as effec-tive as preammoniation (and at pH 6 was nearly as effective as free residual chlori-nation). Both concurrent addition and preammoniation were more effective thanprereaction of the chlorine, except at pH 8, where all three modes behaved in a sim-ilar manner (Ward et al., 1984).

In pilot plant studies, Means et al. (1986b) found that concurrent and sequentialmethods (chlorine at rapid mix and ammonia at end of flocculation tank) gave bet-ter performance at removing m-SPC bacteria than preammoniation (but poorerthan free chlorination). Concurrent addition gave about as low a TTHM value aspreammoniation.

In the presence of concentrations of organic nitrogen similar to ammonia, pre-reacted chloramines may give better performance than dynamically formed chlor-amines from preammoniation due to the favorable competition for chlorine bymany organic N compounds (and their low biocidal potency) relative to inorganicnitrogen (Wolfe and Olson, 1985).

Characterizing and Improving Contact Tank Hydraulics. The presence of back-mixing—deviations from ideal plug flow behavior—will reduce the disinfection effi-ciency of chlorine contactors. In serpentine contactors, the degree of backmixing(i.e., the dispersion) can be estimated from the geometry of the tank, in particularthe length-to-width ratio (Stover et al., 1986).

When a finished water reservoir is used as a contact tank, severe backmixing andalso stratification may occur (Boulos et al., 1996; Grayman et al., 1996). Varioustypes of baffles may be used to counteract these tendencies, as shown in Figure 14.8(Grayman et al., 1996). In addition, both hydraulic scale models and computational

DISINFECTION 14.39

fluid dynamic models may be useful in assessing potential improvements in flow pat-terns from alternative baffling arrangements (Boulos et al., 1996; Grayman et al.,1996;Taras, 1980). Similarly, chlorine contact chambers may be upgraded in terms ofhydraulic performance (increasing t10 and decreasing dispersion) by the installationof baffles (Bishop et al., 1993).


FIGURE 14.8 Schematic of baffling arrangements for reservoirs. (Source:Grayman et al., 1996.)

Chlorine Dioxide

It is necessary to generate chlorine dioxide on a continuous basis for use as a disinfec-tant.Although a few European potable water treatment plants have been reported touse the acid-chlorite generation process (Miller et al., 1978), the most common syn-thesis route for disinfectant ClO2 generation is the chlorine-chlorite process.

Chemistry of Generation. Theoretically, chlorine dioxide may be produced byeither the oxidation of a lower-valence compound or reduction of a more oxidizedcompound of chlorine. Chlorites (ClO2

−) or chlorous acid (HClO2) may be oxidizedby chlorine or persulfate to chlorine dioxide, or may undergo autooxidation (dis-proportionation) to chlorine dioxide in solutions acidified with either mineral ororganic acids. Chlorates (ClO3

−) may be reduced either by use of chlorides, sulfuricacid, sulfur dioxide, or oxalic acid, or electrochemically to form chlorine dioxide(Masschelein, 1979a).

For practical purposes in water treatment, chlorine dioxide is generated exclusivelyfrom chlorite inasmuch as the reductive processes using chlorate as a starting materialare capital intensive and competitive only at larger capacities (Masschelein, 1979b).

In the acid-chlorite process, sodium chlorite and hydrochloric acid react accord-ing to Equation 14.43:

5 NaClO2 + 4 HCl = 4 ClO2 + 5 NaCl + 2 H2O (14.43)

The resulting chlorine dioxide may be evolved as a gas or removed in solution.Mechanistically, this process occurs via a series of coupled reactions, some of whichmay involve the in situ formation of chlorine, catalysis by chloride, and the oxidationof chlorite by chlorine (Gordon et al., 1972; Masschelein, 1979a; Noack and Doeff,1979). In addition, the yield of the reaction as well as the rate of the process areimproved by low pH values in which formation of both gaseous chlorine and chlorousacid is favored. Under these favorable conditions, the reaction proceeds in the order ofminutes; however, to achieve these conditions, excess hydrochloric acid is required.

During the acid-chlorite reaction, the following side reactions result in chlorineproduction:

5 ClO2− + 5 H+ = 3 ClO3

− + Cl2 + 3 H+ + H2O (14.44)

4 ClO2− + 4 H+ = 2 Cl2 + 3 O2 + 2 H2O (14.45)

4 HClO2 = 2 ClO2 + HClO3 + HCl + H2O (14.46)

if equal weights of the reactants are added, than close to 100 percent of the conver-sion of chlorite may occur; this means a final pH below 0.5 (Masschelein, 1979b).

Alternatively, chlorine dioxide may be produced by the oxidation of chlorite withchlorine gas according to Equation 14.47:

2 NaClO2 + Cl2 = NaCl + 2 ClO2 (14.47)

As in the previous case, low pH accelerates the rate of this process, as does excesschlorine gas. However, if chlorine gas is used in stoichiometric excess, the resultantproduct may contain a mixture of unconsumed chlorine as well as chlorine dioxide.

The rate of the direct reaction between dissolved Cl2 and chlorite has been mea-sured (Aieta and Roberts, 1985), with a forward second-order rate constant given bythe following:

kf = 1.31 × 1011 exp(−4800/T) 1/M-s (14.48)

DISINFECTION 14.41

In the chlorine-chlorite process, sodium chlorite is supplied as either a solid pow-der or a concentrated solution. A solution of chlorine gas in water is produced by achlorinator-ejector system of design similar to that used in chlorination.The chlorine-water solution and a solution of sodium chlorite are simultaneously fed into a reactorvessel packed with Raschig rings to promote mixing (Miller et al., 1978). Equation14.46 shows that 1 mole of chlorine is required for 2 moles of sodium chlorite—or0.78 parts of Cl2 per part of NaClO2 by weight. However, for this reaction to proceedto completion, it is necessary to reduce the pH below that provided by the typicallyacidic chlorine-water solution produced by an ejector. At 1:1 feed ratios by weight,only 60 percent of the chlorite typically reacts (Miller et al., 1978).

To provide greater yields, several options exist. First, it is possible to producechlorine-water solutions in excess of 3500 mg/L using pressurized injection of gas. Inthis case, however, there will be an excess of unreacted chlorine in the product solu-tion, and the resulting disinfectant will consist of a mixture of chlorine and chlorinedioxide. The second option consists of addition of acid to the chlorine-chlorite solu-tion. For example, a 0.1 M HCl/M chloride addition enabled the production of a dis-infectant solution of 95 percent purity in terms of chlorine dioxide, and achieved a90 percent conversion of chlorite to chlorine dioxide (Jordan, 1980).A third process,developed by CIFEC (Paris, France), involves recirculation of the chlorinator ejec-tor discharge water back to the ejector inlet to produce a strong (5000–6000 mg/L)chlorine solution, typically at pH below 3.0, and in this manner to increase the effi-ciency of chlorite conversion (Miller et al., 1978). It has been reported that this lastoption is capable of producing 95 to 99 percent pure solutions of chlorine dioxide.The intricacies of chlorine dioxide reactions and by-products necessitate carefulprocess monitoring during operation of the generator (Lauer et al., 1986).

EXAMPLE 14.6 A water utility has a chlorination capacity of 1000 tons/day (454kg/day) and is considering a switchover to chlorine dioxide to be generated using thechlorine-chlorite process. If the existing chlorination equipment is to be used, whatis the maximum production capacity of chlorine dioxide, and how much sodiumchlorite must be used under these conditions? Assume ideal stoichiometry and noexcess chlorite or chlorine requirements.

1. From Equation 14.47, 2 moles of sodium chlorite (NaClO2) react with 1 mole ofchlorine to produce 2 moles of chlorine dioxide. Because chlorine has a molecularweight of 70, the current chlorinators have a capacity of 454,000/70 = 6486 mol/dayof chlorine. Therefore, 12,971 mol/day of sodium chlorite are required, and theresult would be an equal number of moles of chlorine dioxide. The molecularweights are:

Chlorine dioxide 35 + 2(16) = 67Sodium chlorite 23 + 35 + 2(16) = 89

Therefore, the sodium chlorite required is 12,971(89) = 1.15 × 106 g/day (2,540lb/day). The chlorine dioxide produced would be 12,971(67) = 0.87 × 106 g/day(1,914 lb/day).

Application. The use of chlorine dioxide is limited by two factors. First, the maxi-mum residual that does not cause adverse taste and odor problems is 0.4 to 0.5 mg/las ClO2 (Masschelein, 1979b). Second, the chlorite produced by reduction of chlorinedioxide as demand is exerted has been found to cause certain types of anemia, andtherefore the maximum chlorine dioxide dose must be 1 mg/l to minimize this effect.

Augenstein (1974) suggested that chlorine dioxide residuals have moderate sta-bility during distribution. However, it is not clear whether the analytical methods


used in that study could adequately differentiate chlorine dioxide from its reactionproducts.

Insofar as chlorine dioxide is produced free of chlorine (in the acid-chlorite pro-cess, or in “optimized” chlorine-chlorite processes), the reactions with organic mate-rial to produce chlorinated by-products appear less significant than with chlorine(Aieta and Berg, 1986; Lykins and Griese, 1986).

Ozonation

Generation. The use of ozone in water treatment in the United States has beenconfined to preoxidation (Glaze, 1987). Additional information on the use of ozoneas an oxidant is provided in Chapter 12.

Due to its instability, ozone is produced from gas-phase electrolytic oxidation ofoxygen, either using very dry air or pure oxygen. The ozone-enriched gaseous phaseis then contacted with the water to be treated in a bubble contactor (either diffusedair or turbine mixed) or in a countercurrent tower contactor. Due to the cost ofozone, it is highly desirable to maximize the efficiency of transfer of ozone from gasto liquid.

Most ozone generators used in water treatment use one of two designs (Glaze,1987).The most common for large plants is a bank of glass tube generators as shownin Figure 14.9. Small plants may use this type of generator on a smaller scale, or aplate-type generator in which the ozone is generated between ceramic plates. Cool-ing the tubes or plates increases the efficiency of ozone production. Ozone genera-tors may use pure oxygen, oxygen-enriched air, or air as the feed gas. If air is used,the most economical operation gives a product stream that contains about 2 percentozone by weight. Enhancement of the amount of oxygen in the stream increases theeconomical yield of ozone; for example, pure oxygen can generate a stream contain-ing 5 to 7 percent ozone economically. In any case, the gas stream must have a very

DISINFECTION 14.43

FIGURE 14.9 Large-scale tube-type generator for production of ozone from air or oxy-gen by cold plasma discharge.

low dew point (−50°C) and must be free of organic vapors. Figure 14.10 shows a flowdiagram for a plant that utilizes oxygen enriched air.

Ozone generator designs that utilize different electrode configurations, such assurface discharge models, are also available. Also, ozone may be generated by irra-diation of air with high-energy ultraviolet radiation (with wavelengths less than 200 nm). Photochemical generators are not yet capable of producing as much ozoneas plasma generators, but may be particularly useful for small-scale applicationssuch as swimming pool disinfection.

Contactors. After generation, ozone is piped to a contactor where it is transferredinto the water. The most common type of contactor is the countercurrent spargedtank with diffuser (Figure 14.11). In this reactor, ozone-containing gas forms smallbubbles as it is passed through a porous stone at the bottom of the tank.As the bub-bles rise through the tank, ozone is transferred from the gas phase into water accord-ing to the rate equation

Rate of transfer [mol/(m3)(s)] = KLa (C* − C) (14.49)

where C (mol/m3) is the prevailing concentration of ozone in the liquid, C* is theconcentration at saturation, and KLa is the overall transfer coefficient (s−1). Thevalue of C* depends on the percentage of ozone in the gas and may be calculatedfrom the equation

C* = (14.50)Pgas�H


FIGURE 14.10 Flow diagrams for air and oxygen purification for ozone production from oxygen-enriched air. The air purification unit may be omitted when pure oxygen is used, or it may be usedwithout oxygen enrichment.

where Pgas is the partial pressure of ozone in the gas phase (atm) and H is theHenry’s law constant for ozone (0.082 atm-m3/mol at 25°C). At a partial pressure ofozone of 0.02 atm (corresponding to 3 weight percent of ozone in air), Equation14.36 predicts that the concentration of ozone in water at saturation will be 0.24mol/m3 (12 mg/L). More details on theory and application of liquid to gas transfermay be found in Chapter 5.

Ozone contactors such as that shown in Figure 14.11 may be sized to transferozone on a small or very large scale. Alternately, other designs such as spargedstirred tank reactors and venturi injectors may be used. If injected under a largehead of water pressure, ozone may be transferred at higher rates because the satu-ration concentration of ozone, the C* term in Equation 14.49, is higher.

Ozone contactors must have systems for the collection of ozone off-gas. Ozone istoxic and must be kept within OSHA allowable limits (ACGIH, 1994) within thetreatment plant and surrounding areas. In some regions of the United States, ozonedischarge from the treatment plants may be regulated. As a consequence of theserequirements, thermal and catalytic ozone destroyers are routinely used in plantsemploying ozonation.

Ozone in water is so highly corrosive that only certain construction materialsmay be used in plants that generate and utilize ozone for water treatment. Metal incontact with ozone should be 304 stainless steel, and gasket materials should besome form of inert polymer such as a fluorocarbon. Concrete is a typical material forconstruction of basins, but joints must be caulked with an inert material.

Unfortunately, the characteristics that promote efficient gas-liquid mass transfer,particularly the desirability of intense agitation, lead to hydraulic characteristics thatdecrease disinfection efficiency (short-circuiting). Therefore, laboratory results on

DISINFECTION 14.45

FIGURE 14.11 Diagram of typical countercurrentsparged column for transfer of ozone from gas to liquidphase.

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(<2/100 ml) were obtained at a dose of 1.29 mg/l (transferred dose of 1.04 mg/l). Aneconomic analysis showed that the GAC-ozone process was cheaper than a processemploying full conventional treatment with chlorination.

The contact time required for ozone inactivation of microorganisms at com-monly employed ozone doses (several mg/L) is typically quite short—seconds toseveral minutes—rather than the characteristically longer disinfection times com-monly used for chlorine or chloramines.

The lack of a persistent residual following ozone treatment has generally neces-sitated application of an additional terminal disinfection process such as chlorami-nation.

Ozone Contactor Hydraulics. The role of hydraulics in ozone contactors has twoinfluences on performance. It influences ozone transfer efficiency, and also, by virtueof the influence of dispersion on performance, it influences the effectiveness withwhich a given combination of residual and time can be effective. The complex inter-actions involved are illustrated in several papers (Lev and Regli, 1992; Martin et al.,1992; Roustan et al., 1991). The designer is faced with a trade-off in which multiplestages, each approximating a completely mixed system, are used to transfer ozone.As the number of stages increases, the overall system hydraulics and also potentiallythe transfer efficiency may improve. However, this also would significantly increaseconstruction and operating costs for the ozonation units.

UV Radiation

UV radiation can be effectively produced by the use of mercury vapor (or morerecently, antimony vapor) lamps. While to date there has been limited experiencewith UV disinfection in water treatment, there is a substantial operational data baseassociated with UV disinfection of wastewater effluents (Stover et al., 1986; WaterPollution Control Federation, 1984).

There are a variety of physical configurations in which UV has been employed.In two of these, the UV lamp(s) are surrounded by quartz sheaths and the jacketedlamps are immersed in the flowing water. The flow may be in a closed or open ves-sel, and may be either parallel or perpendicular to the axes of the lamps. In the thirdconfiguration, the water flows through Teflon tubes (which are relatively transpar-ent to UV radiation) surrounded by UV lamps.

There are a number of newer UV technologies emerging for disinfection appli-cation. These include (O’Brien et al., 1996):

● Use of medium-pressure UV lamps with higher-intensity emissions and potentialfor longer lamp lifetimes

● High-frequency pulsed UV disinfection

For design purposes, it is necessary to ensure that there is turbulence (to allow allelements of fluid to come sufficiently close to the lamp surfaces) and yet to minimizethe degree of transverse mixing (short-circuiting). The latter must be explicitly con-sidered in design calculations (Stover et al., 1986).


The contact times for UV disinfection systems can be quite short, generally under1 min. Therefore, the space required for UV disinfection units is relatively small.However, since there is no residual, some additional terminal disinfection processwould usually be required. The potential for UV reactions to produce organic by-products is minor since the intensities required for UV disinfection are less thanthose needed to cause photochemical effects. Operationally, it is essential to employan effective cleaning program to periodically removal biological and chemical foul-ing materials from lamp jacket or teflon tube surfaces (Stover et al., 1986;Water Pol-lution Control Federation, 1984). In wastewater disinfection systems, lamp sizing istypically (to within a factor of 2) 1 kW per million gallons (3780 m3) peak designflow (Scheible, 1993).

Role of Hydraulics. The role of hydraulics in UV disinfection has two effects onperformance. In addition to a negative influence of backmixing, due to the phe-nomenon of dispersion noted earlier, the movement of microorganisms relative tolamp surfaces may dramatically influence the integrated exposure to radiation dur-ing contacting. The effect of turbulence to decrease the potential for stratification(defined as travel of a given microorganism at a fixed distance from lamp surfaces)is desirable to increase UV disinfection efficiency (Haas and Sakellaropoulous,1979).

There is an interaction between the residence time distribution and the distribu-tion of intensities of exposure that microorganisms experience in transiting a UVdisinfection unit.The use of computational hydraulic models may assist in providinga more precise prediction of the effect of flow rate and water quality parameters(absorbency) on the inactivation expected in full-scale UV units (Chiu et al., 1999;Lyn et al., 1999).

USE OF MULTIPLE DISINFECTANTS

In view of the reexamination of disinfection practices that has occurred over the past20 years, a number of case studies of multiple disinfectants have been described.These might include a relatively reactive primary disinfectant (such as O3 or ClO2)followed by a secondary disinfectant that is used to maintain a chemical residual inthe distribution system.

For example, at the Louisville Water Company (Hubbs et al., 1980), there was aplant-scale examination of a process in which ClO2 is added between the coagula-tion basin effluent and the softening basin influent along with excess ammonia toconvert the free chlorine to combined chlorine (predominantly monochloramine).Some additional chlorine is added after filtration to provide a monochloramineresidual in the distribution system. There was some regrowth in the filters; however,bacterial quality in the distribution system was not altered, and TTHMs werereduced from 30 ppb (free chlorination) to less than 5 ppb.

The use of multiple disinfectants may be advantageous in terms of achieving dis-infection efficiency, since possible synergistic effects may occur—although furtherstudy is necessary to delineate the exact extent of such phenomena. For example, theexposure of E. coli to mixtures of free chlorine and chloramine resulted in substan-tially greater inactivation than would be predicted by their individual effectiveness(Kouame and Haas, 1991). Similarly, the following combinations of disinfectantsmay offer some level of inactivation of the highly resistant Cryptosporidium oocysts(Finch et al., 1998):

DISINFECTION 14.47

● Free chlorine followed by monochloramine● Ozone followed by either free or combined chlorine● Chlorine dioxide followed by either free or combined chlorine

RELATIVE COMPARISONS

The dominant disinfection technology in the United States is presently free residualchlorination. However, chloramination, chlorination with partial dechlorination, andchlorine dioxide treatment are viable alternative primary disinfectants. In addition,ozone, or possibly UV, when supplemented with a chemical that can produce a last-ing residual in the distribution system, have application.

Table 14.10 summarizes important aspects of the major technologies and theirtechnical pros and cons.

DIAGNOSTICS AND TROUBLESHOOTING

Since real-time measurements of microorganisms surviving disinfection processes arenot presently possible, the operation of disinfection processes relies on indirect assess-ment of performance. With chlorine contacting systems, control is most frequentlybased on the use of residual measurements on a batch or continuous basis. A varietyof control schemes are available. The feed of chlorine may be controlled to automati-cally provide a relatively constant residual.The presence of excessive chlorine demandmay be symptomatic of a process upset or malfunction (Stover et al., 1986).

With ozone systems, it is possible to use either off-gas ozone monitors or mea-surement of dissolved residual for control of ozonation. However, the degree of


TABLE 14.10 Applicability of Alternative Disinfection Techniques

Consideration Cl2 O3 ClO2 UV

Equipment reliability Good Good Good Fair to good

Relative complexity Simple Complex Moderate Moderateof technology

Safety concerns Yes Moderate Yes Moderate

Bactericidal Good Good Good Good

Virucidal Moderate Good Moderate Good

Efficacy against protozoa Fair Moderate Fair Fair to moderate

By-products of possible Yes Some Some None knownhealth concern

Persistent residual Long None Moderate None

Reacts with ammonia Yes No No No

pH dependent Yes Slight Slight No

Process control Well developed Developing Developing Developing

Intensiveness of operations Low High Moderate Highand maintenance

automation in control systems is not as readily developed as in the case of chlorina-tion systems.

For UV disinfection, light absorption may be monitored continually, and lampsmay be controlled (by either turning on or off sets of lamps, or varying power) toproduce a given delivered energy intensity. This can also allow for compensation offouling of lamp and jacket surfaces with operations.

BIBLIOGRAPHY

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ACGIH. “Threshold Limit Values for Chemical Substances and Physical Agents and Biologi-cal Exposure Indices.” American Conference of Governmental Industrial Hygienists, Cincin-nati, OH, 1994.

Aieta, E., and J. Berg. “A Review of Chlorine Dioxide in Drinking Water Treatment.” JournalAWWA, 78(6):62, 1986.

Aieta, E., and P. Roberts. “The Chemistry of Oxo-Chlorine Compounds Relevent to ChlorineDioxide Generation,” In Water Chlorination: Environmental Impact and Health Effects, vol. 5(R. Jolley, ed.). Boca Raton, FL: Lewis Publishers, 1985.

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14. disinfection

Documents

axial dispersion model

henrys law constant

finished water reservoir

human enteric viruses

chlorine contact chambers

series event model

free chlorine residuals

chlorine dioxide residuals