Chlorine and Chlorine Compounds
Nancy A. Falk
Marisa Macnaughtan
Atefeh Taheri
William McCormick
Although chlorine is one of the most widely distributed elements on earth, it is not found in a free state in nature. Instead, it exists mostly as a salt, whereas chloride is the anion in combination with sodium, potassium, calcium, and magnesium cations. Elemental chlorine (Cl2) can be synthesized in the laboratory and is a heavy, greenish-yellow gas with a characteristic irritating and penetrating odor. It is likely that chlorine and its compounds must have been known to alchemists for many centuries but only in 1809 did Sir Humphrey Davy conclude that chlorine gas (Cl2) was an element.1 Elemental chlorine is a highly reactive compound. It will react readily with oxygen-containing species to form a variety of chlorine oxides. Elemental chlorine will react reversibly with water to form hypochlorous acid (HOCl) and hydrochloric acid (HCl) (equation [1]).
Observations of the bleaching properties of chlorine gas in a water solution led to its first practical application in textile bleaching, with subsequent commercial production in 1785. Development of sodium and calcium hypochlorites (chloride of lime, chlorinated lime) for more convenient use followed shortly thereafter. The reactions to form these compounds from elemental chlorine are shown in equations (2) and (3).2
In 1846, Semmelweis used chloride of lime to combat and control puerperal fever in his clinic in Vienna. As early as 1854, chlorinated lime was applied in the treatment of sewage in London. Finally, in 1881, a German bacteriologist, Koch, demonstrated under controlled laboratory conditions that pure cultures of bacteria may be destroyed using hypochlorites. Five years later, the American Public Health Association issued a favorable report on the use of hypochlorites as disinfectants.3 Chloride of lime was first introduced to the North American continent by Johnson in 1908 for purification of water. Within a short time, many plants throughout the United States installed the chlorination process for water purification, alone or in combination with filtration, so that by 1911 an estimated 800 000 000 gallons of water were purified by the chlorination process.4
The number of waterborne diseases that have been prevented through public water disinfection in the United States and in Europe is enormous and the list of diseases is extensive. These diseases can be protozoan (amoebiasis, giardiasis, or microsporidiosis), bacterial (botulism, cholera, dysentery, typhoid fever), viral (hepatitis A, polyomavirus infections), or algal (Desmodesmus infections) in nature.5,6,7,8,9 Today, it is rare to find municipal water that is not treated by some form of chlorination. The introduction of elemental chlorine as a commercial product supplemented existing hypochlorite processes for the treatment of water and sewage. The use of chlorine as a disinfectant gained wide acceptance later in other industries. It is important to note that following common practice in industrial applications, the words chlorine or bleach is often broadly used to signify “active chlorine compounds” and are used interchangeably. Generally, what is intended is an “aqueous solution of active chlorine compounds, consisting of a mixture of HOCl, OCl–, Cl2, and other active chlorine compounds.” Where elemental chlorine (Cl2) is intended, it will be referred to as “chlorine gas” or “elemental chlorine.” Note that in some cases, “bleach” or “nonchlorine bleach” may be used to refer to some hydrogen peroxide (H2O2)-based chemistries.
Wide use of bleach as an antiseptic began during World War I when Dakin10 introduced a 0.45% to 0.50% sodium hypochlorite (NaOCl) solution for disinfection of open and infected wounds. Treatment of wounds with hypochlorite
necessitated information regarding solvent action, irritation, and toxicity as well as the rate of reaction on the necrotic tissue. Various toxicity studies were reported.11,12,13,14 More recent innovations in this field are discussed in this chapter.
necessitated information regarding solvent action, irritation, and toxicity as well as the rate of reaction on the necrotic tissue. Various toxicity studies were reported.11,12,13,14 More recent innovations in this field are discussed in this chapter.
CHEMISTRYElemental Chlorine
One of the most important commercial preparations is elemental chlorine (Cl2 gas). Today, over 65 million short tons of elemental chlorine are produced globally and 14 million short tons in the United States.15 The main method of manufacture is through electrolysis of a salt solution (either sodium or potassium chloride), known as the chloralkali process. Useful by-products of the process include hydrogen gas and sodium hydroxide. Although elemental chlorine is a gas (approximately 2.5 times as heavy as air), it is supplied through compression and cooling as an amber liquid (1.5 times as heavy as water) and shipped in steel cylinders or tank cars. When released to atmospheric conditions, liquid chlorine reverts immediately to a gaseous form. The largest use of elemental chlorine is in the manufacturing of polyvinylchloride plastics and other major polymers such as polyurethanes, followed closely by the synthesis of solvents, and other organic and inorganic products. Approximately 5% of the elemental chlorine production goes to the pulp and paper industry, whereas another 5% is used for water disinfection.16 Elemental chlorine is highly reactive in the presence of moisture and possesses a great tendency to combine with organic compounds. Another characteristic of chlorine and chlorine compounds is their unique ability to displace bromine or iodine from their respective salts by metathesis. This chemical mechanism is frequently used in practice for the controlled release of iodine and bromine in solution. Both elemental chlorine and sodium hydroxide produced through electrolysis are used as raw materials in the manufacturing of sodium hypochlorite solutions.2
TABLE 15.1 Commercially produced chlorine compounds | |||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||
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Hypochlorite Salts
The most common liquid form is hypochlorite, particularly in aqueous solution, first discovered by Berthollet17 in 1789 by contacting chlorine gas with potash lye. Hypochlorite salts are the oldest and most widely used of the active chlorine compounds in the field of chemical disinfection. They are (1) proven and powerful antimicrobials controlling a wide spectrum of microorganisms, (2) deodorizers, (3) being nonpoisonous to humans at use concentrations, (4) free of poisonous residuals, (5) colorless and nonstaining, (6) easy to handle, and (7) economical to use.18 Hypochlorite salts are available in solid or liquid form, with details shown in Table 15.1.
Calcium hypochlorite, lithium hypochlorite, and sodium hypochlorite combined with hydrated trisodium phosphate (chlorinated TSP) make up the solid forms; however, sodium and potassium hypochlorite are sold in solution form. Calcium hypochlorite dihydrate is a powder, granulated material, or tablet with a strong chlorine odor.19 It may be blended with compatible inorganic diluents to produce lower available chlorine compounds. Calcium hypochlorite products are soluble in water and fairly stable on prolonged storage. Lithium hypochlorite is a free-flowing, white, granulated material, also with a strong chlorine odor. It is readily soluble in water and is also stable. On exposure to air, powdered hypochlorites attract moisture and become less stable. The sodium hypochlorite solutions range in concentration, with lower available chlorine products used for domestic applications and stronger solutions sold for industrial uses. Sodium hypochlorite in water solutions are less stable, especially in products with higher chlorine concentrations. Potassium hypochlorites are available as liquids, with a higher available chlorine content than the sodium counterparts. Chlorinated TSP is a fine, white crystalline material. In addition to producing hypochlorite ion for sanitizing properties, it also contains alkaline phosphate for detergency; it finds usage in disinfecting equipment used for winemaking and brewing but is subject to phosphate regulations in applicable jurisdictions.
Calcium hypochlorite, lithium hypochlorite, and sodium hypochlorite combined with hydrated trisodium phosphate (chlorinated TSP) make up the solid forms; however, sodium and potassium hypochlorite are sold in solution form. Calcium hypochlorite dihydrate is a powder, granulated material, or tablet with a strong chlorine odor.19 It may be blended with compatible inorganic diluents to produce lower available chlorine compounds. Calcium hypochlorite products are soluble in water and fairly stable on prolonged storage. Lithium hypochlorite is a free-flowing, white, granulated material, also with a strong chlorine odor. It is readily soluble in water and is also stable. On exposure to air, powdered hypochlorites attract moisture and become less stable. The sodium hypochlorite solutions range in concentration, with lower available chlorine products used for domestic applications and stronger solutions sold for industrial uses. Sodium hypochlorite in water solutions are less stable, especially in products with higher chlorine concentrations. Potassium hypochlorites are available as liquids, with a higher available chlorine content than the sodium counterparts. Chlorinated TSP is a fine, white crystalline material. In addition to producing hypochlorite ion for sanitizing properties, it also contains alkaline phosphate for detergency; it finds usage in disinfecting equipment used for winemaking and brewing but is subject to phosphate regulations in applicable jurisdictions.
Because of their wide acceptance as disinfectants in many industries, hypochlorite solutions serve as standards for testing of other disinfectants. Today, hypochlorites are used as disinfectants in most households, hospitals, schools, and public buildings. They are also widely used for microbial control in restaurants, soda fountains, and other public eating places and for sanitizing food processing plants, dairies, canneries, breweries, wineries, and beverage bottling plants. Hypochlorites are sold for treatment of pool and drinking water, sewage, and wastewater effluents. In addition, hypochlorite can be combined with surfactants, salts, fragrances, colorants, and polymers to create products that disinfect with better cleaning efficacy and with improved use experience. Details on forms and choices are discussed later in this chapter.
Chlorine Dioxide
Chlorine dioxide has received more attention in recent years. This chlorinated compound is used with greater frequency for drinking water disinfection, for wastewater treatment, for slime control in cooling tower waters, and for disinfection of spaces containing a complex mixture of surfaces (also see chapter 27). It has a unique ability to break down phenolic compounds and remove phenolic tastes and odors from water. Another favorable feature is its lack of reaction with ammonia. Finally, chlorine dioxide does not form trihalomethanes or chlorophenols, a characteristic of great importance to humans and the environment especially in large-scale municipal water treatment plants. There are some health risks associated with its use. Chlorine dioxide and its inorganic reaction products, such as chlorite, may present a high risk of toxicity.20 Recommended minimum exposure levels for chlorine dioxide gas are an Occupational Safety and Health Administration (OSHA) permissible exposure limit (PEL) of 0.3 mg/m3 (0.1 ppm) as a time-weighted average (TWA) and the short-term exposure limit (STEL) is set at 0.9 mg/m3 (0.3 ppm). For elemental chlorine gas, the OSHA PEL is 1.5 mg/m3 (0.5 ppm) TWA and the STEL is 2.9 mg/m3 (1 ppm), indicating that chlorine dioxide gas is more toxic to humans than chlorine gas.21,22
Chlorine dioxide is used in the chlorination of drinking water as well as in wastewater and for elimination of cyanides, sulfides, aldehydes, and mercaptans. The oxidation capacity of ClO2 in terms of available chlorine to be approximately 2.5 times that of chlorine.23 Its oxidation-reduction potential is close to that of chlorine. The analysis for chlorine dioxide is like that of chlorine but with some modification. As in the case of chlorine detection, chlorine dioxide can be continuously measured by chlorine dioxide analyzer and transmitter systems based on amperometric sensing devices. Chlorine dioxide at room temperature is soluble in water at 2.9 g/L at 30 mm Hg partial pressure. In aqueous solution, the product is decomposed by light. It undergoes a valence change of five in the reduction reaction to chloride and does not pass through a hypochlorous acid phase during this reduction but rather follows a different reaction path.
Chlorine dioxide activity is considered at least equal to that of chlorine.24 Bactericidal efficiency was found to be unaffected at pH levels of 6 to 10. In further work, it was found that chlorine dioxide’s bactericidal activity decreases with lowering of temperature.25 In later work with spores, the authors demonstrated greater sporicidal activity for chlorine dioxide than for chlorine. The greater sporicidal activity of chlorine dioxide is explained by greater utilization of its oxidation capacity, involving a full change of five electrons. Vegetative bacteria do not activate this full oxidation potential.26 Harakeh et al27 evaluated solutions of chlorine dioxide against Yersinia enterocolitica and Klebsiella pneumoniae. The authors showed when these were grown in a natural aquatic environment of low temperature and lower nutrient contents that they were more resistant to chlorine dioxide than the same microorganisms grown under optimum laboratory conditions.27 Berman and Hoff28 showed good antiviral activity by chlorine dioxide at pH 10 in less than 15 seconds. It also has been shown to be very effective in inactivating Cryptosporidium parvum oocysts in drinking water.29
Chlorine dioxide is an extremely reactive compound and consequently cannot be manufactured and shipped in bulk but is prepared at the place of consumption. In practice, this consists of mixing a solution of chlorine with
a solution of sodium chlorite; the product is then applied to a water supply. The chlorine dioxide is formed according to equation (4):
a solution of sodium chlorite; the product is then applied to a water supply. The chlorine dioxide is formed according to equation (4):
Chlorine dioxide may also be produced by the following reactions: (1) acidification of chlorates with hydrochloric or sulfuric acid; (2) reduction of chlorates in acid medium; (3) reacting acids with chlorites; and (4) through electrolysis, using sodium chloride, sodium chlorite, and water. By combining separate solutions of sodium chlorite and acid, Alliger30 developed a disinfectant composition that produces chlorine dioxide at the site of application. Chlorine dioxide decahydrate may be prepared commercially, but it must be kept refrigerated because it decomposes at room temperature and can be dangerously explosive under some conditions. de Guevara31 prepared stable antiseptic solutions using inorganic boron compounds such as sodium tetraborate, boric acid, and sodium perborate to stabilize chlorine dioxide in aqueous solutions, forming a labile complex. When the stabilized Anthium Dioxcide® ClO2 formula is generated, all traces of chlorine are removed by passage through a column of sodium carbonate peroxide. The final product contains 5% of stabilized ClO2 complex. To make this material biologically active, it is necessary to release the ClO2 in solution by either acidification or introduction of chlorine. The manufacturer’s instructions for use recommend that an adjustment of pH be made with either acetic acid, citric acid, phosphoric acid, or by adding a provided activator a solution.32 Today, this type of stabilized ClO2 is used in the paper industry for removing slimes from paper mill white water systems. The use of chlorine dioxide has been extended to the food industry to provide sanitation for different food products under a variety of situations.33,34,35
Inorganic Chloramines
When ammonia combines with chlorine in an aqueous solution, it forms monochloramine (NH2Cl), dichloramine (NHCl2), nitrogen trichloride (NCl3), or nitrogen (N2). The residual amine tends to suppress the release of hypochlorous acid. In monochloramine, the chlorine is not sufficiently active to demonstrate any antimicrobial activity. It has a low hydrolysis constant, which is too low for hypochlorous acid to be released in sufficient amounts. The activity of chloramines depends on the pH of the solution (Figure 15.1). Because they are highly unstable, inorganic chloramines are not commercialized as such. The instability of NHCl2 is important in breakpoint chlorination used for water and sewage treatment.37
Inorganic chloramine treatment of water supplies was used in the 1930s and early 1940s to improve tastes and odors of water. The most efficient ratio of chlorine to ammonia is 2:1 by weight.4 Use of ammonia in water chlorination provided prolonged stability of chlorine in water distribution systems; however, inferior inactivation has been described when combined available chlorine was compared with free available chlorine.38 It takes approximately 25 times as much chloramine as free available chlorine to effect a rapid bactericidal action; for chloramines, the contact time is approximately 100 times longer than that required for the same residual of free available chlorine to produce the same inactivation. Chloramines are being considered for chlorination of water to prevent trihalomethane formation and for inactivation of Cryptosporidium.29,39 In addition, Kereluk and Borisenok,40 working with 5 to 10 ppm monochloramine solutions, demonstrated antimicrobial activity against bacteria and fungi.
 FIGURE 15.1 Relationships between monochloramine (NH2Cl), dichloramine (NHCl2), and pH. After Baker.36 Copyright © 1959 American Water Works Association. Reprinted by permission of John Wiley & Sons, Inc. |
Organic Chloramines
Organic chloramines are produced by the equilibrium reaction of hypochlorous acid with an amine, amide, imine, or imide, as shown in equation (5).
The organic chloramines are the N-chloro derivatives of the following four groups: (1) sulfonamides: chloramine-T (Figure 15.2; 8, R = CH3), dichloramine-T
(Figure 15.2; 7, R = CH3), chloramine-B (Figure 15.2; 8, R = H), halazone (Figure 15.2; 7, R = COOH); (2) heterocyclic compounds with nitrogen in the ring: hydantoin, dichloroisocyanurate, trichloroisocyanurate, trichloromelamine (Figure 15.2; 3 and 4, 2, 1, and 6, respectively); (3) condensed amines from guanidine derivatives (chloroazodin; Figure 15.2; Specialty); and (4) anilides.37 Once formed, some N-chloramine derivatives can be isolated as solids and are used as solid bleach precursors. When the N-chloramine is redissolved in water, the N-chloramine can undergo a hydrolysis reaction and release hypochlorous acid. Each N-chloramine will have unique properties depending on its water solubility and its hydrolysis constant. These two main factors control the amount of the timing of hypochlorous acid release. There are thousands of options for organic chloramines, and the wide variety of compounds also leads to a wide variety of properties such as solubility, chlorine release profile (hydrolysis constant), and toxicity.
(Figure 15.2; 7, R = CH3), chloramine-B (Figure 15.2; 8, R = H), halazone (Figure 15.2; 7, R = COOH); (2) heterocyclic compounds with nitrogen in the ring: hydantoin, dichloroisocyanurate, trichloroisocyanurate, trichloromelamine (Figure 15.2; 3 and 4, 2, 1, and 6, respectively); (3) condensed amines from guanidine derivatives (chloroazodin; Figure 15.2; Specialty); and (4) anilides.37 Once formed, some N-chloramine derivatives can be isolated as solids and are used as solid bleach precursors. When the N-chloramine is redissolved in water, the N-chloramine can undergo a hydrolysis reaction and release hypochlorous acid. Each N-chloramine will have unique properties depending on its water solubility and its hydrolysis constant. These two main factors control the amount of the timing of hypochlorous acid release. There are thousands of options for organic chloramines, and the wide variety of compounds also leads to a wide variety of properties such as solubility, chlorine release profile (hydrolysis constant), and toxicity.
 FIGURE 15.2 Selected N-chloramines and N-halamines. |
Common uses for N-chloramine compounds include disinfection and cleaning in swimming pools and spas, toilets, laundry, dishwashing detergent, hard surface cleaning, textile bleaching, egg decontamination, water cooling system disinfection, pulp and paper bleaching, and wastewater treatment. The most common commercialized products today include sodium trichloroisocyanuric acid (trichloro-s-c), sodium dichloroisocyanurate, 1,3-dichloro-5,5-dimethylhydantoin, and 1-bromo-3-chloro-5,5-dimethylhydantoin (Figure 15.2; 1, 2, 3, and 4, respectively). Other N-halamines are available as specialty chemicals and therefore only manufactured in small quantities. Historically, N-halamines such as chloramine-T and others were commonly manufactured but no longer have active registration status in the United States.
PRINCIPLES, MECHANISMS, AND OTHER ASPECTS OF CHLORINE DISINFECTIONDefinitions of Chemical Terms
Available chlorine may be defined as a measurement of oxidizing capacity and is expressed in terms of the equivalent amount of elemental chlorine. The concentration of
hypochlorite (or any other oxidizing disinfectant) may be expressed as available chlorine by determining the electrochemical equivalent amount of Cl2 to that compound. By equation (6), 1 mole of elemental chlorine is capable of reacting with two electrons to form inert chloride:
hypochlorite (or any other oxidizing disinfectant) may be expressed as available chlorine by determining the electrochemical equivalent amount of Cl2 to that compound. By equation (6), 1 mole of elemental chlorine is capable of reacting with two electrons to form inert chloride:
From equation (7), it can also be noted that 1 mole of hypochlorite (OCl–) may react with two electrons and two protons to form chloride and water:
Hence, 1 mole of hypochlorite is equivalent (electrochemically) to 1 mole of elemental chlorine and may be said to contain 70.91 g of available chlorine (identical to the molecular weight of Cl2). Because calcium hypochlorite (Ca[OCl]2) and NaOCl contain 2 and 1 moles of hypochlorite per mole of chemical, respectively, they also contain 141.8 g and 70.91 g available chlorine per mole. The molecular weights of Ca(OCl)2 and NaOCl are, respectively, 143 and 74.5, so that pure preparations of the two compounds contain 99.2 and 95.8 weight percent available chlorine; hence, they are effective means of supplying chlorine for disinfection purposes.
Chlorine demand. In chlorination of water, a certain part of this chlorine is consumed by water impurities, and any unconsumed chlorine remains as residual available chlorine. The difference between the chlorine applied and the chlorine remaining in the water may be referred to as the chlorine demand of this water. Because of its electronic configuration, chlorine possesses a strong tendency to acquire extra electrons, changing to inorganic chloride ions. This affinity for electrons makes chlorine a strong oxidizing agent. Hence, chlorine in water reacts quickly with (1) inorganic-reducing substances, such as ferrous iron (Fe2+), manganous manganese (Mn2+), nitrites (NO2-), and hydrogen sulfide (H2S), and (2) organic materials (other than amines), wherein the chlorine atom loses its oxidizing properties by reduction to chloride and is lost as a disinfectant.41 The reactions with the inorganic-reducing substances are quite rapid and stoichiometric, whereas those with organic material are in general slow and depend largely on the concentration of the free available chlorine.
Free and combined available chlorine. When chlorine survives the chlorine demand of the water, the chlorine measured may be reported as free, combined, or total residual chlorine depending on the analytic method used. The term free available chlorine is usually applied to three forms of chlorine that may be found in water: (1) elemental chlorine (Cl2), (2) HOCl, and (3) hypochlorite ion (OCl–).41 These forms may be found in water, provided there is no ammonia or other nitrogenous compounds to form chloramines and there is enough chlorine to satisfy the organic and inorganic water demands. During a chlorination process, a certain portion of chlorine combines with ammonia and other nitrogenous compounds present in natural water to form chloramines or N-chloro compounds. This combination of chlorine with ammonia or with other nitrogenous compounds is referred to as combined available chlorine. The free and combined available chlorine, when present in the water, are collectively described as total residual (available) chlorine.
Breakpoint chlorination refers to when a sufficient amount of chlorine is applied to satisfy the initial water demand, and an extra quantity of chlorine is added to provide a slight residual of free available chlorine. Before establishing this free residual, additions of chlorine (or hypochlorites) oxidize all the inorganic and organic material until at some so-called breakpoint, demand is fully satisfied. Any further additions of available chlorine result in a constant rise of free available chlorine in proportion to the dose (Figure 15.3). A detailed review of breakpoint chlorination and its application in water treatment is given by White.43
Marginal chlorination refers to the addition of chlorine to water sufficient just to overcome the chlorine demand consumption and to obtain an initial level of available chlorine regardless of the type of residual produced. Superchlorination is a further step in which chlorine is added beyond the level that is needed to yield an initial residual without regard to the type of residual produced. If too much chlorine is added, and a lower chlorine level is desired, a dechlorinating chemical (eg, sodium thiosulfate) may be added, and this process is referred to as dechlorination. The terms prechlorination and postchlorination generally relate to a position of chlorination before to or after filtration, respectively.44
 FIGURE 15.3 Ideal residual chlorine curve (ammonia solution). After Butterfield.42 Copyright © 1948 American Water Works Association. Reprinted by permission of John Wiley & Sons, Inc. |
Analysis of Available Chlorine
There are several methods to determine the available chlorine in solution or in products. In the iodometric method, the free chlorine liberates iodine in the acidified test solution containing potassium iodide (KI), and the liberated iodine is titrated with a standard sodium thiosulfate solution to a starch endpoint. In the sodium arsenite method, free chlorine is titrated with a standard sodium arsenite solution using KI-starch paper as the external indicator. In the orthotolidine (OT) method, when added to dilute chlorine solution (at the ppm level), the colorless OT reagent turns to yellow-orange-red depending on the chlorine concentration. The intensity of color determines the amount of available chlorine present and is compared with previously prepared color standards.45 The Palin DPD method uses N,N-diethyl-p-phenylenediamine (DPD) reagent, and the dilute chlorine solution turns the reagent pink to red, depending on the concentration of chlorine (at the ppm level).46 There are modifications in the OT and Palin methods to distinguish free available chlorine, total available chlorine, combined available chlorine, and chloramines. The DPD method is the basis for chlorine test strips used to quickly check chlorine levels in swimming pools and disinfectant solutions. The amperometric method consists of electrometric titration in which the current passes through a titration cell containing a dilute chlorine solution as the oxidizing agent and the standard phenylarsine oxide as titrating reducing agent. The volume of phenylarsine oxide consumed determines the free available chlorine (at the ppm level) in solution, and the endpoint is indicated electrically. This titrating detection unit is composed of an indicator electrode, a reference electrode, and a microammeter.43
Chlorine measurements can be accurately made using polarographic membrane techniques. The plastic probe contains an anode and electrolyte and is terminated by a membrane-covered noble metal cathode. When the probe is immersed in a chlorinated solution, the chlorine is reduced to chloride at the cathode; the generated current, linear with chlorine concentration, is displayed on the meter of the analyzer. The probe lead is the only connection between the analyzer and the sample being measured with no need for reagents of any kind. This probe can be installed in a pipeline or any other desirable location and requires virtually no maintenance. One of the more important features of probe analysis is the capability to distinguish between different forms of chlorine that may be present, thus yielding maximum efficiency at minimum cost. This application is widely accepted by various industries because of its safety, performance, and economy.43
Stability of Chlorine in Solution
The stability of free available chlorine in solution depends largely on the following factors: (1) chlorine concentration, (2) presence and concentration of catalysts or reducing agents, (3) pH of the solution, (4) temperature of the solution, (5) presence of organic material, (6) ultraviolet irradiation, and (7) ionic strength. Any of these factors, alone or in combination, may greatly affect the stability of free available chlorine in solution. Iron and aluminum seem to have only a slight effect on the stability of chlorine in solution, whereas copper, nickel, and cobalt are powerful catalysts of decomposition. The most stable free available chlorine solutions are those having the following characteristics: (1) low chlorine concentration; (2) absence of copper, cobalt, nickel, or other catalysts; (3) high alkalinity; (4) low temperature; (5) absence of organic material; (6) storage in dark, closed containers (ie, shielded from ultraviolet light); and (7) low ionic strength. The stability of chlorine in solution or products may be rated by its half-life, which denotes the number of days required for the available chlorine content to be reduced to half its initial value.47
Organic chloramines are considerably more stable in solution than free chlorine compounds because they release chlorine rather slowly into solution, with delayed bactericidal action. Solutions of chloramine-T (see Figure 15.2; 8, R = CH3) are quite stable, and a moderate exposure to high temperature, sunlight, or organic material does not seem to cause any appreciable decomposition.48 To ensure free chlorine stability in solution, chlorine stabilizers are often used that combine with chlorine to form N-chloro compounds, prolonging the life of chlorine considerably but at the same time producing a slower germicidal effect.
Factors Affecting Chlorine Microbicidal Activity
A long history and wide use of chlorine compounds have yielded much laboratory and field evaluation data, mostly concerning hypochlorites, but with application to all active chlorine compounds to some extent. Disinfection effectiveness largely depends on the concentration of free, undissociated hypochlorous acid in water solution and the relationship between pH and the degree of dissociation of HOCl, as shown in Figure 15.4. In addition to pH, various other environmental factors, alone or in combination, determine the antimicrobial action of chlorine. A full understanding of these environmental factors and manipulation thereof enables the user of chlorine compounds to make proper adjustments for best results.
Effect of pH
The pH has perhaps the greatest influence on the antimicrobial activity of chlorine in solution. An increase in pH substantially decreases the biocidal activity of chlorine, and a decrease in pH increases this activity. Early works in 1921 and 1934 showed this pH dependency on hypochlorite effectiveness.49,50 In 1937, Charlton and Levine,51
using Bacillus metiens and calcium hypochlorite solutions, showed that 100 ppm available chlorine at pH 8.2 exhibits approximately the same kill of spores as a 1000 ppm solution at pH 11.3, demonstrating the controlling effect of pH. Later, the effect of pH on 25 ppm available chlorine solution to produce 99% kill of B metiens spores was shown.52 The results were 2.5 minutes for pH 6, 3.6 minutes for pH 7, 5 minutes for pH 8, 19.5 minutes for pH 9, 35.5 minutes for pH 9.35, 131 minutes for pH 10, and 465 minutes for pH 12.86. The authors attributed the striking changes in killing time to changes in concentrations of undissociated hypochlorous acid and concluded that the concentration of HOCl is closely related to the speed of inactivation by hypochlorites in solution.
using Bacillus metiens and calcium hypochlorite solutions, showed that 100 ppm available chlorine at pH 8.2 exhibits approximately the same kill of spores as a 1000 ppm solution at pH 11.3, demonstrating the controlling effect of pH. Later, the effect of pH on 25 ppm available chlorine solution to produce 99% kill of B metiens spores was shown.52 The results were 2.5 minutes for pH 6, 3.6 minutes for pH 7, 5 minutes for pH 8, 19.5 minutes for pH 9, 35.5 minutes for pH 9.35, 131 minutes for pH 10, and 465 minutes for pH 12.86. The authors attributed the striking changes in killing time to changes in concentrations of undissociated hypochlorous acid and concluded that the concentration of HOCl is closely related to the speed of inactivation by hypochlorites in solution.
 FIGURE 15.4 Relationships between hypochlorous acid (HOCl), hypochlorite (-OCl), and pH. After Baker.36 Copyright © 1959 American Water Works Association. Reprinted by permission of John Wiley & Sons, Inc. |
Mercer and Somers,53 using Bacillus macerans spores, showed that 15 ppm hypochlorite solution effected 99% reduction of organisms within 8.5 minutes at pH 6, and that approximately 42 minutes were required at pH 8 for the same reduction. They also found no significant difference in sporicidal activity with chlorine gas, sodium hypochlorite, and calcium hypochlorite.53 Friberg and Hammarstrom,54 in their work with bacteria and viruses, concluded that the viricidal effect of free available chlorine is affected by the pH in much the same manner as bactericidal action. Watkins et al55 reported on the viricidal activity of NaOCl solution, which at 12.5 ppm available chlorine completely inactivated phage of Streptococcus cremoris within a 30-second interval; as pH was lowered from 9 to 4.4, progressively faster phage destruction occurred. This increased activity at lower pH levels was somewhat similar to results obtained with hypochlorites against bacterial spores and non-spore-forming bacteria.
Bactericidal activity of chloramines is also influenced by pH. With a decrease in pH, there is a corresponding increase in dichloramine formation, which is a more effective bactericide than the monochloramine. Work using cysts of Entamoeba histolytica verified this point by showing that dichloramine was a considerably more powerful cyst-penetrating agent than monochloramine, attributed to faster penetrating power to one of the hydrolysis products of dichloramine (ie, HOCl).56
Effect of Concentration
It is logical to assume that an increase in concentration in available chlorine in a solution brings a corresponding increase in antibacterial activity. This supposition may hold true if other factors, such as pH, temperature, and organic content, are held constant. Experiments with Staphylococcus aureus at a constant pH value of 9 showed that by increasing the available chlorine in the hypochlorite solutions from 0.3 to 0.6, 1.2, to 2.0 ppm, the killing time was shortened or the bactericidal rate increased. The 2 ppm available chlorine produced complete kill in 5 minutes and 1.2 ppm in 10 minutes, whereas 0.3 ppm did not completely kill the organism even in 30 minutes.57 Others tested hypochlorite solutions at concentrations of 25, 100, and 500 ppm of available chlorine at a constant pH of 10 and temperature of 20°C. The times required to provide the 99.9% kill of resistant B metiens spores were 31 minutes for 500 ppm, 63.5 minutes for 100 ppm, and 121 minutes for 25 ppm available chlorine solutions. It was concluded that a 4-fold increase in the concentration of hypochlorite solution results in a 50% reduction in killing time and a 2-fold increase in only a 30% reduction.52,58
Effect of Temperature
The effect of temperature was demonstrated with Mycobacterium tuberculosis; using 50 ppm available chlorine hypochlorite solution at pH 8.35, complete kill was obtained in 30 seconds at 60°C, in 60 seconds at 55°C, and in 2.5 minutes at 50°C.59 Under the same test conditions, 200 ppm available chlorine solutions at pH 9 destroyed the organism in 60 seconds at 50°C and in 30 seconds at 55°C.
A temperature effect on the bactericidal activity of Ca(OCl)2 solution was observed at 20°C, 30°C, 35°C, and 50°C. The 25 ppm hypochlorite solution at a constant pH of 10 killed in 121, 65, 38.7, and 9.3 minutes, respectively. A 60% to 65% reduction in killing time was
observed with a 10°C rise in temperature.52 Later, in work with hypochlorite solutions at 25 ppm available chlorine and three different pH levels (10, 7, 5), a rise of 10°C produced a reduction of 50% to 60% in killing time and that a drop of 10°C increased the necessary exposure time by approximately 2.1- to 2.3-fold. This work also revealed that temperature coefficients were only slightly affected by pH.58 Work with Pseudomonas fragi showed that Ca(OCl)2 at 3 ppm available chlorine produced a 99.99% kill in 4 minutes at 21°C, but in approximately 10 minutes at 4.4°C.60 The effects of temperature on the bactericidal action of free available chlorine are especially evident at a pH higher than 8.5 and also when the chlorine residuals are low (0.02-0.03 ppm available chlorine).61
observed with a 10°C rise in temperature.52 Later, in work with hypochlorite solutions at 25 ppm available chlorine and three different pH levels (10, 7, 5), a rise of 10°C produced a reduction of 50% to 60% in killing time and that a drop of 10°C increased the necessary exposure time by approximately 2.1- to 2.3-fold. This work also revealed that temperature coefficients were only slightly affected by pH.58 Work with Pseudomonas fragi showed that Ca(OCl)2 at 3 ppm available chlorine produced a 99.99% kill in 4 minutes at 21°C, but in approximately 10 minutes at 4.4°C.60 The effects of temperature on the bactericidal action of free available chlorine are especially evident at a pH higher than 8.5 and also when the chlorine residuals are low (0.02-0.03 ppm available chlorine).61
With respect to temperature effect on chloramines, it has been reported that 2.5-fold higher concentrations of chlorine and 9-fold longer exposure times are necessary to produce the same kill at 3°C as at 20°C.41 From all this work, it is evident that an increase in temperature increases bactericidal activity.
Effect of Organic Material
Organic material in chlorine solution consumes available chlorine and reduces its capacity for bactericidal activity; this is evident especially in solutions with low levels of chlorine. It has been reported that hypochlorites are selective in their attack on various types of organic material. There seems to be a difference of opinion among various workers on this subject. Among different sugars, only levulose consumed chlorine and the chlorine loss with other nonnitrogenous substrates (lipids and alcohols) was negligible. Sodium hypochlorite was also more reactive with organic substrates than either chloramine-T or azochloramide.62 Prucha63 reported early on the effect of 1% to 5% skim milk on chlorine losses in solution. Loveless64 studied the amount of available chlorine loss in hypochlorite solution in the presence of 1.5% whole milk at a temperature range of 21°C to 100°C for a period of 60 minutes, with all solutions showing some loss in available chlorine at 21°C and that the rate of loss increased with rising temperatures. The control solutions with no milk did not seem to lose any available chlorine during the 60 minutes, except for a small loss at 100°C. But the presence of milk in the tests with hypochlorite solutions did not seem adversely to affect the bactericidal action of chlorine.65
If the organic matter contains proteins, the chlorine reacts and forms chloramines, retaining some of its antibacterial activity even though the available chlorine levels are reduced considerably. This explains some questionable results in the early literature regarding the disappearance of anthrax spores from chlorinated tannery wastes in the absence of measurable free available chlorine, or that hypochlorite solution of 130 ppm of available chlorine completely killed Salmonella pullorum in the presence of 5% organic matter in the form of chicken manure.66 Similar results were reported with Salmonella typhosa and human feces.67
Johns,68 using a modification of glass slide technique and Escherichia coli and S aureus as test organisms, reported no evidence of germicidal reduction due to the presence of skim or whole milk in the freshly prepared hypochlorite solutions. This may be significant for the use of hypochlorites as sanitizers on milk farms and in dairy plants because minute amounts of milk may be encountered with no particular adverse effect on the activity of hypochlorites. However, Lasmanis and Spencer,69 in their work with hypochlorite solutions using strains of staphylococci (coagulase positive and negative), found that with 3% of skim milk, they did not obtain complete kill of organisms, although smaller amounts of milk exhibited progressively lesser effects on the bactericidal action.
It appears that sugars and starches may affect the germicidal activity of chlorine. Shere70 reported that 500 ppm of alkyl aryl sulfonate did not exhibit any slowing action on the germicidal effectiveness of the hypochlorite solutions. Other organic materials, such as tyrosine, tryptophan, cystine, egg albumin, peptone, body fluids, tissues, microbes, and vegetable matter, when present in a disinfecting solution, can consume chlorine to satisfy the organic water demand; in these cases, the chlorine may lose its function as a germicidal agent unless it forms chloramines or unless the chlorine dosage is adjusted to overcome this demand. This loss of chlorine due to organic matter may be significant in cases in which minute amounts of chlorine are used. Higher levels of chlorine tend to produce a safety reserve for performing the desired bactericidal action.
Effect of Hardness
Water hardness components such as Mg2+ and Ca2+ ions do not exhibit any impact on the antibacterial action of hypochlorite solution. In studies with 5 ppm available chlorine, sodium hypochlorite solution at 0 and 400 ppm hardness at 20°C, a complete kill of bacteria at two examined levels of hardness was obtained, indicating that raising the hardness from 0 to 400 ppm did not have any inhibitory action.42
Effect of Addition of Ammonia or Amino Compounds
The bactericidal activity of free chlorine is considerably diminished when chlorine is added to water containing ammonia or amino compounds and the concentration of chlorine is plotted against residual chlorine (see Figure 15.3). Chlorine can react immediately with ammonia to form monochloramines and dichloramines. As more chlorine is added to the ammonia solution, to a ratio
of chlorine to ammonia of 5:1, formation of chloramines continues until all the ammonia has been converted. Up to this point, chlorine remains in the form of combined available chlorine. After the so-called “hump” has been obtained, added amounts of chlorine oxidize the chloramines, slowly reducing the residual chlorine and ammonia, until they both drop practically to zero. Increase of chlorine beyond this point (breakpoint) produces an increase in free available chlorine.42 The available chlorine curves for different N-chloro compounds formed from amino acid or proteins vary because of variability of reactions and varying stabilities of the products of the reactions.
of chlorine to ammonia of 5:1, formation of chloramines continues until all the ammonia has been converted. Up to this point, chlorine remains in the form of combined available chlorine. After the so-called “hump” has been obtained, added amounts of chlorine oxidize the chloramines, slowly reducing the residual chlorine and ammonia, until they both drop practically to zero. Increase of chlorine beyond this point (breakpoint) produces an increase in free available chlorine.42 The available chlorine curves for different N-chloro compounds formed from amino acid or proteins vary because of variability of reactions and varying stabilities of the products of the reactions.
If ammonia concentrations are less than one-eighth of the total available chlorine added, the ammonia was found to be destroyed and the excess chlorine can remain as free available chlorine, exhibiting fast bactericidal action.58 If the concentration of ammonia is greater than one-fourth that of free chlorine, the available chlorine will exist in the form of chloramines and thus will have slow bactericidal activity. Water temperature influences the antibacterial action of the ammonia-chlorine treatment, with efficiency decreasing with lower temperatures. An excess of ammonium salt in the presence of high levels of organic material enhanced the bactericidal effectiveness of hypochlorite solution.71,72
From the information available, it appears that the killing time of chlorine is extended considerably in chloramines or N-chloro compounds, and the higher the concentration of ammonia or nitrogenous compounds the greater the lag in bactericidal time.
Effect of Addition of Iodine or Bromine (Halogen Mixture)
There is considerable evidence that small additions of bromine or iodine to chlorine solutions greatly enhance the bactericidal activity of chlorine. Improvement of bactericidal results in solutions containing chlorine with a small amount of ammonia salt and bromide ions has been reported.72 The addition of sodium bromide (NaBr) to hypochlorite solution resulted in a 33% to 1000% increase in bactericidal effectiveness against a variety of bacteria at pH 11.73 Chlorine-bromine mixtures at various ratios increased the germicidal activity in purified and natural waters containing low and high amounts of nitrogenous growth-promoting material in a pH range of 5.4 to 8.6. Also chlorine reinforced by 5% to 10% bromide was effective in decreasing the number of chlorine- and bromine-resistant bacteria.74 Kamlet,75 using equimolecular mixtures of bromine and chlorine, obtained superior bactericidal effects against E coli compared with either chlorine or bromine alone. Others claimed an advantage in using germicidal mixtures of iodine and chlorine.76
Paterson77 demonstrated that a halogen-substituted mixture, such as N-bromo-N-chlorodimethylhydantoin, exhibited bactericidal activity against test bacteria superior to that obtained for either N,N-dibromomethyl or N,N-dichlorodimethylhydantoin. Zsoldos78 produced a germicidal mixture by introducing dichlorodimethylhydantoin and KI into aqueous solution, thereby generating a hypoiodous acid and chloramine combination. Table 15.2 summarizes the biocidal effect of free available chlorine for a number of microorganisms.
MECHANISMS OF ANTIMICROBIAL ACTIONHypochlorous Acid Formation in Cells
The HOCl is an example of a reactive oxygen species (ROS). The ROS are chemically reactive chemical species that assist mammals during host defense. During the process, macrophages and neutrophils produce high concentrations of H2O2, superoxide, and hypochlorous acid to kill invading microorganisms.103 The buildup of ROS, a condition called oxidative stress, is suspected to be linked to the cause of different human diseases including Parkinson disease, Alzheimer disease, Huntington disease, and multiple sclerosis.104 A series of reactions occur in the space between an ingested bacterium and the membrane of the phagosome and as a result the ingested bacteria within the phagosome are killed (shown in Figure 15.5). The NADPH oxidase system reduces molecular oxygen to the superoxide radical. The influx of protons (H+) or other cations compensate the charge transfer. Nevertheless, the pH in the phagosome rises to about pH 8 that indicates that other cations such as potassium ions (K+) may enter the phagosome instead of protons.105 The protons are used to reduce superoxide to H2O2, which can be broken down to oxygen and water in a catalase-dependent reaction. Alternatively, H2O2 can combine with chloride to form HOCl in a reaction catalyzed by myeloperoxidase (MPO). Different peroxidases vary in their substrate specificity, and only MPO can generate HOCl.106 Initial studies indicated that the halide requirement could be met by iodide, bromide, or chloride or by the pseudohalide, thiocyanate.107,108,109,110,111 Additionally, it is shown nitrite might substitute halide in the MPO-mediated antimicrobial system in vitro112,113; however, it is not clear that nitrite is formed in sufficient amounts to contribute significantly to the microbicidal activity of the MPO system.114
MPO forms three different complexes on reaction with products of the respiratory burst of phagocytes: compounds I, II, and III.114 H2O2 reacts rapidly with the iron of MPO (which is normally in the ferric form) to form a complex compound I, which has an oxygen bound by a double bond to the heme iron.115 Compound I can also be formed by the reaction of MPO with HOCl. Compound I, the primary catalytic complex of MPO,
reacts with a halide in a two-electron reduction to form the corresponding hypohalous acid and regenerating the native Fe3+-MPO. The reaction of compound I with excess H2O2 results in the formation of compound II, which is inactive with respect to the oxidation of chloride. Compound II can be reduced to the active, native enzyme by oxygen radical anion or another reducing agent. Dioxygen radical anion can also react directly with native MPO to form compound III, an oxyperoxidase, which, like oxyhemoglobin, has oxygen attached to the heme iron. Compound III is unstable, decaying to native MPO with a half-decay time of several minutes at room temperature.
reacts with a halide in a two-electron reduction to form the corresponding hypohalous acid and regenerating the native Fe3+-MPO. The reaction of compound I with excess H2O2 results in the formation of compound II, which is inactive with respect to the oxidation of chloride. Compound II can be reduced to the active, native enzyme by oxygen radical anion or another reducing agent. Dioxygen radical anion can also react directly with native MPO to form compound III, an oxyperoxidase, which, like oxyhemoglobin, has oxygen attached to the heme iron. Compound III is unstable, decaying to native MPO with a half-decay time of several minutes at room temperature.
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