Chlorination

Chlorination Chlorination has been used on an emergency basis from as far back as 1850. Continuous chorination of public water supplies began in 1904 in England, and was followed by the U.S. around 1909. After the facilities for gaseous chlorine were developed in 1912, chlorination grew rapidly. In that time, infectious waterborne diseases have decreased dramatically. Chlorine may have saved more lives than antibiotics. Nevertheless, chlorine has had its share of controversy. Chlorine use was delayed until 1909 in the U.S. because of a lawsuit arguing it was unsafe. Indeed, concentrated chlorine gas is extremely toxic. As recently as the mid-1990's, the Clinton Administration proposed a ban on chlorine (the "Chlorine Zero Discharge Act"). In many ways it is understandable, because chlorine helps makes up dioxins and PCB's, not to mention the CFC's role in stratospheric ozone depletion. However, the law was not passed when a group of scientists pointed out that chlorine is already found in nature [salt (sodium chloride) comes to mind...]. Moreover, chlorine plays a huge role in the safety of our water supplies. To better understand the issues, we need to take a closer look at chlorine chemistry. Chlorine can be found in gas, liquid, and solid forms as shown below. The key compunds are shown in red, and are further discussed below. 1. Gas: Cl2 C12 (Elemental Cl gas) + H20 <--> HCL + HOCl (Hypochlorous Acid) <--> OCL+ H+ (Hypochlorite Ion) 2. Liquid: sodium hypochlorite (bleach) NaOCl (Sodium Hypochlorite) + H20 <--> Na+ + OCL(Hypochlorite Ion) < H+> HOCl (Hypochlorous Acid) 3. Solid: calcium hypochlorite (as powder or tablets) Ca(OC1)2 (Calcium Hypochlorite) + H20 <--> Ca++ + 2 OCL(Hypochlorite Ion) < H+> HOCl (Hypochlorous Acid) HCl (hydrochloric acid) is actually ineffective in killing microbes at the concentrations used in water treatment. HOCL (hypochlorous acid) is the active form that kills the microbes (concentrations of about 0.2-1.2 ppm are needed for drinking water supplies). OCl- (hypochlorite ion) also is an active killing ingredient, and along with HOCl makes up what we refer to as "free available chlorine." HOCL is roughly a eighty times more effective in killing power than OCl-, largely because the uncharged HOCL is more effective in penetrating cell walls. How much chlorine do we need to add? It depends... on the quality of the water. Not all of the chlorine is converted to free available chlorine (i.e., the HOCl and OCl- that makes up the active killing ingredients). The chlorine can participate in a variety of reactions, yielding one of the strangest curves we see in our field: For a graph like the one above, we normally expect a straight line (i.e., "what you add is what you get" in the form of a residual). But a number of reactions are acting to reduce the free chlorine residual, and we can analyze them by referring to each of the letters in the graph (a, b, c, and d): a. Fe++ --> Fe+++ Mn++ --> Mn+++ NO2- --> NO3- Because of these three reactions, chlorine is reduced to HCl. The more iron, manganese, and nitrites in the water, the more of these reactions that occur. b. Combined chlorine b1. Chloro-organics In the first part of this curve (labelled b1), chlorine is combining with various organics in the water. b2. HOCL + NH3 --> NH2Cl (monochloramine) HOCL + NH2Cl --> NHC12 (dichloramine) HOCL + NHC12 --> NC13 (trichloramine) In the second half of this curve (labelled b2), chlorine is combining with ammonia in the water, forming what we collectively refer to as combine chlorine, or chloramines. c. --> --> --> HCl + N2 A variety of reactions occur starting with letter c on the curve, ultimately releasing HCl and nitrogen gas. We sometimes refer to this as "burning off the chloramines." d. Free available chlorine Once we have "burned off the chloramines," any additional chlorine is converted directly into free available chlorine. Letter d is also referred to as the "breakpoint", and this type of chlorination is referred to as "breakpoint chlorination."