Introduction:
Chlorination of public water supplies and polluted waters serves primarily to destroy or deactivate disease-producing microorganisms. Disinfection with chlorine is widely practiced. Chlorination may produce some adverse effects including taste and odor problem. in recent years, chlorination has been found to produce trihalomethanes (THMs) and other organics of health concern (THMs are suspected human carcinogens). Thus, use of alternative disinfectants, such as chlorine dioxide and ozone that do not cause this particular problem, is increasing.
Theory:
Disinfectant capabilities of chlorine depend on its chemical form in water, which in turn is dependent on pH, temperature, organic content of water, and other water quality factors. Chlorine is used in the form of free chlorine [e.g., chlorine gas] or as hypochlorites [e.g., NaOCl and Ca(OC1)2]. Chlorine applied to water either as free chlorine or hypochlorite initially undergoes hydrolysis to form free chlorine consisting of aqueous molecular chlorine, hypochlorous acid and hypochlorite ion.
Chlorine gas rapidly hydrolyzes to hypochlorous acid according to:
Cl2 + H2O = HOCl + H+ +Cl–
Aqueous solutions of sodium or calcium hypochlorite hydrolyze too:
14.1
Ca(OCl)2 + 2H2O = Ca2+ + 2HOCl + 2OH. 14.2
NaOCl + H2O = Na+ + HOCl + OH–
Hypochlorous acid is a weak acid and will disassociate according to: 14.3
HOCl H+ +OCl– 14.4
The two chemical species formed by chlorine in water, hypochlorous acid (HOCl) and hypochlorite ion (OCl–), are commonly referred to as “free” or “available” chlorine.
Figure 14.1: Distribution of Chlorine species at 250C
Figure 13.1 shows that Cl2 can be significantly at low pH values (below pH 2); while HOCl is dominant between pH 3 and 6. Between pH 6 and 9, the relative fraction of HOCl decrease, while the corresponding fraction of OCl- increases. In waters with pH between 6.5~8.5, the reaction is incomplete and both species (HOCl and OCl-) will be present. Hypochlorous acid is the more germicidal of the two, especially at short contact time. The dissociation of HOCl is also temperature dependent. The effect of temperature is such that at a given pH, the fraction of HOCl will be lower at higher temperatures.
Reactions of Chorine with Impurities in Water:
Reactions with Ammonia:
Free chlorine reacts readily with ammonia and certain nitrogenous compounds to form what are collectively known as "combined chlorine". The inorganic chloramines consist of three species: monochloramine (NH2CI), dichloramine (NHCl2) and trichloramine or nitrogen trichloride (NCI3). The presence and concentrations of these combined forms depend on a number of factors including the ratio of chlorine to ammonia-nitrogen, chlorine dose, temperature, pH and alkalinity.
NH3 + HOCI = NH2C1 + H2O; pH 4.5 to 8 14.5
NH2CI + HOCI = NHCl2 + H2O; pH 4.5 to 8 14.6
NHCl2 HOCI = NCI3 + H2O; pH < 4.5 14.7
In addition to chlorinating ammonia, chlorine also reacts to oxidize ammonia to chlorine-free products (e.g., nitrogen gas and nitrate) as shown below.
3 Cl2 + 2 NH3 = N2 (g) + 6H+ + 6 CI- 14.8
4C12 + NH3 + 3H2O = 8C1- + NO3- + 9H+ 14.9
The mono- and dichloramines have significant disinfecting power and are therefore of interest in the measurement of chlorine residuals. Combined chlorine in water supplies may be formed in the treatment of raw waters containing ammonia; chlorinated wastewater effluents, as well as certain chlorinated industrial effluents normally contain only combined chlorine.
Reactions with Other Impurities:
Chlorine combines with various reducing agents and organic compounds thus increasing the chlorine demand which must be satisfied before chlorine is available to accomplish disinfection.
Fe2+, Mn2+, NO2-, and H2S are examples of inorganic reducing agents present in water supplies that will react with chlorine. Chlorine can react with phenols to produce mono-, di-, or trichlorophenols, which can impart tastes and odors to waters, Chlorine also reacts with humic substances present in water to form trihalomethanes (THMs, e.g., chloroform, brornoform, etc.) which are suspected human carcinogens (Note: According to USEPA, maximum allowable level of THMs in drinking water is 100 µg/L).
Break Point Chlorination
If chlorine is added to water containing reducing agents and ammonia (either naturally present or added to water to produce combined chlorine), a hump-shaped breakpoint curve is produced as shown in following figure. The different segment of the curve is described as follows :
If the water is free of ammonia and other compounds that may react with chlorine, the application of chlorine will yield free available chlorine residual of same concentration. This is denoted by the ‘no demand line’ or the "zero demand line" (see Fig.).
Chlorine first reacts with reducing agents such as H2S, Fe-2+, Mn2+ and develops no measurable residual as shown by the portion of the curve from Origin up to point A.
Figure 14.2: Generalized curve obtained during breakpoint chlorination of water sample containing ammonia
Addition of chlorine beyond point A results in forming mainly mono- and dichloramines. With mole ratios of chlorine to ammonia up to 1:1 [i.e., C12:NH3-N = 1:1], both mono and di-chloramines are formed. Chloramines thus formed are effective disinfectants and are shown as combined available chlorine residual in figure (From A to B).
Further increase in the mole ratio of chlorine to ammonia result in formation of some trichloramine and oxidation of part of ammonia to N2 and NO3-. These reactions are essentially complete when 1.5 mole of chlorine has been added for each mole of ammonia nitrogen originally present in water [i.e., C12:NH3-N = 1.5:1]. This is represented by the portion of the curve from B to C.
Addition of chlorine beyond point C would produce free chlorine residuals and is referred to as "breakpoint chlorination". In other words, chlorination of water to the extent that all ammonia is converted to N2 or higher oxidation state is referred to as "breakpoint chlorination'.
Addition of chlorine beyond point C would produce free chlorine residuals and is referredto as "breakpoint chlorination". In other words, chlorination of water to the extent that all ammonia is converted to N2 or higher oxidation state is referred to as "breakpoint chlorination".
Environmental Significance:
Breakpoint chlorination is required to obtain a free chlorine residual for better disinfection if ammonia is present in a water supply. While free chlorine residuals have good disinfecting powers, they are usually dissipated quickly in the distribution system. For this reason, final treatment with ammonia if often practiced to convert free chlorine residuals to longer-lasting combined chlorine residuals. The difference between the amount of chlorine added to the water and the amount of residual chlorine (i.e., free and combined available chlorine remaining) at the end of a specified contact period is termed as "chlorine demand'.
Apparatus:
Erlemeyer flask (250 mL)
Bottle
Beaker (250 mL)
Measuring cylinder
Dropper
Stirrer
Reagents:
Starch Indicator
Standard 0.025 N Sodium thiosulfate
Potassium Iodine crystal
Concentrated Acetic Acid
Chlorine water
Procedure:
1.) Place 200-mL portion of the water to be chlorinated in each of six 250-mL flasks.
2.) Add required quantity (as instructed by your teacher) of "chlorine water" (stock solution of bleaching powder in water) in each of the flasks. The chlorine content of the "chlorine water" (determined earlier in the laboratory) would be provided to you by your teacher. Calculate the chlorine dose for each of the six flasks.
3.) Shake each flask gently and allow to stand for 30 minutes.
4.) Determine residual chlorine of water from each flask by the starch-iodine method as described below:
Starch-Iodine Method:
The starch-iodine method is based on the oxidizing power of free and combined chlorine residuals to convert iodide ion into free iodine at pH 8 or less, as shown below.
Cl2 + 2I- = I2 + 2 Cl-
In the starch-iodine method, the quantity of chlorine residuals is determined by measuring the quantity of iodine by titration with a reducing agent sodium thiosulfate (Na2S2O3). The end point of titration is indicated by the disappearance of blue color, produced by the reaction between iodine and starch (which is added as indicator during the titration),
I2 + 2 Na2S2O3 = Na2S4O6 + 2 Nal or, I2 + 2S2O32- = S4O62- + 2I- I2 +
starch = blue color
(Qualitative test for the presence of iodine/chlorine)
The titration is carried out at pH 3 to 4, because the reaction with thiosulfate is not stoichiometric at neutral pH due to partial oxidation of the thiosulfate to sulphate.
Procedure for determination of residual chlorine concentration:
1.) Place 200 mL of the sample in an Erlenmeyer flask.
2.) Add 'about 1g of potassium iodide (estimated on a spatula) and 2 mL of concentrated Acetic acid to the water.
3.) Add 0.025 N sodium thiosulfate drop by drop from a burette until the yellow color almost disappears.
4.) Add 1 mL of starch solution to the water.
5.) Continue addition of standard sodium thiosulfate (Na2S2O3) solution until the blue color just disappears.
6.) Record the quantity (in mL) of sodium thiosulfate (Na2S2O3) solution used.
Calculation:
Residual chlorine (mg/L) = mL of 0.025N sodium thiosulfate used x M.F.
M.F
Table:
