Drinking Water Treatment

Clean drinking water is the most important public health factor. But well over 2 billion people worldwide do not have adequate supplies of safe drinking water. Worldwide, between 15 to 20 million babies die every year from water-borne diarrheal diseases such as typhoid fever, dysentery, and cholera. Contaminated water supplies and poor sanitation cause 80% of the diseases that afflict people in the poorest countries. The development of municipal water purification in the last century has allowed cities in the developed countries to be essentially free of water-carried diseases. Since the introduction of filtration and disinfection of drinking water in the U.S., water-borne diseases, such as cholera and typhoid, have been virtually eliminated.

However, in 1974, it was discovered that water disinfectants react with organic compounds that are naturally occurring in water and form unintended disinfection byproducts (DBPs) that may cause health risks.3 724 Trihalomethane DBPs were regulated by the EPA in 1979.10

Since then, several DBPs (bromodichloromethane, bromoform, chloroform, dichloroacetic acid, and bromate) have been shown to be carcinogenic in laboratory animals at high doses. Some DBPs (bromodichloromethane, chlorite, and certain haloacetic acids) also can cause adverse reproductive or developmental effects in laboratory animals. In the belief that DBPs present a potential public health risk, the EPA published guidelines for minimizing their formation3839 and established standards in 1998 for drinking water concentrations of DBPs and disinfectant residuals (see Appendix A). The goal of EPA disinfectant and disinfection byproduct regulations is to balance the health risks of pathogen contamination, normally controlled by water disinfection, against DPB formation.

Water Sources

Drinking water supplies come either from surface waters or groundwaters. In the U.S., groundwater sources or wells supply about 53% of all drinking water and surfacewater sources, such as reservoirs, rivers, and lakes, supply the remaining 47%. Groundwater comes from underground aquifers into which wells are drilled to recover the water. Wells range from tens to hundreds of meters deep. Generally, water in deep aquifers is replaced by percolation from the surface very slowly over hundreds to thousands of years. Water in the deep Ogallala aquifer in the Great Plains region of the U.S. is estimated to be thousands of years old and is called "fossil water." Replenishment of such aquifers occurs over thousands of years, and it is easy to withdraw water from them at a rate that greatly exceeds replacement. Such aquifers are essentially nonrenewable resources in our lifetime. The Ogallala aquifer has been depleted significantly over the past several decades, principally by agricultural irrigation.

Groundwater tends to be less contaminated than surface water. It is normally more protected from surface contamination and, because it moves more slowly, organic matter has time to be decomposed by soil bacteria. The soil itself acts as a filter so that less suspended matter is present.

Surface waters come from lakes, rivers, and reservoirs. It usually has more suspended materials than groundwater and requires more processing to make it safe to drink. Surface waters are used for purposes other than drinking and often become polluted by sewage, industrial, and recreational activities. On most rivers, the fraction of "new" water diminishes with distance from the head waters, as the water becomes more and more used. On the Rhine river in Europe, for example, communities near the mouth of the river receive as little as 40% "new" water in the river. All the other water has been previously discharged by an upstream city or originates as nonpoint source return flow from agricultural activities. Water treatment must make this quality of river water fit to drink. Filtration through sand was the first successful method of municipal water treatment, used in London in the middle 1800s. It led to an immediate decline in the amount of water-borne diseases.

Water Treatment

Major changes are occurring in the water treatment field driven by increasingly tighter water quality standards, a steady increase in the number of regulated drinking water contaminants (from about 5 in 1940 to around 100 in 1999), and new regulations affecting disinfection and disinfection byproducts. Municipalities are constantly seeking to refine their water treatment and provide higher quality water by more economical means. A recent development in water treatment is the application of membrane filtration to drinking water treatment. Membrane filters have been refined to the point where, in certain cases, they are suitable as stand-alone treatment for small systems. More often, they are used in conjunction with other treatment methods to economically improve the overall quality of finished drinking water.

Basic Drinking Water Treatment

The purpose of water treatment is (1) to make water safe to drink by ensuring that it is free of pathogens and toxic substances, and (2) to make it a desirable drink by removing offensive turbidity, tastes, colors, and odors.

Conventional drinking water treatment addresses both of these goals. It consists of four steps:

1) Primary settling

2) Aeration

3) Coagulation and filtration

4) Disinfection

Not all four of the basic steps are needed in every treatment plant. Groundwaters, in particular, usually need much less treatment than surface waters. Groundwaters may need no settling, aeration, or coagulation. For clean groundwaters, only a little chlorine 0.16 ppm) is added to protect the water while in the distribution system. The relatively new treatment technology of membrane filtration is increasingly being used in conjunction with the more traditional treatments and as a stand-alone treatment.

Primary Settling

Water, which has been coarsely screened to remove large particulate matter, is brought into a large holding basin to allow finer particulates to settle. Chemical coagulants may be added to form floc. Lime may be added at this point to help clarification if pH < 6.5. The floc settles by gravity, removing solids larger than about 25 microns.


The clarified water is agitated with air. This promotes oxidation of any easily oxidizable substances — for example those which are strong reducing agents. Chlorine will be added later. If chlorine were added at this point and reducing agents were still in the water, they would reduce the chlorine and make it ineffective as a disinfectant.

Ferrous iron, Fe2+, is a particularly troublesome reducing agent. It may arise from the water passing through iron pyrite (FeS2) or iron carbonate (FeCO3) minerals. With dissolved oxygen present, Fe2+ is oxidized to Fe3+, which precipitates as ferric hydroxide, Fe(OH)3, at any pH greater than 3.5. Fe(OH)3 gives a metallic taste to the water and causes the ugly red-brown stain commonly found in sinks and toilets in iron-rich regions. The stain is easily removed with weak acid solutions, such as vinegar.

Coagulation and Filtration

The finest sediments, such as pollen, spores, bacteria, and colloidal minerals, do not settle out in the primary settling step. For the finished water to look clear and sparkling, these fine sediments must be removed. Hydrated aluminum sulfate, Al2(SO4)318 H2O, sometimes called alum or filter alum, applied with lime, Ca(OH)2, is the most common filtering agent used for secondary settling.

At pH = 6-8, Al(OH)3(s) is formed as a light, fluffy, gelatinous flocculant having an extremely large surface area that attracts and traps small suspended particles, carrying them to the bottom of the tank as the precipitate slowly settles. In this pH range, Al(OH)3 is near its minimum solubility and very little Al3+ is left in solution. Additonal filtration with sand beds or membranes may be used in a final polishing step before disinfection.


Killing bacteria and viruses is the most important part of water treatment. Proper disinfection provides a residual disinfectant level that persists throughout the distribution system. This not only kills organisms that pass through filtration and coagulation at the treatment plant, it prevents reinfection during the time the water is in the distribution system. In a large city, water may remain in the system for 5 days or more before it is used. Five days is plenty of time for any missed microorganisms to multiply. Leaks and breaks in water mains can permit recontamination, especially at the extremities of the system where the pressure is low. High pressure causes the flow at leaks to always be from the inside to the outside. But at low pressure, bacteria can seep in.

As a result of concerns about DBPs, the EPA and the water treatment industry are placing more emphasis on the use of disinfectants other than chlorine, which at present is the most commonly used water disinfectant. Another approach to reducing the probability of DBP formation is by removing DBP precursors (naturally occurring organic matter) from water before disinfection. However, use of alternative disinfectants has also been found to produce DBPs. Current regulations try to balance the risks between microbial pathogens and DBPs. DBPs include the following, not all of which pose health risks:

• Halogenated organic compounds, such as trihalomethanes (THMs), haloacetic acids, haloketones, and other halogenated compounds that are formed primarily when chlorine or ozone (in the presence of bromide ion) are used for disinfection.

• Organic oxidation byproducts, such as aldehydes, ketones, assimilable organic carbon (AOC), and biodegradable organic carbon (BDOC). The latter two DBPs result from large organic molecules being oxidized to smaller molecules, which are more available to microbes, plant, and aquatic life as a nutrient source. Oxidized organics are formed when strong oxidizing agents (ozone, permanganate, chlorine dioxide, or hydroxyl radical) are used.

• Inorganic compounds, such as chlorate, chlorite, and bromate ions. These are formed when chlorine dioxide and ozone disinfectants are used.

Disinfection Procedures

Most disinfectants are strong oxidizing agents that react with organic and inorganic oxidizable compounds in water. In some cases, the oxidant is produced as a reaction byproduct — hydroxyl radical is formed in this way. In addition to destroying pathogens, disinfectants are also used for removing disagreeable tastes, odors, and colors. They also can assist in the oxidation of dissolved iron and manganese, prevention of algal growth, improvement of coagulation and filtration efficiency, and control of nuisance water organisms such as Asiatic clams and zebra mussels.

The most commonly used water treatment disinfectant is chlorine. It was first used on a regular basis in Belgium in the early 1900s. Other disinfectants sometimes used are ozone, chlorine dioxide, and ultraviolet radiation. Of these, only chlorine and chlorine dioxide have residual disinfectant capability. With chlorine or chlorine dioxide, adding a small excess of disinfectant maintains protection of the drinking water throughout the distribution system. Normally, a residual chlorine or chlorine dioxide concentration of about 0.2 to 0.5 mg/L is sought. Disinfectants that do not provide residual protection are normally followed by a low dose of chlorine in order to preserve a disinfection capability throughout the distribution system.

Part of the disinfection procedure involves removing DBP precursors, mainly total organic carbon (TOC), by coagulation, water softening, or filtration. A high TOC concentration (greater than 2.0 mg/L) indicates a high potential for DBP formation. Typical required reduction percentages of TOC for conventional treatment plants are given in Table 6.3.


Required Percentage Removal of Total Organic Carbon by Enhanced Coagulationa for Conventional Water Treatment Systems1

Source Water TOC

Source Water Alkalinity (mg/L as CaCO3)

a Enhanced coagulation is defined, in part, as the coagulant dose where an incremental addition of 10 mg/L of alum (or an equivalent amount of ferric salt) results in a TOC removal to below 0.3 mg/L.

b Applies to utilities using surface water and groundwater impacted by surface water.

Disinfection Byproducts and Disinfection Residuals

The principal precursor of organic DBPs is naturally occurring organic matter (NOM). NOM is usually measured as total organic carbon (TOC) or dissolved organic carbon (DOC). Typically, about 90% of TOC is in the form of DOC (DOC is defined as the part of TOC that passes through a 0.45 |jm filter). Halogenated organic byproducts are formed in water when NOM reacts with free chlorine (Cl2) or free bromine (Br2). Free chlorine may be introduced when chlorine gas, chlorine dioxide, or chloramines are added for disinfection. Free bromine is a product of the oxidation by disinfectants of bromide ion already present in the source water.

Reactions of strong oxidants with NOM also form nonhalogenated DBPs, particularly when nonchlorine oxidants such as ozone and peroxone are used. Common nonhalogenated DBPs include aldehydes, ketones, organic acids, ammonia, and hydrogen peroxide.

Bromide ion (Br) may be present, especially where geothermal waters impact surface and groundwaters, and in coastal areas where saltwater incursion is occurring. Ozone or free chlorine oxidizes Br- to form brominated DBPs such as: bromate ion, bromoform, cyanogen bromide, bromopicrin, and brominated acetic acid.

Strategies for Controlling Disinfection Byproducts

Once formed, DBPs are difficult to remove from a water supply. Therefore, DBP control is focused on preventing their formation. Chief control measures for DBPs are

• Lowering NOM concentrations in source water by coagulation, settling, filtering, and oxidation

• Using sorption on granulated activated carbon (GAC) to remove DOC

• Moving the disinfection step later in the treatment train, so that it comes after all processes that decrease NOM

• Limiting chlorine to providing residual disinfection, following primary disinfection with ozone, chlorine dioxide, chloramines, or ultraviolet radiation

• Protection of source water from bromide ion

Table 6.4 is a list of the cancer classifications assigned by the EPA for disinfectants and DBPs as of January 1999.


EPA Cancer Classifications for Disinfectants and DBPs38

Compound Cancer Classification"









Monochloroacetic acid

Dichloroacetic acid


Trichloroacetic acid

















Chloral hydrate


Cyanogen chloride










Hypochlorous acid



Chlorine dioxide


a The EPA classifications for carcinogenic potential of chemicals are37 A: Human carcinogen; sufficient evidence in epidemiologic studies to support causal association between exposure and cancer. B: Probable human carcinogen; limited evidence in epidemiologic studies (B1) and/or sufficient evidence from animal studies (B2). C: Possible human carcinogen; limited evidence from animal studies and inadequate or no data in humans. D: Not classifiable; inadequate or no animal and human evidence of carcinogenicity. E: No evidence of carcinogenicity for humans; no evidence of carcinogenicity in at least two adequate animal tests or in adequate epidemiologic and animal studies. Note: Not all of the EPA cancer classifications are found among the listed disinfectants and DBPs. The EPA is in the process of revising these cancer guidelines.

Chlorine Disinfection Treatment

At room temperature, chlorine is a corrosive and toxic yellow-green gas with a strong, irritating odor. It is stored and shipped as a liquefied gas. Chlorine is the most widely used water treatment disinfectant because of its many attractive features:

• It is effective against a wide range of pathogens commonly found in water, particularly bacteria and viruses.

• It leaves a residual that stabilizes water in distribution systems against reinfection.

• It is economical and easily measured and controlled.

• It has been used for a long time and represents a well-understood treatment technology. It maintains an excellent safety record despite the hazards of handling chlorine gas.

• Chlorine disinfection is available from sodium and calcium hypochlorite salts, as well as from chlorine gas. Hypochlorite solutions may be more economical and convenient than chlorine gas for small treatment systems.

In addition to disinfection, chlorination is used for

• Taste and odor control, including destruction of hydrogen sulfide.

• Color bleaching.

• Controlling algal growth.

• Precipitation of soluble iron and manganese.

• Sterilizing and maintaining wells, water mains, distribution pipelines, and filter systems.

• Improving some coagulation processes.

Problems with chlorine usage include

• Not effective against Cryptosporidium and limited effectiveness against Giardia lamblia protozoa.

• Reactions with NOM can result in the formation of undesirable DBPs.

• The hazards of handling chlorine gas require special equipment and safety programs.

• If site conditions require high chlorine doses, taste and odor problems may arise.

Chlorine dissolves in water by the following equilibrium reactions:

At pH values below 7.5, hypochlorous acid (HOCl) is the dominant dissolved chlorine species. Above pH 7.5, chlorite anion (OCl) is dominant (see Figure 6.3). The formation of H+ means that chlorination reduces total alkalinity.

The active disinfection species, Cl2, HOCl, and OCl-, are called the total free available chlorine. All these species are oxidizing agents, but chloride ion (Cl-) is not. HOCl is about 100 times more effective as a disinfectant than OCl-. Thus, the amount of chlorine required for a given level of disinfection depends on the pH. Higher doses are needed at a higher pH. At pH 8.5, 7.6 times as much chlorine must be used as at pH 7.0, for the same amount of disinfection. HOCl is more effective than OCl- because, as a neutral molecule, it can penetrate cell membranes of microorganisms more easily than OCl- can.

When chlorine gas is added to a water system, it dissolves according to Equations 6.11-6.13. All substances present in the water that are oxidizable by chlorine constitute the chlorine demand. Until oxidation of these substances is complete, all the added chlorine is consumed, and the net dissolved chlorine concentration remains zero as chlorine is added. When no chlorine-oxidizable matter is left, for example when the chlorine demand has been met, the dissolved chlorine concentration (chlorine residual) increases in direct proportion to the additional dose (see Figure 6.4).

If chlorine demand is zero, residual always equals the dose, and the plot is a straight line of slope = 1, passing through the zero. Chlorine is supplied as the bulk liquid under pressure, the boiling point of chlorine gas is -35° C at 1 atmosphere pressure. The total time of water in the

_ /—

_ /HOCl

\ /ocr



r^-—L 1 1

^r 1 1 1

FIGURE 6.3 Distribution diagram for dissolved chlorine species. Free chlorine molecules, Cl2, exist only below about pH = 2. At pH = 7.5, [HOCl] = [OCl-].

FIGURE 6.3 Distribution diagram for dissolved chlorine species. Free chlorine molecules, Cl2, exist only below about pH = 2. At pH = 7.5, [HOCl] = [OCl-].

Chlorine Dose

FIGURE 6.4 Relations among chlorine dose, chlorine demand, and chlorine residual.

chlorine disinfection tank is generally about 20-60 minutes. A typical concentration of residual chlorine in the finished water is 1 ppm or less.


In addition to chlorine gas, the active disinfecting species HOCl and OCl- can be obtained from hypochlorite salts, chiefly sodium hypochlorite (NaOCl) and calcium hypochlorite (Ca(OCl)2). The salts react in water according to Equations 6.14 and 6.15.

Notice, that while adding chlorine gas to water lowers the pH, Equations 6.11-6.13, hypochlo-rite salts raise the pH.

Sodium hypochlorite salts are available as the dry salt or in aqueous solution. The solution is corrosive with a pH of about 12. One gallon of 12.5% sodium hypochlorite solution is the equivalent of about 1 lb of chlorine gas. Unfortunately, sodium hypochlorite presents storage problems. After one month of storage under the best of conditions (low temperature, dark, and no metal contact), a 12.5% solution will have degraded to about 10%. On-site generation of sodium hypochlorite is accomplished by passing low voltage electrical current through a sodium chloride solution. On-site generation allows smaller quantities to be stored and makes the use of more stable dilute solutions (0.8%) feasible.

Calcium hypochlorite is commonly available as the dry salt which contains about 65% available chlorine. 1.5 lbs of calcium hypochlorite are equivalent to about 1 lb of chlorine gas. Storage is less of a problem with calcium hypochlorite; normal storage conditions result in a 3-5% loss of its available chlorine per year.


Chlorine dose: the amount of chlorine originally used. Chlorine residual: the amount remaining at time of analysis.

Chlorine demand: the amount used up in oxidizing organic substances and pathogens in the water, for example the difference between the chlorine dose and the chlorine residual. Free available chlorine: the total amount of HOCl and ClO- in solution. (Cl2 is not present above pH = 2.)

Drawbacks to Use of Chlorine: Disinfection Byproducts (DBPs) Trihalomethanes (THMs)

The problem of greatest concern with the use of chlorine is the formation of chlorination byproducts, particularly trihalomethanes (CHCl3, CHBrCl2, CHBr2Cl, CHBr3, CHCl2I, CHBrClI) and carbon tetrachloride (CCl4) as possible carcinogens. It was once thought that THMs were formed by chlorination of dissolved methane. It is now known that they come from the reaction of HOCl, with acetyl groups in NOM, chiefly humic acids. Humic acids are breakdown products of plant materials like lignin. There is no evidence that chlorine itself is carcinogenic.

In addition to the general strategies for controlling DBPs listed earlier, another option is available with chlorine use. Addition of ammonia with chlorination forms chloramines (see Breakpoint Chlorination to Remove Ammonia). Chloramines are weaker oxidants than chlorine and are useful for providing a residual disinfectant capability with a lower potential for forming DBPs.

Chlorinated Phenols

If phenol or its derivatives from industrial activities are in the water, taste and color can be a problem. Phenols are easily chlorinated, forming compounds with very penetrating antiseptic odors. The most common chlorinated phenols arising from chlorine disinfection are shown in Table 6.5, with their odor thresholds, several of which are in the ppb (|Jg/L) range. At the ppm level, chlorinated phenols can make water completely unfit for drinking or cooking. If phenol is present in the intake water, treatment choices are to employ additional nonchlorine oxidation for removing phenol, to remove phenol with activated charcoal, or to use a different disinfectant. The activated charcoal treatment is expensive and few communities use it.

Example 6.3

Water has begun to seep into the basement of a home. The home's foundation is well above the water table and this problem had not been experienced before. The house is located about 50 ft


Odor Thresholds of Phenol and Chlorinated Derivatives from Drinking Water Disinfection With Chlorine

Odor Threshold in Water Phenol Compound Chemical Structure (ppb)










downgradient from a main water line and one possibility is that a leak has occurred in the pipeline. The water utility company tested water entering the basement for the presence of chlorine, thinking that if the water source was the pipeline, the chlorine residual should be detected. When no chlorine was found, the utility company concluded that they were not responsible for the seep. Was this conclusion justified?

Answer: No. Water would have to travel at least 50 ft through soil from the pipeline to the house. The chlorine residual should not exceed 4 mg/L (see Appendix A) and would almost certainly come in contact with enough oxidizable organic and inorganic matter in the soil to be depleted below detection. A better water source marker would be fluoride, assuming the water supply is fluoridated. Although fluoride might react with calcium and magnesium in the soil to form solid precipitates, it is more likely to be detectable at the house than is chlorine. However, neither test is conclusive. The simplest and best test would be to turn off the water in the pipeline long enough to observe any change in water flow into the house. This, however, might not be possible. Another approach would be to examine the water line for leaks, using a video camera probe or soil conductivity measuring equipment.


Many utilities use chlorine for disinfection and chloramines for residual maintenance. Chloramines are formed in the reaction of ammonia with HOCl from chlorine — a process that is inexpensive and easy to control. The reactions are described in the section on breakpoint chlorination. Although the reaction of chlorine with ammonia can be used for the purpose of destroying ammonia, it also serves to generate chloramines, which are useful disinfectants that are more stable and longer lasting in a water distribution system than is free chlorine. Thus, chloramines are effective for controlling bacterial regrowth in water systems although they are not very effective against viruses and protozoa. The primary role of chloramines is their use as a secondary disinfectant to provide residual treatment — an application which has been practiced in the U.S. since about 1910. Being weaker oxidizers than chlorine, chloramines form far fewer disinfection byproducts. However, they are not useful for oxidizing iron and manganese. When chloramine disinfection is the goal, ammonia is added in the final chlorination step. Chloramines are always generated on site.

Optimal chloramine disinfection occurs when the chlorine:ammonia-nitrogen (Cl2:N) ratio by weight is around 4, before the chlorination breakpoint occurs. Under these conditions, monochloramine (NH2Cl) and dichloramine (NHCl2) are the main reaction products and the effective disinfectant species. The normal dose of chloramines is between 1 and 4 mg/L. Residual concentrations are usually maintained between 0.5 and 1 mg/L. The maximum residual disinfection level (MRDL) mandated by the EPA is 4.0 mg/L.

Chlorine Dioxide Disinfection Treatment

Chlorine dioxide (ClO2) is a gas at temperatures above 12°C with high water solubility. Unlike chlorine, it reacts quite slowly with water, remaining mostly dissolved as a neutral molecule. It is a very good disinfectant, about twice as effective as HOCl from Cl2 but also about twice as expensive. ClO2 was first used as a municipal water disinfectant in Niagara Falls, NY in 1944. In 1977, about 100 municipalities in the U.S. and thousands in Europe were using it. The main drawback to its use is that it is unstable and cannot be stored. It must be made and used on site, whereas chlorine can be delivered in tank cars.

Much of its reactivity is due to being a free radical. ClO2 cannot be compressed for storage because it is explosive when pressurized or when it is at concentrations above 10 percent by volume in air. It decomposes in storage and can decompose explosively in sunlight, when heated or agitated suddenly. So it is never shipped and is always prepared on site and used immediately. Typical dose rates are 0.1-1.0 ppm.

Sodium chlorite is used to make ClO2 by one of three methods:

5 NaClO2 + 4 HCl o 4 ClO2(g) + 5 NaCl + 2 H2O. (6.16)

Sodium chlorite is extremely reactive, especially in the dry form, and it must be handled with care to prevent potentially explosive conditions. If chlorine dioxide generator conditions are not carefully controlled (pH, feedstock ratios, low feedstock concentrations, etc.), the undesirable byproducts chlorite (ClO2) and chlorate (ClO3) may be formed.

Chlorine dioxide solutions below about 10 g/L will not have sufficiently high vapor pressures to create an explosive hazard under normal environmental conditions of temperature and pressure. For drinking water treatment, ClO2 solutions are generally less than 4 g/L and treatment levels generally are between 0.07 to 2.0 mg/L.

Since ClO2 is an oxidizer but not a chlorinating agent, it does not form trihalomethanes or chlorinated phenols. So it does not have taste or odor problems. Common applications for ClO2 have been to control taste and odor problems associated with algae and decaying vegetation, to reduce the concentrations of phenolic compounds, and to oxidize iron and manganese to insoluble forms. Chlorine dioxide can maintain a residual disinfection concentration in distribution systems. The toxicity of ClO2 restricts the maximum dose. At 50 ppm, ClO2 can cause breakdown of red corpuscles with the release of hemoglobin. Therefore, the dose of ClO2 is limited to 1 ppm.

Ozone Disinfection Treatment

Ozone (O3) is a colorless, highly corrosive gas at room temperature, with a pungent odor that is easily detectable at concentrations as low as 0.02 ppmv — well below a hazardous level. It is one of the strongest chemical oxidizing agents available, second only to hydroxyl free radical (HO-), among disinfectants commonly used in water treatment. Ozone use for water disinfection started in 1893 in the Netherlands and in 1901 in Germany. Significant use in the U.S. did not occur until the 1980s. Ozone is one of the most potent disinfectants used in water treatment today. Ozone disinfection is effective against bacteria, viruses, and protozoan cysts, including Cryptosporidium and Giardi lamblia.

Ozone is made by passing a high voltage electric discharge of about 20,000 V through dry, pressurized air.

Equation 6.19 is endothermic and requires a large input of electrical energy. Because ozone is unstable, it cannot be stored and shipped efficiently. Therefore, it must be generated at the point of application. The ozone gas is transferred to water through bubble diffusers, injectors, or turbine mixers. Once dissolved in water, ozone reacts with pathogens and oxidizable organic and inorganic compounds. Undissolved gas is released to the surroundings as off-gas and must be collected and destroyed by conversion back to oxygen before release to the atmosphere. Ozonator off-gas may contain as much as 3000 ppmv of ozone, well above a fatal level. Ozone is readily converted to oxygen by heating it to above 350° C or by passing it through a catalyst held above 100° C. OSHA currently requires released gases to contain no more than 0.1 ppmv of ozone for worker exposure. Typical dissolved ozone concentrations in water near an ozonator are around 1 mg/L.

The dissolved ozone gas decomposes spontaneously in water by a complex mechanism that includes the formation of hydroxyl free radical, which is the strongest oxidizing agent available for water treatment. Hydroxyl radical essentially reacts at every molecular collision with many organic compounds. The very high reaction rate of hydroxyl radicals limits their half life in water to the order of microseconds and their concentration to less than about 10-12 mol/L. Both ozone molecules and hydroxyl free radicals play prominent oxidant roles in water treatment by ozonation.

Ozone concentrations of about 4-6% are achieved in municipal and industrial ozonators. Ozone reacts quickly and completely in water, leaving no active residual concentration. Decomposition of ozone in water produces hydroxyl radical (a very reactive short-lived oxidant) and dissolved oxygen, which further aid in disinfection and diminishing BOD, COD, color, and odor problems. The air-ozone mixture is typically bubbled through water for a 10-15 minutes contact time. The main drawbacks to ozone use have been its high capital and operating costs and the fact that it leaves no residual disinfection concentration. Since it offers no residual protection, ozone can be used only as a primary disinfectant. It must be followed by a light dose of secondary disinfectant, such as chlorine, chloramine, or chlorine dioxide for a complete disinfection system.

Several ways to assist ozonation are adding hydrogen peroxide (H2O2), using ultraviolet radiation (UV), and/or raising the pH to around 10-11. Hydrogen peroxide decomposes to form the reactive hydroxyl radical, greatly increasing the hydroxyl radical concentration above that generated by simple ozone reaction with water. Reactions of hydroxyl radicals with organic matter cause structural changes that make organic matter still more susceptible to ozone attack. Adding hydrogen peroxide to ozonation is known as the Advanced Oxidation Process (AOP) or Peroxone process. UV radiation dissociates peroxide, forming hydroxyl radicals at a rapid rate. Raising the pH allows ozone to react with hydroxyl ions (OH-, not the radical HO-) to form additional hydrogen peroxide. In addition to increasing the effectiveness of ozone oxidation, peroxide and UV radiation are also effective as disinfectants. The use of these ozonation enhancers is known as the AOP process.

The equipment for ozonation is expensive, but the cost per gallon decreases with large scale operations. Generally, only large cities use ozone. ClO2 is not as problem free as ozone, but it is cheaper to use for small systems.

In addition to disinfection, ozone is used for

• DBP precursor control

• Protection against Cryptosporidium and Giardi

• Taste and odor control, including destruction of hydrogen sulfide

• Color bleaching

• Precipitation of soluble iron and manganese

• Sterilizing and maintaining wells, water mains, distribution pipelines, and filter systems

• Improving some coagulation processes

Ozone DBPs

Although it does not form the chlorinated disinfection byproducts that are of concern with chlorine use, ozone can react to form its own set of oxidation byproducts. When bromide ion (Br) is present — where geothermal waters impact surface and groundwaters or in coastal areas where saltwater incursion is occurring — ozonation can produce bromate ion (BrO3), a suspected carcinogen, as well as brominated THMs and other brominated disinfection byproducts. Controlling the formation of unwanted ozonation byproducts is accomplished by pretreatment to remove organic matter (activated carbon filters and membrane filtration) and scavenge BrO3- (pH lowering and hydrogen peroxide addition).

When bromide is present, the addition of ammonia with ozone forms bromamines — by reactions analogous to the formation of chloramines with ammonia and chlorine — and lessens the formation of bromate ion and organic DBPs.

Potassium Permanganate

Potassium permanganate salt (KMnO4) dissolves to form the permanganate anion (MnO4), a strong oxidant effective at oxidizing a wide variety of organic and inorganic substances. In the process, manganese is reduced to manganese dioxide (MnO2), an insoluble solid that precipitates from solution. Permanganate imparts a pink to purple color to water and is, therefore, unsuitable as a residual disinfectant.

Although it is easy to transport, store, and apply, permanganate generally is too expensive for use as a primary or secondary disinfectant. It is used in drinking water treatment primarily as an alternative to chlorine for taste and odor control, iron and manganese oxidation, oxidation of DPB precursors, control of algae, and control of nuisance organisms, such as zebra mussels and the Asiatic clam. It contains no chlorine and does not contribute to the formation of THMs. When used to oxidize NOM early in a water treatment train that includes post-treatment chlorination, permanganate can reduce the formation of THMs.

Peroxone (Ozone + Hydrogen Peroxide)

The peroxone process is an advanced oxidation process (AOP). AOPs employ highly reactive hydroxyl radicals (OH-) as major oxidizing species. Hydroxyl radicals are produced when ozone decomposes spontaneously. Accelerating ozone decomposition by using, for example, ultraviolet radiation or adding hydrogen peroxide, elevates the hydroxyl radical concentration and increases the rate of contaminant oxidation. When hydrogen peroxide is used, the process is called peroxone.

Like ozonation, the peroxone process does not provide a lasting disinfectant residual. Oxidation is more complete and much faster with peroxone than with ozone. Peroxone is the treatment of choice for oxidizing many chlorinated hydrocarbons that are difficult to treat by any other oxidant. It is also used for inactivating pathogens and destroying pesticides, herbicides, and volatile organic compounds (VOCs). It can be more effective than ozone for removing taste- and odor-causing compounds such as geosmin and 2-methyliosborneol (MIB). However, it is less effective than ozone for oxidizing iron and manganese. Because hydroxyl radicals react readily with carbonate, it may be necessary to lower the alkalinity in water with a high carbonate level in order to maintain a useful level of radicals. Peroxone treatment produces similar DBPs as does ozonation. In general, it forms more bromate than ozone under similar water conditions and bromine concentrations.

Ultraviolet (UV) Disinfection Treatment

Ultraviolet radiation at wavelengths below 300 nm is very damaging to life forms, including microorganisms. Low-pressure mercury lamps, known as germicidal lamps, have their maximum energy output at 254 nm. They are very efficient, with about 40% of their electrical input being converted to 254 nm radiation. Protein and DNA in microorganisms absorb radiation at 254 nm, leading to photochemical reactions that destroy the ability to reproduce. UV doses required to inactivate bacteria and viruses are relatively low, of the order of 20-40 mWs/cm2. Much higher doses, 200 mWs/cm2 or higher, are needed to inactivate Cryptosporidium and Giardia lamblia.

Color or high levels of suspended solids can interfere with transmission of UV through the treatment cell and UV absorption by iron species diminishes the UV energy absorbed by microorganisms. Such problems may necessitate higher UV dose rates or pretreatment filtration. To minimize these problems, UV reaction cells are designed to induce turbulent flow, have long water flow paths and short light paths (around 3 inches), and provide for cleaning of residues from the lamp housings. Wherever used, usually in small water treatment systems, UV irradiation is generally the last step in the water treatment process, just after final filtration and before entering the distribution system. UV systems are normally easy to operate and maintain although severe site conditions, such as high levels of dissolved iron or hardness, may require pretreatment.

UV does not introduce any chemicals into the water and causes little, if any, chemical change in water. Therefore, overdosing does not cause water quality problems. UV is used mostly for inactivating pathogens to regulated levels. Since it leaves no residual, it can serve only as a primary disinfectant and must by followed by some form of chemical secondary disinfection, generally chlorine or chloramine. UV water treatment is used more in Europe than in the U.S. Small-scale units are available for individuals who have wells with high microbial levels.

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