The Disinfection of Drinking Water

The goal of disinfection of public water supplies is the elimination of the pathogens that are responsible for waterborne diseases. The transmission of diseases such as typhoid and paratyphoid fevers, cholera, salmonellosis, and shigellosis can be controlled with treatments that substantially reduce the total number of viable microorganisms in the water.

While the concentration of organisms in drinking water after effective disinfection may be exceedingly small, sterilization (i.e., killing all the microbes present) is not attempted. Sterilization is not only impractical, it cannot be maintained in the distribution system. Assessment of the reduction in microbes that is sufficient to protect against the transmission of pathogens in water is discussed below.

Chlorination is the most widely used method for disinfecting water supplies in the United States. The near universal adoption of this method can be attributed to its convenience and to its highly satisfactory performance as a disinfectant, which has been established by decades of use. It has been so successful that freedom from epidemics of waterborne diseases is now virtually taken for granted. As stated in Drinking Water and Health national Academy of Sciences, 1977), "chlorination is the standard of disinfection against which others are compared."

However, the discovery that chlorination can result in the formation of trihalomethanes (THM's)) and other halogenated hydrocarbons has prompted the reexamination of available disinfection methodology to determine alternative agents or procedures (Morris, 1975).

The method of choice for disinfecting water for human consumption depends on a variety of factors (Symons et al., 1977). These include:

its efficacy against waterborne pathogens (bacteria, viruses, protozoa, and helminths);

the accuracy with which the process can be monitored and controlled;

its ability to produce a residual that provides an added measure of protection against possible posttreatment contamination resulting from faults in the distribution system;

the aesthetic quality of the treated water; and

the availability of the technology for the adoption of the method on the scale that is required for public water supplies.

Economic factors will also play a part in the final decision; however, this study is confined to a discussion of the five factors listed above as they apply to various disinfectants.

The propensity of various disinfection methods to produce by-products having effects on health (other than those relating to the control of infectious diseases) and the possibility of eliminating or avoiding these undesirable by-products are also important factors to be weighed when making the final decisions about overall suitability of methods to disinfect drinking water. The subcommittee has not attempted to deal with these problems since the chemistry of disinfectants in water and the toxicology of expected by-products have been studied by other subcommittees of the Safe Drinking Water Committee, whose reports appear in Chapter III of this volume (Chemistry) and Chapter IV toxicity) of Drinking Water and Health, Vol.3.


The general considerations noted in the immediately following material should be borne in mind when considering each method of disinfection. Available information on the obvious major candidates for drinking water disinfection-chlorine, ozone, chlorine dioxide, iodine, and bromine-is then evaluated for each method individually in the following sections. Other less obvious possibilities are also examined to see if they have been overlooked unjustly in previous studies or if it might be profitable to conduct further experimentation on them. Disinfection by chloramines is dealt with in parallel with that effected by chlorine because of the close relationship the former has to chlorine disinfection under conditions that might normally be encountered in drinking water treatment.

The evaluations in this report are not exhaustive literature reviews but, rather, are selections of the studies that, in the judgment of the committee, provide the most accurate and relevant information on the biocidal activities of each method of disinfection. The analytical methods that are described in this report are those that are most likely to be used by persons involved in disinfection research or water treatment. A review of all existing analytical methods, some of which may be more sophisticated than those described below, would be impractical within the constraints of time and space available and is not within the scope of this document.

After the methods of disinfection are examined individually, their major characteristics and biocidal efficacy are compared by means of summary tables and c t (concentration, in milligrams per liter, times contact time, in minutes) values required for similar inactivations under identical conditions. The conclusions of the study are then recorded on the basis of this evidence.


In any comparison of disinfection methods, certain considerations should be discussed at the outset since they are relevant to most, if not all, methods. The quality of the raw water (i.e., its content of solids and material that will react with the disinfectant), treatment of the water prior to disinfection, and the manner in which the disinfectant is applied to the water will directly affect the efficacy of all disinfectants. Equally applicable to all methods are appropriate standards for verifying the adequacy of disinfection, differences in response to disinfectants between organisms that were obtained directly from the field and those that have been acclimated to laboratory culture, and the maintenance of potability from treatment plant to the consumer's tap. The use of chlorination as presented in examples in the following pages does not imply that it is necessarily the method of choice. Rather, this method has been studied more thoroughly than other methods.

Raw Water Quality

In addition to potential pathogens, raw water may contain contaminants that may interfere with the disinfection process or may be undesirable in the finished product. These contaminants include inorganic and organic molecules, particulates, and other organisms, e.g., invertebrates. Variations among these contaminants arise from differences in regional geochemistry and between ground- and surface-water sources.


Many inorganic and organic molecules that occur in raw water exert a "demand," i.e., a capacity to react with and consume the disinfectant. Therefore, higher "demand" waters require a greater dose to achieve a specific concentration of the active species of disinfectant. This demand must be satisfied to ensure adequate biocidal treatment.

Ferrous ions, nitrites, hydrogen sulfide, and various organic molecules exert a demand for oxidizing disinfectants such as chlorine. The bulk of the nonparticulate organic material in raw water occurs as naturally derived humic substances, i.e., humic, fulvic, and hymatomelanic acids, which contribute to color in water. The structure of these molecules is not yet fully understood. However, they are known to be polymeric and to contain aromatic rings and carboxyl, phenolic, alcoholic hydroxyl, and methoxyl functional groups. Humic substances, when reacting with and consuming applied chlorine, produce chloroform (CHCl3) and other THM's. Water, particularly surface waters, may also contain synthetic organic molecules whose demand for disinfectant will be determined by their structure. Ammonia and amines in raw water will react with chlorine to yield chloramines that do have some biocidal activity, unlike most products of these side reactions. If chlorination progresses to the breakpoint, i.e., to a free-chlorine residual, these chloramines will be oxidized causing more added chlorine to be consumed before a specific free-chlorine level is achieved. This phenomenon is discussed more fully below.

The nature of the demand reactions varies with the composition of the water and the disinfectant. Removal of the demand substances leaves a water with a lower requirement for a disinfectant to achieve an equivalent degree of protection against transmission of a waterborne disease.


Various treatments applied to raw water to remedy undesirable characteristics, e.g., color, taste, odor, or turbidity, may affect the ultimate microbiological quality of the finished water. Microorganisms may be physically removed or the disinfectant demand of the water altered.

Presedimentation to remove suspended matter, coagulation with alum or other agents, and filtration reduce the organic material in the raw water and, thus, the disinfectant demand. Removal of ferrous iron similarly reduces the demand for oxidizing disinfectants as will aeration, which eliminates hydrogen sulfide. Prechiorination to a free chlorine residual is practiced early in the treatment sequence as one method to alter taste- and odor-producing compounds, to suppress growth of organisms in the treatment plant, to remove iron and manganese, and to reduce the interference of organic compounds in the coagulation process. The necessity for these treatments or others is determined by the characteristics of the raw water. The selection of one of the various methods to achieve a particular result will be based upon cost-effectiveness in the particular situation. When chlorination is used, the application or point of application in the treatment sequence of some of the above-mentioned procedures can affect the undesirable THM content of the finished water.

Reduction of precursors in raw water by coagulation and settling prior to chlorination reduces final THM production (Hoehn et al., 1977; Stevens et aL, 1975). The Louisville Water Company reduced THM concentrations leaving the plant by 4O%-5O% by shifting the point of chlorination from the presedimentation basin to the coagulation basin (Hubbs et al., 1977). The available information on these variations is limited, and a universally applicable procedure cannot be recommended in view of the diverse treatments required for different raw waters.

Particulates and Aggregates

To inactivate organisms in water, the active chemical species must be able to reach the reactive site within the organism or on its surface. Inactivation will not result if this cannot occur. Microorganisms may acquire physical protection in water as a result of their being adsorbed to the enormous surfaces provided by clays, silt, and organic matter or to the surfaces of solids created during water treatment, e.g., aluminum or ferric hydrated oxides, calcium carbonate, and magnesium hydroxide. Viruses, bacteria, and protozoan cysts may be adsorbed to these surfaces. Such particles, with the adsorbed microorganisms, may aggregate to form clumps, affording additional protection. Organisms themselves may also aggregate or clump together so that organisms that are on the interior of the clump are shielded from the disinfectant and are not inactivated. Organisms may also be physically embedded within particles of fecal material, within larger organisms such as nematodes, or, in the case of viruses, within human body cells that have been discharged in fecal material.

To disinfect water adequately, the water must have been pretreated, when necessary, to reduce the concentration of solid materials to an acceptably low level. The primary drinking water turbidity standard of I nephelometric turbidity unit (NTU) is an attempt to assure that the concentration of particulates is compatible with current disinfection techniques. Where it is possible to obtain lower turbidities, this is desirable.

Disinfection studies in which the complications of adsorbed organisms, aggregation, or embedment were thought to occur were excluded from this study. The conclusions in this report should not be extrapolated to such situations as the disinfection of turbid or colored waters.

The Importance of Residuals

Water supplies are disinfected through the addition or dosage of a chemical or physical agent. With a chemical agent, such as a halogen, a given dosage should theoretically impart a predetermined concentration (residual) of the active agent in the water. From a practical point of view, most natural waters exert a "demand" for the disinfectant, as discussed above, so that the residual in the water is less than the calculated amount based on the dosage. The decrease in residual, which is caused by the demand, is rapid in most cases, but it may be prolonged until the residual eventually disappears. In addition, the chemical agent may decompose spontaneously, thereby yielding substances having little or no disinfection ability and exerting no measurable residual. For example, ozone not only reacts with substances in water that exert a demand, but it also decomposes rapidly. To achieve microbial inactivation with a chemical agent, a residual must be present for a specific time. Thus, the nature and level of the residual, together with time of exposure, are important in achieving disinfection or microbial inactivation. Because the nature of the dosage-residual relationship for natural waters has not been and possibly cannot be reliably defined, the efficacy of disinfection with a chemical agent must be based on a residual concentration/time-of-exposure relationship.

Residual measurements are important and useful in controlling the disinfection process. By knowing the residual-time relationship that is required to inactivate pathogenic or infectious agents, one can adjust the dosage of the disinfecting agent to achieve the residual that is required for effective disinfection with a given contact time. Thus, the effectiveness of the disinfection process can be controlled and/or judged by monitoring or measuring the residual.

Following disinfection of a water supply at a treatment plant, the water is distributed to the consumers. A persistent residual is important for continued protection of the water supply against subsequent contamination in the distribution system. Accidental or mechanical failures in the distribution system may result in the introduction of infectious agents into the water supply. In the presence of a residual, disinfection will continue and, as a result, offer continued protection to the users. Physical agents such as radiation may provide effective disinfection during application, but they do not impart any persistent residual to the water.

The dosage of a chemical agent that is used to effect microbial inactivation should not be so great that it imparts a health hazard to the water consumer. From another point of view, the aesthetic quality of the finished water should not be impaired by the dosage of the chemical agent or the residual that is required for effective disinfection. These qualities might include discoloration of water from potassium permanganate (KMnO4) or iodine or problems of taste and odor from excessive chlorine.

Application of the Disinfectant

Optimum inactivation occurs when the disinfectant is distributed uniformly throughout the water. To disperse the chemical disinfectant when it is added to the water, it must be mixed effectively to assure that all of the water, however small the volume, receives its proportionate share of the chemical. Additions of a disinfectant at points in a flowing water stream, e.g., from submerged pipes. is seldom adequate to assure uniform concentration. In such cases, mechanical mixing devices are needed to disperse the disinfectant throughout the water. Disinfection by radiation treatment also requires good mixing to bring all of the water within the effective radiation distance.

Microbiological Considerations1

Comparison of the biocidal efficacy of disinfectants is complicated by the need to control many variables, a need not realized in some early studies. Halogens in particular are significantly affected by the composition of the test menstruum and its pH. temperature, and halogen demand. For very low concentrations of halogen to be present over a testing period, halogen demand must be carefully eliminated. Different disinfectants may have different biocidal potential. In earlier work,

1Nomenclature in this report follows that recommended in the Eighth Edition of Bergey's Mast of Determinative Bacteriology (Buchanan and Gibbons, 1974). Thus, the name of an organism mentioned in the text may not be that used by the author of the work cited

analytical difficulties may have precluded defining exactly the species present, but new techniques allow the species to be defined for most disinfectants. Information on the species of disinfectant actually in the test system should be included in future reports on disinfection studies.

Investigators studying efficacy have usually adopted one of two extremes. Some have conducted carefully designed laboratory experiments with controls for as many variables as possible. Certain of these investigators have reduced the temperature to slow the inactivation reactions. Although these experiments yield good basic information and can be used to determine which variables are important, they often have little quantitative relationship to field situations. The other extreme, a field study or reconstruction of field conditions, is difficult to control. Moreover, their results are often not repeatable.

In addition to the variables noted above, prereaction of chemicals in the test system, the culture history of the organism being used, and the "cleanup" procedures applied to it may also affect the observed results. Despite these problems, there have been some attempts to standardize efficacy testing.


A major factor that influences the evaluation of the efficacy of a particular disinfectant is the test microorganism. There is a wide variation in susceptibility, not only among bacteria, viruses, and protozoa (cyst stage), but also among genera, species, and strains of the microorganism. It is impractical to obtain information on the inactivation by each disinfectant for each species and strain of pathogenic microorganism of importance in water. In addition, interpretation of the data would be confounded by the condition and source of the test microorganism (e.g., the degree of aggregation and whether the organisms were "naturally occurring" or laboratory preparations), the presence of solids and particulates, and the presence of materials that react with and consume the disinfectant.

The overwhelming majority of the literature on water disinfection concerns the inactivation of model microorganisms rather than the pathogens. These disinfectant model microorganisms have generally been nonpathogenic microorganisms that are as similar as possible to the pathogen and behave in a similar manner when exposed to the disinfectant. The disinfectant model systems are simpler, less fastidious, technically more workable systems that provide a way to obtain basic information concerning fundamental parameters and reactions. The

information gained with the model systems can then be used to design key experiments in the more difficult systems. The disinfection model microorganism should be clearly distinguished from the indicator organism. The indicator microorganism, as defined in Drinking Water and Health national Academy of Sciences, 1977), is a "microorganism whose presence is evidence that pollution (associated with fecal contamination from man or other warm-blooded animals) has occurred." Following are criteria for the indicator microorganism (Fair and Geyer, 1954):

1. The indicator should always be present when fecal material is present and absent in clean, uncontaminated water.

2. The indicator should die away in the natural aquatic environment and respond to treatment processes in a manner that is similar to that of the pathogens of interest.

3. The indicator should be more numerous than the pathogens.

4. The indicator should be easy to isolate, identify, and enumerate.

Only a restrictive application of the second criterion is necessary for a disinfection model. The response of the test microorganism to the disinfectant must be similar to that of the pathogen that it is intended to simulate. The disinfection model is not meant to function as an indicator microorganism.

During the latter part of the nineteenth century, investigators recognized the presence of a group of bacteria that occured in large numbers in feces and wastewater. The most significant member of this group (currently called the coliform group) is Escherichia coli. Since the late nineteenth century, this coliform group has served as an indicator of the degree of fecal contamination of water, and E. coli has been used routinely as a disinfection model for enteric pathogens. Butterfield and co-workers (Butterfield and Wattie, 1946; Butterfield et aL, 1943; Wattie and Butterfield, 1944) provided information on the inactivation of E. coli and other enteric bacterial pathogens with chlorine and chloramines. At pH values above 8.5, all strains of E. coli were more resistant to free chlorine than were Salmonella typhi strains. At pH values of 6.5 and 7.0, strains of S typhi were more resistant. Only slight differences between the two genera were found when chloramines were used as the disinfectant. The bactericidal activity of chloramine was noticably less than that of free chlorine.

Bacteria of the coliform group, especially E. coli, have proved useful as an indicator and disinfection model for enteric bacterial pathogens but

are poor indicators and disinfection models for nonbacterial pathogens. E. coil has been observed to be markedly more susceptible to chlorine than certain enteric viruses and cysts of pathogenic protozoa (Daling et al., 1972; Kruse', 1969).

The bacterial viruses of E. coil have received increased attention as possible disinfection models and indicators of enteric viruses in water and wastewater. At present, the data to justify the bacterial viruses as indicators for enteric viruses are limited and inconsistent However, there is a growing body of knowledge on the utilization of bacterial viruses as disinfection models.

Hsu (1964) and Hsu et al. (1966) first reported the use of the f2 virus as a model for disinfection studies with iodine. They showed that inactivation of both the f2 virus and poliovirus 1 were inhibited by increasing concentrations of iodide ion and that both f2 RNA and poliovirus 1 RNA were resistant to iodination.

Dahling et al. (1972) compared the inactivation of two enteric viruses (poliovirus 1 and coxsackievirus A9), two DNA phages c12 and T5), two RNA phages (f2 and MS2), and E. coil ATCC 11229 under demand4ree conditions with free chlorine at pH 6.0. They found enteric viruses to be most resistant to free chlorine followed by RNA phages, E. coil, and the T phages.

Shah and McCamish (1972) compared the resistance of poliovirus 1 and the coliphages f2 and T2 to 4 mg/liter combined residual chlorine. The f2 virus was shown to be more resistant to this form of chlorine than poliovirus 1 and Ti coliphage.

Cramer et al. (1976) compared the inactivation of poliovirus 3 (Leon) and f2 with chlorine and iodine in buffered wastewater. Both viruses were treated together in the same reaction flask, thereby eliminating any inherent differences due to virus preparations and replicate systems. In wastewater effluent at pH 6.0 and 10.0 with a 30 mg/liter dosage of halogen under prereacted halogen added to wastewater, allowed to react, viruses added at zero time) and dynamic (viruses added to wastewater, halogen added at zero time) conditions, f2 was, in each case, at least as or more resistant to chlorine and iodine than poliovirus 1. The f2 virus appears to be more sensitive to free chlorine but more resistant to combined chlorine than poliovirus 1 is.

Neefe et al. (1945) observed that the agent of infectious hepatitis was inactivated by breakpoint chlorination (free chlorine) but not completely inactivated by combined chlorine.

Engelbrecht et al. (1975) reported that the use of a yeast (Candida parapsilosis) and two acid-fast bacteria (Mycobacterium fortuitum and

Mycobacterium phlei) may provide suitable disinfection models. They observed that the yeast was more resistant to free chlorine than were poliovirus 1 and the enteric bacteria under all conditions tested. The acid-fast bacilli were most resistant.

There is no generally accepted disinfection model for protozoan cysts. In disinfection studies for protozoan diseases, investigators have used the pathogen or its cysts. Work with such Systems is, however, generally difficult.

The use of disinfection models provides useful information that is helpful to the comparison of the relative efficiencies of various disinfectants in the laboratory and in controlled field investigations. Strains of E. coli have been used extensively as models for enteric pathogenic bacteria. While not as widely accepted, the bacterial viruses of E. coli are used as disinfection models for enteric viruses. The difficulty of available methods has limited the number of disinfection studies with protozoan cysts.


The resistance or sensitivity to disinfectants of some bacteria (e.g.,

F. coli) in the laboratory may bear very little resemblance to their responses in nature. This is true in spite of the fact that standardized procedures govern the conditions under which cells are grown, harvested, washed, etc., when they are used as inocula. Examples of such differences range from Gram-negative bacteria and their comparative resistance to disinfectants in general (Carson et al., 1972; Favero et aL, 1971, 1975) to Gram-positive bacterial spores and heat resistance (Bond et aL, 1973) and to halogen resistance of Entamoeba histolytica cysts from simian hosts as opposed to those grown in in-vitro systems (Stringer et al., 1975). Presumably, the mechanisms creating this phenomenon among these three groups vary widely.

The comparative resistance to disinfectants among Gram-negative bacteria varies greatly. A good example of this is the study of Favero and Drake (1966). They first applied the term "naturally occurring" to certain Gram-negative bacteria with the potential for rapid growth in water. They observed that Pseudomonas alcaligenes, a common bacterial contaminant in iodinated swimming pools, could grow well in swimming pool waters that had been sterilized by membrane filters and rendered free of iodine or chlorine. Starting with contaminated swimming pool water that contained a variety of bacteria, they isolated a pure culture of P. akaligenes by an extinction-dilution technique in which filter-steri

lized swimming pool water was used as the diluent and growth medium. Since these cells had been isolated in pure culture without exposure to conventional laboratory culture media, they were referred to as "naturally Occurring" P. akaligenes. Subsequent tests showed that these naturally occurring cells were significantly more resistant to free iodine than were cells of the same organism that had been subcultured one time on trypticase soy agar. In fact, standard disinfectant tests using the cells that had been subcultured on an enriched laboratory medium suggested that P. akaligenes should never be found in pools that had been disinfected even minimally with iodine. This was obviously an erroneous assumption. The discovery that naturally occurring cells were extremely resistant to iodine explained the relatively high concentrations of P. akaligenes that accumulated in pool water that had been iodinated for several weeks.

Subsequently, Favero et al. (1971, 1975) and Carson et al. (1972) published a series of papers showing that Pseudomonas aeruginosa could grow rapidly in distilled water, which they obtained from hospitals, and could reach high concentrations of cells that remained stable for a long time. Naturally occurring cells that were grown in distilled water reacted quite differently to chemical and physical stresses than did cells grown on standard laboratory culture media. For example, naturally occurring cells of P. aeruginosa were significantly more resistant to chlorine, quaternary ammonium compounds, and alkaline glutaraldehyde than were subcultured cells.

In halogen-disinfected waters, naturally occurring bacteria can be from one to two orders of magnitude more resistant to the disinfectant than cells of the same Organism that had been subcultured on conventional laboratory culture media. Since standard disinfectant testing necessarily employs subcultured and washed bacterial cells, a false sense of confidence may be created if these data are used as an absolute criterion for the dilution of a disinfectant. These results could explain the frequent discrepancies between tests that are performed under laboratory conditions and those that are performed under field conditions.

If bacteria could be used in their naturally occurring state, one might explore the possibility of bridging the gaps between laboratory and field conditions by using this experimental system. The ability of some Gram-negative bacteria to grow in water makes it possible to produce and control large numbers of cells for such studies.

More difficult to answer is the more basic question of why naturally Occurring cells of Gram-negative water bacteria become more sensitive

to disinfectants when grown in a minimal medium than the same strain when grown in water. One would expect the reverse to occur. Milbauer and Grossowicz (1959) showed that cells of E. coli were much more sensitive to chlorine when grown on a medium of glucose mineral salts than when grown on nutrient agar. Since Favero and Drake (1966) reported that filter-sterilized dehalogenated swimming pool water could be considered a minimal medium, one would expect that P. alcaligenes cells that were grown in this environment would be less resistant to iodine than those grown in trypticase soy broth. This phenomenon has not been explained. Evidently it is not primarily a genetic response since the extreme difference in iodine resistance occurs with one subpassage on trypticase soy agar.

Over the years various investigators have tried without success to "train" bacteria to become more resistant to chlorine and/or iodine (Favero, 1961; Favero et al., 1964; Kruse', 1969). This failure is not surprising, because if halogens are truly a general cytoplasmic poison that affects primarily the sulfhydryl groups of enzymes (see pp. 36-39), it would be very difficult for an organism to modify its physiology to the extent that it becomes resistant, very unlike the situation with antibiotics and bacteria. Consequently, the extreme resistance or differing resistances of naturally occurring bacteria can be attributed only to "environmental" factors and, perhaps, to the different compositions of cell walls and membranes. However, there have been no data to substantiate this hypothesis.

Despite the questions that have been raised by differences in the behavior of organisms under both laboratory and field conditions, valuable comparative information can be obtained from studies of disinfectants that are conducted in similar laboratory Systems.


Chlorine is a strong oxidizing disinfectant that has been used to treat drinking water supplies for more than 60 yr. The gas was named "chlorine" after the Greek word for green, "chloros," because of its characteristic color. About 1800, chlorine gas was used as a general disinfectant in both France and England. In the United States, electrolytically produced chlorine was first used directly for water disinfection for only a week or two in 1896 at the Louisville Experimental Station in Kentucky. The first continuous municipal application of chlorine (as sodium hypochlorite [NaOC]) to water in the United States

occurred in 1908 at Jersey City, New Jersey. This was followed in 1912 by the first full-scale use of liquid chlorine for water disinfection at Niagara Falls, New York, where solution-feed equipment was used. This use of chlorine successfully eliminated recurring outbreaks of typhoid fever. In 1913, improved solution-feed equipment was developed to measure chlorine gas, dissolve it in water, and apply the solution to the water supply. This equipment was first installed at Boonton, New Jersey, where it replaced the use of sodium hypochiorite. Other equally significant historical occurrences (Laubush, 1971; White, 1972) led to the eventual addition of chlorine to drinking water in most of the United States for disinfection to destroy or inactivate pathogenic microorganisms (see Drinking Water and Health, National Academy of Sciences, 1977).

Chemistry of Chlorine in Water

Chlorine has an atomic number of 17, a melting point of -1020C, a

boiling point of -350C, and an oxidation potential of -1.36 V at 250C

(i.e., 2 Cl C12 + 2 e-). It is a green-yellow gas at room temperature. When chlorine is added to water, the following chemical reactions


C12 + H20 HOCI + H+ + Cr (1)

HOCI H+ +OCr (2)

Extremely little molecular chlorine (C12) is present at pH values greater

than pH 3.0 and total chlorine concentrations of less than 1,000

mg/liter (White, 1972). The hypochiorous acid (HOCl) that is produced further ionizes to form hypochlorite ion (OCl-) and hydrogen ion (H+) (Reaction 2). The dissociation of hypochlorous acid is dependent chiefly upon pH and, to a much lesser extent, temperature, with almost 100% hypochlorous acid present at pH 5 and almost 100% hypochlorite ion present at pH 10 (Figure Il-I). Free available chlorine refers to the concentration of hypochlorous acid and hypochlorite ion, as well as any molecular chlorine existing in a chlorinated water.


During chlorination of a water supply for disinfection, chlorine will react with any ammonia NH3) in the water to form inorganic chloramines.

The Disinfection of Drinking Water 19

100 - 0

80 - 20




20 C


4C 60

20 - 80



5 6 7 8 9 10 ii



FIGURE 11.1 Effect of pH on quantities of hypochlorous acid (HOQ) and hypochlorite ion (OCI-) that are present in water. Data from Fair eta'., 1948.

Furthermore, ammonia is sometimes deliberately added to chlorinated public water supplies to provide a combined available chlorine residual, i.e., inorganic chloramines. Chlorine will also react with organic amines. The organic chloramines that are produced (see below) are considered encompassed in the term "combined available chlorine."

Although inorganic chloramines are less effective oxidizing and disinfecting agents than hypochlorous acid and hypochlorite ion, they are more stable. Consequently, they will produce a residual in water that will persist for a longer time (Symons et al., 1977).

Inorganic chloramines are formed when hypochlorous acid reacts with ammonia:

NH3 + HOCl NH2Cl + H20 (3)

NH2Cl + HOCI NHCl2 + H20 (4)

The chloramine that is formed in the reaction depends upon the ratio of ammonia to hypochlorous acid and the pH of the system. Dichloramine (NHCl2) is the predominant form of chloramine at a 1:1 molar ratio of ammonia to chlorine at pH values of 5 and below, whereas at pH values of 9 and above, monochloramine NH2CL) predominates. Figure 11-2 shows the proportions of monochloramine and dichloramine formed for pH values of 4 to 9 and temperatures of 00C, 100C, and 250C (Morris, 1978, personal communication).


Chlorine is also known to combine slowly with Organic or albumenoid nitrogen (amines) to form organic chloramines (Taras, 1953):

R-NH2 + HOCl R-NHCI + HOH (5)

Although the reaction between organic amines and chlorine is generally considered to be slow, organic chloramines may be formed, thereby producing a stable combined available chlorine residual after many hours of contact. It is generally accepted that most organic chloramines have little disinfecting capability, i.e., less than the inorganic chloramines (Feng, 1966; Nusbaum, 1952).


Hypochlorous acid and other chlorine compounds having disinfecting ability by virtue of their being oxidizing agents will oxidize sulfites (SO32-), sulfides (S-), and ferrous (Fe2+) or manganous (Mn2+) ions. The disinfecting species are reduced, and the products have no disinfecting activity. All of the interfering compounds that destroy the disinfecting ability of the added chlorine exert a "chlorine demand," which may be defined as the difference between the amount of chlorine applied and the quantity of free or combined available chlorine residual measured in the water at the end of a specified contact period. When chlorine is added to water with no chlorine demand, a linear relationship is established between the chlorine dosage and the free chlorine residual (Figure 11-3).

1.c 0



s 5


z z

z h z


z 0.E - 0 0.4 ZW

0 0


w 10 w

LL 0.6 ~

0 0. 4 a

z 25 z

0 0


0.2 -

4 5 6 8 9


FIGURE 11-2 Proportions of Mono- and dichloramine (NH2CL and NHC12) in water chlorination with equimolar concentrations of chlorine and ammonia. Data from J.C. Morris, personal communication.

However, when increasing amounts of chlorine are added to water containing reducing agents and ammonia, the so-called breakpoint phenomenon occurs. The breakpoint is that dosage of chlorine that produces the first detectable amount of free available chlorine residual.

When chlorine is added to water, it reacts with any reducing agents and ammonia that are present. It is believed that chlorine reacts first with


0 1.0 2.0

z z

z 10 - z

2.0 -



w w

w w

z z

0 0

z 5


U, U)


U, U,

w w


0 0


FIGURE 11-3 Diagrammatic representation of completed breakpoint reaction. From Morris, 1978, personal communication.

the reducing agents. Since the chlorine is destroyed, no measurable residual is produced. Following the oxidation of these reducing agents, e.g., sulfides, sulfites, nitrites NO2-), and ferrous ions, the chlorine will react with ammonia to form inorganic chloramines. The quantity of monochloramine and dichloroamine that is formed is determined primarily by the pH of the water and the ratio of chlorine to ammonia. When the ratio by weight is less than 5:1, or the molar ratio is less than 1:1, and the pH is in the range of 6.5 to 8.5, the combined available chlorine residual is probably due primarily to monochloramine (Reaction 3). With additional chlorine, the ratio of chlorine to ammonia changes with the result that the monochloramines are converted to dichloramines (Reaction 4). When all of the ammonia has been reacted, a free available chlorine residual begins to develop. As the concentration increases, the previously formed chloramines are oxidized to nitrous oxide (N20), nitrogen trichloride (NCl3), and nitrogen N2) The reactions leading to the formation of these oxidized forms of nitrogen destroy the combined available chlorine residual so that the measurable residual in the water actually decreases. Upon completion of the

oxidation of all the chloramines, the addition of more chlorine creates the breakpoint phenomenon. At the breakpoint dosage, some resistant chloramines may still be present, but at such small concentrations that they are unimportant.

As pointed out by Morris (1970), the occurrence of reactions giving rise to the "breakpoint" is most rapid in the pH range 7.0 to 7.5. At greater and lesser pH values, it becomes slower and less distinct, e.g., at pH's < 6 or > 9 the concept of "breakpoint" is not significant. In the pH range 7.0 to 7.5 the "breakpoint" is about half developed within 10 mm at 150C to 200C and is then substantially completed within about 2 hr.

Analytical Methods and Their Evaluation

Standard Methods (1976) lists six acceptable methods for the determination of chlorine residuals in natural and treated waters: iodometric methods, amperometric titration, the stabilized neutral orthotolidine (SNORT) method, the ferrous diethyl-p-phenylenediamine (DPD) method, the DPD colorimetric method, and the leuco crystal violet (LCV) method.

The amperometric, LCV, DPD, and SNORT methods are unaffected by dichloramine concentrations in the range of 0 to 9 mg/liter (as C12) in the determination of free chlorine. If nitrogen trichioride is present, it reacts partially as free available chlorine in the amperometric, DPD, and SNORT methods. Nitrogen trichloride does not interfere with the LCV procedure for free chlorine. The sample color and turbidity may interfere with all colorimetric procedures. Thus, a compensation must be made. Also, Organic contaminants in the sample may produce a false4ree chlorine reading in most colorimetric methods.

Standard Methods (1976) contains data on the precision and accuracy of the methods used in the measurement of chlorine. These data were obtained from participating laboratories by the Analytical Reference Service (1969, 1971), which then operated in an agency that preceded the Environmental Protection Agency. However, as noted in Standard Methods (1976), these results are valuable only for comparison of the methods tested, and many factors, such as analytical skill, recognition of known interferences, and inherent limitations, determine the reliability of any given method. Moreover, some oxidizing agents, including free halogens other than chlorine, will appear quantitatively as free chlorine. This is also true of chlorine dioxide. Also, some nitrogen trichloride may be measured as free chlorine. The actions of interfering substances should be familiar to the analyst because they affect a particular method.

Although orthotolidine (i.e., orthotolidine and orthotolidine arsenite) methods have been widely used in many disinfection studies, they are omitted from the 14th edition of Standard Methods (1976) primarily because of their inaccuracy and high overall (average) total error in comparison with other available methods.

Research studies on disinfection are restricted by the limitations that are inherent in the methods themselves or by poor selection of methods by the investigator. The chemical conditions of the test water have not always been well defined. The types of titratable chlorine, i.e., free (11ypochlorous acid or hypochlorite ion) or combined (mono or dichloramine) in the chlorinated water, have not always been differentiated, and the rates of microbial destruction or inactivation have not always been studied in experimental systems with little or no chlorine demand. In fact, reports prior to the 1940's have been especially difficult to interpret, because reliable test methods for distinguishing between free and combined chlorine, between hypochlorous acid and hypochlorite ion, and between mono- and dichloramine in solution were not developed until the 1950's. For example, many earlier researchers claimed to have tested mono and dichloramine by controlling the pH and the ratio of chlorine to nitrogen. They used methods such as the orthotolidine or thiosulfate titrations to determine total chlorine residual. Much of this early work is now questionable, since it was not possible to detect free chlorine contamination in their chloramine solutions or the quantitative ratios between the mono and dichloramine tested. In addition, these earlier studies had high chlorine demand in the test systems.

Some more contemporary studies have lacked quantitated information on chlorine residual and/or types of chlorine present in the test systems.

Biocidal Activity

In the absence of reducing agents, inorganic ammonia, and organic amines, the addition of chlorine to municipal water supplies will result in free available residual chlorine, represented by the hypochlorous acid or hypochlorite ion. The pH determines the relative amounts of each species. However, inorganic chloramines will be formed if the background level of ammonia in the water supply is significant or if ammonia is intentionally added during treatment. If such is the case, monochloramine would predominate due to the alkaline pH of most finished water (see Figure 11-2).

In 1966, Feng proposed that the active forms of chlorine would exhibit disinfection properties in the following descending order:

C12> HOCI > OCr > NHCl2 > NH2CI > R-NHCI

Butterfield et al. (1943) published the first treatise on the use of chlorine demand-free water for studies of water disinfection. They proposed that to study the disinfectant capacity of any chlorine species, the test medium must meet certain exacting criteria. It must be nontoxic to bacteria except for the variables under study such as chlorine and pH, well buffered at the desired pH, free of all ammonia and Organic matter capable of forming chlorine-addition products, free of background chlorine, and of such a nature that calculated additions of chlorine are recoverable after 5 min without a loss in residual and that free chlorine must still be present several hours after contact.

Most studies of combined chlorine have dealt with poorly defined mixtures of mono- and dichloramine. Also, test conditions have often been inadequately defined, poorly controlled, or both.


Free Chlorine (HOCI and Ocl- Butterfield et al. (1943) studied percentages of inactivation as functions of time for E. col4 Enterobacter aerogenes, Pseudomonas aeruginosa, Salmonella typh4 and Shigella dysenteriae. They used different levels of free chlorine at pH values ranging from 7.0 to 10.7 and two temperature ranges-20C to 50C and 200C to 250C. Their work is of great importance, since very few other studies have been conducted that dealt with the action of disinfectants on pathogens. Generally, they found that the primary factors governing the bactericidal efficacy of free available chlorine and combined available chlorine were:

the time of contact between the bacteria and the bactericidal agent, i.e., the longer the time, the more effective the chlorine disinfection process;

the temperature of the water in which contact is made, i.e., the lower the temperature, the less effective the chlorine disinfecting activity; and

the pH of the water in which contact is made, i.e., the higher the pH, the less effective chlorination.

Thus, the test bacteria will be killed more rapidly at lower pH values and at higher temperatures. Since hypochlorous acid would predominate

at lower pH's figure 11-1), the data of Butterfield et aL show that it is a better bactericide than the hypochlorite ion. For example, to produce a 100% inactivation of an initial inoculum of 8 x 105 E. coil in 400 ml of sterile chlorine demand-free water (2,000/mi) at 200C-250C with a chlorine level of 0.046 to 0.055 mg/liter, 1.0 min was required at pH 7.0, but at pH 8.5, 9.8, and 10.7, between 20 and 60 min of exposure were needed. At higher concentrations of chlorine, i.e., from 0.1 to 0.29 mg/liter, exposure of 1.0 min was required at pH 7.0, 10 min at pH 8.5, 20 min at 9.8, and 60 min at 10.7. A similar pH effect was noted for 5. typhi.

Unfortunately, Butterfield et al. (1943) lifted their cells from agar slants but failed to wash them in demand-free water. The cells probably carried trace amounts of albumenoid nitrogen from the slants to the test flasks, thereby creating the small chlorine demand that the investigators had tried so carefully to avoid. The effect of such a trace amount of chlorine demand would be most apparent in test solutions with very low chlorine levels. In studies using approximately 0.1 mg/liter or less free chlorine at decreasing pH values, Butterfield et al. (1943) observed that the disinfection of the organisms required a very long time. This might indicate interference at the low levels due to the formation of combined chlorine.

Under very exact controlled test conditions of pH and temperature and using chlorine demand-free buffer systems, Scarpino et aL (1972) observed that at 50C E. coil was 99% destroyed by 1.0 mg/liter hypochlorous acid at pH 6 in less than 10 s and at pH 10 by 1.0 mg/liter hypochlorite ion in about 50 5. Their studies, which totally eliminated any form of combined chlorine from the test solutions, indicated that hypochlorous acid was approximately 50 times more effective than the hypochlorite ion as a bactericide. Fair et al. (1948) and Berg (l%6) analyzed the data of Butterfield et al. (1943) on the destruction of E. coil. They reported that hypochlorous acid was 70-80 times as bactericidal as hypochlorite ion.

Engelbrecht et al. (1975) investigated new microbial indicators of disinfection efficiency. In their chlorination studies, they noted the following decreasing order of resistance to free chlorine at pH values of 6, 7, and 10 and at 50C and 200C: acid-fast bacteria > yeasts > poliovirus > Salmonella typhimurium> E. coil.

Monochloramine (NH2Cl) In 1948, Butterfield summarized previous results on the bactericidal properties of chloramines (and free chlorine) in water at pH values ranging from 6.5 to 10.7 and in two temperature ranges-20C to 50C and 200C to 250C. The test bacteria included strains

of Escherichia coli, Enterobacter aerogenes, Pseudomonas aeruginosa, Salmonella typhi and Shigella dysenteriae. Although he admitted that adequate tests for separate determination of free and combined chlorine forms were not used in these studies, the solutions were vigorously prepared to ensure the exclusions of free chlorine. Chloramines were determined using orthotolidine; readings made after 10 to 30 s at 200C gave free chlorine levels, and those after standing for 10 min at 200C were recorded as total residual chlorine. Since no free chlorine was reported (and should not have been found, according to the authors), the 10-mm readings of total residual chlorine were also those of total chloramine levels. No distinction could be made between monochloramine and dichloramine. However, he estimated that at pH 6.5, 7.0, 7.8, 8.5, 9.5, and 10.5 the chloramines were present as monochloramine at

35%, 51%, 84%, 98%, 100%, and 100%, respectively. The balance was believed to be dichloramine.

Butterfield and his associates (Butterfield, 1948; Butterfield et aL, 1946) found that chloramine disinfection was always slower than that for free chlorine. For example, in order to achieve a 100% inactivation of the initial number of bacteria tested after 60 mm of contact time at 200 C, 0.6 mg/liter chloramine was required at pH 7.0 and 1.2 mg/liter at pH 8.5. At 40C, a 100% inactivation required 1.5 mg/liter chloramine at pH 7.0 and 1.8 mg/liter at pH 8.5. However, only 0.03 to 0.06 mg/liter free chlorine was needed at pH ranges of 7.0 and 8.5 at either 40C or 220C to achieve 100% inactivation in 20 mm.

The bactericidal effects of monochloramine alone were confirmed by Siders el al. in 1973. They found that at 150C E. coli was 99% destroyed in approximately 20 mm using 1.0 ppm monochloramine in pH 9 borate buffer. Since E. coli was less resistant to monochloramine than were the animal viruses tested, Siders et al. questioned the validity of using E. coli as an indicator organism for measuring the viral quality of a chlorinated water supply. Chang (1971) had previously calculated that 4 mg/liter monochloramine would be needed to give 99.999% reduction of E. coli bacterium in 10 nun at 250C.

Dichloramine (NHCI2) Chang (1971) calculated that 1.2 mg/liter dichloramine would be needed to give 99.999% reduction of enteric bacteria in 10 mm at 250C.

In carefully conceived studies, Esposito et al. (1974) examined the destruction rates of test organisms in contact with dichloramine in demand-free phthalate buffer at pH 4.5 and 150C. Figure 114 shows the comparisons that were made among enteroviruses (poliovirus 1 and

100 u

Poliovirus 1

-A coxsackie



10 e




z A


1.0 : ~X 174

E. coli AT~~ 11229


1.0 10 100 1000


FIGURE 114 Inactivation of various rnicroorganisms with dichloramine (NHCl2) at pH 4.5 and 150C. From Esposito, 1974.

coxsackievirus A9), the bacteriophage ~X-l74, and E. coli (ATCC 11229).

From a review of the literature and an analysis of the data, Chang (1971) calculated that the relative bactericidal efficiency of dichloramine to monochloramine was 3.3:1. However, Esposito (1974) and Esposito et al. (1974) showed experimentally that Chang's estimate was conservative. They found that dichloramine was 35 times more bactericidal than was monochloramine, not 3.3. At 150C, poliovirus 1 was 17 times more resistant than coxsackievirus A9, 83 times more resistant than ~X-174, and 1,700 times more resistant than E. coli to dichloramine (Figure IIA). They observed that dichloramine was a better bactericide than monochloramine.

Organic Chloramines These chlorine derivatives exhibit some bactericidal activity, but markedly less than either free chlorine or the inorganic chloramines (Feng, 1966; Nusbaum, 1952).

In summary, the bactericidal efficiency of hypochlorous acid, the hypochlorite ion, monochloramine, and dichloramine have been accu

rately defined in recent years by investigators using rigidly controlled test conditions. The order of disinfection efficiency presented by Feng (1 %6) has been confirmed. Comparative c t values are shown in Table 11-3 at the end of the section on chlorine.


In reviewing disinfection of enteroviruses in water, Clarke and Chang (1959) excluded all studies on the inactivation of viruses by chlorine that were conducted before 1946. Their justification for this exclusion was the failure of these studies to differentiate between free and combined chlorine. Furthermore, they attributed the irregular virucidal results of some studies to the use of animal inoculation methods for assaying virus concentrations. For these same reasons, those studies have been omitted in this report. The advent of viral propogation techniques using tissue cultures (Enders el al., 1949) enabled research of a more exacting nature to be performed, resulting in more precise virus inactivation data.

Free Chlorine (HOCl and OCl-) Generally, enteroviruses are more resistant to free chlorine than are the enteric bacteria (Chang, 1971; Clarke and Kabler, 1954; Scarpino et aL, 1972). For example, in what was probably the first well-defined study, Clarke and Kabler (1954) used purified coxsackievirus A2 to investigate viral inactivation in water by free chlorine. They carefully controlled their free chlorine residuals with a modified form of the orthotolidine test to determine total chlorine and an orthotolidine-arsenite method for free chlorine. (Combined chlorine was then calculated as the difference between "total" and "free" chlorine readings.) They measured virus recoveries by using suckling mice and the LD50 quantitation procedure. Their results indicated that inactivation times for the virus increased with increasing pH (6.9 to 9.0), decreasing temperatures (270C-290C to 3cC-6oC), and decreasing total chlorine concentration. They estimated that approximately 7 to 46 times as much free chlorine was required to obtain comparable inactivation of coxsackievirus A2 as was required for a suspension of E. coli cells (Butterfield et al., 1943). For instance, Butterfield et al. (1943) found that at pH 7.0 and at 20C to 50C, 99.9% of £. coli cells were inactivated in S mm with 0.03 mg/liter of free chlorine. At approximately the same pH and temperature ranges, Clarke and Kabler (1954) observed 99.6% inactivation of coxsackievirus A2 in 5 mm with 1.4 mg/liter of free chlorine, i.e., 46 times as much free chlorine as that required to inactivate

E. coli cells. At a pH of 8.5 at 250C, 99.9% of E. coli cells were inactivated in 3 min with 0.14 mg/liter of free chlorine (Butterfield et al.,

1943), while at a pH of 9.0 at 270C to 290C, 99.6% of the virus was inactivated in 3 min by 1.0 mg/liter of free chlorine. Thus, Clarke and Kabler's work showed that 7 times as much free chlorine was required to inactivate the test coxsackievirus compared to the time necessary to kill the bacterium E. coli. In a subsequent study, Clarke et al. (1956) found that adenovirus type 3, E. col4 and Salmonella typhi were all inactivated or destroyed at approximately the same concentration of free chlorine.

In 1958, Weidenkopf reported his studies on the rate of inactivation of poliovirus I as a function of free available chlorine (HOCl and OC1-) and pH at 00C. His results showed that 99% inactivation of poliovinis 1 was obtained with 0.10 mg/liter free chlorine in 10 min at pH 6.0 and of 00C. At pH 6, most of the free chlorine should have been present as hypochlorous acid. Increasing the pH to 7.0 increased the time required for the same degree of inactivation by approximately 50%. At that pH, the free chlorine should have been a mixture containing predominantly hypochlorous acid and significant levels of hypochlorite ion. Both Weidenkopf (1958) and Clarke et al. (1956) indicated that an increase in pH (from 7.0 to 8.5, Weidenkopf; from 8.8 to 9.0, Clarke et al.) increased the inactivation time about sixfold. At these pH's, the free chlorine should have been present as mixtures of both hypochlorous acid and hypochlorite ion, but predominantly as the ion.

After comparing these studies, Clarke et aL (1964) concluded that at 00C to 60C poliovirus 1 and coxsackievirus A2 were considerably more resistant to hypochiorous acid than was E. coli, while adenovirus type 3 was more sensitive (Figure 11-5). For 99% destruction of E. coli, a 99-s contact time was required when the system was dosed with 0.1 mg/liter free chlorine as hypochiorous acid. The same percentage of the adenovirus was inactivated in approximately one-third of that time by the same concentration of hypochlorous acid. Under the same conditions, 8.5 mm was required to inactivate 99% of the poliovirus, i.e., approximately S times the contact time required for E. coli. Coxsackievirus required a contact time for 99% inactivation in excess of 40 min, more than 24 times that required for E. coli.

Clarke and Kabler (1954) and Clarke et al. (1956) reported that the time required for inactivation of coxsackievirus A2 and adenovirus increased with increasing pH, decreasing total chlorine concentrations, and decreasing temperatures. Clarke and Chang (1959) concluded from data in the literature that a 100C increase in temperature increased the rate of virus inactivation by a factor of 2 to 3.

Kelly and Sanderson (1958) reported that each of six enteric viruses possessed a different sensitivity to chlorine. Their results suggested that the inactivation of enteric viruses in water at pH 7.0 and 250C required a

1.0 lililil I II II ( I IIIII~


a +0


C,, 10



a '9


w .010





.001 I 1111111 I II I II I I I II I

.1 1.0 2 10 20 100


FIGURE II-5 Concentration-time relationship for 99% destruction of £. coli and several viruses by hypochlorous acid at O-0C. From Clarke et aL, 1964

minimum free residual chlorine concentration of 0.3 mg/liter with a contact time of at least 30 min. With combined chlorine in water, a concentration of at least 9.0 mg/liter was necessary for a 99.7% inactivation of poliovirus at 250C and a pH of 7.0. Poliovirus 1 (strain MK 500) was the most resistant strain tested and coxsackievirus B5 the most sensitive. Poliovirus I (Mahoney strain), poliovirus 2, coxsackievirus B I, and poliovirus 3 were intermediate in resistance. The virucidal efficiency of hypochlorous acid was more than 50 times greater than that of the chloramines (Kelley and Sanderson, 1958, 1960).

Liu et al. (1971) studied the manner in which 20 strains of human enteric viruses responded to free chlorine. They used Potomac River water that had been partially treated by coagulation with alum and filtration through sand. Chlorine was added to the water at one dosage, 0.5 mg/liter. The final pH was 7.8. They stored the sample at 20C. There was a wide range of resistance to chlorine by the viruses. The most sensitive virus was reovirus type 1 which required 2.7 min for inactiving 4 logs (99.99%) of the virus with 0.5 mg/liter of free chlorine. The most resistant, as judged by extrapolating the experimental data, was

TABLE 11-1 Time Required for 99% Inactivation by Free Residual Chlorine at 5.00C t O.2oCa

Concentration Minutes

of Free Chlorine, for 990/o Rank

p11 mg/literb Virus Strain Inactivation Ordering

6.00 0.46-0.49 Coxsackie A9 (Griggs) 0.3

6.00 0.48-0.49 Echo I (F&ouk) 0.5 2

6.00-6.02 0.48-0.51 Polio2 (Lansing) 1.2 3

6.00-6.03 0.38-0.49 Echo S (Noyce) 1.3 4

6.00 0.47-0.49 Polio I (Mahoney) 2.1 5

6.00-6.06 0.51-0.52 Coxsackie B5 (Faulkner) 3.4 6

Coxsackie A9 (Griggs) NDC

7.81-7.82 0.47-0.49 Echo I (Farouk) 1.2 I

Polio 2 (Lansing) NDC

7.79-7.83 0.48-0.52 EchoS (Noyce) 1.8 3

7.80-7.84 0.46-0.51 Polio I (Mahoney) 1.3 2

7.81-7.82 0.48-0.50 Coxsackie B5 (Faulkner) 4.5 4

10.00-10.01 0.48-0.50 Coxsackie A9 ((1riggs) 1.5 1

10.00-10.40 0.49-0.51 Echo I (Farouk) 96.0 6

9.89-10.03 0.48-0.50 Polio2 (Lansing) 64.0 4

9.97-10.02 0.49-0.51 EchoS (Noyce) 27.0 3

9.99-10.40 0.50-0.52 Polio I (Mahoney) 21.0 2

9.93-10.05 0.50-0.51 Coxsackie B5 (Faulkner) 66.0 S

a Data from Fngelbrecht et al., 1978.

b Range of measured free chlorine residual in the "test" reactor at the termination of each of three separate experiments.

~ ND not determined.


poliovirus 2, which required 40 min for the same degree of inactivation. Using actual experimental data, the most resistant virus was echovirus 12, which required a contact time of greater than 60 min for 99.99% inactivation. Liu et al. (1971) concluded from their extrapolated values that the reoviruses were the least resistant to chlorine treatment, that both adenoviruses and echoviruses were less resistant, and that the polioviruses and coxsackieviruses were the most resistant. However, assuming a 20-mm contact time, most of the viruses tested at pH 7.8 and 20C would have been 99.99% inactivated with a free chlorine residual of

0.5 mg/liter.

Using six of the same virus strains studied by Liu et al. (1971), Engelbrecht et al. (1978) investigated the effect of pH on the kinetics of chlorine inactivation at 5.0 + 0.20C. The suspending medium was buffered, chlorine demand-free, distilled-deionized water. Each virus stock was also prepared so as to be chlorine demand-free. Table 11-1 summarizes the results, giving the chlorine levels used and the time required for two logs (99%) inactivation of the viruses at pH 6.0 and 7.8 in phosphate buffer and at pH 10.0 in borate buffer. Because of the use of two different buffer systems, i.e., at pH 6.0 and 10.0, virus inactivation was determined at pH 7.8 with each of the two buffer systems. The kinetics of inactivation of poliovirus 1 at pH 7.8, using both the phosphate and borate buffer, were the same. The results shown in Table

11-1 indicate that there is a significant difference in the time required for two logs inactivation for the various viruses at pH 6.0 and 10.0. In every case, the rate of inactivation at pH 10.0 was significantly less than at pH 6.0. The rank ordering in Table 11-1 shows that there is also a wide range of sensitivity of related viruses to chlorine disinfection. For example, at pH 10.0, coxsackie BS was 40 times more resistant than coxsackie A9. There are several cases in which the relative sensitivity to chlorine was altered (rank ordering) between pH 6.0 and 10.0, suggesting important effects of pH on the virion as well as on the chlorine species, i.e., hypochlorous acid versus. hypochlorite ion. This observation can be seen more clearly in Table 11-2 in which the time required for two logs of inactivation of the various viruses and the ratio of inactivation times at pH 6.0 and 10.0 are compared. Even at pH 7.8, differences in relative sensitivity appear when ranked and compared to results at pH 6.0 or 10.0 (Table lI-I).

Combined Chlorine (NH2CI and NHCl2) Viral inactivation rates with chloramines have been found to be much slower than with free chlorine. For example, Kelly and Sanderson (1958) studied the effects of chlorine on several enteric viruses. They reported that at pH 7 at 250C-280C, 0.2-

TABLE 11-2 Comparison of Virus Inactivation by Free Residual Chorine at pH 6.0 and 10.0, at 5.00C +- 0.2oCb

Minutes for 99% Inactivation

Virus Strain pH 6.0 pH 10.0 ratioC

Coxsackie A9 (Griggs) 0.3 1.5 S

Echo 1 (Farouk) 0.5 96.0 192

Polio2 (Lansing) 1.2 64.0 53

echo 5 (Noyce) 1.3 27.0 21

Polio 1 (Mahoney) 2.1 21.0 10

Coxsackie BS (Faulkner) 3.4 66.0 19

a Chlorine concentrations as noted in Table 11-1 at pH 6.0 and 10.0.

b Data from Engelbrecht et al., 1978. Time required at pH 10.0

Time required at pH 6.0

0.3 mg/liter free chlorine inactivated 99.9% of all test viruses in 8 mm. At the same temperature and pH, combined chlorine at 0.7 mg/liter and at least 4 hr of contact time were needed to achieve 99.7% inactivation of the test viruses.

Although most viral inactivation studies with chloramines have not differentiated between mono and dichloramine (Kelly and Sanderson, 1958; Lothrop and Sproul, 1969), Kelly and Sanderson (1958, 1%0) noted that viral inactivation by chloramines proceeds more rapidly at pH 6-7 than at pH 8-l0. This tendency indicates that dichloramine may be more virucidal than monochloramine since its proportion increases with increased hydrogen ion concentration.

In 1971, Chang proposed that 5.0 mg/liter dichloramine or 20 mg/liter monochloramine would be needed to inactivate enteroviruses by 99.99% in 10 mm at 25 CC. Subsequently, Siders et al. (1973) presented evidence of the first comparative studies on viral inactivation due solely to monochloramine. At pH 9 at lSCC, poliovirus 1 (Mahoney) was 10 times more resistant to monochloramine than was E. coli. Similarly, coxsackievirus A9 was approximately 4 times more resistant than E. coli. Siders' data can be compared to Chang's theory on the disinfectant capacity of monochloramine. Assume that poliovirus inactivation has a temperature coefficient for a 100C rise (Qio) of 3. Based on Chang's (1971) calculations that 20 mg/liter monochloramine at 25CC would be needed

to achieve 99.999% reduction of enterovirus in 10 min, one would extrapolate that 3 x 20 or 60 mg/liter monochloramine would be needed at I 5 C to achieve an equivalent reduction of poliovirus in 10 min. However, examination of the data of Siders et al. reveals that approximately 18 mm were needed to achieve only 99% inactivation using 60 ppm monochloramine. Therefore, it appears that Chang slightly overestimated the virucidal capacity of monochloramine.


Some studies, mostly in the field of wastewater treatment, have shown that ova and larvae of the helminth parasites that affect humans and that could occur in U.S. water supplies are resistant to current chlorination procedures. They can survive concentrations and exposure periods considerably in excess of those used in the treatment of municipal water supplies. In studies of various free-living nematodes, Chang et al (l%0) observed that 2.5 to 3.0 mg/liter of free chlorine for a l20-min contact period and 15 to 45 mg/liter of free chlorine for I mm were not lethal. Free chlorine residuals as high as 95 to 100 mg/liter for S min killed only 40%-50% of the nematodes. Thus, it may be speculated that all the helminths, including their larvae, may approach the degree of resistance to chlorine that had been demonstrated by the free-living nematodes.

There have been a number of studies on the effectiveness of chlorine in destroying or inactivating cysts of the protozoan parasite, Entamoeba histolytica, in water, especially during the early 1940's (Brady et al., 1943; Chang, 1944b; Chang and Fair, 1941). Varied results reflect primarily the different experimental conditions and techniques that were used. The presence of organic matter, pH, and temperature, as well as the concentration and form of chlorine and exposure period, have been shown to exert an influence on disinfection. However, the consensus is that, compared with bacteria, these cysts are rather resistant to current chlorination procedures, but are much less resistant than helminths.

Brady et al. (1943) conducted field-simulated studies with cysts in raw water that had been treated with calcium hypochlorite (CaOCl), resulting in pH levels ranging from approximately 7.5 to 8.0. They found that at temperatures of 230C-260C, exposures of 20 min and longer to residuals of 3 to 4 mg/liter were required to produce an estimated 99% cyst destruction as judged via a culture technique. Chang (1944b), also using a culture technique, studied the cysticidal effectiveness of calcium hypochlorite solution, chloramines, and gaseous chlorine in tap water as well as the effects of pH and organic matter on the biocidal activity. At contact periods of up to 30 mm, gaseous chlorine was the most powerful,

hypochlorite solution slightly less so, and chloramines the least. Increase in pH and organic matter reduced cysticidal efficacy. For comparison with Brady et al. (1943), the "lethal" residual concentration in tap water at 180C and pH 6.8-7.2 ranged from 2.8 to 3.2 mg/liter at 15 min, and from 1.8 to 2.2 mg/liter at 30 min (as estimated from a graph in Chang, 1944b).

Recently, Stringer et al. (1975) reported on comparative studies of the cysticidal efficacy of chlorine, bromine, and iodine as disinfectants. Using chlorine gas bubbled into buffered distilled water as stock, they obtained 99.9% cyst inactivation (as measured by excystment capability) after 15 min exposure to 2 mg/liter free chlorine in "clean water" at pH 6. However at pH 8 a contact time exceeding 60 min was required to achieve 99% mortality. In "secondary treated sewage effluent," Stringer et al. considered 13.7 mg/liter chlorine at pH 8 to be ineffectual as a cysticide.

In keeping with these findings, it is unlikely that the chlorine residuals generally maintained in distribution systems provide much protection against E. histolytica cysts in the event of contamination because of cross-connections, seepage, etc.

During the past 10 yr, a number of outbreaks of waterborne infections from Giardia lamblia (another intestinal protozoan) have been reported national Academy of Sciences, 1977). Most incidents in the United States that were traced to municipal water supplies involved surface water sources where disinfection appeared to be the only treatment. The cysts of this parasite are thought to be as resistant to chlorine as those of

E. hystolytka. However, there seem to be no studies of the resistance of this parasite to chlorine or other disinfectants.

Mechanism of Action

One of the earliest references to the mechanism of inactivation of microorganisms by chlorine resulted from the work of Chang (1944a,b). While studying the inactivation of E. histolytica cysts by chlorine, he observed greater uptake of chlorine and less survival at low pH than at high pH. This observation was associated with the increased inactivation efficiency of the undissociated hypochlorous acid. Supportive evidence for the hypothesis that permeability of the uncharged chlorine species is important in determining sensitivity to chlorine has been provided by Skvortsova and Lebedeva (1973), Kaminski et al. (1976), and Dennis (1977). Chang (1 944a) also noted that the inactivation of amoebic cysts was accompanied by microscopic damage to the cell nucleus, which was dependent on chlorine penetration.

The importance of penetration and/or damage to the permeability barrier of the cell membrane as a result of exposure to chlorine has been observed by several investigators. In 1945, Rahn suggested that the inactivation of bacteria by chlorine was due to multiple injuries to the cell surface. From their work with bacterial spores, Kulikovsky et al. (1975) implicated permeability damage as a mechanism of chlorine inactivation. Studies with E3cherichia coli have shown that chlorine causes leakage of cytoplasmic material, first protein, then RNA and DNA, into the suspending menstruum. It also inhibits the biochemical activities that are associated with the bacterial cell membrane Venkobachar, 1975; Venkobachar et al, 1977). Friberg (1957) observed that E. coli also loses nondialyzable phosphorus following exposure to chlorine. In a recent study, Haas (1978) demonstrated that chlorine caused certain bacteria and yeast to release Organic matter or UV-absorbing material, presumably protein or nucleic acid or their precursors. This investigator also noted that chlorine affected the uptake and retention of potassium by these same microorganisms.

Green and Stumpf (1946) and Knox et al. (1948) indicated that destruction of bacteria by chlorine was caused by an inhibition of the mechanism of glucose oxidation. Specifically, they suggested that chlorine affected the aldolase enzyme of E. coli by oxidizing the sulfhydryl group that is associated with the enzyme. Venkobachar et al. (1977) recently reported that chlorine significantly inhibits both oxygen uptake and oxidative phosphorylation. The latter effect was attributed to inhibition of the respiratory enzyme rather than to a deficiency in phosphate uptake. However, it is unclear whether free or combined chlorine was used in these studies. Haas (1978) also observed chlorine to affect the respiration of bacteria as well as the rate of synthesis of protein and DNA. Others have also noted that chlorine affects the nucleic acids or physically damages DNA (Bocharov, 1970; Bocharov and Kulikovskii, 1971; Fetner, 1962; Rosenkranz, 1973; Shih and Lederberg, 1976a,b).

It appears that chlorine, having penetrated the cell wall, encounters the cell membrane and alters its permeability. Simultaneously or subsequently, the chlorine molecules may enter the cytoplasm and interfere with various enzymatic reactions. It should be noted that permeases and respiratory enzymes are associated with the cytoplasmic membrane of bacteria.

Chang (1971) supported the hypothesis that the rapid destruction of vegetative bacteria by chlorine was due to the extensive destruction of metabolic enzyme systems. He also addressed the subject of virus inactivation, commenting that viruses are generally more resistant to

chlorine than bacteria. He associated this observation with the fact that viruses completely lack a metabolic enzyme system. He speculated that inactivation of viruses by chorine probably result from the denaturation of the capsid protein. Furthermore, since protein denaturation is more difficult to achieve than destruction of enyymatic R-S-H bonds by oxidizing agents, it is understandable why greater levels of chlorine are required to inactivate viruses than bacteria. However, from their experimental work with the bacterial virus f2, Olivieri et al. (1975) concluded that chlorine caused initial lethal damage to the viral genome and that the capsid protein was affected after the virus was inactivated. Dennis (1977) reported that the incorporation of chlorine into the f2 bacterial virus is dependent on pH and that the higher rates of incorporation occur at lower pH values.

There is limited information in the literature on the mechanism of inactivation of microorganisms by chloramines. Nusbaum (1952) proposed that since low levels of inorganic chloramines were effective in inactivating bacteria, the mechanism of action must be essentially the same as that of hypochlorous acid on enzymies. Ingols et al. (1953) showed that monochloramine was not able to immediately and irreversibly oxidize sulfydryl groups. Such oxidation would have resulted in the rapid inactivation of the bacteria. They hypothesized that since monochloramine required higher concentrations and longer contact times to destroy bacteria completely and could not readily and irreversibly oxidize the sulfhydryl groups of the glucose oxidation enzymes, its ability to inactivate microorganisms should be attributed to changes in enzymes that may not be involved in the inactivation of the organism by hypochlorous acid. Thus, while the sulfhydryl group may be the most vulnerable to a strong oxidant like hypochlorous acid, changes in other groups produced by the weaker oxidant, monochloramine, may lead also to microbial inactivation. More recent information indicates that the destructive effects of chloramine might be associated with the effects of chloramine on nucleic acids or DNA of cells (Fetner, 1962; Shih and Lederberg, 1976a,b).

Nusbaum (1952) suggested that the disinfective activity of dichloramine occurs by a mechanism similar to monochloramine, but there do not appear to be any data to support this contention.

Considering the mechanism of destruction or inactivation of microorganisms by chlorine and associated compounds, it is interesting to note that Fair et al. (1948) speculated that there might be three or four "targets" or points of attack and that perhaps all must be affected before there is death. This "multiple hit" concept supported the observation

that monochloramine must alter groups other than the sulfhydryl group to be effective in the destruction of microorganisms.

Thus, the action of chlorine on microbes such as bacteria and amebic cysts may involve some or all of the steps in the following sequence:

penetration of the disinfectant through the cell wall followed by attack on the cell membrane (the site of cellular respiration in bacteria) and disruption of permeability of the cell membrane, which leads to a loss of cell constituents, thereby disrupting metabolic functions within the cell including those involving nucleic acids. Changes in viability may result from this process. Experimental studies on virus (Olivieri et al., 1975) demonstrated that chlorine caused initial damage on the viral nucleic acid while leaving the capsid protein unaffected until after the virus was inactivated.


Chlorine is the most widely used water supply disinfectant in the United States. Depending upon the predominant species of chlorine, hypochlorous acid, and/or hypochlorite ion, disinfection with chlorine can achieve greater than 99.9% destruction of bacteria. For example, a chlorine residual of 0.2 to 1.0 mg/liter and a contact time of 15 to 30 min will inactivate 99.9% of E. coli (Walton, 1969). According to Walton, a properly designed, constructed, and operated water treatment plant, consisting of chemical coagulation, sedimentation, filtration, and disinfection, can remove or destroy more than 99.999% of the coliform bacteria that are present. Although most investigations on the removal or destruction of bacteria have used E. col4 there is evidence that the bacterial pathogens, e.g., Salmonella typhi, respond somewhat similarly to E. coli.

Laboratory studies have demonstrated that that there is limited virus inactivation after the added chlorine has reacted with any ammonia that is in the water. Most inactivation probably occurs in the first few seconds before the chlorine has completed its reaction with ammonia (Olivieri et al., 1971).

Table 11-3 displays c t values for F. coli and poliovirus inactivation for the various species of free and combined chlorine.

Research Recommendations

Recent reports of enhanced chlorine resistance of certain viral and bacterial strains should be investigated and the mechanism of increased resistance elucidated, if the reports are corroborated.


TABLE 11-3 Dosages of Various Chlorine Species Required for 990/o Inactivation of Escherichia Coli and Poliovirus 1

Concen- Contact

Test Disinfecling tration, Time, Tempera-

Microorganism Agent mg/liter mm c-t0 pH ture,0C References

L coli Hypochlorous

acid (HOCi)


ion (OC1~)



0.1 0.4 0.04 6.0 S Scarpino el al., 1974

1.0 0.92 0.92 10.0 S Scarpino el (*1., 1974

1.0 175.0 175.0 9.0 S

1.0 64 64.0 9.0 IS

1.2 33.5 40.2 9.0 25

Dichloramine 1.0 5.5 5.5 4.5 IS


Siders el al., 1973

Dom, 1974

Siders el al., 1973

Dorn, 1974

Siders el al., 1973

Dorn, 1974

Esposito, 1974

Esposito el al., 1974

1)()Ii()vira~ I Ilypochlorous

~cid (lIOCI)


ion (OCI~)

Monochlorarnine 10


1.0 1.0

0.5 2.1

1.0 2.1






I.0 6.0 0

1.05 6.0 5

2.1 6.0 S

1.0 6.0 15

10.5 10.0 S

3.5 3.5 10.0 IS

90 900 9.0 IS

10 32 320 9.0 25

Dichloramine 100 140 14,000 4.5 s


100 So 5,000 4.5 15

Wei(1cnkopf, I 958

IRngelbrecht ~ ~jI., 1978

S~~rpino el aI~, 1974

Itrigmo el al., 1978

Lngelbrechi ci ul., 1978

Itrigano ci al., 1978

Siders Cl al., 1973

Porn, 1974

Siders ci 'ii., 1974

Porn, 1974

[~posito, 1974

Lsposito CI (ii., 1974

1~posito, 1974

Lsposito ei aI~ 1974

Other recommendations, applicable to other agents as well as to chlorine, are included after the evaluations of the other methods of disinfection.


Chemistry of Ozone in Water

Ozone has the molecular formula 03, a molecular weight of 48 g/mol, and a density, as a gas, of 2.154 g/liter at 00C and 1 atm. It is approximately 13 times more soluble in water than is oxygen. At saturation in water at 200C, a 2% weight mixture of ozone and oxygen contains about 11 mg of ozone and 40 mg of oxygen per liter.

Ozone has a half-life in pure distilled water of approximately 40 min at pH 7.6, but this decreases to 10 min at pH 8.5 (Stuum, 1958) at 14.60C. Rising temperatures increase the rate of decomposition. Because its half-life is so short, ozone must be generated on the site where it is to be used.

Ozone is a powerful oxidant that reacts rapidly with most organic and many inorganic compounds. It does not convert chloride to chlorine under test conditions (U.S. Environmental Protection Agency, 1976), nor does it react extensively with ammonia (NH3) (Singer and zilli, 1975). However, bromide and iodide are oxidized to bromine and iodine. Singer and Zilli (1975) reported that oxidation of ammonia was pH-dependent. At pH 7.0, 9% of a 29 mg/liter ammonia nitrogen NH3-N) solution was oxidized in 30 min, but at pH 9.070% of a 24.4 mg/liter solution was oxidized in the same time. During disinfection, only minor amounts of ammonia are oxidized when ozone is used. Ozone's limited reaction with ammonia is desirable, but its fast reaction rate with most organic and many inorganic compounds further shortens its persistence in water.

Production and Application of Ozone

Ozone is produced on site from a stream of clean dry air or oxygen by passing an electrical discharge between electrodes that are separated by a dielectric. Approximately twice the percent of ozone by weight is obtained if oxygen, rather than air, is used as the feed stream. Power requirements are about 13 to 22 kWh/kg of ozone that is generated from air and approximately half that when oxygen is used (Rosen, 1972). Compressors and dryers may increase these requirements by 20% to 50%. Other factors affecting efficiency are the rate of gas flow, applied voltage,

and the temperature of the gas. The heat that is produced during the process must be removed by cooling with either air or water.

The ozone gas stream must be fed into the water to effect the transfer of ozone. The usual methods are to inject the ozone gas stream through an orifice at the bottom of a co or countercurrent contact chamber or to aspirate the gas into a contact chamber where it is mixed with the water mechanically. Successful design and operation of the contactor system is necessary to minimize costs of the operation.

Commercial equipment is available in a wide range of capacities-from a few grams of ozone per day to more than 40 kg/day. Larger capacities are obtained by adding additional units. Successful delivery of ozone to the water to be treated requires a dependable power supply and reasonably maintenance-free ozonization equipment.

Ozone has been used in a great number of water treatment plants throughout the world. However, in small institutions and private residences, its use appears limited, because it requires dependable power supplies and, usually a second disinfectant to furnish a disinfecting residual in the system. The maintenance and repairs that are required for the specialized ozone generation equipment provide further barriers against the use of ozone by small institutions.

Analytical Methods

The disinfection process is usually controlled in one of two ways: by the dosage of a specified amount of ozone or by the maintenance of a specified rninimum residual for a given time. Residual measurements in both the gas stream and water are sometimes required. Standard Method (1976) contains descriptions of the measurement of ozone in water by the iodometric, orthotolidine-manganese sulfate, and orthotolidine-arsenite methods. Of these methods the iodometric method, which is subject to the fewest interferences, is the method of choice. Determinations must be made immediately since ozone decomposes rapidly.

In all three methods, the oxidant compounds that result from the reaction of ozone with contaminants in water may react with the test reagents, thereby indicating a higher concentration of ozone than is actually present. This is particularly true in the presence of organic matter, which results in the formation of organic peroxides. In the iodometric method this interference and others are minimized by stripping the ozone from the sample with nitrogen or air and absorbing it from this gas stream in an iodide solution. These and other methods are described in Standard Methods.

Schecter (1973) developed a UV spectrophotometric method to

measure the triiodide that is formed by the oxidation of iodide by ozone. She reported a better sensitivity at low ozone concentrations (0.01 to 0.3 mg/liter) than achieved with the normal titration method. The effects of interferences on the direct measurement of ozone without sparging the ozone to a separate iodide solution were not indicated. These effects are noted in Standard Methods.

Analytical determination of ozone in water in the presence of other oxidants is poor. Considerable work in this area is needed.


As discussed above, ozone is unstable in water with a half-life of approximately 40 min at pH 7.6 and 14.60C. Many regard the half-life in water supplies at higher ambient temperatures to be 10 to 20 nun.

In a recent review of the literature, Peleg (1976) concluded that the possible species to be found in an ozone solution were ozone, hydroxyl radicals (.OH), hydroperoxyl radicals (H02.), oxide radicals (0.), ozonide radicals (03.), and, possibly, free oxygen atoms (.0.). Hydrogen peroxide (H202) may also be present by dimerization of the hydroxyl radicals. There have been no studies on the disinfecting activities of these individual species except for those on hydrogen peroxide, which is a poor biocide (when compared to chlorine). Peleg concluded that evidence indicates that the dissociation species are better disinfectants than ozone.

Biocidal Activity


Wuhrrnann and Meyrath (1955) found that 99% of £scherchia coli were inactivated in 21 s with 0.0125 mg/liter residual ozone, in 62 s with 0.0023 mg/liter residual ozone, and in 100 s with 0.0006 mg/liter residual ozone at pH 7.0 and 120C. Spores of a Bacillus species were much more resistant, requiring 2 min with 0.191 mg/liter residual ozone and S min with 0.049 mg/liter residual ozone for 99% inactivation. These data were obtained at pH 7.2 and 220C. The observed ozone residuals were reported as being constant throughout the test periods.

Katzenelson et al. (1974) found that an initial residual ozone concentration of 0.04 mg/liter in demand-free water inactivated approximately 3 logs of E. coli in 50 5, while a concentration of 1.3 mg/liter achieved the same degree of inactivation in S 5. They obtained their inactivation data from experiments in which the loss of initial ozone

residuals did not exceed 20%. Their work was conducted at 1 0C and pH 7.2. The ozone was determined by the method of Schecter (1973).

Using washed cells of E. col4 Bacillus megaterium, and Bacillus cereus, which were suspended in deionized water with 10 6 cells/ml, and different concentrations of ozone up to 0.71 mg/liter, Broadwater et al. (1973) demonstrated that the inactivation of all three test organisms gave an "all-or-none" response. They measured the ozone by stripping it into iodide at the end of the contact period. Consequently, any demand should have used up whatever ozone was needed. They did measure what they presumed to be ozone and not some breakdown product. Thus, they appear to have eliminated the ozone demand problem as much as is possible with present techniques. However, it is possible that their results reflect some effect or artifact not yet understood. With the constant contact time of 5 mm, no inactivation of vegetative cells of E. coli or B. megaterium occurred until the initial residual concentration of ozone was 0.19 mg/liter, when the density of viable cells of both species decreased to "near zero." With the vegetative cells of B. cereus, an initial ozone residual of 0.12 mg/liter was required before any inactivation occurred, and then it was nearly complete. Spores of B. megaterium and

B. cereus were not inactivated until the ozone residual reached approximately 2.29 mg/liter, when the spores of both organisms showed an "all-or-none" response. All of the inactivation experiments were performed at 280C, but the pH was not reported.

Burleson et al. (1975) determined that E. col4 Staphylococcus aureus, Salmonella typhimurium, Shigella flexneri, Psuedomonus fluorescens, and Vibrio cholerae were all reduced by 7.5 logs in 15 s after exposure to approximately 0.5 mg/liter of ozone at 250C in phosphate-buffered saline. The pH was not reported. In this study, the bacteria were placed in unozonized water (no initial residual), and the ozone was then sparged into the water. After 15 s the ozone concentration in solution reached approximately 0.5 mg/liter. This technique does not render quantitative data. Ozone was determined by spectrophotometric measurement of iodine that was released from iodide without stripping of the ozone.

The inactivation of E. coli with initial counts of about S x 105/mi in buffered demand-free water at 11 0C was measured by Ross et al. (1976). For 99% inactivation with 0.1 mg/liter of initial ozone residuals, contact times of 16.5 and 21 s were required at pH 6 and 10, respectively. Ozone was measured by the diethyl-p-phenylenediamine (DPD) method.

Farooq (1976) and Farooq et aL (1977a,b,c) studied the effects of pH and temperature in ozone inactivation studies with Mycobacterium fortuitum and candida parapsilosis (a yeast). They concluded that if the ozone residual remains constant, the disinfection capability will not be

affected by a change in pH. They also demonstrated that for a given dosage a rise in temperature increases the rate of inactivation, even though the ozone residual was decreased. (Ozone was less soluble at higher temperature). Vorchinskii (1963) concluded that whereas the bactericidal dose of ozone at 40C to 60C was unity, the corresponding dose at 180C to 210C was 1.6 and at 360C to 380C was 3.6.

The work of Farooq and his colleagues is in agreement with that of Morris (1976), who observed that the disinfection capability of ozone does not change significantly with pH, at least over the normal pH range (6 to 8.5) of water supplies.


Katzenelson et al. (1974) investigated the inactivation of poliovirus 1 at 50C and pH 7.2. A 99% inactivation of the virus occurred in less than 8 s with 0.3 mg/liter of initial residual ozone. Their experimental methods were the same as those described above for bacteria. They also observed that inactivation resulted from two distinct stages (rates) of action. The first stage of inactivation, less than 8 s long, produced a virus inactivation of from 99% to 99.5%, depending upon the ozone residual. The second stage, which lasted from 1 to S min, still left some viruses infective. Additional work showed that the slower second-stage inactivation apparently involved the inactivation of viruses that were clumped together. The single virus particles were inactivated during the first stage. After ultrasonic treatment, 99.5% of the virus was inactivated within 8 s and more than 99.99% within 3 min at an initial ozone residual of 0.1 mg/liter.

Coin et al. (1964) investigated the inactivation of poliovirus 1 (Mahoney) in a batch system in which a measured amount of ozone was introduced into a virus suspension in distilled and filtered river water. In distilled water, nearly 1 log of the virus was inactivated in 4 min at a 4- min ozone residual of approximately 0.23 mg/liter. In river water, the inactivation was in excess of 99.99% in 4 min when the ozone concentration, also measured at 4 min, exceeded 0.3 mg/liter. The ozone was measured by the iodide titration without stripping of the ozone. Rather large ozone demands existed in this system. An initial S mg/liter residual ozone concentration in the distilled water decreased to about 0.6 to 0.8 mg/liter in 4 mm. The same initial residual in river water decreased to between 0.2 to 0.6 mg/liter. Temperature and pH were not reported.

Keller et al. (1974) studied ozone inactivation of viruses in a water supply by using both batch tests and a pilot plant with a 38 liter/min (10

TABLE 11-4 Ozone Required for Inactivation of Viruses in 10 Minutes at pH 7.0 and 25oCa

Ozone Required, mg/liter

99.90/o 99%

Virus Inactivation Inactivation

Coxsackie B3 0.6 0.095

Polio 3 0.22 0.082

Polio 1 0.095 0.042

Echo 1 0.086 0.044

Coxsackie BS 0.076 0.053

Polio 2 0.052 0.039

a Data from Evison, 1977.

gal/min) flow rate. Inactivation of poliovirus 2 and coxsackievirus B3 in 5 mm was greater than 99.9% in the batch tests when the ozone residual was between 0.8 and 1.7 mg/liter at the end of the 5-mm contact period. Initial ozone residuals varied from 1.6 to 2.8 mg/liter. At the pilot plant, greater than 99.999% inactivation of coxsackievirus B3 could be achieved with an ozone dosage of 1.45 mg/liter, which provided an ozone residual of 0.28 mg/liter after 1 min in lake water. Temperature and pH conditions were not reported. Ozone was measured by the iodide method without prior stripping of the ozone.

Burleson et al. (1975) obtained inactivations of >99.9% for vesicular stomatitis virus, >99.99% for encephalomyocarditis virus, and 99.99% for GD VII virus in 15 s with an ozone residual of approximately 0.5 mg/liter of ozone. The experiment was conducted in the same manner as that described above for their studies with bacteria. Evison (1977) reported data for the inactivation of a number of viruses in buffered water at pH 7.0 and 25CC (Table 114).

The reported ozone concentrations were evidently those measured initially and maintained by the addition of ozone during the experiment. Ozone was measured by a colorimetric version of Palm's DPD technique. The Evison data show that more ozone or a longer time are required for inactivation than do the data of other workers. This may have resulted from the virus purification used. Her viruses were purified by low-speed centrifugation and filtration through an 0.2-um membrane filter. These cleanup procedures are neither as complete nor as thorough as those used by other investigators. The unremoved cell debris and organic matter offer protection to the virus. Either higher ozone residuals

or longer contact times would be required to inactivate such preparations to the same extent as clean virus. Other data showed that the inactivation of coliphage 185 was relatively unaffected by pH's ranging from 6.0 to 8.0 at ozone concentrations exceeding 0.05 to 0.10 mg/liter. Evison (1977) also concluded that the rate of inactivation of the coliphage 185 by ozone was much less affected by temperature than was the inactivation by chlorine.

Farooq (1976) observed a 99% inactivation of ml poliovirus 1 in

30 s in distilled water at an initial ozone residual concentration between

0.23 and 0.26 mg/liter at 240C and pH 7.0. In this study, ozone was

measured by the Schecter (1973) method.

Sproul et al. (1978) reported total inactivation in 10 s of poliovirus 1 (Sabin), which had initial concentrations of 1.3 x 102 to 2.4 x 103 plaque-forming units (PFU)/ml. The initial ozone residual of 0.012 to

0.085 mg/liter decreased by approximately one-third in 40 5. The experiments were conducted with the Sharpe dynamic reactor. Ozone was measured by the Schecter (1973) method.


No studies on the efficacy of ozone as an antihelmintic agent appear to have been reported.

Ozone may have application as an antiparasitic agent in the treatment of water supplies but only limited information is available. Newton and Jones (1949) reported that ozone, with 5-min residuals as low as 0.3 mg/liter, inactivated from 98% to >99% of Entamoeba histolytica cysts that were suspended in water. Initial ozone residuals that were required to obtain 5-mm residuals of 0.3 mg/liter varied from 0.7 mg4iter to 0.9 mg/liter. With the ozone concentrations used, the cysticidal action was not affected by temperatures from 100C to 270C nor by pH's between 6.5 to 8.0. Ozone was measured by titration of iodine, which was released from iodide directly in the reactors without removal of ozone by sparging.

Mechanism of Action

Investigations of the inactivation of bacteria by ozone have centered on the action of ozone on the cell membrane. Scott and Lesher (1%3) concluded from their work with E. coli that the primary attack of ozone occurred on the double bonds of the fatty acids in the cell wall and

membrane and that there was a consequent change in the cell permeability. Cell contents then leaked into the water. This was confirmed by Smith (1967). Prat et al (1968), examining extracts from ozonized E. coli, reported that thymine was more sensitive to ozone attack than were cytosine and uracil.

Riesser et al. (1977) showed that ozone attacked the protein capsid of poliovirus 2 in such a manner that the virus was not taken up into susceptible cells. An electrophoretic study showed complete loss of viral proteins in a poliovirus 2 sample that had showed an inactivation of 7 logs in 20 min.

Summary and Conclusions

Inactivation with ozone at specified ozone residuals is relatively insensitive to pH's between 6.0 and 8.5. Moreover, ozone does not react with ammonia over this same range when short detention times are used. The data on temperature are not sufficiently firm to permit conclusions concerning its effect on disinfection. Ozone must be generated on site, and the process is relatively energy intensive. To make economic comparisons of ozone with other disinfectants, the cost of local power must be ascertained.

Available kinetic data on ozone inactivation are presented in Table IIS The c t product for 99% inactivation of E. coli appears to be approximately 0.006 at near-neutral pH values and at temperatures of <100C. There was less consistency for the c t products of poliovirus 1. These values varied from <0.005 to 0.42 at pH 7 and at temperatures from 50C to 250C.

The c t products vary over a broad range. These variations illustrate the difficulty of doing quantitative experimentation with ozone and microorganisms in water. Among other reasons, these difficulties are caused by undetected ozone demand in the water, poor analytical techniques for residual ozone, and nonuniformity of microorganisms from one laboratory to another. As an example of the latter, different strains of poliovirus with different inactivation rates are used, but the inactivation data are frequently not reported as strain-specific. Further-more, viruses frequently exist in an undetected clumped state rather than in the presumed single discrete particle state. Higher c t products are required for the clumped viruses than for the unclumped ones.

Because of ozone's relatively short half-life in water, another disinfectant must be added to maintain a disinfection capability in the

TABLE 11-5 Concentration of Ozone and Contact Time Necessary for 990/o Inactivation of. Eschenchia coli

andPoIiol Virus

Test Ozone, Contact Temperature,

Microorganism mg/liter Time, min c.ta pH 0C References

E. coli 0.07 0.083 0.006 7.2 1 Katzenelson el al., 1974

0.065 0.33 0.022 7.2 1 Katzenelson el al., 1974

0.04 0.50 0.02 7.2 1 Katzenelson el al., 1974

0.01 0.275 0.027 6.0 ii Rossefal., 1976

0.01 0.35 0.035 6.0 II Rossetal., 1976

0.0006 1.7 0.001 7.0 12 Wuhnnann and Meyrath, 1955

0.0023 1.03 0.002 7.0 12 Wuhrmann and Meyrath, 1955

0.0125 0.33 0.004 7.0 12 Wuhrmann and Meyrath, 1955

Polio 1 <0.3 0.13 <0.04 7.2 S Katzenelsoneial., 1974

0.245 0.50 0.12 7.0 24 Farooq, 1976

0.042 10 0.42 7.0 25 Evison, 1977

<0.03 0.16 <0.005 7.0 20 Sprouleial., 1978

distribution system. The most effective disinfectant, its optimum concentration, and method of addition must be determined. The disinfection process with ozone will probably be controlled by specifying the ozone residual at the beginning and end of a given contact time.

Research Recommendations

Future research studies with ozone should be conducted to:

provide data on the inactivation of enteric pathogens;

provide more data on the inactivation of protozoan cysts, especially those of Giardia lamblia;

provide analytical methods that are specific for ozone;

provide additional definitive data on bacteria and viruses, and eliminate discrepancies; and

provide operating data from full-scale drinking water plants to demonstrate reliability of operation, operating costs, ozone dosages, residuals, contact times, and disinfection results.


Chlorine dioxide (C102) was first prepared in the early nineteenth century by Sir Humphrey Davey (1811). By combining potassium chlorate (KClO3) and hydrochlonc acid (HCl), he produced a greenish-yellow gas, which he named "euchlorine." Later, this gas was found to be a mixture of chlorine dioxide and chlorine. The bleaching action of chlorine dioxide on wood pulp was recognized by Watt and Burgess


Large quantities of chlorine dioxide are produced each day in the United States. Although its primary application has been the bleaching of wood pulp, it is also used extensively for bleaching and dye stripping in the textile industry and for bleaching flour, fats, oils, and waxes (Gall, 1978).

In the United States, chlorine dioxide was first used in 1944 at the water treatment plant in Niagara Falls, New York, to control phenolic tastes and odors arising from the presence of industrial wastes, algae, and decaying vegetation (Synan et al, 1945). Granstrom and Lee (1958) surveyed water treatment plants believed to be using chlorine dioxide.

The majority of respondents (956 plants) were using it for taste and odor control. Other uses reported were algal control (7 plants), iron and manganese removal (3 plants), and disinfection (15 plants).

Sussman (1978) compiled a partial listing of plants using chlorine dioxide. He reported that the compound is used primarily to control taste and odor in the United States. In England, Italy, and Switzerland, it is used for disinfection of water supplies.

The Chemistry of Chlorine Dioxide in Water

Chlorine dioxide reacts with a wide variety of organic and inorganic chemicals under conditions that are usually found in water treatment systems (Stevens et al., 1978). However, two important reactions do not occur. Chlorine dioxide per se does not react to cause the formation of trihalomethanes CFHM's) (Miltner, 1976). However, THM's will be formed if the chlorine dioxide is contaminated with chlorine. Such a situation may occur when chlorine is used in the preparation of chlorine dioxide.

Chlorine dioxide does not react with ammonia, but will react with other amines (Rosenblatt, 1978). The amine structure determines reactivity. Tertiary amines are more reactive with chlorine dioxide than secondary amines, which, in turn, are more reactive than primary amines.

Production and Application

Chlorine dioxide condenses to form an unstable liquid. Both the gas and liquid are sensitive to temperature, pressure, and light. At concentrations above 10% in air, chlorine dioxide may be explosive, and at 4% in air, it can be detonated by sparks (Gall, 1978; Sussman, 1978). As a result, the preparation and distribution of chlorine dioxide in bulk have not been deemed practical. It has been generated and used on site.

Sodium chlorate (NaClO3) or sodium chlorite NaClO2) may be used to generate chlorine dioxide. The method of production will depend upon the amount of chorine dioxide that is required. The reduction of sodium chlorate is the more efficient process and is generally used when large volumes and high concentrations of chlorine dioxide are needed. Commercial processes that are used in North America for large-scale production of chlorine dioxide are based on the three reactions listed below. To reduce the sodium chlorate, each process uses a different

agent: sulfur dioxide (S02), methanol (CH3OH), and the chloride ion (Cl-).

2NaClO3 + H2S04 + 502 2C102 + 2NaHSO4 (6)

2NaCIO3 + CH3OH + H2504

2C102 + HCHO + Na2SO4 + 2H20 (7)

NaCIO3 + NaCI + H2S04 C102 + C12 + Na2SO4 + H20 (8)

All of these processes are used in the pulp and paper industry. They can also be used to prepare chlorine dioxide for the large waterworks that might require several metric tons per day. Small quantities of chlorine are formed during the side reactions and intermediate reactions in these processes. A more detailed review of the chemistry that is involved in the production of chlorine dioxide from chlorate is given by Gall (1978) and Gordon et al. (1972).

Chlorine dioxide can be prepared from chlorine and sodium chlorite through the following reactions:

C12 + H20 HOCl + HCl (9)

HOCI + MCI + 2NaClO2 2C102 + 2NaCl + H20 (10)

The theoretical weight ratio of sodium chlorite to chlorine is 1.00:0.39 (Dowling, 1974). With available sodium chlorite (80%), the weight ratio is 1:0.30. In practice, Gall (1978) recommended a chlorite to chlorine ratio of 1:1. The excess chlorine lowers the pH, thereby increasing the reaction rate and optimizing the yield of chlorine dioxide. Dowling (1974) reported that the maximum theoretical yield of chlorine dioxide was produced when the ratio was normally maintained at a minimum of


Alternatively, chlorine dioxide may be prepared from sodium hypochlorite @4aOCl) and sodium chlorite. The sodium hypochlorite is acidified to yield hypochlorous acid (HOCl), and the chlorine dioxide is generated according to Reaction 10. Each of the methods produces a solution containing both chlorine and chlorine dioxide.

Chlorine dioxide may also be prepared by the addition of a strong acid, such as sulfuric acid (H2504) or hydrochloric acid, to sodium chlorite as shown in the following reactions:

lONaClO2 + 5H2S04 8C102 + 5Na2SO4 + 2HCl + 4H20 (11)

SNaCIO2 + 4HCl 4C102 + 5NaCl + 2H20 (12)

Although some investigators have claimed that this method produces chlorine-free chlorine dioxide, Feuss (1964) and Schilling (1956) reported that chlorine is also formed. Dowling (1974) indicated that chlorine was formed even when sulfuric acid was used.

Analytical Methods

Chlorine dioxide is one of the few stable nonmetallic inorganic free radicals (Rosenblatt, 1978). It does not contain available chlorine in the form of hypochlorous acid or hypochlorite ion (OCl-). However, concentrations of chlorine dioxide are often reported in terms of available chlorine. The chlorine atom in chlorine dioxide has a valance of + 4. A reduction to chloride results in a gain of five electrons. In terms of available chlorine, chlorine dioxide has 263% or more than 2.5 times the oxidizing capacity of chlorine.

Se x 35A5

67.45 x100 = 263% (13)

The weight ratio of chlorine dioxide to available chlorine is 67.45 to 35.45 or 1.9. However, in water treatment practices this increased oxidizing capacity is rarely realized. The reduction of chlorine dioxide depends heavily on pH and the nature of the reducing agent. At neutral or alkaline pH, chlorine dioxide is reduced to chlorite, a net gain of one electron. Thus, only one-fifth or 20% of its oxidizing capacity is utilized. At low pH, the chlorite (C102-) is reduced to chloride (Cl-) releasing the remaining four available electrons.

Chlorine dioxide may be determined iodometrically (Standard Methods, 1976), amperometrically (Hailer and Listek, 1948; Standard Methods, 1976), spectrophotometrically (Gordon et al., 1972), and colorinietrically (Aston, 1950; Hodgen and Ingols, 1954; Masschelein, 1966; Palin, 1957, 1960, 1967, 1970, 1974, 1975; Post and Moore, 1959; Standard Methods, 1976).

Several studies contain comparisons of various analytical methods and procedures for the measurement of chlorine dioxide. Adams el al. (l%6) compared the H-acid tyrosine, amperomeiric, and diethyl-p-phenylenediamine (DPD) procedures for free chlorine, chlorine dioxide, and chlorite. They reported DPD to be the most reliable, as did Dowling (1974), Mutner (1976), and the U.S. Public Health Service (Mutner, 1976). Myhrstad and Samdal (1969) noted that the DPD method (Palin, 1957, 1960) yielded consistently higher residual measurements for chlorine dioxide than those that are produced with other analytical methods. Mter analysis with acid chrome violet K (Masschelein, l%6), chlorine dioxide was not observed in the water of the distribution system; however chlorite was found. The residuals that were previously interpreted as chlonne dioxide were apparently due to chlorite.

Recently, more sophisticated procedures were suggested for the determination of chionne dioxide. Moffa et al. (1975) reported the use of electron spin resonance and Issacsson and Wettermark (1976) described a chemilurninescent method for active chlorine compounds. Stevenson et al. (1978) reviewed electrochemical methods and presented preliminary results for a membrane amperometric probe. Under development is a sensor that shows a linear response region from about 0.5 to 10 mg/liter. The response to hypochiorous acid and chloramines was low, and the sensor does not measure chlorite or other ionized species.

No one procedure appears to possess the necessary sensitivity, selectivity, and simplicity to permit reliable determinations in the treatment of water. Each of the titration methods are prone to error because of volatilization. They are time-consuming and particularly complex when differentiation of chlorine and oxychloro species are necessary. The colorimetric procedures require strict control of pH, temperature, and reaction times and will be affected by turbidity. In addition, the selectivity of the indicators for chlorine dioxide is questionable. The direct spectrophotometric determination of chlorine dioxide at 360 nm is selective and rapid but is not sufficiently sensitive for use in water. Limited experience with the more recent procedures (chemiluminescence and membrane amperometric probe) does not permit an evaluation.

In practice, the principal distinction that must be made is that between the active biocidal species (hypochlorous acid, the hypochloride ion, and chlorine dioxide), the moderately biocidal species (monochloramine [NH2CIJ, dichloramine [NHCI2], and nitrogen trichloride [NCl3]), and the relatively nonbiocidal species (chlorite and chlorate [C103-) ions). This is imperative when the primary purpose for the addition of chlorine dioxide is the inactivation of microorganisms. Furthermore, when the

formation of THM's is to be considered, the distinction between free chlorine and chlorine dioxide becomes important.

Bioeidal Activity


Experimental data on the efficacy of chlorine dioxide as a disinfectant became available in the early 1940's. McCarthy (1944, 1945) reported that chlorine dioxide was an effective bactericide in water with a low organic content. When the levels of organic material in the water were high, chlorine dioxide was less effective.

Ridenour and Ingols (1947) reported that chlorine dioxide was at least as effective as chlorine against Escherichia coli after 30 min at similar residual concentrations. Both chlorine and chlorine dioxide residues were determined by the orthotolidine-arsenite (OTA) method. The bactericidal activity of chlorine dioxide was not affected by pH values from 6.0 to 10.0. Ridenour and Armbruster (1949) extended their observation to other enteric bacteria. The common waterborne path~ gens were similarly inactivated with chlorine dioxide. They also reported that the efficiency of chlorine dioxide decreased as the temperature decreased from 200C to 50C.

Ridenour et aL (1949) found chlorine dioxide to be more effective than chlorine ~ased on OTA residuals) against bacterial endospores. They indicated that less weight of chlorine dioxide than chlorine is required to inactivate the spores of Bacillus mesentericw~ in either demand-free water or in waters contaii~ing ammonia. In the waters containing ammonia, chlorine had to be applied beyond breakpoint before efficient sporicidal activity was observed.

The work of Ridenour and colleagues is not discussed in depth because the small amounts of free chlorine that are produced during the generation of chlorine dioxide are not distinguished from the chlorine dioxide by the OTA method that they used for both stock solutions and residual determination. In their 1947 paper (Ridenour and Ingols, 1947), the survival measurements were + and not quantitative. In their 1949 paper (Ridenour and Armbruster, 1949), the survival measurement was quantitative, but only one contact time was used, 5.0 mm. Since chlorine dioxide is a rapidly acting disinfectant, c t products may be misleading using this contact time.

Russian investigators (Bedulevich et al., 1953; Trakhtman, 1946) found chlorine dioxide to be a more effective or at least as effective a bactericide as chlorine. Additional data were reported for Bacillus

anthracis. They observed that the efficiency of chlorine dioxide decreased as the pH of the system containing the B. anthracis increased.

Early studies on disinfectant activity are difficult to interpret because the methods of preparing chlorine dioxide invariably included the addition or production of chlorine. Analytical procedures were not sufficiently advanced to differentiate between chlorine dioxide and other oxychloro species. Thus, the quantitative analyses of stock solutions and reports of dose and residual chlorine dioxide may be in error. This suggests that the initial and residual concentrations of chlorine dioxide were probably lower than reported values and that the comparative bactericidal efficiency would suffer accordingly. In addition, the older investigations did not take into account the volatility of chlorine dioxide. Depending upon concentration and length of exposure, losses of 7% to 30% can occur within an hour.

Many of the difficulties that were encountered during the early studies were overcome in a series of studies reported by Benarde and co-workers during the mid- 1960's. Their work on disinfection was based heavily on the improved methods of preparing and analyzing chlorine dioxide, which were reported by Granstrom and Lee (1958). They prepared chlorine dioxide by oxidizing sodium chlorite with persulfate (S2O8~ under acid conditions. The resulting chlorine dioxide was swept to a collection vessel by high purity nitrogen gas. Chlorine dioxide was measured spectrophotometrically at 357 nm.

Bernarde el al. (1965) compared the bactericidal effectiveness of chlorine with that of chlorine dioxide at pH 6.5 and 8.5 in a demand-free buffered system. At pH 6.5, both chlorine and chlorine dioxide inactivated a freshly isolated strain of E. coli in less than 60 5. Chlorine was slightly more effective at the lower dosages at the lower pH. At pH

8.5, chlorine dioxide was dramatically more effective than chlonne. Greater than 99% inactivation of E. coli was observed in 15 s with 0.25 mg/liter chlorine dioxide, while, under the same conditions, chlorine required almost S min for similar inactivation. Chlorine dioxide was significantly more efficient than chlorine in the presence of high levels of organic and nitrogenous material. Bernarde et al. (1967a) also reported that temperature affects the rate of inactivation of bacteria with chlorine dioxide. A decrease in disinfectant activity was observed as temperature decreased from 300C to 50C. For 99% inactivation of E. coli with 0.25 mg/liter of chlorine dioxide, 190, 74, 41, and 16 s were required at 50C, 100C, 200C, 300C, respectively. More recent work by Cronier (1977) in a clean system also demonstrated the excellent bactericidal activity of chlorine dioxide. Results of both studies are shown in Table 11-6.

TABLE 11-6 Concentrations of Chlorine Dioxide and Contact Times Necessary for 99% Inactivation of Escherichia coli

Chlorine Contact Temper

Test Dioxide, Time, ature,

~eroorganism rng(liter C.ta pH 0C

E. coli 0.25 1.8 0.45 6.5 S

(freshly 0.50 0.83 0.41 6.5 S

isolated 0.75 0.50 0.38 6.5

fromfeces)b 0.25 1.2 0.30 6.5 10

0.50 0.47 0.24 6.5 10

0.75 0.3 0.23 6.5 10

0.25 0.68 0.17 6.5 20

0.50 0.35 0.18 6.5 20

0.75 0.25 0.19 6.5 20

0.25 0.27 0.07 6.5 32

0.50 0.22 0.11 6.5 32

0.75 0.15 0.11 6.5 32

E. coli 0.30 1.8 0.54 7.0 S

(ATCC 11229)C 0.50 0.98 0.49 7.0 S

0.80 0.58 0.41 7.0 S

0.30 1.3 0.39 7.0 15

0.50 0.75 0.38 7.0 15

0.80 0.47 0.38 7.0 15

0.30 0.98 0.29 7.0 25

0.50 0.55 0.28 7.0 25

0.80 0.35 0.28 7.0 25

a Concentration of chlorine dioxide times oontact time. b Data from Bernarde et aL, 1967a.

Data from Cronier, 1977.


In the mid-i 940's, there were also investigations on the virucidal activity of chlorine dioxide. Ridenour and Ingols (1946) reported that chlorine dioxide was as effective as chiorme against a mouse-adapted strain of poliovirus (Lansing). Again, their comparison was based upon OTAdetermined residual levels of each disinfectant. Hettche and Ehlbeck (1953) found chlorine dioxide to be more effective against poliovirus than either chlorine or ozone. In addition to the difficulties that are associated with the preparation and measurement of chlonne dioxide

and chlorine, the early virus studies were also saddled with difficult and time-consuming virus test systems.

The first definitive work on the virucidal activity of chlorine dioxide was done in G6teborg, Sweden. Warriner (1967) showed that the rate of inactivation of poliovirus 3 increased with increasing pH at pH values of 5.6 to 8.5 (see Table 11-7). Similar to the action on bacteria, the viral inactivation occurred rapidly. When chlorine and chlorine dioxide were combined, the inactivation was Synergistic. The inactivation with the two chlorine species together was more efficient at lower pH, but the presence of chlorine dioxide enhanced inactivation by chlorine at the higher pH.

Cronier (1977) compared the inactivation of E. COli (ATCC 11229), poliovirus 1, and coxsackievirus A9. All of the microorganisms that were tested were sensitive to low concentrations (<I mg/liter) of chlorine dioxide. Poliovirus I and coxsackievirus A9 were more resistant than E. coli. Similar to its bactericidal activity, chlorine dioxide was more effective as a virucide at higher pH. Cronier also reported that on a weight basis, it was similar to hypochlorous acid and better than hypochlorite ion, monochloramine, and dichloramine.

The effect of bentonite (Al03 45i02 . H20) (added to the test as a model for turbidity) on disinfection with chlorine dioxide was studied by Scarpino et al. (1977). At turbidity levels below 2.29 nephelometric turbidity units (NTU) at 250C, no protection was afforded to poliovirus 1 by the bentonite. However, at 3.22 and 14.1 NTU's, poliovirus inactivation was noticeably decreased. The bentonite at these levels appeared to offer protection to the virus.


The data that are available on the efficacy of chlorine dioxide on helminths or protozoan cysts do not appear to be suitable for comparison with the action of other disinfectants.


The data presented in Tables 11-6 and 11-7 were collected for demand-free systems within the last 15 yr, when relatively reliable chemical methods and/or quantitative biocidal assay procedures were used. Table

11-6 shows the concentration of chlorine dioxide necessary for 99% inactivation of E. coli strains. Even at 50C and chlorine dioxide concentrations of 0.25 to 0.30 mg/liter, less than 2 min was required for 99% inactivation. Increases in chlorine dioxide concentration and/or

TABLE 11-7 Concentrations of Chlorine Dioxide and Contact Times Necesary for 99% Inactivation of Viruses

Chlorine Contact Temper-

Test Dioxide, Time, ature,

Microorganism mg/liter min C.ta pH oC

P0h.0virus3b 0.5 5.0 2.5 5.6 20

0.5 5.0 2.5 7.2 20

0.5 0.25 0.125 8.5 20

1.6 5.0 8.00 5.6 20

1.6 1.0 1.60 7.1 20

1.6 0.25 0.40 8.0 20

Poliovirusic 0.3 16.6 5.0 7.0 S

0.5 12.0 6.0 7.0 S

0.8 6.8 5.4 7.0 S

0.3 4.2 1.3 7.0 15

0.5 2.5 1.25 7.0 15

0.8 1.7 1.4 7.0 15

0.3 3.6 1.08 7.0 25

0.5 2.0 1.0 7.0 25

0.8 1.5 1.2 7.0 25

Coxsackievirus 0.3 1.2 0.4 7.0 15

A9 0.5 0.05 0.34 7.0 15

0.8 0.25 0.20 7.0 15

Concentration of chlorine dioxide times contact time. b Data from Wairiner, 1967~

Data from Cronier, 1977.

temperature markedly reduced the contact time that was necessary. Table 11-7 presents similar data for viruses. Increased chlorine dioxide concentration, temperature, and pH decreased the contact time that was required to produce 99% inactivation of the viruses. The amount of inactivation depended on the virus that was tested.

Mechanism of Action

There is little information concerning the mode of action by which chlorine dioxide inactivates bacteria and viruses. Ingols and Ridenour (1948) suggested that the bactericidal effectiveness of chlorine dioxide is due to its adsorption on the cell wall with subsequent penetration into the cell where it reacts with enzymes containmg sulfhydryl groups.

Benarde et al. (1967b) demonstrated that chionne dioxide abruptly inhibited protein synthesis. The incorporation of 14C-labeled amino acids into protein by whole cells stopped within a few seconds affer the addition of chlorine dioxide. Subsequently, Olivieri (1968) reported a dose-response in the inhibition of protein synthesis in bacteria that had been treated with chlorine dioxide. The site of action was localized in the soluble portion (enzymes) of the cell extracts of treated cells without affecting the integrity of the ribosomes' function in protein synthesis.


Chlorine dioxide is an effective bactericide and virucide under the pH, temperature, and turbidity that are expected in the treatment of potable water. It should be noted that the U.S. Environmental Protection Agency (1978) has set an interim Maximum Contaminant Level of 1 mg/liter on chlorine dioxide (U.S. Environmental Protection Agency, 1978) because of the unresolved questions on its health effects (Symons et aL, 1977).

Research Recommendations

A simple, selective, and sensitive test for chlorine dioxide should be developed to monitor residual concentration.

Chlorine dioxide is currently used at several plants. A review of plant records and field studies on the stability and effectiveness of chlorine dioxide in the distribution system should be undertaken.

Studies should be directed toward evaluating the inactivation of protozoan cysts.

More information should be obtained on the mode of action by which chlorine dioxide inactivates bacteria, viruses, and cysts.


The use of iodine as a biocide has had a long history, primarily as an antiseptic for skin wounds and mucous surfaces of the body and, to a lesser degree, as a powerful sanitizing agent in hospitals and laboratories (Gershenfeld, 1977). The use of iodine as a disinfectant of drinking and swimming pool water has not been extensive, mainly because of the costs and problems that are involved in applying the dosage.

Aside from the emergency iodination of small volumes for field and emergency drinking water and limited experience with swimming pool

TABLE 11-8 Hydrolysis of Iodine at 2S~C Showing the Percentage of Iodine Species at Different PH'sa

Percentage of Iodine Species by Concentration of Iodine

0.5 mg/liter 5.0 mg/liter 50.0 mg/liter

pH 12 HOI or 12 HOI OF 12 HOI OP

S 99 1 0 100 0 0 100 0 0

6 90 10 0 99 1 0 100 0 0

7 52 48 0 88 12 0 97 3 0

8 12 88 0.005 52 48 0 86 14 0

a Data from chang, 1958.

disinfection (Black et aA, 1959), the only substantial experience with iodine disinfection of piped water system is that of Black et al. (l%5, 1968). Two water systems serving three prisons in the state of Florida were disinfected satisfactorily with 1 mg/liter of iodine. A persistent residual was maintained throughout the distribution system despite a finished water pH of 8.0 to 9.5. No adverse effects on health were observed among those consuming the water.

Chemistry of Iodine in Water

Iodine is the only common halogen that is a solid at room temperature, and it possesses the highest atomic weight (126.91). Of the four common halogens, it is the least soluble in water, has the lowest standard oxidation potential for reduction to halide, and reacts least readily with orgamc compounds.

'2 + H20 HOI + H+ + I- (14)

Diatomic iodine (12) reacts with water to form hypoiodous acid (HOI) and iodide ion (I-). The effect of pH on this reaction is shown in Table


The distribution of chemical species of iodine given in Table 11-8 was taken from the calculations that were made by Chang (1958) from the equilibrium expression:

- Kh



The value of Kb, the hydrolysis constant, is given by Wyss and Strandskov (1945) as 3 X 10-12 at 250C. With iodine residuals at 0.5 mg/liter, which are expected in water Systems, and a pH of 5, approximately 99% of the total iodine residual is present as iodine and only 1% as hypoiodous acid. At pH 7, the two forms are present in almost equal concentrations. At pH 8, only 12% is present as elemental iodine and 88% as hypoiodous acid, which can be converted to hypoiodite ion (OI~).

HOI H+ + or (16)

fH+~[OI~] K = 45xlO~13

[HOI] a (17)

[H+] - 'Ca [HOI] (18)


[HOVl H+

[OI~] 'Ca


The dissociation constant, Ka~ of hypoiodous acid (at 200C) is 4.5 X 10-13. Consequently, the dissociation of hypoiodous acid, which occurs at high pH's, is not important for practical purposes. However, as confirmed by the field studies in Florida, hypoiodous acid can form iodate ion by autooxidation at pH values above 9.

3H01 + 20H H013 + 2H20 + 2r (20)

The iodate ion possesses no disinfecting ability (Marks and Strandskov, 1950).

Production and Application

Iodine may be added to a municipal water supply by several procedures. One method is to employ nonhazardous solvents and solubil~ng agents such as ethyl alcohol (C2H5OH) and potassium iodide (KI) to overcome the low concentration of aqueous iodine stock for solution feeders. Another method produces the required concentration of iodine by passing water through a bed of crystalline iodine (saturator). This has

been used in many small, semipublic and private home water Systems. Since the maximum concentration is limited by solubility to 2O(~~300 mg/liter at the ambient temperatures that are expected for drinking water, some physical complications would accompany the introduction of this method into the large waterworks system in view of the large saturation beds required. The iodinated anion exchange resin bed and the vaporization technique are not sufficently developed to be considered for use in public water supplies.

In certain circumstances, potassium iodide might be combined with an oxidation reaction to release iodine. The chemistry is simple, and the persistence of the iodine that is generated may be much better than chlorine. This method has been used in swimming pools and for dechlorination purposes where a chlorine residual may be exchanged for an iodine residual, or an iodine residual may be provided where ozone is the primary disinfectant.

Analytical Methods

Both amperometric titration and leuco crystal violet (LCV) colorimetric methods give acceptable results when used to measure free iodine in drinking water. Waters containing oxidized forms of manganese interfere with the LCV method. Also, when iodide ion exceeds 50 mg/liter and chloride ion exceeds 200 mg/liter, the amperometric method is preferred. Under unusual situations, where mixtures of chlorine, bromine, and ozone occur along with iodine, the problem of separation is difficult (Stand£~rd Methods, 1976).

Biocidal Activity

Table 11-9 shows the relative resistance of bacteria, viruses, and cysts to inactivation by iodine.


Chang and Morris (1953), summarizing the development of a universal water disinfectant tablet for the military, concluded that iodine concentrations of S to 10 mg/liter were effective against bacterial pathogens. Iodine was less dependent than chlorine upon the pH, temperature, contact time, and secondary reactions with nitrogenous impurities in the water. As a cysticide, iodine was poor in water with a high pH. Consequently, the tablet that was formulated contained an acid buffer to lower the pH of the water. The tablet, called "Globaline," released 8

TABLE 11-9 Comparative Values from Confirmed

Experiments on Disinfecting Water with Iodine at 230C to

300C, at pH 7.0~


Iodine for 99~/o

Organism mg/liter Inactivation c.tb

Coliform 0.4 1 0.4


Poliovirus 1 20 1.5 30.0

f2 Virus 10 3.0 30.0

Simiancvsts 15 10~0 150.0

a References to these values appear in the text.

b Concentration of iodine times contact time.

mg/liter of iodine, which is extremely high compared to the amount that is possibly needed for public water supplies. The tablet is not widely accepted, since color, taste, and odor problems are fairly common (O'Connor and Kapoor 1970).

Chambers el a!. (1952) investigated the effect of iodine concentration, pH, exposure time, and temperature on 13 enteric bacteria in a clean system. Their results are not completely quantitative in that the reported iodine concentration is that required to inactivate all test organisims plated out affer 1, 2, and s mm of contact. This was equivalent to approximately 99.9% activation with the procedure that they followed. There was definitely an observed pH effect. At 20C to 50C, some bacterial species required 3 to 4 times as much iodine for similar inactivation at pH 9.0 as was required at pH 7.5. For S min of exposure, the required dosage for >99.9% bacterial inactivation in high pH water at 20QC~26cC was always less than I mg/liter. A summary of their work with E. coli is given in Table 11-10. This is the best available information on iodine as a bactericide.


Studies on the efficacy of iodine on viruses have shown that viruses are more resistant to disinfection than are vegetative cells of bacteria. Kruse (1969) compared virus and bacterial inactivation. At an iodine dose of 10 mg/liter in 0.048 mM potassium iodide at room temperature and pH 7.0,

TABLE 11-10 Concentrations of Iodine and Contact

Times Necessary for 99% Inactivation ofEscherichia coli~

Contact Temper

Iodine, Time, ature

mg/liter mm c.tb pH 0c

1.3 1 1.3 6.5 2-5

0.9 2 1.8 6.5

1.3 1 1.3 7.5

0.7 2 1.4 7.5

0.8 1 0.8 7~5

0.6 2 1.2 7.5

0.8 1 0.8 8.5

0.9 2 1.8 8.5

1.8 1.8 9.1

1.2 2 2.4 9.1

0.35 0.35 6.5 2O~25

0.20 2 0.40 6.5

0.45 1 0.20 7.5

0.30 2 0.60 7.5

0.45 1 0.45 8.5

0.40 2 0.80 8.5

0.45 1 0.20 9.1

0.30 2 0.60 9.1

a Data from chambers e: a'., 1952.

b Concentration of iodine times oontact time.

E. coli was inactivated more rapidly than ~ virus but both were reduced by 4 logs in less than 1 mm. However, when the potassium iodide concentration was raised to 0.5 M, bacterial inactivation was 4 logs in 1 min while virus inactivation was only 0.5 log in 1 hr ~igure 11-6).

Berg et al. (1964) measured the dynamics of survival for poliovirus, coxsackievirus, and echovinis that were iodinated at pH 6.0 and at temperatures of 50C, 150C, and 250C. The work of the Johns Hopkins group from 1962 to 1969 (Kuuse', 1969) was primarilv with the virus model f2, although some comparative poliovirus 1 work was done. Cramer et al. (1976) have shown that the mode of inactivation of these two viruses by iodine is similar. Data on the inactivation of virus by iodine are summarized in Table 11-11 for both polio and f2. The results of inactivation studies by the various groups compare very favorably. At pH 6.0, iodine approaches the order of magnitude of virucidal activity of

The Disinfection of Drinking Water 67


0.04 mM Iodine

0.048 mM K). pH 7.0

0.04 mM Iodine

0.5 M K), pH 6.3 1.0 mM PCMB. pH 8.0





- Phage

-- E.

Sc 10~ C,,





1 2 3 0 10 20 30 40 50 60 C







I I I I I I 10 20 30 40

50 -~ ~60

CONTACT TIME (minutes)

Figure 11-6 Survival of E. coli and f2 bacteriophage reacted with 0.04 mM iodine containing (a) 0.048 mM potassium iodide (KI) at 370C and ~) 0.5 TnAg potassium iodide at 00C, and survival obtamed with (c) 1 mM of a Su'fhydryl reacting agent, ~ chioromercurihenzoic acid (C1H~H4COOH). From Knise', 1%9, with permission.

hypochiorous acid (HOCi). At the pH's likely to be maintained in the distribution system (pH 8.0), iodine is a vastly more effective virucide than combined chlorine and is not far removed from the activity range of free chlorine.


Many studies on cyst inactivation have been reported, but there are discrepancies in their results due mainly to differences in the test Systems that were used. The dose, pH, and temperature that were used in many of the studies are in doubt as are the different sources, cleaning methods for cysts, number of cysts used per test, and determination of viability. Stringer (1970) compared the resistance to iodine disinfection of Entamoeba histolytica cysts obtained from in-vitro culturing with those harvested from a human camer and mixed amoebic cysts from monkeys. In water at room temperature at pH 6.5, 3 times the iodine dose was required in 10 mm of exposure for 99% inactivation with wild (naturally formed) cysts compared to the cultured cysts. Large numbers of amoebic cysts from simian hosts were more readily available than cysts from

TABLE Il-il Concentrations of Iodine and Contact Times Necessary for 990/o Inactivation of Polio and f~ Viruses

with Flash Mixing

Test Iodine, Contact Tempera-

Microorganism mg/liter Time, min C.ta p11 ture,0C References

f2 Virus 13 10 130 4.0 s Kruse', 1969

f2 Virus 12 10 120 5.0 Kruse, 1969

f2 Virus 7.5 10 75 6.0 Kruse, 1969

f2 Virus s 10 50 7.0 Knise', 1969

f2 Virus 3.3 10 33 8.0 Kruse, 1969

f2 Virus 2.7 10 27 9.0 Kruse, 1969

f~ Virus 2.5 10 25 10.0 Kruse', 1969

f2 Virus 7.6 10 76 4.0 25-27 Kruse', 1969

Poliovirus 1 30 3 90 4.0 Kruse', 1969

f2 Virus 64 10 64 5.0 Kruse', 1969

f2 Virus 4.0 10 40 6.0 Kruse', 1969

Poliovirus 1 1.25 39 49 6.0 Berg et al., 1964

Poliovirus 1 6.35 9 57 6.0 Berg etal., 1964

Poliovirus 1 12.7 s 63 6.0 Berg et al., 1964

Poliovirus 1 38 1.6 60 6.0 Bergelal., 1964

Poliovirus I 30 2.0 60 6.0 Cramer el al., 1976

f2 Virus 3.0 10 30 7.0 Knise', 1969

Poliovirus 1 20 1.5 30 7.0 Kruse, 1969

f2 Virus 2.5 10 25 8.0 Knise', 1969

f2 Virus 2.0 10 20 9.0 Knise', 1969

f2 Virus 1.5 10 15 10.0 Knise', 1969

Poliovirus 1 30 0.5 15 10.0 Cramer el al., 1976

human stools. They served as a reliable model for human E. histolytica. The most extensive and reliable data on the cysticidal properties of

iodine are found in the work of Chang and Stringer and their co-workers (Change and Baxter, 1955; Chang el al., 1955; Stringer, 1970; Stringer and Kiuse, 1971; Stringer et aL, 1975). The only quantitative comparative study of the cysticidal properties of chlorine, bromine, and iodine believed to be published is that of Stringer, who reported a 99.9% inactivation. The earlier studies of others involved "total kills" and are quite dependent on the cyst density that was used.

All investigators more or less agree that iodine is an excellent cysticide in low pH waters (<pH 4.0). suggesting that molecular iodine (12), rather than hypoiodous acid, is the active agent. Chang (1958) observed that when the titratable iodine dosages exceed 20 mg/liter in the presence of iodide ion, the triiodide ion (l~ has a cysticidal efficiency that is equal to 1/11, 1/8, and 1/7 that of 12 at 60C, 250C, and 350C. He further calculated relative cysticidal efficiency of hypoiodous acid to be one-third that of 12 at 60C and one-half that of 12 at 200 C.

There may be the problem of ineffective iodine residual in the distribution system where cross-connection introduction of cyst contarnination is a possibility. This is due to the practice (for corrosion control) of maintaining water in the distribution system at approximately pH 8.0. At the low halogen residual levels that are usually maintained, the most cysticidal form, 12, would be less prevalent than hypoiodous acid. Water at 250C and pH 8.0 with a total iodine residual of 0.5, 1.0, and 2.0 mg/liter would contain only 0.06, 0.2, and 0.62 mg/liter of molecular iodine. Table 11-12 shows the time in rninutes required for 90% inactivation of simian cysts with residuals of bromine, chlorine, and iodine in water at pH 6.0, 7.0, and 8.0 and a temperature of 300C. At pH

8.0, iodine is inferior to free bromine and chlorine. However, in the presence of excess ammonia, with which the halogens could react, bromine appears to be more effective at pH 8.0 than do similar dosages of iodine or chlorine. Residual iodine in water at pH 8.0 has the ability to persist.

The kinetics of iodine as a cysticide in water at 200C to 300C have

been assembled and compared. The cultured E. histolytica cyst data of

Chang and co-workers (Chang and Baxter, 1955; Chang and Morris,

1953; Chang et aL, 1955) and the results of Stringer et al. (1975) present a

consistent picture. For cultured cysts in water at pH S the c t values for

99% inactivation calculated by Chang were 200 at 30C, 130 at 100C, and

65 at 230C (Chang and Morris, 1953). Stringer et al. (1975) extended the

cultured cyst data at 300C for a range of pH given below:


TABLE 11-12 Contact Times Necessary for Low Residual Bromine, Chlorine, and Iodine in Water at 300C to Effect 990/0 Inactivation of Cysts from Simian Stoolsa

Contact Time Required, mm, by 10.min Ilalogen Residual Concentration, mg/liter

0.5 mg/liter 1.0 mg/liter 2.0 mg/liter

p11 Br2 C12 12 Br2 C12 12 Br2 C12 12

Bulferedwater 6.0 10 10 20 4 4 10 3 3 s

7.0 12 14 40 S 12 20 4 s 7

8.0 15 20 NDb 10 15 80 S 10 20

Bufferedwater 6.0 10 65 20 S 35 10 4 22 S

inpresenceof 7.0 30 120 40 10 55 20 7 35 7

excessammonia 8.0 35 NDb NDb 13 80 80 9 50 20

a Data from Stringer et al., 1975. The proportions of the molecular species present in the test system depend upon the halogen, the pH, and the presence or absence of ammonia.

b ND not determined.

pH c t for 99% inactivation

6.0 40

7.0 50

8.0 100

8.5 200

These c t values from Stringer et al. (1975) were approximately one-half the values they obtained with E. histolytica cysts from human stools and for mixed cysts from simian stools.

Mechanism of Action

The failure of a strong commercial iodine disinfectant to inactivate poliovirus (Wallis et al., 1963) led to interest in the responsible mechanism, especially the role of pH and iodide ion. Berg et aL (1964) claimed that iodine inactivation of coxsackievirus resulted from biomolecular reaction with a single iodine molecule and that clumping of virions played a role in resistance to disinfection. Fraenkel-Conrat (1955) pointed out inconsistencies in the literature regarding the virucidal properties of iodine. Hsu (1964) and Hsu et aL (1%5) clarified the mechamsm of iodine on cells and virus. Hsu extracted fully active transforming DNA from iodine-inactivated Haemophil~ parainfluenzae cells and just as much infectious RNA from f2 bacterial virus that had been inactivated (5 logs) by iodine as from noniodinated controls. Iodine inactivation, unlike the action of chlorme, appears to be attributable to a reaction with vital amino acids in proteins. Further experiments were conducted to determine whether sulfhydryl, tryptophanyl, histidyl, or tyrosyl groups were involved. With bacterial cells, there was a striking similarity between the kinetics of inactivation with iodine and the application of p-chloromercuribenzoic acid, a sulfhydryl reacting agent. While the active chemical species of iodine is not known, Hughes (1957), Allen and Keefer (1955), Bell and Gelles (1951), and Hsu (1964) hypothesized that the hydrated cationic iodine species (H201 +) attacks the base, first against the sulfhydroxyl groups, and is not materially affected by the presence of iodide ion as was tyrosine. When sullhydroxyl groups are the site of inactivation, a low pH should favor the reaction. With the viruses, the sulfliydroxyl group is not involved. Evidence of tyrosine's involvement came from parallel experiments showing similar patterns of inactivation curves between iodination of f2 virus and Ltyrosine. Li (1942, 1944) had shown that the presence of iodide ion and low pH iodine (12) inhibited the iodination of tyrosine. Both poliovirus and f2 virus inactivation with iodine was inhibited by the iodide ion

(Cramer et al., 1976). At pH 4.0, both polio and f2 viruses survived jodination whereas at pH 10 inactivation was complete. Although iodine decomposes rapidly to jodate (IO3~ and iodide at pH 10, flash mixing of iodine ~ielding hypoiodous acid) overcomes this difficulty effectively, and the virus inactivation is complete in less than I min (Longley, 1964). Therefore, there is evidence that iodine action, with little or no iodide present, is effected by the modification of protein without destroying DNA or RNA (Brammer, 1963; Hsu, 1964). The mode of action of iodine in cyst peneiration and inactivation has not been studied.


Iodine has many features that are comparable to free chlorine and bromine as a water disinfectant, but iodamines are not formed. Free iodine is an effective bactericide over a relatively wide range of pH. Field studies on small public water systems have shown that low levels of 0.5 to I mgliter of free iodine can be maintained in distribution systems and that the magnitude of residual is sufficient to produce safe drinking water with no adverse effects on human health. Like other halogens, the effectiveness of iodine against bacteria and cysts is significantly reduced by high pH, but unlike bromine and chlorine it is much more effective agamst viruses because of the enhanced iodination of tyrosine. Currently, its use is restricted primarily to emergency disinfection of field water supplies because of its high cost and because it is difficult to apply to large systems. The possible adverse health effects of increased iodide intake for susceptible individuals in the population must also be considered.

Research Recommendations

Studies should be conducted to determine the consequences for human health of the long-term consumption of iodine in drinking water with special regard for more susceptible subgroups of the population.


Chemistry of Bromine in Water

Bromine was first applied to water as a disinfectant in the form of liquid bromine (Br2) (Henderson, 1935), but it can also be added as bromine

chloride gas (BrC1) (Mills, 1975) or from a solid brominated ion exchange resin (Mills, 1969). Oxidation of bromide (Br~) to bromine can also be accomplished either chemically or electrochemically. Oxidation with aqueous chlonne gives either bromine or hypobromous acid (HOBr), depending on the ratio of chlorine to bromide. Both bromine (Liebhafsky, 1934) and bromine chloride (Kanyaev and Shilov, 1940) hydrolyze to hypobromous acid:

Br2 + H20 HOBr + + Br (21)

B~l + H20 HOBr + + C~ (22)

Molecular bromine exists in water at moderately acid pH and liigh bromide concentrations since the equilibrium constant of Rea~tion 21 is 5.8 x 10~ at 250C. Like chionne, bromine chionde has a much higher hydrolysis constant than this, so it does not exist as the molecular form in appreciable concentrations under conditions of water treatment. The ratio of molecular bromine to hypobromous acid depends on both pH and bromide concentration. From the equilibrium expression for Reaction 21:


log (HOBr) = 1og(Br~) - pH +8.24 (23)

Thus, for a solution containing 10 mg/liter bromide and a pH of 6.3, 1% of the bromine is Br2, while at lower bromide concentrations or higher pH aqueous bromine occurs almost entirely as hypobromous acid, which is a very weak acid with a dissociation constant (for Reaction 24) of 2 x l0~ at 250C (Farkas and Lewin, 1950).

HOBr~H+ + OBr (24)

The hypobromite ion (OBr~) becomes the major form of bromine above pH 8.7 at 250C. Lower temperature decreases both of the above equilibrium values, thereby increasing the pH range where hypobromous acid is the major chemical form of bromine in water.

Bromine and bromine chloride also react with basic nitrogen compounds to form combined bromine or bromamines (Galal-Gorchev and Morris, 1965; Johannesson, 1958; Johnson and Overby, 1971):

HOBr + NH3 NH2Br + H20 (25)

NH2Br + HOBr NHBr2 + H2O (26) NHBr2 + HOBr NBr3 + H20 (27)

The observed breakpoint for ammonia ~H3) solutions that have been treated with bromine is similar to that seen with chlorine. For bromine, this point corresponds to 17 mg/liter bromine for 1 mg/liter of ammonia nitrogen ~H3-N). At this point, a minimum of bromarnine stability occurs (Inman el al., 1976) as ammonia nitrogen is oxidized to mtrogen gas:

3HOBr + 2NH3 N2 + 3HBr + 3H2O (28)

At bromine-t~ammonia ratios higher than thls and in the acid pH range, nitrogen tribromide ~Br3) is stable. It is the most abundant bromamine in such aqueous solutions. At lower bromine concentrations, dibromamine ~HBr2) can be present, but it is quite unstable. Only at alkaline pH values and very high ammonia concentrations such as those found in wastewater are significant quantities of monobromamine ~H2Br) formed. Organic bromamines are also formed, but there is little information on their forms or stability.

Production and Application

Bromine is produced by oxidation of bromide-rich brines (that contain between 0.05% and 0.6% bromide) with chlorine. Bromine is then stripped with steam or air and is collected as liquid Br2.

Bromine chloride is produced by m]"wung equal molar quantities of pure bromine and chlorine ~ills, 1975). It condenses to liquid bromine chloride below 50C at 1 atm pressure or above 30 psig at 250C. In the liquid phase, more than 80% of the liquid is bromine chloride, and the remainder is Br2 and C12. In the gas phase, 40% of the bromine chloride dissociates to Br2 and C12 at 250C.

Bromine has been applied to water as liquid Br2. The difficulties that are encountered when handling the liquid and its corrosive nature, especially when wet, have encouraged the use of bromine chloride. Liquid bromine chloride is removed from cylinders under moderate pressure. It is then vaporized, and the gas is metered in equipment that is

similar to that used for chlorine. Like chlorine, bromine chloride is shipped as the dry liquid in steel containers. Gas feeders must be made of Teflon, Kynar, or Viton plastics because bromine cliloride is more reactive than chlorine with polyvinylchloride plastics.

Analytical Methods

Bromine concentrations can be measured iodometrically by procedures that are identical to those used to measure total chlorine residuals (Standard Methods, 1976). None of these procedures is capable of distinguishing free bromine (Br2, HOBr, OBr~) from combined bromine Q)romamines) or other oxidants that are capable of reacting with iodide under the slightly acid conditions used in these procedures. However, bromate does not interfere except at low pH (Kolthoff and Belcher, 1957). UV spectroscopy can be used to measure the brurnammes selectively in the presence of one another and without interference from free bromine because none of the free bromine forms except Br3 has strong UV absorptivities (Johnson and Overby, 1971).

Biocidal Activity

Since residual bromine was rarely measured in the studies cited here, c t (concentration, mg/liter, times contact time, mm) products have generally been calculated using the dosage of bromine added to the system.


The efficacy of bromine inactivation of bacteria has been summarized by Farkas-Hinsley (1966). Vegetative cells are readily inactivated by bromine, but reports often leave the type of bromine compound and its residual concentration as uncertain. Consequently, it is difficult to inake quantitative comparisons with other disinfectants. Tanner and Pitner (1939) reported that the concentrations of hypobromous acid that are required to give "complete kill" in 30 min are 0.15 mg/liter for Escherichia Coli and 0.6 mg/liter for Salmonella typhi. Spores of Bacillus subtilis required more than 150 mg/liter of bromine. Kruse et al. (1970) found that 4 mg of bromine per liter as hypobromous acid gave S logs of disinfection of E. coli in 10 mm at 00C. At pH 4.5, where significant Br2 was present, 2 mg/liter bromine at 00C gave 4 logs of E. Coli disinfection in 3 mm. However, Kiuse' also reported that 0.1 M bromide decreased E. coli disinfection to 2 logs of inactivation from 4.5 logs at 0.001 M

bromide, using 4 mg/liter bromine at pH 7.5 and 00C in each case. This is in conflict with observations made at low pH, since high bromide should also produce bromine. The data at pH 4.5 showed 4 logs inactivation with a c . t of 6 compared to S logs inactivation from a c t of 40 for hypobromous acid. Even these data, the best available, are based on the concentrations of bromine added to the solution. It is difficult to compare the various studies on disinfection by bromine, since the chemical species of bromine present in most instances have not been measured or reported.

The effect of the formation and decomposition of the bromamines on the efficacy of bacterial disinfection was first discussed by Johannesson (1958). He reported that 0.28 mg/liter of monobromamine expressed as bromine resulted in 99% inactivation in less than 1 min, while the same concentration of N-bromodimethylamine [N(CH3)2Br] required 12 miii for 99% inactivation of E. COli. The measurements of the halogen remainin~ after the experiments showed very litfie loss.

The efficacy of bromine inactivation of spores of Bacillus metiens and B. subtilis has been studied by Marks and Strandskov (1950) and Wyss and Stockton (1947). Both pairs of investigators found that the activity was markedly pH-dependent at low pH values where molecular bromine predominates. The activity increased rapidly from pH 4 to pH 3, the lowest that was tested. The c t required for 99% inactivation was 2~25 at 250C and pH 3.0. At pH 6 to 8, the c t was 225 to 375 (Wyss and Stockton, 1947). At pH 9, where the hypobromite ion starts to predominate, the c t increased to more than 500.

Marks and Strandskov (1950) also determined that at pH 7 and 250C monobromamine was one-half as effective as hypobromous acid against B. metiep~~ spores. N-bromosuccimmide was 1/17, and N-bromopiperidine was 1/800 as effective. Wyss and Stockton (1947) showed the effect of changing chemical form and concentration by increasing the concentration of ammonia that was added to a 20 mg/liter bromine solution at pH 7 containing B. subtilis spores. Table 11-13 shows the measured residual and the time required for 99% inactivation. Before the ammonia is added to the solution, hypobromous acid is the major chemical form. When the ammonia nitrogen reaches 1 mg/liter, nitrogen tribromide predominates. At 2 mg/liter ammonia nitrogen, the breakpoint occurs with rapid loss of combined bromine residual. At 10 and 30 mg/liter ammonia nitrogen, dibromamine is present and is less stable and effective than hypobromous acid. At I ,000 mg/liter ammonia nitrogen, monobromamine has a c t value for 99% inactivation that is as effective as nitrogen tribromide and hypobromous acid.

The effect of temperature on the efficacy of hypobromous acid

TABLE 11-13 The Effect of Ammonia on Bromine

Concentrations Required for 99% Inactivation of Bacillus subtilis Spores at 20 mg/liter, pH 7, 25oCa

10-mm Contact

Ammonia, Residual, Time,

mg NH3.N/Iiter mg/Iiterb min

0 20 14 280

1 16 19 304

2 2 >100 >200

10 8 85 680

30 9 70 630

100 14 18 252

1,000 19 14 266

a Data from Wyss and Stockton, 1947.

b Initial bromine ooncentration 20 mg/iiter.

Concentration of bromine times oontact time.

inactivation of B. subtilis spores is shown in Table 11-14 (Wyss and Stockton, 1947). The activity increases approximately 2.3 times for each 100C rise in temperature. After measuring washed vegetative cells of B. subtilis, Wyss and Stockton also found them to be 500 times less resistant. They reported no temperature data for their vegetative cell studies.


Kruse' et al. (1970) reported that 4 mg/liter of bromine as hypobromous acid at pH 7 and 00C gave 3.7 logs of inactivation of f2 E. coli phage virus in 10 mm. At higher bromide concentrations and lower pH, where Br2 becomes the principal form of bromine, the rates of inactivation increased. A C t of 40 was required for 3.7 logs inactivation at pH 7.0 and 00C compared to a c t of I for 4.5 logs inactivation at 00C and pH 4.8. At pH 7.5 and with a concentration of 0.1 M bromide at 00C, a c. t of 12 yielded 5.5 logs inactivation. In these studies the residual concentrations of bromine were probably not significantly different from those in solutions that did not contain added salts. In the presence of added bromide, amines were difficult to quantify. Mter 2 hr, Kruse' and his colleagues found no viable virus in a solution of 4 mg/liter bromine, starting from an initial titer of 1.7 x 109 plaque-forming units (PFU)/ml at 00C and pH 7.5, when excess arrirnonia was present. However, the

TABLE 11-14 The Effect of Temperature on Bromie Concentrations Required for 99% Inactivation of Bacillus subtilis Spores with 25 mg/lite~ Bromine at pH 7.0~

Temperature, Contact Time,

cc min c.tb

35 4 100

25 10 250

20 16 400

15 21 525

10 37 925

s 54 1,350

a Data from Wyss and Stockton, 1947.

b Conoentration of bromine times contact time.

addition of bromine to phage solutions that contained excess methylamine (CH3NH2) and glycine (H2NCH2COOH) essentially stopped disinfection under these conditions. This may be due either to the formation of the N-bromo compounds or to the reduction of bromine.

Taylor and Johnson (1974) used ~x174 E. coli phage and measured concentrations of the disinfecting species in constant residual solutions. They reported that 0.32 mg/liter hypobromous acid required only 1.1 mm or a c. t of 0.35 to inactivate 99% of this phage at 00C. Compared to hypobromous acid, molecular bromine was approximately 3 times as fast, requiring Br2 at c t of 0.1. Nitrogen tribromide was less potent, requiring a c - t of 1.0 for 99% inactivation of this phage at 00C.

The effect of temperature on ~xl74 inactivation was measured for 0.16 mg/liter of hypobromous acid expressed as bromine at pH 7. The Arrhenius plot of ln k against I /T (where T is temperature) from 2730K to 3030K gave a slope equivalent to 37 kcal/mol or an increase in rate of inactivation of 1.9 times for a rise in temperature from 150C to 250C.

Sharp's group (Floyd et al., 1976, 1978; Sharp et al., 1975, 1976) has made careful studies with measured, controlled residual concentrations. They demonstrated that even under these controlled chemical conditions the degree of aggregation or clumping had a marked effect on the apparent, observed inactivation rates with reovirus and type 1 poliovirus (Sharp et al., 1975). Reoviruses, as single particles, required only 1 s for 3 logs of inactivation at pH 7 and 20C with hypobromous acid at 0.46 mg/liter bromine. Aggregated samples gave much slower rates, especially at high levels of inactivation. At 3 logs of inactivation, the time

required doubled for the same concentration conditions (Sharp el al., 1976). Poliovirus I was also rapidly inactivated by hypobromous acid when single virus particles were studied (Floyd et al., 1976). At 20C and pH 7, only 3.5 mg/liter bromme or a C t of 0.2 was required for 2 logs of inactivation. However, the rate of inactivation did not increase linearly with concentration at this temperature. Higher concentrations were much less efficient than longer exposures. At 1 mg/liter hypobr~ mous acid, 7 s were required for 1 log of inactivation or a c t of 0.23 for 2 logs of inactivation at 20C and pH 7 with hypobromous acid. The inactivation rate increased only slightly with temperature at this concentration. At 100C a c t of 0.21 for 2 logs of inactivation and at 200C a c t of 0~06 was required for 99% inactivation with hypobromous acid.

The inactivation of single poliovirus particles in buffered, distilled water and constant residual concentrations for the other major bromine chemical forms have also been studied by Floyd el al. (1978). Table 11-15 gives the calculated c t required to yield 99% inactivation for these different forms near 1 mg/liter expressed as bromine. The values in Table 11-15 depend on concentration except for Br2, for which time and concentration were inversely related as normally assumed from the Watson-Chick relationship (Morris, 1975). Floyd et aL (1978) also demonstrated that dibromamine, nitrogen tribromide, and hypobromous acid were less efficient at higher concentrations while the hypobromite ion became more effective as the concentration increased. The rate for the hypobromite ion is rapid compared to the other bromine compounds. This is interesting in light of the fact the hypochlorite ion (OCl-) is generally considered to be a much poorer disinfectant than hypochlorous acid (HOCl).


There appears to be no information regarding the effectiveness of bromine as a disinfectant against eggs or larvae of helminth parasites in water.

With protozoa that are important to public health, attention appears to have been devoted primarily to the efficacy of bromine in destroying cysts of Entamoeba histolyzica. Most relevant information has been provided by Stringer et al. (1975) as a result of their studies on the comparative cysticidal efficacies of various halogens (shown in Table II12 in the section on iodine). They used bromine stock solutions that were prepared by bubbling nitrogen through bottles of elemental bromine. Concentrations were determined by a colorimetric bromocresol purple

TABLE 11-15 Expositre (c.t) a to Various Brornine

Compounds Required for 99% Inactivation of Poliovirus

1, Mahone~


Chemical Form c.t 0C pH

Dibromamine 1.2 4 7.0


Nitrogen tribromide 0.19 5 7.0


Bromine 0.03 4. 5.0


Hypobromite ion 0.01 2 10.0


Rypobromous acid 0.24 2 7.0


Rypobromousacid 0.21 10 7.0

Hypobromous acid 0.06 20 7.0

a Conoentration of oomnound times oontact time.

b Data from Floyd et aL, 1978.

test. Cyst survival was evaluated by the excystation method that these investigators had developed.

In dose-response experiments in distilled water for exposures of 10 miii and at pH levels from pH 4 to pH 10, bromine was found by Stringer et al. (1975) to be the most effective and fastest acting of the halogens tested over the widest pH range. At pH 4, 99.9% cyst mortality was obtained with 1.5 mg/liter of free bromine residual; whereas 2 mg/liter of chlorine and S mg/liter of iodine were required to attain the same mortality. Furthermore, increases in pH seemed to have less effect on the cysticidal efficacy of bromine as compared with the other halogens. At pH 10, 99.9% mortality was obtained with residuals of 4 mg/liter of bromine, 12 mg/liter of chlorine, and 20 mg/liter of iodine.

In studies more nearly simulating usual water treatment procedures, "flash mixing" of halogens at a dosage of 2 mg/liter with the cyst suspensions was used. In buffered distilled water, bromine again proved to be the most effective: 99.9% inactivation was obtained at pH 6 and pH 8 after 15 mm of contact. Longer exposures at the lower pH were required for the other halogens in order to provide even less inactivation than this.

An interesting aspect of the study of Stringer el al. (1975) was that

ammonia bromamines were nearly as cysticidal as free bromine except in waters of high pH (Table 11-12).

Thus, under conditions likely to be found during the treatment of natural water supplies, free and combined bromine appear to be a practical, effective cystide, at least as far as cysts of E. histolytica are concerned.

Mechanisms of Action

The pattern of bromine disinfection appears to be similar to that of chlorine. After comparing the activity of chlorine, bromine, and iodine against spores Marks and Strandskov (1950) noted that Br2 was 9 times more effective than hypobromous acid and that the hypobromite ion and tribromide ion (Br3~) had very low activity. They noted that a high degree of polarity contributed to the inactivity of the ionic forms and the reduced activity of the hypohalous acid compared to the free molecular halogen. They also found that "the killing rates of bacterial spores for the hypohalous acids and probably for the molecular halogens decrease in the [following] order: chlonne, bromine, iodine."

Many workers attribute this decrease to the greater oxidation potential of the halogens with lower molecular weight. However, the polarity and perhaps the sizes of the halogens may be important in getting the disinfectants to the vital site.

Olivieri et al. (1975) demonstrated that bromine was effective in inactivating both naked viral RNA and intact virus. The primary site of inactivation of f2 phage more likely involves the reaction of bromine with the protein coat of the virus, because inactivation of RNA that was prepared from bromine-treated virus lagged significantly behind the inactivation of intact virus. The mechanism of inactivation with bromine, as with the other halogens, involves moving the disinfectant to the vital site, mass transport, and the reactivity of the bromine with that site, oxidation. The effectiveness of bromamines as disinfectants may be explained by their relatively low polarity and high reactivity. The effectiveness of the hypobromite ion as a virucide may be due to the fact the high pH's of the solutions make the virus more sensitive to the ion. Or, it could be attributed to Olivieri's observation that bromine acts primarily on the protein coat of viruses.

Summary and Conclusions

Laboratory tests show that bromine is an effective bactericide and virucide. It is more effective than chlorine in the presence of ammoma.

As a cysticide, it is highly active. Bromine is active over a relatively wide range of pH, and it retains some of its effectiveness as hypobromous acid to above pH 9.

The major disadvantage of bromine as a disinfectant is its reactivity with ammonia or other amines that may seriously limit its effectiveness under conditions that are encountered in the treatment of drinking water. Data on the effectiveness of bromine against bacteria are complicated by this reactivity and the lack of characterization of the residual species in disinfection studies.

Research Recommendations

Further research is required to quantitate the effectiveness of the various species of bromine and the bromaniines, both organic and inorganic, against bacteria, viruses, and cysts, particularly with bacteria.

The reactivity of bromine and the dosages that are required to maintain effective residual in natural water systems should be investigated.

The effectiveness of various technologies for application of bromine or bromine chloride to intended drinking water should be evaluated.


Ferrates are salts of ferric acid (}12FeO4) in which iron is hexavalent. Fremy (1841) first synthesized potassium ferrate (K2FeO4) in the mid-nineteenth century. Since then, a wide variety of metallic salts have been prepared. However, only a few of the preparations yield ferrates of sufficient purity and stability for use in the treatment of water. Ferrates are strong oxidizing agents that have a redox potential of -2.2 V or 0.7 V in acid and base, respectively (Wood, 1958).

The Chemistry of Ferrate in Water

Aqueous solutions of potassium ferrate are unstable and decompose to yield oxygen (02), hydroxide (OH-), and insoluble hydrous iron oxide


The hydrous iron oxide is a coagulant that is commonly used in water treatment. Initial ferrate concentration, pH, temperature, and the surface character of the resulting hydrous iron oxide affect the rate of decomposition (Schreyer and Ockerman, 1951; Wagner et al., 1952; Wood, 1958). Ferrate solutions are most stable in strong base (>3 M or at pH 10 to 11). Schreyer and Ockerman (1951) reported that dilute aqueous solutions of ferrate are more stable than concentrated solutions. The presence of other inorganic ions also affects ferrate stability in aqueous solutions.

Ferrate reacts rapidly with reducing agents in solutions (Murmann, 1974). It will also oxidize ammonia ~H3). The rates for this reaction increase with pH (optimum pH range 9.5-11.2), molar ratio of ferrate to ammonia, and temperature (Strong, 1973). The oxidation of ammonia by ferrate below pH 9 is markedly slower than that observed for chlorine. Some information is available on the reaction of ferrates with various organic compounds (Audette et aL, 1971; Becarud, 1966; zhdanov and Pustovarova, 1967).


Concentrated solutions of sodium ferrate ~a2F~4) may be prepared electrochemically from the more common iron forms. Scrap iron can be converted to ferrate iron with a 40% efficiency (Murmann, 1974). Potassium ferrate may be prepared by wet oxidation of Fe(III) with potassium permanganate (KMnO4). Subsequent recrystallization yields a crystalline solid with greater than 90% purity (Schreyer et al., 1950). Ferrate salts are not commercially available. Consequently, sufficient quantities for pilot or full-scale testing would have to be especially prepared.

Analytical Methods and Residuals

Aqueous ferrate solutions have a characteristic violet color that is similar to permanganate and a wavelength maximum at 505 nm in the visible portion of the electromagnetic spectrum. The molar extinction coefficient at 505 nm in l0~~ M sodium hydroxide ~aOH) was 1 070 _ 30 1 ~7 cm~1 (Wood, 1958). Since ferrates are unstable, residual cannot be

Biocidal Activity


Gilbert el al. (1976) reported the inactivation of a pure culture of Escherichia coli at pH 8.0. Increased ferrate concentrations from 1.2 to 6.0 mg/liter yielded increased rates of inactivation of E. coli. They observed 99% inactivation at 6.0 mg/liter, pH 8.0, and 270C in approximately 8.5 mm. The rates of inactivation of E. coli with ferrate appear to be of the same order of magnitude as monochioranune ~H2Cl). Waite (l978a,b) extended Gilbert's disinfection studies to include entenc pathogens and Gram-positive bacteria and evaluated the effects of pH and temperature. The rate of E. coli inactivation by ferrate increases as pH decreases from pH 8.0 to pH 6.0. At the lower pH value, the rate of ferrate decomposition is increased, and little inactivation was observed after S min. Several inconsistencies at low ferrate concentrations (0.12 mg/liter) were observed, but low-level inactivations (<50% inactivation) are difficult to interpret. At lower temperatures, the biocidal activities of ferrate appear to increase as temperature increases. At higher temperatures, the ferrate decomposition becomes an important factor. Salmonella typhimurium and Shigellaftexneri were inactivated in a manner similar to that of E. coli. However, considerably higher dosages of ferrate (>12 mg/liter) were necessary for inactivation of Streptococcus faecalis. Ferrate concentrations of 12 mg/liter and 60 mg/liter required S min and 15 miii, respectively, for 99% inactivation at pH 8.0. Other Gram-positive bacteria tested (Bacillus cereus, Streptococcus bovis, and Staphylococcus aureus) were noticeably more resistant to ferrates than were the enteric bacteria. Higher doses of ferrate (15-20 mg/liter) were also necessary to inactivate 99% to 99.9% of the bacteria in the presence of organic material in wastewater.


Waite (1978a,b) reported that the inactivation of the RNA bacterial f2 virus with ferrate appeared to be more dependent on pH than was bacterial inactivation. Considerably more viral inactivation was oh-served at pH 6.0 than at pH 8.0. At 1.2 mg/liter ferrate, just under 4 min were required for 99% inactivation, while for similar conditions 13 min were necessary at pH 8.0.

The disinfection efficacy of ferrate is sumrarized in Table 11-16. The ferrate dose is given in the table, since ferrate residuals were not reported. Ferrate decomposes rapidly in aqueous solution, and the dose

TABLE 11-16 Concentrations of Ferrate (FeO42-) and Contact Times

Necessary for 990/0 Inactivation of Escherichia coli, Streptococcusfaecalis, and f2 Virusa

Ferrate Contact Tempera-

Test Dosage, Time, ture,

Microorganism mg/liter rnin c.t~ pH 0C

~chenchia 1.2 4.6 5.52 7.0 20

coli 6.0 2.5 15.0

Streptococcus 1.2 365 438 7.0 20

faecalis 6.0 4.0 24

f2Virus 1.2 3.7 4.4 6.0 20

6.0 1.5 9

a Data from Waite, 1978b.

b Concentration of ferrate times contact time.

represents an overestimate of ferrate concentration. Nevertheless, the c t products (concentration, mg/liter, times contact time, mm) indicate the difference in efficacy relative to the test microorganism and provide a crude idea of the c t product for ferrate.

Mechanism of Action

No studies have elucidated the mechanism of inactivation caused by ferrate.


While not as biocidal as the free halogens, chlorine dioxide, or ozone, ferrate appears to be similar to or slightly better than the chloramines as a bactericide and more active as a virucide than the chloramines. The combined application of ferrate as a disinfectant and a coagulant makes it an attractive alternate biocide. Far more information on the feasibility of large-scale preparation of ferrate and its biocidal activity (particularly under water treatment conditions) are necessary before ferrate can be given consideration as a water disinfectant for public systems.


High pH values are obtained when drinking water is softened by the commonly used precipitation method, which removes calcium (Ca2+) and magnesium (M~+) ions that cause water hardness. In such cases, calcium hydroxide [Ca(OH)2] is the usual source of hydroxide that is used to raise the pH. In this process, pH values as high as 10.5 are reached and maintained up to 6 hr. Low pH conditions are not used in water treatment processes. Therefore, while low pH's are lethal to most microorganisms, they are of less interest in this report than high pH conditions, which might have potential for disinfection through modification of current practices.

Attainment of Elevated pH

High pH values in water are obtained by using either calcium hydroxide or sodium hydroxide ~aOH). Sodium hydroxide is a by-product of the production of chlorine by the electrolysis method. Calcium hydroxide is prepared by reacting calcium oxide (CaO) with water. The calcium oxide is produced by heating calcium carbonate (CaC03) to drive off the carbon dioxide (C02). The calcium oxide is normally prep&ed off the site of the water treatment plant. However, it is often slaked on site at all but the smallest plants, which purchase calcium hydroxide and use it direcdy. A few, very large water treatment plants produce their own calcium hydroxide by calcining their water-softening sludge.

Hydroxide (OH~ is added to water as a water slurry containing calcium hydroxide or as a water solution of sodium hydroxide. In the usual water-softening technique, the water is flocculated for I~30 min to promote the precipitation reactions. The precipitates are then removed by sedimentation with retention times from 1 to 6 hr.

Disinfection with high pH is readily accomplished, but the required pH values are so high that the pH must be reduced before the water is consumed. In a full-scale water treatment plant, this is accomplished in the recarbonation stage, which is a normal component of the watersoftening process. However, for small institutions or private residences the pH reduction requirement would mitigate against the use of this disinfection technique.

Analytical Methods and Process Control

The pH can be reliably measured with the glass electrode at pH values of approximately 7 and at room temperatures. Special care must be taken to obtain accurate values at other temperatures and at the elevated pH values that are necessary for disinfection with this process.


The hydroxide that is used to obtain the necessary pH values enters into few side reactions in natural water after the initial formation of metallic hydroxides and the reactions with carbon dioxide. The amounts of hydroxide for these reactions can be readily computed after the water has been chemically analyzed. There is a slow additional loss of hydroxide from the water by reaction with carbon dioxide entering from the atmosphere. However, this is not a serious problem.

Biocidal Activity

The microbial inactivation under consideration in this chapter is that caused solely by high pH. When the pH of water is raised, the opportunity exists for precipitation of many compunds associated with carbonates, oxides, and hydroxides, among many others. Even in distilled water, calcium carbonate may precipitate if calcium hydroxide is used to raise the pH and if sufficient carbon dioxide enters the water during the experiment. These precipitates provide opportunities for adsorption and coagulation of microorganisms and will cause an increased removal over that obtained from the pH effect alone without the precipitate. The research cited below, unless otherwise noted, was conducted with procedures that obviated this problem or that made its effect insignificant.


Wattie and Chambers (1943) determined the times required to obtain 100% inactivation of several bacterial species at initial concentrations of approximately 1,500 organisms/ml in 2~C to 250C dechlorinated tapwater, using calcium hydroxide for pH adjustment (see Table 11-17). The inactivation was significantly less at each pH for each organism between 00C to I 0C. At a pH of 11.5, the time for 100% inactivation of Escherichia coli was increased from 210 mm between 2OCC to 250C to 355 mm between 00C to 10C. The time for 100% inactivation of



TABLE 11-17 Contact Times Necessary for High pH Conditions to Effect 1000/0 Inactivation of Initial Concentrations of 1,500 Organisms/mi0

Contact Time Required for 1000/0 Inactivation by High pH Conditions, mm

Esclierlchia Enierobacier Pseudomonas Salmonella Shigella

p11 coli aerogenes aeruginosa lyphi dysenierlae

9.01-9.5 >540 - - >540

9.51-10.0 >600 >600 420 >540 >300

10.01-10.5 >600 >600 300 >540 >300

10.51-11.0 600 >540 240 240 180

11.01-11.5 300 >600 120 120 75

TABLE ll- 18 Inactivation of Escherichia coli by High pH Conditions at 250C in a Dilution Mediuma

Contact time,

pH min Inactivation, %

11.55 50 10

11.70 50 90

11.81 35 96

12.01 10 99.8

12.04 6 97

a Data from Berg and Berman, 1967.

Salmonella typhi was increased from 75 min to 270 min for the same temperature change. No mention was made of difficulties from the formation of precipitates.

Riehi et al. (1952) reported F. coli inactivations of 95% in 8 hr at pH

10.5 and 50C, 100% in 2 hr at pH 10.5 and 150C, and 100% in approximately 30 mm at pH 10.5 and 250C in distilled water. Calcium hydroxide was used for pH adjustment. Inactivations of 50% and 55% in 10 hr were noted for Salmonella montevideo and 5. ~yphi, respectively, at

20C and pH 10.6 in distilled water. Initial concentrations of organisms were approximately 1,000/mi. They also observed that higher temperatures gave more inactivation at the same pH and contact time and that the composition of the water did not markedly influence the bacterial survival.

Berg and Berman (1967) determined the inactivation of F. coli at 250C in the dilution water medium that is described in the 11th edition of Standard Methods (1960) (see Table 11-18). Sodium hydroxide was used to adjust the pH.


Wentworth et al. (1968) reported that little or no poliovius 1 was inactivated in distilled water at pH 11.2 at room temperature in 90 inin when calcium hydroxide or sodium hydroxide was used to increase the pH. For poliovirus I in distilled water without added salts, Sproul et al. (1970) showed that no inactivation occurred in 90 min at pH values of 10.5 or less when calcium hydroxide and sodium hydroxide furnished the hydroxide. Their work was done at 21 0C to 220C. Sproul (1975) obtained

TABLE 11-19 Inactivation of Echovirus 7 by High pH Conditions at 25~C in a Dilution Mediuma

Contact time,

pH mm inacLivation, %

11.23 7 99.98

11.49 4 99.99

11.79 4 99.5

11.92 1.5 99.9

a Data from Berg and Berman, 1967.

poliovirus 1 inactivations in 30 min of 7% at pH 11.5, 59% at pH 11.9 in 30 mm, 94% at pH 12.1 in 20 min, and 99.83% at pH 12.5 in S 'nin. He used distilled water with 100 mg/liter of sodium chloride ~aCl) at 220C to 230C.

Berg and Berman (1967) showed the inactivation of echovirns 7 AGKP8Al at 250C in a dilution water medium (Stan~~rd Method~, 1960). They used sodium hydroxide to adjust pH (see Table 11-19). From his work with the f2 bacteriophage, Donovan (1972) suggested that a calcium-virus complex was formed and that a large part of the observed decrease in titer at pH 11.5 was caused by an aggregation of the virus. Electron micrographs and chemical evidence supported the observation.


The susceptibility of protozoa or helminths to inactivation by high pH conditions does not appear to have been studied.

Mechanisms of Action

The inactivation mechanism of high pH on bacteria has not been examined. There is ample evidence to show that the poliovirus is inactivated by the disruption of the capsid and a loss of the RNA to the water (Boeye and Van Elsen, 1967; Maizel et aL, 1967; Van Elsen and Boeye, 1966). Boeye and Van Elsen (1%7) suggested that, at pH 10 and at elevated temperatures (>300C), the RNA was released from the capsid in a degraded form or that it was quickly and extensively broken down after release. The applicability of this mechanism to enteroviruses other than poliovirus is unlinown.


The utilization of high pH conditions for disinfection of water is feasible, but higher pH values than are normally used in water treatment and long contact times are required. At pH values of approximately 10.5, bacteria are inactivated in up to 600 min, but viruses are probably unaffected. A pH value of 12.0 to 12.5 and a contact time of approximately 30 min can probably yield 99.0% to 99.9% inactivation of most bacteria and of certain viruses. This process has a drawback: the pH that is necessary for effective disinfection must be reduced before the water can be consumed. Use of the high pH as the sole means of disinfection is not recommended. In many situations where water is softened by the precipitation process and where the water is of poor biological quality, the disinfection potential of this process could be used by increasing the pH above its present values.

Research Recommendations

Although this method has not been used deliberately for disinfection in the past, additional work on this method may be warranted since high pH conditions are attained in many treatment plants. Studies are needed to determine the pH values required for inactivation of protozoans, a broader and more representative group of viruses, and a larger group of enteric bacteria.


Hydrogen peroxide (H202) is a strong oxidizing agent that has been used for disinfection for more than a century. Its instability and the difficulty of preparing concentrated solutions have tended to limit its use. However, by 1950 electrochemical and other processes were developed to produce pure hydrogen peroxide in high concentration, which is known as stabilized hydrogen peroxide (Schumb et aL, 1955). This product has been subjected to increased study and application. Most recently it has been used to disinfect spacecraft (Wardle and Renninger, 1975), foods (Toledo, 1975), and contact lenses (Gasset et al., 1975; Spaulding et al., 1977). Although there has been some interest in using hydrogen peroxide as a disinfectant for wastewater CFaki and Hashim~ to, 1977), it has been used more for control of bulking in the activated sludge waste treatment process (Sezgin et al., 1978). Its use in drink:ing water disinfection appears minimal.

Analytical Methods

Analytical methods for hydrogen peroxide are based on oxidation or reduction with potassium permanganate (KMnO4), potassium iodide (KI), or ceric sulfate [Ce(504)2) (Chadwick and Hoh, 1966). A colorimetric procedure based on the oxidation of titanium sulfate (TiOSO4) (Snell and Snell, 1949) is sensitive to about I mg/liter in the absence of other oxidizing agents.

Production and Application

Commercially produced hydrogen peroxide is available in aqueous solution, usually ranging from 30% to 90%. It is stabilized during manufacture by addition of such compounds as sodium pyrophosphate (Na4P2O7), acetophenetidin (CH3CONH~H4OC2H5), or acetanilide (CH3CONHQH5). As the concentration of hydrogen peroxide is decreased, the concentration of stabilizers is typically increased. Experience in the waterworks industry using hydrogen peroxide is nonexistent. However, hydrogen peroxide is widely used as a bleaching agent in making cotton textiles or in wood pulping (Chadwick and Hoh, 1966). Presumably, a concentrated solution would be diluted and applied with a chemical metering pump. Careful attention to safety in handling would be required because of the possibility of fire or explosion.

Biocidal Activity

The few pre-1965 references on disinfection by hydrogen peroxide have been summarized by Yoshpe-Purer and Eylan (1968). Although its bactericidal activity was indicated, a need for catalytic Fe2+ or Cu2+ was reported. Yoshpe-Purer and Eylan (1968) worked with Escherichia co/i, Salmonella typhi, and Staphylococcus aureus in pure culture and as a mixture of bacteria. They also studied the effect of hydrogen peroxide concentrations (from 30 to 60 mg/liter), contact time (10 to 420 min), and initial concentration of organisms. Without ever measuring residuals, they concluded that bacterial inactivations were relatively slow, that E. coli was more resistant to hydrogen peroxide than S typhi or S. aureus, and that the required inactivation time was increased as the initial concentration of organisms was increased. Regardless of the hydrogen peroxide concentration or the type of organism used , all tests were claimed to be "sterile" in 24 hr. Table 11-20 summarizes some of these data.

Toledo et al. (1973) studied inactivation of S. aureus and spores of

TABLE H-20 Contact Times Necessary for Hydrogen Peroxide (H202) to Effect 99% Inactivation of Various Concentrations of Escherichla Coli, Salmonella typhi, and Staphylococcus aureusa

Contact Time Required

for 99% Inactivation by

Initial Hydrogen Peroxide, mm

Test Conoentration,

Microorganism bacteria/rnl 30 mg/liter 60 mg/liter 90 mg/liter

Esch£nclaa 102 b b 360

coli 104 b b b

106 b b 360

Salmonella 102 60 >30 <45 <10

ryphi 104 >300 >45 <60 30

106 >300 >60 >45<60

&aphylococcus 102 b >15<30

aurew 104 c b 60

106 c b 180

Data from Yoshpe-Purer and Eylan, 1968. Organism grown on nutrient agar (2~hr

culture), suspended in saline, and exposed to varyihg conoentrations of hydrogen peroxide in

1 liter of tap water, pH 6.5, at room temperature (not specified).

b Inactivation in 300 mm was less than 90%. No data reported.

several species of Bacillus and Clostridium with 25.8% (258,000 mg/liter) hydrogen peroxide. The time to reduce £ aureus by 6 logs was 1 ruin, whereas reduction of spores by 99% required from less than 1 miii to approximately 17 mm, depending on the species. Increasing the hydrogen peroxide concentration up to 41% (or 410,000 mg/liter) or increasing temperature from 240C to 760C significantly reduced the required inactivation time.

Several studies on virus inactivation by hydrogen peroxide have been reported. Lund (1963) obtained a 99% inactivation of poliovirus (Saukett strain) in about 6 hr with 0.3% (3,000 mg/liter) hydrogen peroxide. Mentel and Schmidt (1973) worked with rhinovirus (types lA, lB, and

7). They found that while a 1.5% (15,000 mg/liter) concentration required approximately 24 min for 99% inactivation, equivalent inactivation occurred in 4 mm with a concentration of 3% (30,000 mg/liter).

Information is lacking on the effect of hydrogen peroxide on protozoa and helminths in water.

In none of the studies cited above is there any indication that the dosage of hydrogen peroxide was measured other than by the volumetric addition of hydrogen peroxide. No measurements of residual were reported, although Yoshpe-Purer and Eylan (1968) claimed that a residual was present for up to 13 days as indicated by inactivation of added doses of bacteria.

Mechanism of Adion

No studies have specifically identified the mechanism of action of hydrogen peroxide. Spaulding et al. (1977) believed that the hydrogen peroxide molecule itself was not responsible for the action but, rather, that the free hydroxyl radical (HO.) that it produced was the speciiic inactivating agent. They claimed that the catalytic effect of iron or copper ions supported this theory. Yoshpe-Purer and Eylan (1968) also reported that free radicals were important but that sufficient catalyzing metal ions were available either from the tap water that they used or from the bacterial cells themselves.


Because of its relatively high cost and the high concentrations that are required to achieve disinfection in reasonable time, hydrogen peroxide is not a generally satisfactory disinfectant for drinking water.

Research Recommendations

If further research is to be conducted, the stability of hydrogen peroxide and its residual effect in the presence of organic material and other substances in water or distribution systems should be investigated. Parallel studies using chlorine and hydrogen peroxide, separately and in conjunction, should be conducted.


Ionizing radiation may be electromagnetic or particulate. As used for disinfection or sterilization, electromagnetic radiation may be UV, gamma, or X rays, and the particles may be alpha or beta or neutrons, mesons, positrons, or neutrinos. This discussion is limited to gamma rays and beta particles (or electrons).

Although there is extensive literature on the use of ionizing radiation

in the preservation of food and other materials (Silverman and Sinskey, 1977), little information is available on water treatment.


Laboratory studies of destruction of microorganisms are generally conducted by exposing small containers holding the test suspensions to the shielded source of ionizing radiation. However, in a practical application, the shielded source must be designed to permit relatively thin sheets of flowing liquid to be exposed. This is particularly important for high-energy electrons that have a less penetrating power than gamma rays. The MIT study (Massachusetts Institute of Technology, 1977) included a field demonstration in which a plant treating approximately 380,000 liters/day (minimum dosage of 400,000 rads) was designed and operated. The energy source was a 50-kW electron accelerator operated at 850,000 V. Sludge was pumped under the electron beam through a drum system that produced a layer that was 1 .2-m wide and 2-mm thick. The entire operation was conducted in a concrete vault to shield the operators against X rays, which are produced incidentally. Shielding of workers is a major requirement both for gamma ray and high-energy electron sources.

Analytical Methods and Residual

No residual is produced by ionizing radiation. Consequently, dosage has been measured exclusively. The MIT (1977) report and Silverman and Sinksey (1977) summarized analytical methods including calorimetry, photoluminescent dosimetry, colorimetry, and the use of ionization chamber instruments.

Biocidal Activity


The first study of ionizing radiation in the treatment of water was reported by Dunn (1953). In addition to reviewing the literature and providing some discussion on ionizing radiation, he studied the use of a l,()()0~Ci source of cobalt-60 (providing gamma rays) and a Van de Graaff generator (providing high-energy electrons). Working with natural waters and wastewaters that were not chemically or bacteriologically characterized or controlled, he exposed samples to a constant radiation flux (2,000 R/min fr~m the cobalt-60 or electrons from 3-Mv

TABLE 11-21 Dosages of Cobalt-60 Irradiation

Necessary for 99% Inactivation of Microorganisrns in

Distilled Watera

Test Dosage Required,

Microorgaiiism rad

Bacillus subtilis var niger 3.5 x ~

Mycobacierium Smegma(is 1.4 x 1O~

Escherichiacoli 6.5 x 1O~

Micrccoccuspyogenes 'aar aureus 5.8 x ~

E. coliphageT3 3.2 x 1O~

a Data from ~we et al., 1956.

operation of the Van de Graaff generator). Var'ying only the time of exposure, Dunn found that it took 0.125 miii (dosage 250,000 R) with the cobalt-60 source to reach 95.4% to 99.999% inactivation of total initial numbers of bacteria (as measured by a plate count at 370C).

Ridenour and Arnibruster (1956) studied the effect of a 1~kCi source of cobalt-60 (3,000 R/min) on a variety of natural waters, wastewaters, and pure cultures (approximately 2 x l0~ organisms per milliliter) of 10 different organisms. They found that a dosage of 100,000 roentgen equivalent physical (rep) reduced the count of all species tested by

>99%. (In order of increasing resistance, the species were: Enterobacter aerogenes, Escherichia col4 Shigella ftexneri, Salmonella typhi, 5. sonnei, Salmonella sp., Staphylococcus aureus, 5. paratyphi B, Streptococcus faecalis, and Bacillus subtilis.) With river water, a dosage of from 50,000 to 100,000 rep reduced the total bacterial count and the coliform index by at least 99%, but 150,000 rep were required to reduce the streptococcus mdex by 99%. Varying the pH (5.0, 7.0, and 8.5) had no effect on the rate of inactivation.

Also using cobalt-60 (1,100 Ci), Lowe et al. (1956) exposed pure cultures of bacteria that were suspended in double distilled water and sterile settled sewage. The concentration of microorganisms was approximately 1 x lO~/ml. Although they did not report exposure time, they summarized their exposure data in rads (radiation absorbed dose). (See Table 11-21).

There have been more studies on the disinfection of wastewater than on disinfection of drinking water. Ballantine et al. (1969) and Compton et al. (1970) considered ionizing radiation not as good as other available methods but, with increased emphasis on water reclamation (Eliassen

and Trump, 1973) and the use of high-energy electrons (Massachusetts Institute of Technology, 1977; Wright and Trump, 1956), the prospects for practical application seem improved. In the MIT (1977) study, which is the most systematic one conducted to date, washed cells or spores were suspended in 0.067-M phosphate buffer (1 x 108 cells/ml) and exposed to high-energy electrons from a Van de Graaff generator. The approximate dosages, in rads, required for a 99% inactivation of the organisms tested were: Escherichia coli (K12), 38,000; Salmonella typhimurium (LT2), 24,000; Micrococcus sp., 35,000; Aspergillus niger (spores), 78,000; 5. typhimurium (24), 100,000; Streptococcusfecah~~, 300,000; and Clostridjum perftingens (spores), 400,000. The investigators found that the solids in the wastewater had no effect on disinfection.

Of the factors affecting disinfection, oxygen was most important. For example, for E. coli K 12, the dosage required for 99% inactivation was more than doubled when an atmosphere of air was replaced by one of nitrogen. and almost halved when the air was replaced by an oxygen-rich atmosphere.


The inactivation of viruses does not appear to be influenced by the source of the iomzmg radiation (Lea, 1955). It is affected by factors that are similar to those described above for bacterial inactivation (Ballentine et al., 1969). Lowe et al. (1956) found that 32,000 rads were required for 99% inactivation of E. coli phage T3 (see Table 11-21). In the MIT study (1977), from 300,000 to 420,000 rads were required for 99% inactivation of coxsackievirus B3, poliovirus 2, echovirus 7, reovirus 1, and adenovirus S (in order of increasing sensitivity), each suspended in 0.05 M glycine at pH 7.0.


The MIT study (1977) appears to be the only one in which the effect of ionizing radiation on both protozoa and helminths was investigated; however, no quantitative data were reported. Brannan et al. (1975) found that to attain a 90% reduction of embryonation of Ascaris lumbricoides eggs, a dosage of approximately 30,000 rads was required.

Mechanism of Action

In a review of the literature, Silverman and Sinskey (1977) summarized the mode of action of disinfection by ionizing radiation. Two effects are

recognized: the direct effect, in which the primary cellular target is DNA that is damaged by the energy released from the ionizing radiation, and the indirect effect, which is associated with the production of such substances in the cell menstruum as hydrogen peroxide (H202), organic peroxides, and free radicals. Oxygen is important because it reacts with electrons and radicals, which in turn react to form hydrogen peroxide or ozone (03).


Ionizing radiation can disinfect water effectively; however, the large source, shielding, and relatively thin exposure layers that are required create difficult engineering and safety problems. The complex technology may limit application to large facilities that can provide for adequate safeguards. The absence of residual dismfection is also restrictive.

Research Recommendations

Any further research should address practical methods of delivering the ionizing radiation to the water mass. Basic engineering design, construction, and operation data should be developed.


Potassium permanganate (KMnO4) is a strong oxidizing agent, which was first used as a municipal water treatment chemical by Sir Alexander Houston of the London Metropolitan Water Board in 1913. In the United States, it was used in Rochester, New York, in 1927 and in Buffalo, New York, in 1928. Since 1948, it has been used more widely in waterworks as an algicide (Fitzgerald, 1964; Kemp et a'., 1966), as an oxidant to control tastes and odors (Spicher and Skrinde, 1963; Welch, 1963), to remove iron and manganese from solution (American Water Works Association, 1971; Shull, 1962), and, to a limited extent, as a disinfectant (Cleasby et aL, 1964).

The relatively limited information concerning disinfection with potassium permanganate is subject to criticism, because there have been no studies of the effects of organic constituents of the medium (test system) or destruction of a variety of organisms, especially pathogens.

Analytical Methods and Residual

Concentrations of potassium permanganate can be determined readily by iodometric titration using sodium thiosulfate ~a2S2O3) as a titrant or by direct titration with ferrous sulfate. Manganese can be determined by atomic absorption spectrophotometry, or the permanganate ion (MnO4~) can be measured colorimetrically (Standard Methods, 1976). The 'mm-mum detectable concentration of manganese by the colorimetric procedures, using a 100-mi sample, is 50 ~g4iter. A concentration of 0.05 mg/liter or greater can be detected visually by the pink color that is imparted to the water (Welch, 1963).

Production and Application

Crystalline potassium permanganate is highly soluble in water (2.83 g/l00 g at 00C). In waterworks, it is prepared usually as a dilute solution (1% to 4%) and applied with a chemical metering pump (American Water Works Association, 1971). It also may be added as a solid using conventional dry-feed equipment.

Reduction of the permanganate ion produces insoluble manganese oxide (Mn3O4) hydrates. To prevent distribution of turbid water or of water that will cause unsightly staining of plumbing fixtures, potassium permanganate most often is applied as a pretreatment that is followed by filtration. For example, addition of potassium permanganate to a finished water to maintain a residual in a distribution system is unacceptable because of the pink color of the compound itself or the brown color of the oxides.

Biocidal Activity

Cleasby et al. (1964) and Kemp et ~. (1966) summarized the few scientific references to disinfection by potassium permanganate that were published earlier than 1960. Although some bactericidal activity was indicated, no quantitative data were presented, and the early reports disagreed as to its value.

The most systematic study of disinfection by potassium permanganate has been made by Cleasby et al. (1964). They worked exclusively with Escherichia coli (prepared as a lactose broth culture) and studied the effect of potassium permanganate at doses of 1 to 16 mg/liter at pH values of 5.9, 7.4, and 9.2, temperatures of OCC and 200C, and contact times of 4 to 120 miii. They concluded that bacterial inactivation was relatively ineffective, but slightly better at the higher temperature, and

TABLE 11-22 Concentrations and Contact Times Necessary for Potassium Permanganate to Effect 990/0 Inactivation of a

48-Hr Escherichia coli Lactose Broth Culturea

Contact Time Required for 990/0 Inactivation by Potassium Permanganate, mm

1 mg/liter 2 mg/liter 4 mg/liter 8 mg/liter 12 mg/liter pH

10 5.9

b 7.4




16 mg/liter

45 C C C s

b C C C 115


95 45 15 15 s 5.9 20

a c b 80 25 7.4

C C C C b 9.2

a Data from Cleasby el al., 1964.

b Inactivation in 120 min was less than 900/o.

No data reported.




that increasing pH decreased disinfection rates. Table 11-22 surrunarizes in tabular form their data, which were presented originally in graphic fashion.

At least two patent applications involve the use of potassium permanganate as a disinfectant for water in swimming pools (Heuston, 1972; Seidel, 1973). Seidel's patent covered use of a concentration of 0.1 to 0.2 mg/liter in a recirculating system including filtration through quartz gravel. She claimed removal of bacteria and algae, but the specific merit of potassium permanganate in her system cannot be determined from the patent application. Heuston (1972) proposed a tablet contammg 0.001 g of potassium permanganate per 0.2 g tablet to give a dose of 1 mg/liter. Because the tablet includes potassium iodide (KI) and other oxidizing agents, it is impossible to assess the specific killing action of potassium permanganate.

A number of studies on inactivation by potassium permanganate have been reported. However, these dealt largely with the mode of action of potassium permanganate in viral inactivation (Lund, 1963, 1966) or at efforts to control human diseases (Peretts et al., 1960; Schultz and Robinson, 1942; Wagner, 1951), animal diseases (Derbyshire and Arkell, 1971), or plant diseases (Eskarous and Habib, 1972; Hughes and Steindl, 1955). They provide few clear-cut quantitative data on biocidal activity of the compound.

Information is lacking on the effect of potassium permanganate on protozoans and helminths in water.

Mechanism of Action

The mechanism of action of potassium permanganate has not been definitively identified. From the studies that have been conducted (Lund, 1963, 1966) it may be presumed that it exerts its disinfection activities by oxidizing compounds that are involved in essential cellular functions. However, Lund (1963) questioned the suggestion that the oxidative process was the mechanism for poliovirus inactivation, but was unable to exclude it as a possibility.


The relatively high cost, ineffective bactericidal action, and aesthetic unsuitability of maintaining a residual in the distribution system make potassium permanganate a generally unsatisfactory disinfectant for drinking water.

Research Recommendations

Additional studies on potassium permanganate are hardly warranted, but in the unlikely event that further research on its biocidal activities is conducted, it would be desirable to study the effect of organic material on its disinfection and the effect of the compound on organisms other than E. coil. To relate disinfection efficiency to more conventional practices, parallel studies using chlorine should be conducte~


Silver as a metal has been known for millennia, and its use as a water disinfectant dates back to the Persian king Cyrus. The term oligodynamic, which describes the killing effect of small concentrations, was coined by von Naegeli in 1893, but it is unscientific and its use should be rejected. The antibacterial action of silver and silver nitrate (AgNO3) was noted first by Raulin in 1869. Since the late nineteenth century, but especially since World War II, there have been considerable efforts to exploit the use of silver as a disinfectant, particularly for individual (11ome) water systems and swinnning pools. Silver has been used both as a salt, most commonly silver nitrate, or as metallic silver, either bound in filter beds, generated by electrolytic devices, or applied as a colloidal suspension.

The relatively limited information concerning disinfection with silver has been seriously criticized because accurate measurement of low concentrations, i.e., <200 JLg/liter, has been difficult; a suitable neutralizing agent has not always been incorporated in test protocols; and adsorption of silver on surfaces of test vessels has confounded some studies. Despite these limitations, silver has been used in parts of Europe and Japan as a water disinfectant.

Production and Application

In water treatment, silver has been applied principally by dissolving the metal or by incorporating a silver compound in a filter mediurn, often an activated carbon filter. Romans (1954) has described older processes and patents. Davies (1976) reviewed processes for treating swinmiiflg pool water in which silver was added as a soluble salt and the silver ions were kept oxidized by the addition of persulfate; metallic silver was deposited on an activated carbon filter; silver was released from solid silver electrodes, which were alternately made anodic and cathodic; silver was

released from an activated carbon filter containing silver and supplemented by passing the water through anodic silver screens; and silver was released from a pair of silver~opper alloy electrodes by applying a loW reversing voltage.

During the past 6 yr, dozens of patents have been issued in Germany, Spain, South Africa, India, and the United States for devices or Systems for adding silver to home drinking water systems or swimming pools. The home Systems are most frequently combinations of activated carbon to remove tastes and odors and silver to prevent bacterial growth on the filter.

Because of the low solubility of silver, the dose is usually less than 50 ~g/liter. Presumably this is reduced rapidly because of the adsorption of silver to surfaces, but there is no information on the silver residual in treated water. However, the maximum contaminant level (MCL) is 0.05 mg/liter (U.S. Environmental Protection Agency, 1975).

Analytical Methods

Until recendy there have been no satisfactory techniques for measuring silver at the ~g/liter level. Using a dithizone colorimetric method, the minimum detectable quantity of silver is 200 ~~g/h.ter (Standard Methods, 1976). Using an atomic absorption spectrophotometric method, the detection limit is 10 ~g/liter (U.S. Environmental Protection Agency, 1974), and, if a heated graphite furnace is used, the detection limit is reduced to 0.005 ~g/liter (Rattonetti, 1974).

Biocidal Activity

One of the earliest studies was conducted by Just and Szniolis (1936), who found that 100 ~g of silver per liter disinfected water within 3 to 4 hr, that silver nitrate and metallic silver were equally effective, and that histopathological changes occurred in rats that were given water containing 4OO~l,000 ~g silver/liter for lO(~ days. Other systematic studies of silver disinfection have been conducted by Renn and Chesney (195~1956), Wuhrmann and Zobrist (1958), and Chambers and Proctor (1960). These studies have been summarized by Woodward (1963). Table 11-23, adapted from Wuhrmann and Zobrist (1958), shows that in the concentrations used, silver acted slowly, but that there were increases in the rate of bacterial inactivation with increasing temperature and pH. Chambers and Proctor (1960) obtained similar results, but the inactivation rates of Renn and Chesney (l95~l956), who used comparable concentrations of silver, were faster.

TABLE 11-23 Concentrations and Contact Times Necessary for Silver to Effect 99.9% Inactivation ofEscherichia coli in 0.005 M Phosphate Buffe~

Contact Time Required for 99.9% InacLivation

by Silver as Silver Nitrate, mm Tempera


0.01 mg/liter 0.03 mg/liter 0.09 mg/liter 0.27 mg/liter pH ac

1,010 837 156 53 6.3 S

466 214 81 34 7.5

268 109 58 18 8.7

831 344 144 32 6.3 15

316 177 63 21 7.5

216 100 38 13 8.7

1,210 152 68 20 6.3 25

423 86 32 13 7.5

203 40 20 8 8.7

Data from Wuhrmann and Zobrist, 1958.

Both Wuhrmann and Zobrist (1958) and Chambers and Proctor (1960) found that the presence of phosphate (at 60 mg4iter) slowed bacterial inactivation. Because phosphate is typically found at much lower concentrations in drinking water, this observation has litfie practical significance, but it may account for disparate results in some laboratory studies.

Increasing water hardness slowed bacterial inactivation. According to

Wuhrmann and Zobrist, an increase of 3 miii was required to achieve

99.9% bacterial inactivation (at 200C and pH 7.0) for each 10 mg/liter

increase in hardness. Likewise, chlorides interfered with the action of

silver. At 10 mg/liter, they increased the inactivation time by 25% and at

100 mg/liter, by 70%.

The source of silver seemed to be irrelevant to the inactivation rate because silver either added as silver nitrate or dissolved from metallic silver gave approximately equal results.

In most studies of the disinfection of water with silver, Escherichia coli

has been the test organism. Wuhrrnann and Zobrist also tested a

Salmonella species and found it to be at least as sensitive as E. COli.

Yasinskii and Kuznetsova (1973) observed that 90 ~g/liter inactivated

Vibrio comma (1 x 106 cells/ml) in 30 miii or 22.5 ~g/liter in 60 min. Fair (1948) and Harrison (1947) suggested that silver salts were

ineffective against cysts of Entamoeba histolytica, but the data provided in these reports were limited. By contrast, Newton and Jones (1949) observed that electrolytically produced silver in tap water gave greater than 99% cyst inactivation in 1 hr and highly variable residual concentrations that ranged between 17 and 33 mg/liter at pH's of 9.0 to 9.8. When silver nitrate in distilled water at 0.5 mg/liter was used, 4 to 6 hr were required for a 90% to 99% inactivation. At S mg/liter, a >99% inactivation occurred in 3 hr, and at 30 mg/liter, a >99% inactivation occurred in I to 2 hr.

Chang and Baxter (1955) used 150 mg/liter of silver nitrate (95 mg/liter Ag+) with contact lines of 1 to 6 min. They concluded that cysticidal activity was only moderate compared to that of iodine. Apparently, there are no data on silver as a virucidal agent in water except for one study by Lund (1963), who decreased poliovirus activity by 2.5 logs in 4 hr with a concentration of approximately 68 mg/liter.

A good general review of silver and its compounds and a more thorough medical examination of their applications was published by Grier (1977), who concluded that there would be a "significant place for silver compounds in the prevention and treatment of at least some bacterial diseases."

Mechanism of Action

Romans (1954) summarized the mechanism of so-called oligodynamic activity: "However, there is much difference of opinion as to the form of the active principle, as to the mechanism of its action, and the value of the results obtained. Some authorities believe that the active form is a positively charged ion, some think it is a complex ion, others a salt and still others think it acts by formation of proteinates or merely as a catalyst. These are only a few of the ideas that have been expressed with formidable experimental support."

Chang (1970) attributed a direct action to silver in the nonreversible formation of silver-sulfhydryl complexes that could not function as hydrogen carriers:

2R-SH R-S-S-R (normal sulfhydryl) + H2 (30)

R-SH + Ag+ R-SAg (inactive silver complex) + (31)

He considered silver to be bacteriostatic as well as bactericidal. This would explain the relatively long contact times that are required for antibacterial activity at the concentrations that are normally used.

Zimmerman (1952) demonstrated that low concentrations of silver do not enter the cell but are adsorbed onto the bacterial surface just as silver tends to be adsorbed on other surfaces (Chambers and Proctor, 1960). According to Chang (1970), the adsorbed silver ions must immobilize the dehydrogenation process because bacterial respiration takes place at the cell surface membrane.


As a disinfectant, silver may be applicable to home treatment systems and swimming pools. These applications may be more effective in keeping filters free from bacteria than in actually inactivating organisms that are suspended in the water. This is because of the relatively low rate of solution of silver in water and its low biocidal activities.

Woodward (1963) ably summarized the situation: "The high cost of silver will limit it to specialty uses. The fact that silver does not impart taste, odor, or color to water makes it attractive for use. Its slow bactericidal action, although a disadvantage in some situations, may be an advantage in others, particularly where water is stored for long times before use, as on shipboard. Until some of the uncertainties regarding silver are resolved, it would be prudent to use it as a drinking water disinfectant only in situations where substantial factors of safety can be provided and where the bactericidal effectiveness of the procedure can be monitored."

Silver and its compounds are weak, costly disinfectants that are unsuitable for use in municipal drinking water supplies. To achieve acceptable disinfection in a reasonable time would require concentrations exceeding the MCL of 0.05 mg/liter.

Research Recommendations

Now that adequate chemical analytical techniques are available for measuring low concentrations of silver, it would be desirable to conduct disinfection studies in which both dose and residual are measured accurately. To relate disinfection efficiency to more conventional practices, parallel studies using chlorine should be conducted.


Electromagnetic radiation, in wavelengths from 240 to 280 nm, is an effective agent for killing bacteria and other microorganisms in water

(Luckiesh and Holladay, 1944). Conveniently, from a practical point of view, from 30% to more than 90% of the energy emitted by a low-pressure mercury arc, which is enclosed in special UV transmitting glass, is emitted at a wavelength of 253.7 nm (Anonymous, 1960; Childs, l%2; Luckiesh and Taylor, 1946).

Two basically different physical arrangements are commonly used for the application of UV light to water. In one, lamps are placed above the solution to be disinfected at the apex or focus of parabolic or elliptical reflectors. For this purpose, aluminum is preferred because of its high reflectance for the germicidal 253.7-nm wavelength (Luckiest and Taylor, 1946). Although this is an efficient way of applying UV radiation to water, the open nature of the structure can permit contamination. Furthermore, it must operate at atmospheric pressure. Tubular reactors are more common in water treatment because they are sealed and operate under pressure.

To increase intensities and permit higher flow rates, multiple lamp reactors are being designed and put into use. These units generally contain lamps that are positioned parallel to the flow of water through them. A recent patent (Wood, 1974) describes an apparatus using a water film approximately 0.64 cm thick and a flow of liquid that is perpendicular to the lamp in a baffled system. This type of unit is designed for disinfection of fluids with low UV transmittance. For liquids with high transmittance like drinking water, such a design is inefficient, because insufficient depth is available for absorption of UV and a large fraction of the radiant energy is dissipated as heat when the UV is reflected on the walls of the contactor.

In units containing lamps that are surrounded by water, an insulation space must be provided to maintain their efficiency. The maximum efficiency of modern cold cathode lamps is near 400C, dropping off to a 50% output at 240C and 650C. Lamps are normally placed in sleeves that are made of high-silica glass or quartz in order to maximize transmission at the 253.7-nm wavelength. Solarization or opacity can develop in the sleeve as it deteriorates with age. The sleeves must be cleaned regularly for efficient functioning.

Another consideration in the mechanical design of UV contactors is the degree of agitation that is provided and the plug flow characteristics of the contactor. In all disinfection systems that require several orders of magnitude of disinfection, short circuiting or nonplug flow characteristics of contactors, which bypass a part of the flow, limit the efficiency of the process. It was recognized by Cortelyou et al. in 1954 that the degree of agitation in the UV disinfection of water is important in bringing the

target microorganism into close proximity to the UV source where the intensity is highest.


The dose, D, of electromagnetic radiation that is applied to a solution is commonly measured as the intensity of radiant energy input, lo, at the lamp surface or at some given distance from the lamp (Jo being expressed as uW/cm2), multiplied by the time of exposure, 1, in seconds (Luckiesh and Holladay, 1944):

D = 10 t ILW e S/Cm2 (32)

Another measurement of dose is chemical actinometry (Calvert and Pitts, 1966). In this method, a photochemical reaction with a known quantum yield is used to measure the intensity (quanta per second) of light that is absorbed by the actinometry solution. Knowing the time of exposure, the volume of solution, and the volume of sample analyzed, the moles of photochemical reaction per unit of time can be related through the quantum yield to the Einsteins or moles of photons per second that are absorbed by the fluid. From the wavelength of radiation, X, and the area of lamp surface, the average intensity at the lamp surface can be determined.

E = hc/A = 1.2 x 107/X We s/Einstein run (33)

Neither of these measurements of energy input, I~, per unit area of lamp exposed considers the change in intensity through the depth of exposed fluid and the solid angle over which this energy interacts. Recently, attempts have been made to correct the decrease in intensity with depth in studies of wastewater disinfection where this problem is acute (Roeber and Hoot, 1975; Severin, 1978; Venosa et al., 1978). These studies have shown that the usual water quality indices, such as chemical oxygen demand (COD), color, and turbidity, do not adequately predict the loss of intensity through the solution. The direct measurement of UV transmittance has been successful. In spite of these difficulties, Huff et al. (1965) suggested that color at a maximum level of S units, iron at 3.7 mg/liter, and turbidity up to S units did not decrease treatment efficiency below acceptable limits in a unit with a maximum water depth of approximately 7.5 cm.

Biocidal Dose

The biocidal dose of UV energy consists of the intensity of UV energy that is absorbed at the reactive site within the microbe over the time of interaction. The biocidal dose is then a function of the energy input from the UV source into the solution, dispersion of the energy as a function of distance from the source, the depth of the fluid between the organisms and the source as well as its absorptivity, and, finally, the losses and reflection of UV light within the contactor. All of these factors determine the actual intensity of radiation that is available to the microorganism at any one point within the contactor (Childs, 1962; Luckiesh and Holladay, 1944; Luckiesh et al., 1944).

Quite early in the study of the batericidal action of UV light, Gates (1929) found that the relation between the intensity of incident energy and time required for bacterial destruction were not of equal importance in determining experimental disinfection. However, Hoather (1955) found that when one corrects for the transmitted intensity of radiation, taking into consideration the absorptivity of the water and dispersion as a function of distance from the source, the required time of exposure is inversely proportional to the calculated intensity of radiation penetrating the water for a given degree of inactivation. Finally, Oliver and Cosgrove (1975) used chemical actinometry and a pulsed laser to show that their coliform and streptococcal inactivation in wastewater depended only on the total dose that is delivered and not on the dose rate or intensity of light that impinged on the sample. Thus, the same number of

s/cm2 provided the same disinfection regardless of the time and intensity that were used to produce that dosage.

Bacterial spores have also been studied with chemical actinometry. For certain stocks, the rate at which the energy was deposited affected the degree of response (Powers et al, 1974). Powers and his colleagues observed that the effect of the dose rate was not due to geometric factors or intensity variations with depth and solution absorptivity. This explains many earlier results. They did find that dependence on dose rate was seen only for some of their stocks of spores, it was not as dramatic with monochromatic-filtered light of a wavelength of 253.7 nm, and it was observed only at very low doses. The explanations usually given for these low-dose effects are photochemical back reactions, enzymatic repair, and photoreactivation mechanisms (Deering and Setlow, l%3; Harm, 1968).


UV radiation produces no residual. Therefore, monitoring and control of disinfection efficiency are more difficult for UV than for chemical disinfectants.

However, this is not a major problem, because disinfection can be controlled by adjusting the contact time and the UV energy that is transmitted through the solution. The major disadvantage is the lack of a tracer for ensuring the integrity of the distribution system.

Biocidal Activity

UV disinfection follows Chick's Law of kinetics (Luckiesh and Holiaday,



-logNIN0 Q (34)

where N is the number surviving at time t of an original population N0 that has been exposed to an intensity 1. Q is the dose for 1 log survival, one lethal unit exposure, or a lethe.

The variables discussed above interfere with the accurate determination of the average intensity and time of exposure in a practical UV contactor. Consequently, there is considerable disagreement in the literature concerning the absolute magnitude of Q for the microorganisms that have been studied.

Of the Gram-negative bacteria, Escherichia coli is consistently more resistant to disinfection by UV than are the Salmonella and Shigella species that have been studied (Cortelyou et al., 1954; Kawabata and Harada, 1959). The vegetative forms of the Gram-positive bacteria were more resistant than E. coli. Kawabata and Harada found Streptococcus faecalis nearly 3 times more resistant than E. coli, while Bacillus subtilis was 4 times more resistant. The spores of B. subtilis were 6 times more resistant than E. coli.

Luckiesh and Holladay (1944) reported that a dosage of 2,400 'LW . s/cm2 at a wavelength of 254 nm produced one lethe of disinfection of E. coli. Huff et a!. (1965) studied a two-lamp shipboard disinfection unit with an approximately 75-cm long, 20-cm diam. unbaffled stainless steel cylinder. Their dosages were generally quite high, ranging from 3,000 to 11,000 'LW. s/cm2 at a maximum water depth of approximately 19 cm. At the lowest dosage, there was a 0.02% to 0.04% survival of E. coli, which, assuming Chick's Law, gives a dosage of less than 400 'LW. s/cm2/lethe. They also found approximately 2 logs

TABLE 11-24 Ultraviolet Energy Necessary to Inactivate Various Organisms

Test Lethal Dose,

Microorganism ~W.5,cm2)

Escheridua coli 360

&GpJ~YIococcusGlLreus 210

Serraria marcescerts 290

Sarcina luzea 1,250

Badllusglobiggiisrores 1,300

T3 coliphage 160

Poliovirus 780

Vaccinia virus 30

SernlikiForrestvinss 470

EMC virus 650

a Data from Morris, 1972.

less spore inactivation with the same dose as required for vegetative cells. They reported more than 4 logs of inactivation of polio-, echo-, and coxsacltievirus with 4,000 ~W s/cm2. Based on these results, the U.S. Department of Health, Education, and Welfare issued a policy statement on April 1, 1966, stating criteria for the acceptability of UV disinfecting units as a minimum dosage of 16,000 ~W. s/cm2 with a maximum water depth of approximately 7.5 cm.

Morris (1972) has compared the dose of 25~nm UV energy required to inactivate a range of microorganisms suspended in droplets of buffer solution on an aluminum surface. Their results are shown in Table 11-24 expressed as 1£W s/cm2/lethe.

Required dosages reported by Morris (1972) are similar in magnitude to those found by Huff et al. (1965). Both Huff's values and those of Morris (1972) are lower than those of Luckiesh and Holladay (1944), partly because of the additional dose of UV that was provided by reflections from the contactor and aluminum surface, which produce an estimated dosage approximately twice the values given. No studies of the action of UV light on parasitic helminths and protozoa in the treatment of drinking water appear to have been reported.

Mechanism of Action

Inactivation by UV light is believed to act through the direct absorption of UV energy by the microorganism, causing a molecular rearrangement

of one or more of the biochemical components that are essential to the organism's functioning. The major site 9f UV absorption in microorganisms is the purine and pyrimidine components of the nucleoproteins. Since the relative efficiency of disinfection by UV energy as a function of wavelength follows the absorption spectrum of these chromophore groups and because the first law of photochemistry states that the light that is absorbed by a molecule can produce photochemical change, the mechanism of action of UV probably occurs through the progressive biochemical change that is produced primarily in the nucleoproteins. Witkin (1976) has reviewed both UV mutagenesis and inducible DNA repair. The major specific mechanisms that have been suggested for UV damage involve the reversible formation of pyrimidine hydrates and pyrimidine dimers. Breaks in the bonding structure, known as nicking, also occur. The older literature has been reviewed by Reddish (1957).

The repair of damaged nucleoprotein with light of a wavelength longer than that of the damaging radiation is commonly referred to as photoreactivation (Witkin, 1976). This process, originally identified by Kelner (1949), has been studied extensively. It occurs in the visible wavelength range from 300 to 550 nm (Jagger, 1960; Kelner, 1949). In addition, repair of nucleoproteins that have been damaged by UV radiation can also occur in the dark. The mechanisms of repair of damage to DNA caused by UV radiation are generally thought to involve induced enzymes (Witkin, 1976), but repair has also been observed in inactive systems for which photochemical back reactions have been suggested (Powers et al, 1974).


UV light produces no residual. Therefore, if a residual is desired, another disinfectant must be used. Current technology is limited to the use of UV light in small systems. Requirements for equipment maintenance limits widespread adoption of UV light for drinking water disinfection.


Tables 11-25 and 11-26 present significant characteristics of the methods that are considered for disinfection of drinking water.

The biocidal activity of various disinfectants can be compared

conveniently through the numerical value of the product (c . t) of the

concentration of disinfectant (c) multiplied by the time of exposure (t)

that is required to achieve a particular degree of inactivation (e.g., 99%)

under similar conditions of pH, temperature, etc. The lower its c t product, the more effective the disinfecting action of a particular agent.

Table 11-27 displays c t products for the methods that are regarded as the major possibilities for drinking water disinfection. The c t products for similar inactivation under the same conditions obtained for other possible agents (ranked less promising by this or other criteria) are up to several orders of magnitude greater than those shown in Table II-27, e.g., 33,000 and 106 for inactivation of E. coli and poliovirus 1 by hydrogen peroxide, respectively. Therefore, they were not included.


Table 11-28 provides the current status of theoretically possible methods for disinfecting drinking water. To derive the conclusions in the table, the specific biocidal activity data summarized in Table 11-27 were considered as well as information (or lack of it) on the practical application and reliability of the methods.

Chlorination, ozonization, and the use of chlorine dioxide come closest to meeting the criteria desired. At this time none of the other possibilities considered can substitute for techniques presently used to disinfect drinking water.

The ultimate choice between methods will require weighing the characteristics detailed in this evaluation against the nature of byproducts to be expected from the use of the particular method and. the potential toxicities of the by-products. Evaluations of these aspects of drinking water disinfection, also carried out by the Safe Drinking Water Committee, are reported in parallel to this study and should be consulted before final conclusions are drawn.


The lack of data, particularly data on pathogens or data that are amenable to comparison, on the biocidal efficacy of various disinfectants needs to be remedied. Particular emphasis should be placed on the methodology for determining biocidal activity and defining those parameters that need to be controlled for the study to yield data for comparative purposes. Disinfection efficacy studies should be conducted with actual pathogens (bacteria, viruses, and cysts of Giardia lamblia and Entamoeba histolytica) as well as with model systems. Careful control of the character and concentration of the disinfecting species and of the

TABLE 11-25 Summary of Major Possible Disinfection Methods for Drinking Water Efficacy in Demand-Free Systemsb


Disinfection Protozoan of Residual in

Agenta Technological Status Bacteria Viruses cysts Distribution System

Chroninec Widespread use in U.S. drinking

As hypochlorous water

acid (PIOCI)

As hypochlorite

ion (CCI)

Ozonec Widespread use in drinking water outside United States, particularly in France, Switzerland, and the province of Quebec

Chlorine Widespread use for disinfection

dioxidec. (both primary and for distribution system residual) in Europe;

limited use in United States to



No residual possible

NDRd Fair to good (but possible

health effects)


As diatomic iodine (12)

As hypolodous acid (HOI)



counteract taste and odor problems and to disinfect drinking water

No reports of large-scale use in drinking water

Good (but possible health elfects)

+ +++

Limited use for disinfection of Fair

drinking waler

Limited present use on a large scale ++ + Excellent

in U.S. drinking water

a The sequence in which these agents are listed does not constitute a ranking.

b Ratings: ++++, excellent biocidal activity; +++, good biocidal activity; ++ moderate biocidal activity; + low biocidal activity; i of little or questionable value.

By-product production and disinfectant demand are reduced by removal of organics from raw water prior to disinfection.

d Either no data reported or only available data were not free from confounding factors, thus rendering them not amenable to comparison with other data.

MCL 1.0 mg/liter because of health elfects (Symons el al., 1977).

I Poor in the presence of organic material.

TABLE 11-26 Summary of Minor Possible Disinfection Methods for Drinking Water a'

Erneacy in Demand-Free Systemsb


Disinfection Protozoan of Residual in

Agent" Technological Status Bacteria Viruses cysts Distribution System

Ferrate No reports of use in drinking water ++ NDRC Poor

High p11 No reports of large-scale use in + + + ++ + NDRC Feasibility restricted since

conditions drinking water consumption of high pH water

(p11 12-12.5) not recommended

Hydrogen No reports of large-scale use in i I NDRC Poor

peroxide drinking water

Ionizing No reports of use in drinking water ++ ++ NDRC No residual possible


Potassium Limited use for disinfection + NDRC NDRC Good, but aesthetically undesir

permanganate able

Silverd No reports of large-scale use in + NDRC + Good, but possible health elfects

drinking water

UV light Use limited to small systems +++ +++ NDRC No residual possible

"The sequence in which these agents are listed dces not constitute a ranking.

b Ratings: ++++, excellent biocidal activity; +++, good biocidal activity; ++, moderate biocidal activity; +,

tionable value.

low biocidal activity; I, of little or ques

Either no data reported or only available data were not free from confounding factors, thus rendering them not amenable to comparison with other data.

d MCL 0.05 mg/liter because of health elfects (Symons el al., 1977).

Demand-Free Systemsa

F. coli Poliovirus 1 Enta'noeba ~iisioIytica cysts

Disinfection Tempera- Tempera. Tempera-

Agent p11 ture, CC c.tb pH ture, Cc c~ tb p11 ture, CC c.tb

Ilypochlorousacid 6.0 5 0.04 6.0 0 1.0 7 30 20

6.0 S 2.0

7.0 0 1.0

Ilypochlorite ion 10.0 5 0.92 10.5 S 10.5 NDRC

Ozone 6.0 II 0.031 7.0 20 0.005 7.5-8.0 19 1~5d

7.0 12 0.002 7.0 25 0.42

Chlorinedioxide 6.5 20 0.18 7.0 15 1.32

6.5 15 0.38 7.0 25 1.90 NDRC

7.0 . 25 0.28

Iodine 6.5 20-25 0.38 7.0 26 30 7.0 30 80

7.5 20-25 0.40

Bromine NDRC 7.0 20 0.06 7.0 30 18


Monochlorarnine 9.0 15 64 9.0 15 900 NDRC

9.0 25 40 9.0 25 320

Dichlorarnine 4.5 15 5.5 4.5 15 5,000 NDRC

0Conditions closest to pH 7.0 and 200C were selected from studies discussed in the text. Values for other conditions and agents appear in the text along with discussions of the cited studies.

b Concentration of disinfectant (mg/liter) times contact time (mm).

Either no data reported or only available data were not free from confounding factors, thus rendering them not amenable to comparison with other

data. --

d This value was derived primarily from experiments that were conducted with tap water; however, some parallel studies with distilled water showed es- ~

sentially no ditferences in inactivation rates.


TABLE 11-28 Status of Possible Methods for Drinking Water Disinfection

Suitability Suitability for

Disinfection as Inactivating Drinking Water Agent Agent Limitations Disinfectiona

Chlorine Yes Efficacy decreases with increasing p11; affected by ammonia or Yes

organic nitrogen

Ozone Yes On-site generation required; no residual; other disinfectant Yes

needed for residual

Chlorine dioxide Yes On-site generation required; interim MCL 1.0 mg/liter Yes

Iodine Yes Biocidal activity sensitive to pH No

Bromine Yes Lack of technological experience; activity may be pH sensitive No

Chloramines No Mediocre bactericide, poor virucide Nob

Ferrate Yes Moderate bactericide; good virucide; residual unstable; lack of No

technological experience

High pH conditions No Pcor biocide No

Hydrogen peroxide No Pcor biocide No

Ionizing radiation Yes Lack of technological experience No

Potassium permanganate No Poor biocide No

Silver No Poor biocide; MCL 0.05 mg/liter No

UV light Yes Adequate biocide; no residual; use limited by equipment No

maintenance considerations

a This evaluation relates solely to the suitability for controlling infectious disease transmission. See conclusions. b Chlorarnines may have use as a secondrtry disinfectant in the distribution system in view of their persistence.

degree of aggregation and other characteristics of the test organisms (or cysts) are required in these tests.

Studies should be conducted to define the mechanisms resulting in the different responses to disinfectants of laboratory-acclimated cultures and organisms that have been freshly isolated from natural conditions or that are resistant to specific disinfectants for other reasons. Techniques should be developed to facilitate accurate predictions from laboratory data to disinfectant susceptibilities under operating conditions in treatment plants.

There are no practical assays for many human viruses of known or suspected importance, e.g., hepatitis A virus and rotaviruses. Assay models should be developed for such viruses, and studies should be conducted on their susceptibilities to disinfection. Support is also needed for research that would result in better appraisals of the significance to human health of the presence of viruses in water supplies. Such studies should include investigation of the causative agents that are involved in the large number of illnesses presently characterized as "gastroenteritis of unknown etiology" and of the doses of various enteroviruses that are required to establish infections.

In investigations of factors affecting biocidal activity in full-scale treatment plant operations, priority should be given to studying more thoroughly the effects of turbidity (both organically and inorganically caused) on disinfection.

It should be determined if any data are available from treatment plants on the effectiveness of ozone or chlorine dioxide in disease prevention, where these methods of disinfection are not used sequentially or, in the case of ozone, with subsequent chlorination. If such data, free of complications, are available, they should be compared with the records on the efficacy of chlorination disinfection.