Definitions of Chemical Terms
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:
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.
. 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 (H2
S), 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–
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
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
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
generally relate to a position of chlorination before to or after filtration, respectively.44
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
-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
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.
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
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
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
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
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+
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.
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
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
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
demonstrated that a halogen-substituted mixture, such as N
-chlorodimethylhydantoin, exhibited bactericidal activity against test bacteria superior to that obtained for either N
-dibromomethyl or N
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.