Dechlorination of wastewater and cooling water - Environmental

Technol. , 1984, 18 (2), pp 48A–55A. DOI: 10.1021/ ... Publication Date: February 1984. Copyright ... Environmental Science & Technology 1998 32 (4)...
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FEATURES Dechlorination of wastewater and cooling water Questions remain about the environmental effects of dechlorination


George R. Helz Lynn Kosak-Channing a Department of Chemistry University of Maryland College Park, Md. 20742 Two years ago this journal published an article entitled "The Chlorination Question," which reviewed a conference at which the latest research had been presented on topics such as ecosystem damage from discharges of chlorinated sewage and power plant cooling water ( / ) . A growing movement is now under way in the U.S. to install dechlorination facilities, particularly at sewage treatment plants, to alleviate some of the problems discussed at that conference. In California alone, more than two dozen wastewater treatment "Present affiliation: Beltsville Agricultural Research Center, Beltsville, Md. 20705 48A

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plants have dechlorination systems to protect fish and wildlife (2). Some other states have followed California's lead by introducing dechlorination at wastewater treatment plants that discharge into waters supporting commercial or sport fishing (3). In the past several years, roughly 50 wastewater treatment plants in Maryland have installed dechlorination systems, mostly using SO2, in response to the collapse of the shad-fishing industry and similar problems in that state. Until very recently, power plant operators have apparently favored chlorine minimization, rather than dechlorination, as a means of meeting increasingly stringent discharge limitations (4-6). A recent survey of 231 electric power plants in the U.S. identified only one with an automatic dechlorination system and three others with temporary, manual systems (7). However, a number of electric utilities are currently reviewing dechlorination

technology, and the utilities industry may soon adopt dechlorination more widely. The movement toward dechlorination is propelled by a desire to escape the environmental problems associated with both the residual oxidants and halocarbons produced by chlorination. These problems have been discussed extensively during the past decade (e.g.. References 8-16). It is generally believed that dechlorination is beneficial, especially because it reduces acute toxicity (17-21) and even mutagenicity (22) associated with chlorinated waters. In this article, however, we review the chemistry of dechlorination with the purpose of showing that questions remain about the environmental effects of the general use of dechlorination. Types of dechlorinating agents The materials that have been proposed for dechlorination are powdered


© 1984 American Chemical Society

and granulated activated carbon, hy­ drogen peroxide (H2O2), ammonia (NH3), sodium thiosulfate (Na2S 2 0 3 ), and the sulfur(IV) species: sulfur dioxide (SO2), sodium bisulfite ( N a H S 0 3 ) , sodium sulfite ( N a 2 S 0 3 ) , and sodium metabisulfite (Na2S205). Dechlorination with activated car­ bon has been studied in some detail under laboratory and field conditions (23-30). Although it is a feasible method for dechlorination at waste­ water treatment plants, it is not con­ sidered economically attractive if the sole objective is dechlorination (29, 30). Seegert and Brooks found that activated carbon is not 100% efficient for dechlorinating fish tank water (31). Hydrogen peroxide readily re­ duces free chlorine to chloride plus oxygen, but does not react with chloramines at a useful rate (32). There­ fore, its use as a practical dechlori­ nating agent would be extremely lim­ ited. Ammonia, likewise, does not re­ move chloramines. Thiosulfate has been judged unacceptable for dechlo­ rination (33). In drinking water treatment, it has been found to pro­ duce objectionable, "sulfurous" odors (34). The chemical basis of this prob­ lem has not been determined, but it may involve acid hydrolysis, which releases hydrogen sulfide. Of the S(IV) compounds, sulfur dioxide gas is preferred overwhelmingly as a dechlorinating agent. A major factor in its favor is that it can be applied with the same types of equipment and techniques that are used to apply chlorine (29). Comparisons of the toxicity of chlorinated vs. chlor-dechlorinatcd waters uniformly show that dechlori­ nation reduces acute toxicity to aquatic organisms (17-21), but sublethal bi­ ological effects were noted in chlordechlorinated waters and in waters treated only with dechlorinating agents (18, 35, 36). Thiosulfate was used for dechlorination in these studies. It is not known whether these sublethal effects were created by the dechlorinating agent directly, by subsequent reactions of the dechlorinating agent, or by halocarbons or other chlorine by­ products generated before the de­ chlorinating agent was added. Use of excess S(IV) It is of considerable importance that dechlorinating with S(IV) compounds involves replacing residual oxidants with residual S(IV) compounds in the effluent. Although in principle it should be possible to titrate residual oxidants in an effluent exactly to the endpoint, leaving the effluent free of

both residual oxidants and residual S(IV) species, in practice it is neces­ sary to add excess S(IV), partly to compensate for imperfect control ap­ paratus and partly for kinetic rea­ sons—to attain a satisfactory level of dechlorination prior to discharge. For example, Tan et al. analyzed the problem of how to achieve 95% re­ duction of free chlorine (initial con­ centration 0.5 m g / L ) within a 122-m power plant discharge channel in which the flow velocity was 1.9 m / s (37). They concluded that a fourfold molar overdose of S 0 2 would be nec­ essary. They also concluded that a stoichiometric dose would accomplish little chlorine reduction prior to dis­ charge under the assumed treatment conditions. Slower reacting chlor­ amines would presumably require even larger overdoses or longer holdup pe­ riods to meet a requirement of 0.02 m g / L total residual chlorine at the outfall. The chlorine analyzers used for controlling the dechlorination process are reportedly not sufficiently stable to control an exact, titration-to-endpoint dechlorination process (7, 30). Thus, the usual practice is to use a feed-for­ ward control system in which the SO2 dose is controlled by automatic mea­ surements of residual chlorine and flow rate. Alternatively, a simple fixed feed rate is employed (7, 29, 30,38-40). In either case, overdosing is used to achieve what the operators believe to be zero residual chlorine in the ef­ fluent. As will be discussed later, de­ terminations of zero residual chlorine may be an artifact of the analytical method in many cases.

ionize to form sodium ions and sulfite or bisulfite and thus share the same fate as sulfur dioxide. In moderately well buffered waste­ waters or power plant cooling waters, the acidity produced by Reactions 1-3 has a small effect on pH. The doses of S 0 2 required for dechlorination (usually less than 1 0 - 4 M) are small compared to the alkalinity in most cases. In weakly buffered waters, however, acidity is a potential problem. For example, treating cooling water with 5 m g / L Cl 2 introduces up to 0.14 mM of acid. If this is followed by an equivalent dose of SO2, the yield of acidity doubles by the time all the SO2 has been converted to sulfate. The total yield of acid in this example, 2.8 X 10~ 4 M, could overwhelm the alkalinity of certain freshwaters. Generation of acidity can be avoided by chlorinating with hypochlorite salts and dechlorinating with N a 2 S 0 3 .

Changes in pH Sulfur dioxide is about 20 times more soluble than chlorine. The Hen­ ry's law constant for the reaction,


S 0 2 (aq) = S 0 2 (g)


at 25 ° C is 0.81 a t m / M (41). Sulfur dioxide introduced into water rapidly hydrolyzes to form a solution of sul­ furous acid ( H 2 S 0 3 ) . Sulfurous acid, in turn, rapidly ionizes to an equilibrium mixture of bisulfite (HSO3-) and sulfite (SO3 2 -) (41): H 2 S 0 3 = Η+ + H S 0 3 K = 1.74 Χ ΙΟ" 2



H S 0 3 " = H+ + S 0 3 ~ Κ = 6.31 Χ 10" 8


Equimolar concentrations of bisul­ fite and sulfite occur at pH 7.2. When added to water, N a 2 S 0 3 and N a H S 0 3

Reactions with residual oxidants The sulfur(IV) "dechlors," as they are sometimes called, react with both free and combined forms of chlorine. In freshwater, some of the important reactions of sulfite with residual chlo­ rine species are: S 0 3 2 - + O C l - — SO4 2 - + CI-



Rate constant = 2.7 Χ ΙΟ Χ e x p ( - 3 7 7 0 / T ) M - ' s " 1 (42) S 0 3 2 - + NH 2 C1 + H 2 0 - ^ N H 4 + + S 0 4 2 - + C1Rate fast (43) RNHC1 + S 0 3 2 - + H 2 0 -> R N H 2 + S O 4 2 + C I - + H+


Stanbro and Lenkevich measured rate constants for the reaction of sulfite with several organic chloramines (44). The measured rate constants (in L mo)- 1 m i n - 1 · 10" 4 ) for some rep­ resentative organic chloramines at pH 7 and 25 °C are: mcthylamine, 26; ./V-a-acetylysine, 13; alanine, 7; leu­ cine, 14; and 7V-chloroalanylalanylalanine, 0.5. It is frequently stated that the S(IV) dechlorinating agents react rapidly with all residual chlorine fractions (e.g., Reference 33). Stanbro and Lenkevich show clearly that this is not true with respect to some organic chloramines when excess S(IV) con­ centrations are small. For example, in the presence of 10~ 5 M total sulfite at pH 7 and 25 °C, the pseudo-first-order reduction of TV-chloroalanylalanylalaninc occurs with a half-life of 14 Environ. Sci. Technol., Vol. 18, No. 2, 1984


min. At pH 8, the half-life is about three-quarters of an hour (Stanbro and Lenkevich, in preparation). Thus, some organic chloramines may survive in dechlorinated effluents long enough to pose a toxicity problem in receiving waters. The toxicity of organic chloramines varies greatly, but hydrophobic members of this class are similar to HOC1 and OC1- in toxicity (45). With the present standard methods, organic chloramines are not detectable in the presence of excess sulfite. These methods involve conversion of chloramines to an equivalent amount of I2, which is then determined colorimetrically or electrochemically. If sulfite is present even in slight excess, however, the I 2 is likely to be destroyed before it can be detected. Measurements of this type have led to the entirely unjustified conclusion that dechlorination reactions are all "instantaneous." In estuarine and marine waters where chlorine residuals are rapidly replaced by bromine residuals (46), the following reactions are of concern: SO3 2 - + OBr- ~* SO4 2 - + Br" SO3 2 - + NH 2 Br + H 2 0 -* NH 3 + SO4 2 + Br- + H+


RNHBr + S 0 3 2 " + H 2 0 - RNH 2 + SO4 2 + Br- + H+



With respect to the reaction of organic bromamines with sulfite, Stanbro and Lenkevich found that the reaction of TV-bromoalanylalanylalanine with aqueous sulfite proceeds at least an order of magnitude faster than the reaction of the chlorinated analogue (44). Sulfonates The possibility that excess bisulfite reacts with organic components in receiving waters or treated wastewaters to form sulfonates has apparently not been recognized. Most aldehydes, unhindered ketones, and isocyanates react reversibly with bisulfite (4750): \ C = 0 + HSO3" " ^ ^ C(OH)(OS02)-


Aldehyde and ketone functional groups occur naturally in carbohydrates and various biological metabolites. Low-molecular-weight, volatile 50A

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aldehydes and ketones are found in river water and seawater (51, 52). They are probably also part of highmolecular-weight aquatic humus (53). They are furthermore produced by the chlorination process through decomposition of chloro-amino acids (54), through ring-opening reactions involving dihydroxyaromatic compounds (55), and through other related reactions. The significance of the sulfonation reaction can be illustrated by considering, as an example, the sulfonation of benzaldehyde, which has been identified in drinking water (56):

this reaction is too slow to be of environmental significance (47). Nonetheless, the possibility of decomposing the halocarbon products of chlorination is sufficiently interesting to warrant some future research on the rate of this reaction under water treatment conditions. EPA analytical chemists recommended some time ago that ferrous iron, rather than sulfur-containing dechlorinating agents, be used when collecting drinking water samples for halocarbon determination. They believed that halocarbons might be partly decomposed by the sulfurcontaining agents (59).

Ph—CHO + HSO3" =5=*=

Prevention of halocarbons The scientific community first became aware of halocarbons produced by chlorination through the work of Rook (60) and Jolley (10,13). In the past decade, a voluminous literature on this subject has developed, and several reviews are available (8, 61, 62). Trihalomethanes, the principal halocarbon products of chlorination, are common in drinking water (63-65) and power plant effluents (66,67). In chlorinated wastewaters, their production may be suppressed by NH 3 , but chloroform, as an industrial contaminant, is often present nonetheless (46, 68). Existing toxicological information suggests that the trihalomethanes per se are not a critical threat to aquatic organisms (69, 70). Much less is known, however, about the toxicology or yields of the numerous other halocarbon products that have been detected in chlorinated wastewaters and power plant cooling waters (10, 11, 13, 67, 71). Consequently, halocarbon products of chlorination remain a major environmental concern. Even if S(IV) compounds cannot destroy halocarbons at useful rates, they might nevertheless reduce halocarbon levels by preventing halocarbon formation in the first place. Unfortunately, very little research has been done on the kinetics of halocarbon production from chlorination. Thus, it is not possible to determine at present how short the period between chlorination and dechlorination must be before halocarbon yields are significantly reduced. One evaluation indicates that this period must be substantially shorter than 20 min to reduce trihalomethane yields in estuarine water (72). Assorted evidence suggests that trihalomethanes are produced in two stages—an initial, relatively rapid stage in which organic components in the water become halogenated and a slower second stage in which trihalomethanes are released by



The equilibrium constant for this reaction is 1.41 X 104, and the forward rate constant is 245 M - ' s - 1 at 13 °C (57, 58). Assuming a 10 - 4 M concentration of bisulfite (a dose sufficient to remove 7 ppm of chlorine) and assuming a somewhat lower concentration of benzaldehyde, the aldehyde would become 50% sulfonated in less than a minute and after several minutes would reach an equilibrium state in which it was 60% sulfonated. The rate of this reaction is thus slow with respect to removal of free chlorine (Reaction 4) but fast with respect to removal of some organic cloramines (Reaction 6). These numbers are intended only to show that sulfonation reactions are feasible under water treatment conditions; the rate and degree to which different aldehydes become sulfonated varies considerably. Natural pH and temperature variations also affect rates. The potential to form sulfonates during dechlorination is of interest for several reasons. Sulfonates could tie up a portion of the dechlorinating agent, further retarding slow reduction reactions such as those involving organic chloramines. Reversal of the sulfonation reaction upon dilution in a receiving stream could provide a slow release mechanism for bisulfite. Depending on the individual sulfonates involved and on analytical conditions, the sulfonates might or might not be included in assays of excess dechlorinating agent in effluent streams. Another mechanism for forming sulfonates involves alkyl halides (the Strecker reaction): R - X + SO3 2 - -» R - SO3- + X (X = Cl,Br,I) (12) Available information suggests that

hydrolysis (73-75). After the first stage is well advanced, dechlorination may not be effective in reducing ulti­ mate trihalomethane yields. Better information on the kinetics of halocarbon formation is badly needed to determine whether halocarbon yields might be reduced through de­ chlorination. For power plants with once-through cooling systems, de­ chlorination might offer a practical way of controlling halocarbons, be­ cause in many of these systems an ac­ tive chlorine residual is required only for a few minutes, while the cooling water is passing through the plant. On the other hand, wastewater treatment plant operators, who need a contact time long enough to meet disinfection objectives, may not have the flexibility to control halocarbon yields by de­ chlorination. However, as noted above, in amino-nitrogen-rich wastewaters, halocarbon production is less of a problem. Oxygen depletion Most published discussions of de­ chlorination with S(IV) compounds mention the possibility of downstream oxygen deficits from:

s o 3 2 - + >/2o2


Collectively, electric power plants in the U.S. require 210 billion gal/d of cooling water (87). This exceeds all other uses of water including the demand for irrigation. (Potable water demand is more than five times smaller.) Chlorine is used at many of these plants to control fouling in heat exchangers. As dis­ charge limitations on residual chlorine and halocarbons become more stringent, power plant operators may be forced to adopt dechlorination technology.


The rate of this reaction varies with pH and is catalyzed by traces of divalent Mn, Fe, Co, and Cu compounds, while alcohols such as mannitol, ethanol, and glycerol inhibit the rate (28, 76-79). An immense amount of kinetic re­ search has been done on this reaction since the 19th century, but it is still not possible to predict confidently the rate of this reaction in natural waters. In the presence of high concentrations of Cu or Co, this reaction is extremely fast. However, in water treatment sit­ uations, it is probably slow in most cases, judging from reports of coex­ isting oxygen and sulfite in dechlorinated effluents (2, 7, 40). On stoichiometric grounds, one can argue that oxygen depletion by S(I V) species would be minor in well-aerated waters. At normal temperatures, airsaturated water contains more than 2 Χ ΙΟ - 4 Μ oxygen. Since in most treatment situations, this is likely to exceed sulfite doses significantly, and since only half a mole of oxygen is used per mole of sulfite, the sulfite-oxygen reaction could not produce severe oxygen depletion in initially saturated waters even if it were fast. Only when the downstream receiving water is al­ ready suffering oxygen depletion, which unfortunately is often the case, does the depletion caused by 10 - 6 to

ΙΟ - 4 Μ sulfite have the potential to be serious. If the amount of catalytic material in a particular receiving stream is sufficient to make the sulfite-oxygen reaction rapid with respect to the re­ actions of sulfite with residual chlorine species, the dechlorination process could in part be defeated. It is clearly desirable to obtain better information on the rate of this sulfite-oxygen re­ action under water treatment condi­ tions. Trace-metal interactions Apparently no research has been done on the effect of dechlorination on downstream trace-metal chemistry, despite the known toxicity of many metals (80). There are several poten­ tial effects. These can be illustrated by considering the behavior of copper at a hypothetical power plant that uses chlorination followed by dechlorina­ tion. Chlorination decreases the copperbinding capacity of natural organic material, probably releasing free copper(II) to the water (81). This release may be supplemented by other pro­ cesses. For example, in a power plant with copper alloy condenser tubes, corrosion reactions would produce a steady, low-level release of copper

(82). Sigleo et al. found that chlori­ nation of estuarine waters caused copper and other trace metals to ac­ cumulate in colloidal particles, and they suggested that the release of copper by chlorination was largely offset by its coprecipitation in Fe and Mn oxyhydroxides (83). The colloidal oxyhydroxides were presumed to form by chlorine oxidation of Mn(II) and Fe(II) in the water. Dechlorination with sulfur dioxide would have several effects on copper chemistry. First, it would most likely block or even reverse the oxyhydroxide precipitation reaction. The reaction, 2H+ + SO3 2 - + MnO z — Mn 2+ + SO4 2 - + H 2 0 (14) has an equilibrium constant of 10 +45 · 22 (41), so the reductive dissolution of manganese dioxide and related, amorphous Mn oxyhydroxides is strongly favored when sulfite is present at micromolar, or higher, concentra­ tions. The same is true of Fe(III) ox­ ides and oxyhydroxides at pH values below about 9. Sulfite can also be expected to re­ duce and complex the copper rapidly (79, 84). Since Cu(I) complexes are rare in nature, not much is known about their biological effects. How­ ever, because the chemical behavior of Environ. Sci. Technol., Vol. 18, No. 2, 1984


Cu(l) is generally similar to Ag(I), which is highly toxic {80), the poten­ tial toxicity of Cu(I) complexes is a matter of concern. Other reducible metals conceivably are affected by dechlorination in a similar manner. Reductive stripping of sediments and suspended particles under labo­ ratory conditions is known to release significant quantities of trace metals (85, 86). Therefore, if dechlorinated effluents can reduce Fe and Mn oxyhydroxide grain coatings in down­ stream fluvial sediments and on sus­ pended particles being transported past a discharge point by a river, this process would create a major tracemetal source. It is of interest in this regard that accidental overdosing with SO2 in a drinking water distribution system produced an unpleasant taste which was ascribed to ferrous iron, presumably derived from rust in the pipes (34). Summary Based on this review of dechlorina­ tion, we suggest a number of areas in which knowledge is deficient or non­ existent and thus where research is needed. Treatment technology. Experience to date at both power plants and wastewater treatment plants indicates that better SO2 application and control systems are needed (2, 7). Further­ more, analytical methods are needed to detect traces of chloramines in the presence of excess sulfite. Without such methods, it is impossible to eval­ uate the effectiveness of dechlorination and to monitor the process properly. Once these methods have been devel­ oped, systematic studies should be done to determine minimum sulfite overdoses and contact times required to achieve desired reductions in dif­ ferent classes of residual oxidants. The role of water composition, including its oxygen content, on minimum SO2 overdose and contact time must be assessed for these studies to be gener­ ally applicable. Obtaining better ki­ netic data on the reactions of S(IV) species with residual oxidants and components of natural waters is an essential part of this effort. Finally, the importance of S(IV) compounds for reducing halocarbons needs to be de­ termined as a function of the elapsed time between chlorination and de­ chlorination. Downstream chemical effects. The effect of S(IV) compounds on trace metal geochemistry needs to be ex­ amined. Of particular importance is whether S(IV) compounds cause re­ ductive dissolution of the Fe and Mn 54A

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^wca A rudimentary, retrofitted dechlorination system at a small wastewater treatment plant. Dissolved sulfur dioxide is injected into the discharge conduit from the chlorine contact reservoir. Passage of water through the conduit takes only a few minutes, probably not enough time for complete destruction of organic chloramines. However, because of deficiencies in present analytical methods, such chloramines are not detectable in the effluent if sulfite is in excess.

oxyhydroxides in a river's suspended load. Also, studies should be conducted to determine whether metals such as Cu, Cr, and Hg are converted to un­ usual oxidation states. Information should be gathered on the downstream persistence of S(1V) compounds, any organic chloramines that react slowly with dechlorinating agents, and any sulfonates that form during dechlori­ nation. Persistence will be controlled by the nature of the decomposition mechanisms, which should be identi­ fied. Downstream biological effects. The aquatic toxicity research that has been done with S(IV)-containing waters has evaluated mainly short-term, lethal toxicity. Because wastewater treat­ ment plants dcchlorinate continuously, studies are needed of long-term and sublethal effects, such as effects on reproductive and avoidance behavior. Furthermore, if sulfonates, persistent organic chloramines, or metals in un­ usual oxidation states are demon­ strated to be present in dechlorinated waters, the effects of these substances on organisms must be evaluated. Acknowledgment The authors thank Robert Kawaratani and Richard Sugam for supply­ ing lists of relevant literature. Prepa­ ration of this article was supported fi­ nancially by the Electric Power Re­ search Institute, Palo Alto, Calif. Before publication, this article was reviewed for suitability as an ES&T feature by Philip C. Singer, Depart­ ment of Environmental Sciences and

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George R. Helz (I.) is chairman of the Analytical, Nuclear, and Environmental Division of the Chemistry Department at the University of Maryland, College Park, Mil. 20742. He is also chairman-elect of the American Chemical Society's Division of Geochemistry. He has done extensive research on the fate of chlorine in natural waters. Lynn Kosak-Channing (r.) received her PhD in chemistry from the University of Maryland in 1981. She is now a research chemist at the U.S. Department of Agri­ culture, Bellsville Agricultural Research Center, where she is studying tobacco leaf and smoke chemistry. Environ. Sci. Technol., Vol. 18, No. 2, 1984