معالجة مياه الشرب (استخدام الكلور فى معالجة وتطهير المياه)

Chlorine is one of the most versatile chemicals used in water and wastewater treatment. This powerful oxidizing agent is used for:
Disinfection
control of microorganisms
removal of ammonia
control of taste and odor
color reduction
destruction of organic matter
hydrogen sulfide oxidation
iron and manganese oxidation

Physical Properties and Aqueous Reactions concerns.
Chlorine in its gaseous state Chlorine gas is greenish-yellow, and its density is about 2 times that of air.
When condensed, it becomes a clear, amber liquid with a density about 1 times greater than water.
One volume of liquid chlorine yields approximately 500 volumes of chlorine gas, which is neither explosive nor flammable
oxygen, chlorine gas can support the combustion of some substances.
Chlorine reacts with organic materials to form oxidized or chlorinated derivatives.
Some of these reactions, such as those with hydrocarbons, alcohols, and ethers, can be explosive.
The formation of other chlorinated organics, specifically trihalomethanes (THM), poses an environmental threat to public drinking water supplies.
Chlorine gas is also a toxic respiratory irritant. Airborne concentrations greater than 3-5 ppm by volume are detectable by smell, and exposure to 4 ppm for more than 1 hr can have serious respiratory effects.
Because chlorine gas is denser than air, it stays close to the ground when released.
The contents of a 1-ton cylinder of chlorine can cause coughing and respiratory discomfort in an area of 3 square miles.
The same amount concentrated over an area of 1/10 square mile can be fatal after only a few breaths.
Chlorine is generated commercially by the electrolysis of a brine solution, typically sodium chloride, in any of three types of cells: diaphragm, mercury, or membrane.
The majority of chlorine produced is manufactured by the electrolysis of sodium chloride to form chlorine gas and sodium hydroxide in diaphragm cells.
The mercury cell process produces a more concentrated caustic solution (50%) than the diaphragm cell.
Chlorine gas can also be generated by the salt process (which employs the reaction between sodium chloride and nitric acid), by the hydrochloric acid oxidation process, and by the electrolysis of hydrochloric acid solutions.
The gas is shipped under pressure in 150-lb cylinders, 1-ton cylinders, tank trucks, tank cars, and barges.
The four basic categories of chlorine treatment are defined not only by their function but also by their position in a water treatment sequence:
prechlorination
rechlorination
post-chlorination
dechlorination

In chemically pure water, molecular chlorine reacts with water and rapidly hydrolyzes to hypochlorous acid (HOCl) and hydrochloric acid (HCl):
Cl2 + H2O  HOCl + HCl
chlorine water hypochlorous
acid hydrochloric
acid


Both of the acids formed by hydrolysis react with alkalinity to reduce buffering capacity of water and lower pH. Every pound of chlorine gas added to water removes about 1.4 lb of alkalinity.
In cooling water systems, this alkalinity reduction can have a major impact on corrosion rates.
At pH levels above 4.0 and in dilute solutions, the hydrolysis reaction is completed within a fraction of a second.
For all practical purposes the reaction is irreversible.
Hypochlorous acid is a weak acid and dissociates to form a hydrogen ion and a hypochlorite ion.
HOCl « » H+ + OCl
hypochlorous
acid hydrogen
ion hypochlorite
ion


The concentration or distribution of each species at equilibrium depends on pH and temperature. Between pH 6.5 and 8.5, the dissociation reaction is incomplete, and both hypochlorous acid and hypochlorite ions are present.
The equilibrium ratio at any given pH remains constant even if the hypochlorous acid concentration is decreasing.
At constant pH and increasing temperature, chemical equilibrium favors the OCl - ion over HOCl.
The primary oxidizing agents in water are hypochlorous acid and the hypochlorite ion, although hypochlorite has a lower oxidizing potential.
Oxidizing potential is a measure of the tendency of chlorine to react with other materials.
The speed at which these reactions occur is determined by pH, temperature, and oxidation/reduction potential.
As pH increases, the chemical reactivity of chlorine decreases; as temperature increases, reactions proceed more rapidly.
The oxidation reactions of chlorine with such inorganic reducing agents as sulfides, sulfites, and nitrites are generally very rapid.
Some dissolved organic materials also react rapidly with chlorine, but the completion of many organic-chlorine reactions can take hours.
Chlorine Demand.
Chlorine Demand. Chlorine demand is defined as the difference between the amount of chlorine added to a water system and the amount of free available chlorine or combined available chlorine remaining at the end of a specified time period.
The demand is the amount of chlorine consumed by oxidation or substitution reactions with inorganic and organic materials, such as H2S, Fe2+, Mn2+, NH3, phenols, amino acids, proteins, and carbohydrates.
Free available residual chlorine is the amount of chlorine which exists in the treated water system as hypochlorous acid and hypochlorite ions after the chlorine demand has been satisfied.
Free residual chlorination is the application of chlorine to water to produce a free available chlorine residual.
Combined Available Residual Chlorination.
Combined Available Residual Chlorination. Combined available residual chlorine is the chlorine residual which exists in water in combination with ammonia or organic nitrogen compounds.
Combined residual chlorination is the application of chlorine to water to react with ammonia (natural or added) or other nitrogen compounds to produce a combined available chlorine residual.
Total available chlorine is the total of free available chlorine, combined available chlorine, and other chlorinated compounds.
Available Chlorine.
Available Chlorine. "Available chlorine" is an expression of the equivalent weights of oxidizing agents, with chlorine gas as a basis, similar to the expression of alkalinity in terms of calcium carbonate equivalents.
The term originated from the need to compare other chlorine-containing compounds to gaseous chlorine.
Available chlorine is based on the half-cell reaction in which chlorine gas is reduced to chloride ions with the consumption of two electrons.
In this reaction, the equivalent weight of chlorine is the molecular weight of chlorine, 71 g/mole, divided by 2, or 35.5 g/mole.
2e - + Cl2  2Cl -
chlorine chloride ion


The available chlorine of other chlorine-containing compounds is calculated from similar half-cell reactions, the formula weight of the compound, and the equivalent weight of chlorine.
Even though chlorine gas only dissociates into 50% HOCl or OCl -, it is considered 100% available chlorine.
Because of this definition, it is possible for a compound to have more than 100% available chlorine.
The active weight percent chlorine multiplied by 2 indicates available chlorine.
Available chlorine, like oxidation potential, is not a reliable indicator of the occurrence or extent of an oxidation reaction.
It is an even poorer indicator of the antimicrobial efficacy of an oxidizing compound.
For example, the antimicrobial efficacy of hypochlorous acid (HOCl) is much greater than any of the chloramines even though the chloramines have a higher available chlorine.
Chloramine Formation.
Chloramine Formation. One of the most important reactions in water conditioning is the reaction between dissolved chlorine in the form of hypochlorous acid and ammonia (NH3) to form inorganic chloramines.
The inorganic chloramines consist of three species: monochloramine (NH2Cl), dichloramine (NHCl2), and trichloramine, or nitrogen trichloride (NCl3). The principal reactions of chloramine formation are:
NHCl2 + HOCl  NCl3 + H2O
dichloramine hypochlorous
acid trichloramine water

NH3 (aq) + HOCl  NH2Cl + H2O
ammonia hypochlorous
acid monochloramine water

NH2Cl + HOCl  NHCl2 + H2O
monochloramine hypochlorous
acid dichloramine water


The relative amounts of chloramines formed are a function of the amount of chlorine fed, the chlorine/ammonia ratio, temperature, and pH. In general, monochloramine is formed above pH 7 and predominates at pH 8.3.
Dichloramine predominates at pH 4.5. Between these pH values, mixtures of the two chloramines exist. Below pH 4.5, nitrogen trichloride is the predominant reaction product.
The oxidizing potential of monochloramines is substantially lower than that of chloride, and monochloramines are slower to react with organic matter.
These properties reduce the amount of trihalomethanes (THM) formed.
The formation of THM is considered more detrimental in potable water than the reduction of the antimicrobial capabilities of free chlorine.
Therefore, ammonia is often injected into the chlorine feed stream to form chloramines before the chlorine is fed into the potable water stream.
Combined chlorine residuals are ordinarily more chemically stable (less reactive with chlorine demand) than free chlorine residuals.
This property helps maintain stable residuals in outlying pressurized water distribution systems.
However, the lower antimicrobial effectiveness of chloramines compared to free chlorine requires higher combined residuals and/or longer contact times, which are often available in distribution systems.
Breakpoint Chlorination.
is the application of sufficient chlorine to maintain a free available chlorine residual.
The principal purpose of breakpoint chlorination is to ensure effective disinfection by satisfying the chlorine demand of the water.
In wastewater treatment, breakpoint chlorination is a means of eliminating ammonia, which is converted to an oxidized volatile form.
The addition of chlorine to a water that contains ammonia or nitrogen-containing organic matter produces an increased combined chlorine residual.
Mono- and dichloramines are formed, After the maximum combined residual is reached, further chlorine doses decrease the residual.
Chloramine oxidation to dichloramine, occurring, results in a decline in the combined available residuals initially formed.
so the breakpoint: the point at which chlorine demand has been satisfied and additional chlorine appears as free residuals.
so free available residual chlorine increases in direct proportion to the amount of chlorine applied.
Factors that affect breakpoint chlorination are initial ammonia nitrogen concentration, pH, temperature, and demand exerted by other inorganic and organic species.
The weight ratio of chlorine applied to initial ammonia nitrogen must be 8:1 or greater for the breakpoint to be reached.
If the weight ratio is less than 8:1, there is insufficient chlorine present to oxidize the chlorinated nitrogen compounds initially formed.
When instantaneous chlorine residuals are required, the chlorine needed to provide free available chlorine residuals may be 20 or more times the quantity of ammonia present. Reaction rates are fastest at pH 7-8 and high temperatures.
On the initial chlorine dosage produces no residual because of an immediate chlorine demand caused by fast-reacting ions.
As more chlorine is applied, chloramines develop. These chloramines are shown in the total chlorine residual.
At higher chlorine dosages, the slope to breakpoint begins. After the breakpoint, free chlorine residuals develop.
Free chlorine residuals usually destroy tastes and odors, control exposed bacteria, and oxidize organic matter.
Breakpoint chlorination can also control slime and algae growths, aid coagulation, oxidize iron and manganese, remove ammonia, and generally improve water quality in the treatment cycle or distribution system.
OXIDIZING ANTIMICROBIALS IN INDUSTRIAL COOLING SYSTEMS
The oxidizing antimicrobials commonly used in industrial cooling systems are the halogens, chlorine and bromine, in liquid and gaseous form; organic halogen donors; chlorine dioxide; and, to a limited extent, ozone.
Oxidizing antimicrobials oxidize or accept electrons from other chemical compounds.
Their mode of antimicrobial activity can be direct chemical degradation of cellular material or deactivation of critical enzyme systems within the bacterial cell.
An important aspect of antimicrobial efficiency is the ability of the oxidizing agent to penetrate the cell wall and disrupt metabolic pathways.
For this reason, oxidation potential alone does not always correlate directly with antimicrobial efficiency.
The relative microbiological control ability of typical halogens is as follows:
HOCl 3 HOBr 3 NHxBry >>
hypochlorous acid hypobromous acid bromamine

OCl - > OBr - >>> NHxCly
hypochlorite ion hypobromite ion chloramine


Cooling water pH affects oxidizing antimicrobial efficacy.
pH determines the relative proportions of hypochlorous acid and hypochlorite ion or, in systems treated with bromine donors, hypobromous acid and hypobromite ion.
The acid forms of the halogens are usually more effective antimicrobials than the dissociated forms.
Under some conditions, hypochlorous acid is 80 times more effective in controlling bacteria than the hypochlorite ion.
Hypochlorous acid predominates below a pH of 7.6.
Hypobromous acid predominates below pH 8.7, making bromine donors more effective than chlorine donors in alkaline cooling waters, especially where contact time is limited.
Antimicrobial efficacy is also affected by demand in the cooling water system, specifically demand exerted by ammonia.
Chlorine reacts with ammonia to form chloramines, which are not as efficacious as hypochlorous acid or the hypochlorite ion in microbiological control.
Bromine reacts with ammonia to form bromamines.
Unlike chloramines, bromamines are unstable and reform hypobromous acid.
Most microbes in cooling systems can be controlled by chlorine or bromine treatment if exposed to a sufficient residual for a long enough time.
A free chlorine residual of 0.1-0.5 ppm is adequate to control bulk water organisms if the residual can be maintained for a sufficient period of time.
Continuous chlorination of a cooling water system often seems most prudent for microbial slime control.
However, it is economically difficult to maintain a continuous free residual in some systems, especially those with process leaks. In some high-demand systems it is often impossible to achieve a free residual, and a combined residual must be accepted.
In addition, high chlorine feed rates, with or without high residuals, can increase system metal corrosion and tower wood decay.
Supplementing with nonoxidizing antimicrobials is preferable to high chlorination rates.
In once-through systems, free residuals from 0.3 to 0.8 ppm are normally maintained for -2 hr per treatment period.
The rate of recontamination determines the frequency of treatment required.
Open recirculating systems can be treated with a continuous or intermittent halogenation program.
Continuous feed is the most effective and is generally affordable where chlorine gas or hypochlorite is being used and system demand is low.
Free residuals of 0.1-0.5 ppm are manually maintained.
Care should be taken not to feed excessive amounts of halogen that will adversely affect corrosion rates.
Chlorine feed rates should not exceed 4 ppm based on recirculation rate.
The use of halogen donors may be restricted to intermittent feed due to economics, although continuous feed in low-demand systems is effective.
Intermittent feed requires maintaining the same free residual as in the continuous program but for only the last to 1 hr of the chlorine application.
Up to 3 hr of chlorine addition may be required in order to achieve the free residual, depending on system demand, system cleanliness, and the frequency of chlorination
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