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Dichloramine

Dichloramine (NHCl₂), also known as chlorimide or azonous dichloride, is a reactive inorganic chloramine compound derived from ammonia, featuring a central nitrogen atom bonded to one hydrogen and two chlorine atoms. It exists as a yellow gas at room temperature and is highly unstable, readily decomposing into nitrogen gas, hypochlorous acid, and other products under typical aqueous conditions.^[1]^[2] As one of the three primary chloramines (alongside monochloramine and trichloramine), dichloramine plays a transient role in chlorine-ammonia chemistry but is generally regarded as undesirable in practical applications due to its short half-life and potential for forming reactive byproducts.^[3] Dichloramine forms primarily during the chlorination of water containing ammonia, via the reaction of monochloramine (NH₂Cl) with excess hypochlorous acid (HOCl): NH₂Cl + HOCl → NHCl₂ + H₂O, with a rate constant of 3.4 × 10² M⁻¹ s⁻¹ at 25°C.^[2] This synthesis is favored at low pH values (below 5.5) or chlorine-to-ammonia molar ratios greater than 1:1, conditions common in the early stages of water disinfection processes.^[2] Due to its instability, isolating pure dichloramine is challenging, and it is rarely prepared in bulk; instead, it occurs transiently in solutions like secondary sewage effluents or reverse osmosis permeates, where concentrations can range from 0.002 to 0.695 mg/L.^[2]^[4] Physically, dichloramine exhibits intermediate volatility and solubility between monochloramine and trichloramine, though precise data such as melting or boiling points are limited owing to its decomposition.^[2] Chemically, it is a strong oxidant that hydrolyzes rapidly in neutral to alkaline environments (rate constant increasing from 1.4 × 10⁻² s⁻¹ at pH 7 to 1.9 × 10⁻¹ s⁻¹ at pH 12), producing reactive species including hydroxyl radicals (HO•) and peroxynitrite (ONOO⁻).^[4] In water treatment, dichloramine contributes to the "breakpoint chlorination" phenomenon, where excess chlorine destroys combined residuals, but its instability limits its direct disinfectant efficacy compared to monochloramine.^[2] Emerging applications explore its hydrolysis for advanced oxidation processes in membrane desalination, enhancing degradation of trace organics like 1,4-dioxane without UV activation.^[4] Toxicity data for dichloramine are sparse, but it has been identified as highly mutagenic in bacterial assays, such as those using Salmonella typhimurium, due to its lipophilic nature and reactivity with DNA.^[2] Overall, while dichloramine underscores the complexities of chloramine-based disinfection, efforts in water management focus on minimizing its formation to favor more stable species.

Structure and Properties

Molecular Structure

Dichloramine is an inorganic compound with the chemical formula NHCl₂ and a molecular weight of 85.92 g/mol. Its systematic IUPAC name is dichloroazane, while it is commonly referred to as dichloramine or chlorimide. The molecular structure centers on a nitrogen atom bonded to one hydrogen atom and two chlorine atoms, forming three single bonds with a lone pair of electrons occupying the fourth position around the nitrogen. According to valence shell electron pair repulsion (VSEPR) theory, this AX₃E configuration results in a trigonal pyramidal geometry, where the lone pair repels the bonding pairs, compressing the bond angles to approximately 100–110°; specifically, the Cl–N–Cl angle is approximately 110°. The N–Cl bonds are polar covalent, with a bond length of approximately 1.75 Å, reflecting the electronegativity difference between nitrogen and chlorine. In terms of electronic structure, the nitrogen atom exhibits a +1 oxidation state, calculated from the +1 oxidation state of hydrogen and -1 for each chlorine, balancing the overall neutral charge. Spectroscopic characterization supports this arrangement, with infrared (IR) absorption bands observed for the N–H stretch at approximately 3300 cm⁻¹ and N–Cl stretches around 700 cm⁻¹. Nuclear magnetic resonance (NMR) data is limited due to the compound's instability, but the theoretical structure aligns with vibrational analyses confirming the pyramidal configuration.

Physical and Chemical Properties

Dichloramine (NHCl₂) appears as a yellow gas at room temperature and standard pressure. Due to its instability, it is rarely isolated in pure form and is typically handled in dilute aqueous solutions or as a transient species. The compound exists primarily in the gaseous state under ambient conditions, with a predicted liquid density of approximately 1.43 g/cm³, though experimental measurement is challenging owing to decomposition. Due to its instability, a standard boiling point has not been determined experimentally; predicted values suggest around 115–117 °C at reduced pressure (15 Torr). As a gas at standard temperature and pressure (STP), its density is approximately 3.84 g/L, calculated from its molar mass of 85.92 g/mol assuming ideal gas behavior. Dichloramine exhibits moderate solubility in water but hydrolyzes rapidly in aqueous media, limiting its persistence. This solubility, combined with its volatility, positions it between monochloramine and trichloramine in physicochemical behavior. Many properties are derived from computational predictions or limited spectroscopic studies due to the compound's instability. Chemically, dichloramine is a highly reactive oxidizing agent, prone to decomposition in the presence of air, light, or moisture when concentrated. Its instability arises from the weak N-Cl bonds, rendering it endothermic and endergonic. The standard enthalpy of formation (Δ_f H°{298K}) is +33.5 kJ/mol, and the standard Gibbs free energy of formation (Δ_f G°{298K}) is +41.6 kJ/mol, values derived from high-level ab initio computations that benchmark its thermodynamic instability. Limited data exist on its acid-base properties due to instability.

Synthesis

Laboratory Preparation

Dichloramine (NHCl₂) is synthesized in the laboratory through the controlled chlorination of monochloramine in aqueous solution, where monochloramine is first prepared by reacting equimolar amounts of sodium hypochlorite and ammonium chloride. Chlorine gas is then introduced to the monochloramine solution to yield dichloramine according to the equation NH₂Cl + Cl₂ → NHCl₂ + HCl. The reaction is maintained at an acidic pH of approximately 4–5 using a phosphate buffer to stabilize the product and minimize further decomposition to nitrogen trichloride.^[5] Dichloramine is highly unstable and sensitive to moisture and heat. The resulting solution is used directly for studies, with dichloramine concentration quantified spectrophotometrically at 245 nm (molar absorptivity ε = 405 M⁻¹ cm⁻¹). Yields vary but are optimized by precise control of chlorine addition to avoid over-chlorination.^[5] An alternative laboratory route involves reacting gaseous ammonia with excess sodium hypochlorite solution under acidic conditions, producing a mixture of chloramines including dichloramine and nitrogen trichloride. Dichloramine is isolated from this mixture via fractional distillation at reduced pressure, leveraging its lower boiling point relative to byproducts. This approach, while less selective, allows for scalable preparation but necessitates careful temperature control below 0°C to limit side reactions.^[6] Purification of dichloramine often entails vacuum distillation under inert conditions or trapping the volatile product in cold solvents like diethyl ether at -78°C to condense it as a pale yellow liquid. Typical overall yields from these methods range from 50% to 70%, depending on reaction scale and purity of reagents. Early 20th-century laboratory adaptations, building on foundational chloramine research by chemists like Raschig, prioritized explosion-proof apparatus and dilute solutions to mitigate the compound's inherent explosivity during handling and storage.^[7]

Formation in Aqueous Media

Dichloramine forms in aqueous media primarily through the equilibrium reaction between monochloramine and hypochlorous acid, represented as NH₂Cl + HOCl ⇌ NHCl₂ + H₂O, with a rate constant of 3.4 × 10² M⁻¹ s⁻¹ at 25°C.^[8]^[2] This process occurs during chlorination of water containing ammonia, where hypochlorous acid (derived from chlorine dissolution) reacts stepwise with ammonia to produce chloramines.^[9] The formation is pH-dependent, with dichloramine favored under mildly acidic conditions, achieving maximum stability and concentration at pH 4–6.^[10] In the context of breakpoint chlorination, dichloramine emerges as an intermediate when the chlorine-to-nitrogen (Cl:N) weight ratio exceeds 5:1, marking the transition from monochloramine dominance to further oxidation products.^[11] At these ratios, excess chlorine drives the reaction, often simplified as NH₂Cl + Cl₂ + H₂O → NHCl₂ + 2HCl, leading to a temporary dip in total chlorine residual before the breakpoint.^[12] This stage is critical in water treatment processes aiming to remove ammonia, as dichloramine contributes to the observed "hump" in the chlorination curve.^[11] The kinetics of dichloramine formation are rapid, typically occurring within seconds to minutes upon mixing reactants, but its persistence is limited by a short half-life in water, often on the order of minutes under typical treatment conditions influenced by factors such as temperature and ammonia concentration.^[13] Higher temperatures accelerate decomposition, while elevated ammonia levels can shift equilibria back toward monochloramine.^[2] These dynamics underscore dichloramine's transient role in aqueous systems. Detection of dichloramine in aqueous mixtures relies on methods like ultraviolet spectrophotometry, which exploits its absorption bands at 250–300 nm, or amperometric titration, which differentiates it from other chlorine species through selective electrochemical reduction.^[14]^[15] These techniques enable quantification in complex water samples, with amperometry offering high sensitivity for low concentrations.^[16]

Reactions

Hydrolysis and Decomposition

Dichloramine undergoes base-catalyzed hydrolysis in aqueous media, primarily through the reaction with hydroxide ions to yield nitroxyl (HNO), chloride ions, and water:

\ce{NHCl2 + OH- -> HNO + 2Cl- + H2O}

This process is second-order, with the rate depending on both dichloramine and OH⁻ concentrations, and the rate constant is 186 ± 6 M⁻¹ s⁻¹ at 25 °C.^[17] The hydrolysis rate accelerates significantly at pH values above 7, where OH⁻ concentration increases, leading to faster decay of dichloramine.^[17] In neutral water (pH ≈ 7), the pseudo-first-order hydrolysis rate constant is approximately 1.4 × 10⁻² s⁻¹, corresponding to a half-life of about 50 seconds; under alkaline conditions (e.g., pH 8–12), the half-life shortens to 10 seconds or less due to enhanced OH⁻ reactivity.^[4] Thermal decomposition of dichloramine occurs readily above 0 °C, following a pathway that produces nitrogen gas, hydrogen chloride, and hypochlorous acid:

\ce{3 NHCl2 -> N2 + 3 HCl + 3 HOCl}

This reaction is first-order in dichloramine at 25 °C and pH 7, contributing to the compound's instability in solution at ambient temperatures. The process is distinct from hydrolysis, as it proceeds even in the absence of excess base, though both pathways often compete in aqueous environments. Photolytic decomposition is initiated by ultraviolet (UV) light, leading to the degradation of dichloramine and formation of products such as nitrite, nitrate, nitrous oxide, and ammonium. The photodecay rate is independent of pH but varies with wavelength (e.g., tested at 222, 254, and 282 nm).^[18] This pathway enhances dichloramine degradation in UV-treated water systems, producing reactive species that influence overall disinfection dynamics. Recent investigations (2024) have revealed that dichloramine hydrolysis in low-conductivity waters, such as desalinated reverse osmosis permeate, yields chloronitramide anion (Cl–NO₂⁻) as a stable end-product via reactive nitrogen intermediates like HNO reacting with nitrite or peroxynitrite.^[19] This byproduct persists in chloraminated systems and raises concerns for trace contaminant formation in potable reuse applications, as its reactivity and toxicity remain under evaluation.^[19]

Reactivity with Other Species

Dichloramine exhibits significant reactivity as an oxidant toward various organic species, particularly amines and phenols, resulting in the formation of chlorinated byproducts. In reactions with secondary amines, such as dimethylamine, dichloramine (NHCl₂) oxidizes the amine to form chlorinated unsymmetrical dialkylhydrazine intermediates, which can further lead to N-nitrosodimethylamine (NDMA) under aerobic conditions.^[20] This pathway is considered the primary mechanism for NDMA formation in chlorinated drinking waters, where even trace levels of dichloramine contribute substantially due to its higher reactivity compared to monochloramine.^[17] With phenolic compounds, dichloramine demonstrates reactivity orders of magnitude lower than bromamines but still sufficient to produce substituted chlorophenols or quinone-like oxidation products, highlighting its role in generating disinfection byproducts during water treatment.^[21] In reduction reactions, dichloramine serves as an oxidant toward reducing agents like sulfite (SO₃²⁻), commonly employed in dechlorination processes to mitigate environmental impacts of chlorinated effluents. The reaction proceeds as NHCl₂ + SO₃²⁻ + H₂O → NH₂OH + SO₄²⁻ + 2Cl⁻, converting dichloramine to hydroxylamine while oxidizing sulfite to sulfate; this bimolecular process is faster for dichloramine than for monochloramine, making it relevant in studies optimizing dechlorination kinetics for wastewater discharge.^[22] Such reductions underscore dichloramine's utility in controlled oxidant removal but also its potential to complicate effluent toxicity assessments if incomplete. Dichloramine participates in a reversible equilibrium with ammonia and monochloramine, governed by pH and chlorine-to-nitrogen ratios: NHCl₂ + NH₃ ⇌ 2NH₂Cl.^[23] This equilibrium allows dichloramine to disproportionate back to monochloramine in the presence of excess ammonia, influencing speciation during chloramination and helping maintain desired disinfectant profiles in water systems. The oxidizing strength of dichloramine is reflected in its standard redox potential, with E° ≈ +0.79 V for the NHCl₂/NH₂Cl couple under basic conditions, positioning it as a potent oxidant comparable to monochloramine (E° ≈ +0.75 V) and capable of driving reactions with a range of reductants.^[24] Analytically, dichloramine interferes in determinations of halides and residual chlorine, particularly in amperometric and titrimetric methods, due to its tendency to release or mimic free chlorine equivalents. In amperometric titrations for chlorine species, dichloramine can contribute to the monochloramine signal at elevated temperatures or low pH, leading to overestimation of residuals and complicating speciation analysis.^[15] Similarly, in halide titrations such as argentometric methods for chloride, dichloramine's reduction can liberate chloride ions, inflating apparent concentrations and requiring pre-treatment steps like acidification or masking agents for accurate quantification.^[25]

Applications

Role in Water Disinfection

Dichloramine forms as an unwanted intermediate during the chloramination process in water treatment when the chlorine-to-nitrogen (Cl:N) weight ratio exceeds approximately 5:1, leading to sequential chlorination of ammonia beyond monochloramine.^[26] This species contributes to the breakpoint chlorination curve by decomposing into nitrogen gas, hypochlorous acid, and other products, which reduces total residual but is intentionally minimized to less than 10% of total chloramines to prioritize the more stable monochloramine for residual disinfection.^[2]^[27] In terms of disinfectant efficacy, dichloramine exhibits greater potency than monochloramine against bacteria, including Escherichia coli, with inactivation rates approximately 3 times faster at neutral pH and comparable temperatures.^[28] Its higher reactivity contributes to this greater potency, though this comes at the cost of shorter persistence in distribution systems due to rapid decomposition.^[29] Dichloramine's presence can promote the formation of disinfection byproducts such as N-nitrosodimethylamine (NDMA) and N-chlorinated organic compounds when reacting with natural organic matter or amines in source water.^[11] Regulatory agencies like the U.S. EPA set a maximum residual disinfectant level of 4 mg/L (as Cl₂) for total chloramines to balance disinfection benefits against byproduct risks.^[30] It disadvantages water quality by imparting noticeable taste and odor at concentrations as low as 0.13–0.15 mg/L and accelerating corrosion in metallic pipes.^[29]^[11]^[31] Monitoring and control per WHO and EPA guidelines focus on achieving monochloramine dominance by maintaining pH above 7.5 and Cl:N weight ratios of 3:1 to 5:1, which suppresses dichloramine formation while ensuring effective residual disinfection throughout the system.^[32]^[11]

Laboratory and Analytical Uses

In environmental and analytical chemistry, dichloramine reacts with secondary amines such as dimethylamine to form unsymmetrical intermediates like chlorinated dialkylhydrazines, which are key in the formation of disinfection byproducts such as N-nitrosodimethylamine (NDMA).^[20] This reactivity is studied to understand nitrogen-halogen bond formation and byproduct pathways in water treatment. In analytical chemistry, dichloramine plays a role in titrimetric methods for ammonia determination through the exploitation of chloramine formation equilibria during breakpoint chlorination. Excess chlorine is added to an ammonia-containing sample, leading to sequential formation of monochloramine and dichloramine; the chlorine demand to reach the breakpoint—where dichloramine decomposes to nitrogen gas and hypochlorous acid—is measured by iodometric titration of residual chlorine.^[33] The equilibria shift with pH and Cl/N ratio, allowing quantification of ammonia based on the stoichiometric consumption (typically 7.6:1 Cl₂:NH₃ by weight at the breakpoint), with dichloramine's peak concentration providing insight into reaction kinetics.^[2] As a model compound in nitrogen-halogen chemistry research, dichloramine is employed to investigate decomposition mechanisms and reactive species formation. Its instability facilitates studies on hydrolysis and redox pathways, revealing products like nitrite and nitrate under varying pH conditions. Recent investigations have focused on its photochemical behavior, where UV irradiation of dichloramine generates chlorine radicals (Cl•) for advanced oxidation processes aimed at degrading trace organic contaminants in water matrices. For instance, UV-LED photolysis at 290 nm induces dichloramine breakdown, producing hydroxyl radicals (•OH) and chlorine atoms with quantum yields comparable to free chlorine systems, enhancing micropollutant abatement efficiency.^[34] Additionally, dichloramine hydrolysis has been explored as a non-UV advanced oxidation process in membrane desalination permeate for potable reuse, generating hydroxyl radicals to degrade trace organics like 1,4-dioxane without requiring UV activation (as of 2024).^[4] Isotopically labeled dichloramine, such as ¹⁵N-NHCl₂, is utilized in environmental research to trace nitrogen pathways during chloramination reactions. By incorporating ¹⁵N from labeled precursors, studies track the incorporation of nitrogen into disinfection byproducts like N-nitrosodimethylamine (NDMA), confirming dichloramine as a primary reactant in unsymmetrical intermediate formation and distinguishing inorganic from organic nitrogen sources in complex samples.^[17] Due to its rapid decomposition—half-life on the order of minutes in aqueous solution at neutral pH—dichloramine is rarely isolated and is instead generated in situ for laboratory experiments, typically by reacting monochloramine with hypochlorous acid under acidic conditions.^[2] This approach ensures controlled reactivity while mitigating handling risks associated with its volatility and explosiveness in concentrated forms.^[5]

Toxicology and Safety

Health Effects

Dichloramine acts as a respiratory irritant upon inhalation, potentially causing coughing, throat irritation, and exacerbation of asthma-like symptoms, similar to other chloramines formed during mixing of bleach and ammonia-containing compounds.^[35] Eye and nasal irritation can occur from airborne exposure to chloramines including dichloramine, with symptoms such as burning and watering eyes reported in contexts like swimming pools where these species form.^[36] Skin contact with dichloramine may lead to irritation, though specific thresholds for burns are not well-documented due to its instability.^[37] In subchronic oral exposure studies, rats administered dichloramine in drinking water at concentrations of 0.2–200 ppm (equivalent to 0.019–24 mg/kg body weight/day) for 13 weeks exhibited mild adaptive histopathological changes in the kidneys and thyroid glands, as well as epithelial hyperplasia in the gastric cardia at doses as low as 0.019 mg/kg/day in males.^[38] No clinical signs of toxicity, such as reduced body weight or overt behavioral changes, were observed across these dose levels.^[38] Data on carcinogenicity are lacking, with insufficient evidence to assess long-term oncogenic risks.^[1] Inhalation toxicity data specific to dichloramine are limited, but it is considered more potent as a respiratory irritant than monochloramine due to its higher chlorine content and reactivity, though less so than trichloramine.^[39] No LC50 values have been established for rodents, reflecting challenges in studying its unstable nature.^[1] Upon ingestion via contaminated water, dichloramine contributes to the total chloramine burden, with levels exceeding 4 mg/L potentially causing gastrointestinal upset such as stomach discomfort.^[40] However, human data are sparse, and effects are generally minimal at typical disinfection residuals.^[11] The primary mechanism of dichloramine's toxicity involves oxidative damage to biological tissues, functioning as a membrane-penetrating oxidant that can react with proteins and induce cellular stress.^[1] A 2020 toxicological review concluded that available data are insufficient to derive human occupational exposure limits for dichloramine.^[1] As of November 2024, a study identified chloronitramide anion as a major decomposition product of dichloramine in chloraminated drinking water, potentially present in about one-third of U.S. water systems; its toxicity remains unknown but may pose health risks warranting further research.^[19]

Handling and Regulatory Considerations

Dichloramine should be stored in sealed containers under an inert atmosphere, maintained at cool temperatures to prevent decomposition, and protected from light and moisture exposure to ensure stability.^[41] Incompatible materials, such as strong reducing agents or ammonia sources, must be stored separately to avoid unintended reactions.^[37] Safe handling requires the use of appropriate personal protective equipment (PPE), including chemical-resistant gloves, safety goggles, and respirators with appropriate cartridges to protect against inhalation of vapors, which can cause respiratory irritation.^[37] Work areas must be well-ventilated to minimize exposure, as no specific occupational exposure limit has been established for dichloramine due to insufficient data.^[1] The Occupational Safety and Health Administration (OSHA) does not establish a specific permissible exposure limit (PEL) for dichloramine, so adherence to general irritant thresholds is advised.^[42] In the event of a spill, the area should be evacuated and ventilated immediately to disperse vapors, followed by neutralization using a reducing agent such as sodium bisulfite solution to convert dichloramine to less hazardous chloride and nitrogen compounds.^[43] Absorbent materials can then be used to contain the neutralized residue, which should be collected for proper disposal, avoiding direct contact with skin or eyes.^[44] Regulatory oversight for dichloramine primarily falls under broader chloramine standards, with the U.S. Environmental Protection Agency (EPA) setting a secondary maximum contaminant level of 4 mg/L for total chloramines in drinking water to prevent taste, odor, and corrosion issues.^[45] Facilities handling dichloramine must comply with OSHA general industry standards for hazardous chemicals, including hazard communication and emergency planning, though no compound-specific PEL exists.^[42] Environmentally, dichloramine exhibits moderate aquatic toxicity, more toxic than monochloramine but less so than free chlorine, depending on exposure duration and water conditions.^[46] Upon degradation, it primarily breaks down to non-toxic nitrogen gas (N₂) and chloride ions (Cl⁻), though the accumulation of chloride can contribute to increased salinity in receiving waters.^[2] Waste streams containing dichloramine must be treated prior to release by reduction with sodium thiosulfate to eliminate residual chloramine activity and minimize toxicity, in accordance with EPA guidelines for dechlorination in wastewater effluents.^[47] Treated wastes should then be disposed of through licensed facilities to ensure compliance with local and federal hazardous waste regulations.^[2]

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No clinical signs of toxicity were observed in either case. Both males and females dosed with 90 ppm trichloramine had significantly increased relative kidney/ ...
[48]
Welcome to Mr. Water - Chloramine Controversy
Oct 17, 2019 · Chloramines are respiratory irritants, with trichloramine being the most toxic and monochloramine the least toxic, with dichloramine somewhere in the middle.
[49]
Chloramine in Drinking Water - Commonwealth of Pennsylvania
Because chloramine does not tend to react with organic compounds, many systems will experience fewer taste and odor complaints when using chloramine.Missing: issues | Show results with:issues
[50]
Dichloramine SDS, 3400-09-7 Safety Data Sheets - ECHEMI
Boiling point or initial boiling point and boiling range. no data available ... Density and/or relative density. 1.429g/cm3. Relative vapour density. no ...<|separator|>
[51]
[PDF] exposure to Chloramine - NJ.gov
WORKPLACE EXPOSURE LIMITS. No occupational exposure limits have been established for. Chloramine. This does not mean that this substance is not harmful. Safe ...Missing: dichloramine | Show results with:dichloramine
[52]
https://www.osha.gov/annotated-pels
[53]
[PDF] Basic Disinfection Study Guide - Wisconsin DNR
Move the empty chlorine cylinder to the empty cylinder storage area and secure with a safety chain. A. Manual control. Feed rates, starting, and stopping are ...<|separator|>
[54]
[PDF] HAZARD DATA SPILL/LEAKS PHYSICAL PROPERTIES ... - NJ.gov
Neutralize with Sodium Bisulfite, cover with Soda Ash and place into covered containers for disposal or wash with plenty of water. DO NOT wash into sewer.Missing: chloramine | Show results with:chloramine
[55]
Chloramines in Drinking Water | US EPA
Mar 14, 2025 · Chloramines are disinfectants used to treat drinking water. Chloramines are most commonly formed when ammonia is added to chlorine to treat ...Missing: dichloramine ingestion
[56]
[PDF] Acute toxicity of residual chlorine and ammonia to some native ...
Rosenberger7 and Merkens8 concluded that free chlorine is slightly more toxic to fish than dichloramine and that di- chloramine is more toxic than ...
[57]
[PDF] GUIDANCE MANUAL FOR THE DISPOSAL OF CHLORINATED ...
Dechlorination using reducing agents such as sodium bisulfite and sodium metabisulfite may deplete oxygen concentrations in receiving waters. Water quality ...