معالجة مياه الشرب(استخدام مشتقات الكلور(هيبوكلوريت/غاز ثانى اكسيد الكلور)فى معالجة المياه)

. Other sources of halogens and oxidizing agents for microbiological control include:
(hypochlorites (sodium hypochlorite, calcium hypochlorite

(chlorine dioxide)



Hypochlorites
Sodium hypochlorite and calcium hypochlorite are chlorine derivatives formed by the reaction of chlorine with hydroxides.
The application of hypochlorite to water systems produces the hypochlorite ion and hypochlorous acid, just as the application of chlorine gas does.
NaOCl  OCl - + Na+
sodium hypochlorite hypochlorite ion sodium ion

OCl - + Na+ + H2O « HOCl + NaOH
hypochlorite ion sodium ion water hypochlorous acid sodium hydroxide


Ca(OCl)2  2OCl + Ca2+
calcium hypochlorite hypochlorite ion calcium ion


2OCl - + Ca2+ + 2H2O « » 2HOCl + Ca(OH)2
hypochlorite ion calcium ion water hypochlorous acid calcium hydroxide


The difference between the hydrolysis reaction of chlorine gas and hypochlorites is the reaction by-products.
The reaction of chlorine gas and water increases the H+ ion concentration and decreases pH by the formation of hydrochloric acid.
The reaction of hypochlorites and water forms both hypochlorous acid and sodium hydroxide or calcium hydroxide.
This causes little change in pH. Solutions of sodium hypochlorite contain minor amounts of excess caustic as a stabilizer, which increase alkalinity and raise pH at the point of injection.
This can cause hardness scale to form. Addition of a dispersant (organic phosphate/polymer) to the water system is usually sufficient to control this scaling potential.
Alkalinity and pH are significantly changed when sodium or calcium hypochlorite replaces gaseous chlorine.
Gaseous chlorine reduces alkalinity by 1.4 ppm per ppm of chlorine fed; hypochlorite does not reduce alkalinity.
The higher alkalinity of waters treated with hypochlorite reduces the corrosion potential but can increase the deposition potential.
Sodium Hypochlorite.
Sodium Hypochlorite. Sodium hypochlorite, also referred to as liquid bleach, is the most widely used of all the chlorinated bleaches.
It is available in several solution concentrations, ranging from the familiar commercial variety at a concentration of about 5.3 weight percent NaOCl to industrial strengths at concentrations of 10-12%.
The strength of a bleach solution is commonly expressed in terms of its "trade percent" or "percent by volume," not its weight percent: 15 trade percent hypochlorite is only 12.4 weight percent hypochlorite.
Approximately 1 gal of industrial strength sodium hypochlorite is required to replace 1 lb of gaseous chlorine.
The stability of hypochlorite solutions is adversely affected by heat, light, pH, and metal contamination.
The rate of decomposition of 10% and 15% solutions nearly doubles with every 10°F rise in the storage temperature.
Sunlight reduces the half-life of a 10%-15% hypochlorite solution by a factor of 3 to 5.
If the pH of a stored solution drops below 11, decomposition is more rapid. As little as 0.5 ppm of iron causes rapid deterioration of 10-15% solutions.
The addition of concentrated ferric chloride to a tank of sodium hypochlorite causes the rapid generation of chlorine gas.
Normal industrial grades of sodium hypochlorite may be fed neat or diluted with low-hardness water.
The use of high-hardness waters for dilution can cause precipitation of calcium salts due to the high pH of the hypochlorite solution.
"High Test" Calcium Hypochlorite (HTH).
"High Test" Calcium Hypochlorite (HTH).
The most common form of dry hypochlorite is high test calcium hypochlorite (HTH).
It contains 70% available chlorine, 4-6% lime, and some calcium carbonate.
Precipitates form when HTH is dissolved in hard water.
For feeding calcium hypochlorite as a liquid, solutions should be prepared with soft water at 1-2% chlorine concentration.
Care should be exercised in storing granular calcium hypochlorite.
It should not be stored where it may be subjected to heat or contacted by easily oxidized organic material.
Calcium hypochlorite decomposes exothermally, releasing oxygen and chlorine monoxide.
Decomposition occurs if HTH is contaminated with water or moisture from the atmosphere.
Calcium hypochlorite loses 3-5% of its chlorine content per year in normal storage.
All hypochlorites are somewhat harmful to skin and must be handled carefully. Corrosion-resistant materials should be used for storing and dispensing.

CHLORINE DIOXIDE
Chlorine dioxide, ClO2, is another chlorine derivative.
This unstable, potentially explosive gas must be generated at the point of application.
The most common method of generating ClO2 is through the reaction of chlorine gas with a solution of sodium chlorite.
2NaClO2 + Cl2  2ClO2 + 2NaCl
sodium chlorite chlorine dioxide chlorine sodium chloride
Theoretically, 1 lb of chlorine gas is required for each 2.6 lb of sodium chlorite.
However, an excess of chlorine is often used to lower the pH to the required minimum of 3.5 and to drive the reaction to completion.
Sodium hypochlorite can be used in place of the gaseous chlorine to generate chlorine dioxide.
This process requires the addition of sulfuric or hydrochloric acid for pH control.
Other methods used for chlorine dioxide generation include:
5NaClO2 + 5HCl  4ClO2 + 5NaCl + HCl + 2H2O
sodium chlorite hydrochloric acid chlorine dioxide sodium chloride hydrochloric acid water

10NaClO2 + 5H2SO4  8ClO2 + 5Na2SO4 + 2HCl + 4H2O
sodium chlorite sulfuric acid chlorine dioxide sodium sulfate hydrochloric acid water

2NaClO2 + HCl + NaOCl  2ClO2 + 2NaCl + NaOH
sodium chlorite hydrochloric acid sodium hypochlorite chlorine dioxide sodium chloride sodium hydroxide

Rather than hydrolyzing in water as chlorine does, chlorine dioxide forms a true solution in water under typical cooling tower conditions.
For this reason, chlorine dioxide is volatile (700 times more volatile than HOCl) and may be easily lost from treated water systems, especially over cooling towers.
Chlorine dioxide is a powerful oxidant. It reacts rapidly with oxidizable materials but, unlike chlorine, does not readily combine with ammonia.
Chlorine dioxide does not form trihalomethanes (THM) but can significantly lower THM precursors.
In sufficient quantity, chlorine dioxide destroys phenols without creating the taste problems of chlorinated phenols.
It is a good antimicrobial and antispore.
Unlike chlorine, the antimicrobial efficiency of chlorine dioxide is relatively unaffected by changes in pH in the range of 6-9. Chlorine dioxide is also used for the oxidation of sulfides, iron, and manganese.
Complex organic molecules and ammonia are traditional chlorine-demand materials that do not react with chlorine dioxide.
Because chlorine dioxide reacts differently from chlorine, a chlorine dioxide demand test must be conducted to determine chlorine dioxide feed rates.
A residual must be maintained after the chlorine dioxide demand has been met, to ensure effective control of microbiological growth.
The chemical behavior and oxidation characteristics of aqueous chlorine dioxide are not well understood because of the difficulty in differentiating aqueous chlorine-containing species.
Chlorine dioxide is applied to some public water supplies to control taste and odor, and as a disinfectant. It is used in some industrial treatment processes as an antimicrobial.
Chlorine dioxide consumed in water treatment reactions reverts to chlorite ions (ClO2-), chlorate ions (ClO3- ), and chloride ions (Cl -).
There are some concerns about the long-term health effects of the chlorite ion in potable water supplies.
As a gas, chlorine dioxide is more irritating and toxic than chlorine.
Chlorine dioxide in air is detectable by odor at 14-17 ppm, irritating at 45 ppm, fatal in 44 min at 150 ppm, and rapidly fatal at 350 ppm.
Concentrations greater than 14% in air can sustain a decomposition wave set off by an electric spark.
The most common precursor for on-site generation of chlorine dioxide is also a hazardous material: liquid sodium chlorite.
If allowed to dry, this powerful oxidizing agent forms a powdered residue that can ignite or explode if contacted by oxidizable materials.
The hazardous nature of chlorine dioxide vapor and its precursor, and the volatility of aqueous solutions of chlorine dioxide, require caution in the design and operation of solution and feeding equipment.

OTHER USES AND EFFECTS OF CHLORINE
In addition to serving as antimicrobials, chlorine and chlorine compounds are used to reduce objectionable tastes and odors in drinking water; improve influent clarification processes; oxidize iron, manganese, and hydrogen sulfide to facilitate their removal; reduce sludge bulking in wastewater treatment plants; and treat wastewater plant effluents.
Chlorine, along with a coagulant, is often applied to raw water in influent clarification processes
This prechlorination improves coagulation because of the effect of chlorine on the organic material in the water.
It is also used to reduce taste, odor, color, and microbiological populations, and it oxidizes iron and manganese to facilitate removal by settling and filtration.
One part per million of chlorine oxidizes 1.6 ppm of ferrous ion or 0.77 ppm of manganous ion. The addition of 8.87 ppm of chlorine per ppm of sulfide oxidizes sulfides to sulfates, depending on pH and temperature.
Chlorine is a successful activating agent for sodium silicate in the preparation of the coagulant aid, activated silica.
The advantage of this process is that the chlorine used for activation is available for other purposes.
Low-level, intermittent chlorination of return activated sludge has been used to control severe sludge bulking problems in wastewater treatment plants.
Chlorine, injected into sewage and industrial wastes before they are discharged, destroys bacteria and such chemicals as sulfides, sulfites, and ferrous iron. These chemicals react with and consume dissolved oxygen in the receiving body of water
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