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| Names | |
|---|---|
| IUPAC name Sodium hypochlorite | |
Other names
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| Identifiers | |
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3D model (JSmol) | |
| ChEBI | |
| ChemSpider |
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| DrugBank | |
| ECHA InfoCard | 100.028.790 |
| EC Number |
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| KEGG |
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| RTECS number |
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| UNII | |
| UN number | 1791 |
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| Properties | |
| NaOCl | |
| Molar mass | 74.442 g/mol |
| Appearance |
|
| Odor | Chlorine-like and sweetish (pentahydrate)[1] |
| Density | 1.11 g/cm3 |
| Melting point | 18 °C (64 °F; 291 K) (pentahydrate) |
| Boiling point | 101 °C (214 °F; 374 K) (decomposes) (pentahydrate) |
| 29.3 g/(100 mL) (0 °C)[2] | |
| Acidity (pKa) | 7.5185 |
| Basicity (pKb) | 6.4815 |
| Thermochemistry | |
Std enthalpy of formation(ΔfH⦵298) | −347.1 kJ/mol |
| Pharmacology | |
| D08AX07 (WHO) | |
| Hazards | |
| Occupational safety and health (OHS/OSH): | |
Main hazards | corrosive, oxidizing agent[3] |
| GHS labelling: | |
  | |
| Danger | |
| H302,H314,H410 | |
| P260,P264,P273,P280,P301+P330+P331,P303+P361+P353,P304+P340,P305+P351+P338,P310,P321,P363,P391,P405,P501 | |
| NFPA 704 (fire diamond) | |
| Safety data sheet (SDS) | ICSC 1119 (solution, > 10% active chlorine) ICSC 0482 (solution, < 10% active chlorine) |
| Related compounds | |
Otheranions | |
Othercations | |
Related compounds | |
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |
Sodium hypochlorite is analkalineinorganicchemical compound with theformulaNaOCl (also written as NaClO). It is commonly known in a diluteaqueous solution asbleach or chlorine bleach.[4] It is thesodiumsalt ofhypochlorous acid, consisting of sodiumcations (Na+) andhypochloriteanions (−OCl, also written asOCl− andClO−).
The anhydrouscompound is unstable and may decompose explosively.[5][6] It can be crystallized as apentahydrateNaOCl·5H2O, a pale greenish-yellow solid which is notexplosive and is stable if kept refrigerated.[7][8][9]
Sodium hypochlorite is most often encountered as a pale greenish-yellow dilute solution referred to as chlorine bleach, which is ahousehold chemical widely used (since the 18th century) as adisinfectant andbleaching agent. In solution, the compound is unstable and easily decomposes, liberatingchlorine, which is the active principle of such products. Sodium hypochlorite is still the most importantchlorine-based bleach.[10][11]
Its corrosive properties, common availability, and reaction products make it a significant safety risk. In particular,mixing liquid bleach with other cleaning products, such as acids found inlimescale-removing products, will release toxicchlorine gas. A common misconception is that mixing bleach with ammonia also releases chlorine, but in reality they react to producechloramines such asnitrogen trichloride. With excess ammonia andsodium hydroxide,hydrazine may be generated.
Anhydrous sodium hypochlorite can be prepared but, like many hypochlorites, it is highly unstable and decomposes explosively on heating or friction.[5] The decomposition is accelerated bycarbon dioxide at Earth'satmospheric levels - around 4 parts per ten thousand.[6][12] It is a white solid with theorthorhombic crystal structure.[13]
Sodium hypochlorite can also be obtained as acrystallinepentahydrateNaOCl·5H2O, which is not explosive and is much more stable than the anhydrous compound.[6][7] The formula is sometimes given in its hydrous crystalline form as2NaOCl·10H2O.[14] The Cl–O bond length in the pentahydrate is 1.686 Å.[9] The transparent, light greenish-yellow, orthorhombic[15][16] crystals contain 44% NaOCl by weight and melt at 25–27 °C. The compound decomposes rapidly at room temperature, so it must be kept under refrigeration. At lower temperatures, however, it is quite stable: reportedly only 1% decomposition after 360 days at 7 °C.[8][17]
A 1966 USpatent claims that stable solid sodium hypochlorite dihydrateNaOCl·2H2O can be obtained by carefully excludingchloride ions (Cl−), which are present in the output of common manufacturing processes and are said to catalyze the decomposition of hypochlorite intochlorate (ClO−3) and chloride. In one test, the dihydrate was claimed to show only 6% decomposition after 13.5 months of storage at −25 °C. The patent also claims that the dihydrate can be reduced to the anhydrous form by vacuum drying at about 50 °C, yielding a solid that showed no decomposition after 64 hours at −25 °C.[18]
At typical ambient temperatures, sodium hypochlorite is more stable in dilute solutions that contain solvatedNa+ andOCl− ions. The density of the solution is 1.093 g/mL at 5% concentration,[19] and 1.21 g/mL at 14%, 20 °C.[20]Stoichiometric solutions are fairlyalkaline, with pH 11 or higher[8] since the hypochlorite ion is aweak base:
The following species and equilibria are present in NaOCl/NaCl solutions:[21]
The second equilibrium equation above will be shifted to the right if the chlorineCl2 is allowed to escape as gas. The ratios ofCl2, HOCl, andOCl− in solution are also pH dependent. At pH below 2, the majority of the chlorine in the solution is in the form of dissolved elementalCl2. At pH greater than 7.4, the majority is in the form of hypochloriteClO−.[10] Theequilibrium can be shifted by adding acids (such ashydrochloric acid) or bases (such assodium hydroxide) to the solution:
At a pH of about 4, such as obtained by the addition ofstrong acids likehydrochloric acid, the amount of undissociated (nonionized) HOCl is highest. The reaction can be written as:
Sodium hypochlorite solutions combined with acid evolve chlorine gas, particularly strongly at pH < 2, by the reactions:
At pH > 8, the chlorine is practically all in the form of hypochlorite anions (OCl−). The solutions are fairly stable at pH 11–12. Even so, one report claims that a conventional 13.6% NaOCl reagent solution lost 17% of its strength after being stored for 360 days at 7 °C.[8] For this reason, in some applications one may use more stable chlorine-releasing compounds, such ascalcium hypochloriteCa(ClO)2 ortrichloroisocyanuric acid(CNClO)3.[citation needed]
Anhydrous sodium hypochlorite is soluble inmethanol, and solutions are stable.[citation needed]
In solution, under certain conditions, the hypochlorite anion may alsodisproportionate (autoxidize) to chloride andchlorate:[22]
In particular, this reaction occurs in sodium hypochlorite solutions at high temperatures, formingsodium chlorate and sodium chloride:[22][23]
This reaction is exploited in the industrial production of sodium chlorate.
An alternative decomposition of hypochlorite produces oxygen instead:
In hot sodium hypochlorite solutions, this reaction competes with chlorate formation, yielding sodium chloride and oxygen gas:[22]
These two decomposition reactions ofNaOCl solutions are maximized at pH around 6. For example, at 80 °C, withNaOCl andNaCl concentrations of 80mM, over the pH range 5−10.5, both reactions have rate proportional to, decomposition is fastest at pH 6.5, and chlorate is produced with ~95% efficiency.[22] Above pH 11, both reactions have rate proportional to, decomposition is much slower, and chlorate is produced with ~90% efficiency.[24] This decomposition is affected by light[23] and metal ioncatalysts such ascopper,nickel,cobalt,[22] andiridium.[25] Catalysts likesodium dichromateNa2Cr2O7 andsodium molybdateNa2MoO4 may be added industrially to reduce the oxygen pathway, but a report claims that only the latter is effective.[22][failed verification]
Titration of hypochlorite solutions is often done by adding a measured sample to an excess amount of acidified solution ofpotassium iodide (KI) and then titrating the liberatediodine (I2) with a standard solution ofsodium thiosulfate orphenylarsine oxide, usingstarch as indicator, until the blue color disappears.[16]
According to one US patent, the stability of sodium hypochlorite content of solids or solutions can be determined by monitoring theinfrared absorption due to the O–Cl bond. The characteristic wavelength is given as 140.25μm for water solutions, 140.05 μm for the solid dihydrateNaOCl·2H2O, and 139.08 μm for the anhydrous mixed saltNa2(OCl)(OH).[18]
Oxidation ofstarch by sodium hypochlorite, which addscarbonyl andcarboxyl groups, is relevant to the production ofmodified starch products.[26]
In the presence of aphase-transfer catalyst, alcohols are oxidized to the correspondingcarbonyl compound (aldehyde orketone).[27][8] Sodium hypochlorite can also oxidize organicsulfides tosulfoxides orsulfones;disulfides orthiols tosulfonyl halides; andimines tooxaziridines.[8] It can alsode-aromatizephenols.[8]
Heterogeneous reactions of sodium hypochlorite and metals such aszinc proceed slowly to give themetal oxide or hydroxide:[citation needed]
Homogeneous reactions with metalcoordination complexes proceed somewhat faster. This has been exploited in theJacobsen epoxidation.[citation needed]
If not properly stored in airtight containers, sodium hypochlorite reacts withcarbon dioxide to formsodium carbonate:
Sodium hypochlorite reacts with most nitrogen compounds to form volatilemonochloramine,dichloramines, andnitrogen trichloride:
Sodium thiosulfate is an effective chlorine neutralizer. Rinsing with a 5 mg/L solution, followed by washing with soap and water, will remove chlorine odor from the hands.[28]
Potassium hypochlorite was first produced in 1789 byClaude Louis Berthollet in his laboratory on the Quai deJavel inParis, France, by passingchlorine gas through a solution ofpotash lye. The resulting liquid, known as "Eau de Javel" ("Javel water"), was a weak solution of potassium hypochlorite.Antoine Labarraque replaced potash lye by the cheapersoda lye, thus obtaining sodium hypochlorite (Eau de Labarraque).[29][30]
Hence, chlorine is simultaneouslyreduced andoxidized; this process is known asdisproportionation.[citation needed]
The process is also used to prepare the pentahydrateNaOCl·5H2O for industrial and laboratory use. In a typical process, chlorine gas is added to a 45–48% NaOH solution. Some of the sodium chloride precipitates and is removed by filtration, and the pentahydrate is then obtained by cooling the filtrate to 12 °C.[8]
Another method involved the reaction of sodium carbonate ("washing soda") withchlorinated lime ("bleaching powder"), a mixture ofcalcium hypochloriteCa(OCl)2,calcium chlorideCaCl2, andcalcium hydroxideCa(OH)2:
This method was commonly used to produce hypochlorite solutions for use as a hospital antiseptic that was sold after World War I under the names "Eusol", an abbreviation for Edinburgh University Solution Of (chlorinated) Lime – a reference to the university's pathology department, where it was developed.[31]
Near the end of the nineteenth century, E. S. Smith patented thechloralkali process: a method of producing sodium hypochlorite involving the electrolysis ofbrine to producesodium hydroxide and chlorine gas, which then mixed to form sodium hypochlorite.[32][30][33] The key reactions are:
Both electric power and brine solutions were in cheap supply at the time, and various enterprising marketers took advantage of the situation to satisfy the market's demand for sodium hypochlorite. Bottled solutions of sodium hypochlorite were sold under numerous trade names.[citation needed]
Today, an improved version of this method, known as the Hooker process (named after Hooker Chemicals, acquired byOccidental Petroleum), is the only large-scale industrial method of sodium hypochlorite production. In the process, sodium hypochlorite (NaClO) andsodium chloride (NaCl) are formed when chlorine is passed into a cold dilutesodium hydroxide solution. The chlorine is prepared industrially byelectrolysis with minimal separation between theanode and thecathode. The solution must be kept below 40 °C (by cooling coils) to prevent the undesired formation ofsodium chlorate.[citation needed]
Commercial solutions always contain significant amounts of sodium chloride (common salt) as the mainby-product, as seen in the equation above.
A 1966 patent describes the production of solid stable dihydrateNaOCl·2H2O by reacting a chloride-free solution of hypochlorous acid HClO (such as prepared from chlorine monoxide ClO and water), with a concentrated solution of sodium hydroxide. In a typical preparation, 255 mL of a solution with 118 g/L HClO is slowly added with stirring to a solution of 40 g of NaOH in water 0 °C. Some sodium chloride precipitates and is removed by filtration. The solution is vacuum evaporated at 40–50 °C and 1–2mmHg until the dihydrate crystallizes out. The crystals are vacuum-dried to produce a free-flowing crystalline powder.[18]
The same principle was used in a 1993 patent to produce concentratedslurries of the pentahydrateNaClO·5H2O. Typically, a 35% solution (by weight) of HClO is combined with sodium hydroxide at about or below 25 °C. The resulting slurry contains about 35% NaClO, and are relatively stable due to the low concentration of chloride.[34]
Householdbleach sold for use in laundering clothes is a 3–8% solution of sodium hypochlorite at the time of manufacture. Strength varies from one formulation to another and gradually decreases with long storage. Sodium hydroxide is usually added in small amounts to household bleach to slow down the decomposition of NaClO.[10]
Domestic use patio blackspot remover products are ~10% solutions of sodium hypochlorite.
A 10–25% solution of sodium hypochlorite is, according to Univar's safety sheet, supplied with synonyms ortrade names bleach, Hypo, Everchlor, Chloros, Hispec, Bridos, Bleacol, or Vo-redox 9110.[35]
A 12% solution is widely used in waterworks for thechlorination of water, and a 15% solution is more commonly[36] used for disinfection of wastewater in treatment plants. Sodium hypochlorite can also be used for point-of-use disinfection of drinking water,[37] taking 0.2–2 mg of sodium hypochlorite per liter of water.[38]
Dilute solutions (50 ppm to 1.5%) are found in disinfecting sprays and wipes used on hard surfaces.[39][40]
Household bleach is, in general, a solution containing 3–8% sodium hypochlorite, by weight, and 0.01–0.05%sodium hydroxide; the sodium hydroxide is used to slow the decomposition of sodium hypochlorite intosodium chloride andsodium chlorate.[41]
Sodium hypochlorite has destaining properties.[42] Among other applications, it can be used to removemold stains, dental stains caused byfluorosis,[43] and stains on crockery, especially those caused by thetannins intea. It has also been used inlaundry detergents and as a surface cleaner. It is also used insodium hypochlorite washes.
Its bleaching, cleaning, deodorizing, and caustic effects are due tooxidation andhydrolysis (saponification). Organic dirt exposed to hypochlorite becomes water-soluble and non-volatile, which reduces its odor and facilitates its removal.
Sodium hypochlorite in solution exhibits broad-spectrum anti-microbial activity and is widely used in healthcare facilities in a variety of settings.[44] It is usually diluted in water depending on its intended use. "Strong chlorine solution" is a 0.5% solution of hypochlorite (containing approximately 5000 ppm free chlorine) used for disinfecting areas contaminated with body fluids, including large blood spills (the area is first cleaned with detergent before being disinfected).[44][45] It may be made by diluting household bleach as appropriate (normally 1 part bleach to 9 parts water).[46] Such solutions have been demonstrated to inactivate bothC. difficile[44] andHPV.[47] "Weak chlorine solution" is a 0.05% solution of hypochlorite used for washing hands, but is normally prepared withcalcium hypochlorite granules.[45]
"Dakin's Solution" is a disinfectant solution containing a low concentration of sodium hypochlorite and someboric acid orsodium bicarbonate to stabilize the pH. It is effective with NaOCl concentrations as low as 0.025%.[48]
US government regulations allow food processing equipment and food contact surfaces to be sanitized with solutions containing bleach, provided that the solution is allowed to drain adequately before contact with food and that the solutions do not exceed 200 parts per million (ppm) available chlorine (for example, one tablespoon of typical household bleach containing 5.25% sodium hypochlorite, per gallon of water).[49] If higher concentrations are used, the surface must be rinsed with potable water after sanitizing.
A similar concentration of bleach in warm water is used to sanitize surfaces before brewing beer or wine. Surfaces must be rinsed with sterilized (boiled) water to avoid imparting flavors to the brew; the chlorinated byproducts of sanitizing surfaces are also harmful. The mode of disinfectant action of sodium hypochlorite is similar to that of hypochlorous acid.
Solutions containing more than 500 ppm available chlorine arecorrosive to somemetals,alloys, and manythermoplastics (such asacetal resin) and need to be thoroughly removed afterward, so the bleach disinfection is sometimes followed by anethanol disinfection. Liquids containing sodium hypochlorite as the main active component are also used for household cleaning and disinfection, for exampletoilet cleaners.[50] Some cleaners areformulated to be viscous so as not to drain quickly from vertical surfaces, such as the inside of a toilet bowl.
The undissociated (nonionized) hypochlorous acid is believed to react with and inactivate bacterial and viral enzymes.
Neutrophils of the human immune system produce small amounts ofhypochlorite insidephagosomes, which digest bacteria and viruses.
Sodium hypochlorite has deodorizing properties, which go hand-in-hand with its cleaning properties.[42]
Sodium hypochlorite solutions have been used to treat dilutecyanide wastewater, such aselectroplating wastes. In batch treatment operations, sodium hypochlorite has been used to treat more concentrated cyanide wastes, such as silver cyanide plating solutions. Toxic cyanide is oxidized tocyanateOCN−) that is not toxic, idealized as follows:
Sodium hypochlorite is commonly used as abiocide in industrial applications to control slime and bacteria formation in water systems used at power plants, pulp and paper mills, etc., in solutions typically of 10–15% by weight.
Sodium hypochlorite is the medicament of choice due to its efficacy against pathogenic organisms and pulp digestion inendodontic therapy. Its concentration for use varies from 0.5% to 5.25%. At low concentrations it dissolves mainly necrotic tissue; at higher concentrations, it also dissolves vital tissue and additional bacterial species. One study has shown thatEnterococcus faecalis was still present in the dentin after 40 minutes of exposure of 1.3% and 2.5% sodium hypochlorite, whereas 40 minutes at a concentration of 5.25% was effective inE. faecalis removal.[51] In addition to higher concentrations of sodium hypochlorite, longer time exposure and warming the solution (60 °C) also increases its effectiveness in removing soft tissue and bacteria within the root canal chamber.[51] 2% is a common concentration as there is less risk of aniatrogenic hypochlorite incident.[52] A hypochlorite incident is an immediate reaction of severe pain, followed byedema,haematoma, andecchymosis as a consequence of the solution escaping the confines of the tooth and entering the periapical space. This may be caused by binding or excessive pressure on the irrigant syringe, or it may occur if the tooth has an unusually largeapical foramen.[53]
At the variousnerve agent (chemical warfare nerve gas)destruction facilities throughout the United States, 0.5-2.5% sodium hypochlorite is used to remove all traces of nerve agent or blister agent from Personal Protection Equipment after an entry is made by personnel into toxic areas.[54]
0.5-2.5% sodium hypochlorite is also used to neutralize any accidental releases of the nerve agent in the toxic areas.[55]
Lesser concentrations of sodium hypochlorite are used similarly in the Pollution Abatement System to ensure that no nerve agent is released into the furnace flue gas.
Dilute bleach baths have been used for decades to treat moderate to severeeczema in humans,.[56][57] Still, it has not been clear why they work. One of the reasons why bleach helps is that eczema can frequently result in secondary infections, especially from bacteria likeStaphylococcus aureus, which makes managing it difficult. Staphylococcus aureus infection is related to the pathogenesis of eczema and AD. Bleach baths are one method for lowering the risk of staph infections in people with eczema. The antibacterial and anti-inflammatory properties of sodium hypochlorite contribute to the reduction of harmful bacteria on the skin and the reduction of inflammation, respectively.[58] According to work published by researchers at theStanford University School of Medicine in November 2013, a very dilute (0.005%) solution of sodium hypochlorite in water was successful in treating skin damage with aninflammatory component caused byradiation therapy, excess sun exposure or aging inlaboratory mice. Mice withradiation dermatitis given daily 30-minute baths in bleach solution experienced less severe skin damage and better healing and hair regrowth than animals bathed in water. Amolecule callednuclear factor kappa-light-chain-enhancer of activated B cells (NF-κB) is known to play a critical role in inflammation, aging, and response to radiation. The researchers found that if NF-κB activity was blocked in elderly mice by bathing them in bleach solution, the animals' skin began to look younger, going from old and fragile to thicker, with increasedcell proliferation. The effect diminished after the baths were stopped, indicating that regular exposure was necessary to maintain skin thickness.[56][59]
Dilute sodium hypochlorite solutions (as in household bleach) are irritating to mainly the skin and respiratory tract. Short-term skin contact with household bleach may cause dryness of the skin.
It is estimated that there are about 3,300 accidents needing hospital treatment caused by sodium hypochlorite solutions each year in British homes (RoSPA, 2002).
Sodium hypochlorite is a strongoxidizer. Oxidation reactions arecorrosive. Solutions burn the skin and cause eye damage, especially when used in concentrated forms. Despite this, sodium hypochlorite solutions are not classified as oxidizers by theNFPA as "sodium hypochlorite solutions do not readily yield oxygen or other oxidizing gases and do not initiate or promote combustion of combustible materials".[60]
Household bleach and pool chlorinator solutions are typically stabilized by a significant concentration oflye (caustic soda, NaOH) as part of the manufacturing reaction. This additive will by itself cause caustic irritation or burns due todefatting andsaponification of skin oils and destruction of tissue. The slippery feel of bleach on the skin is due to this process.
Contact of sodium hypochlorite solutions with metals may evolve flammable hydrogen gas. Containers may explode when heated due to the release of chlorine gas.[12]
Hypochlorite solutions are corrosive to common container materials such asstainless steel[8] andaluminium. The few compatible metals includetitanium (which however is not compatible with dry chlorine) andtantalum.[10] Glass containers are safe.[8] Some plastics and rubbers are affected too; safe choices includepolyethylene (PE),high density polyethylene (HDPE, PE-HD),polypropylene (PP),[8] somechlorinated andfluorinated polymers such aspolyvinyl chloride (PVC),polytetrafluoroethylene (PTFE), andpolyvinylidene fluoride (PVDF); as well asethylene propylene rubber, andViton.[10]
Containers must allow the venting of oxygen produced by decomposition over time, otherwise, they may burst.[5]
Mixing bleach with some household cleaners can be hazardous.
Sodium hypochlorite solutions, such as liquid bleach, will release toxicchlorine gas when mixed with anacid, such ashydrochloric acid orvinegar.
A 2008 study indicated that sodium hypochlorite and organic chemicals (e.g., surfactants, fragrances) contained in several household cleaning products can react to generate chlorinated organic compounds.[61] The study showed that indoor air concentrations significantly increase (8–52 times for chloroform and 1–1170 times for carbon tetrachloride, respectively, above baseline quantities in the household) during the use of bleach containing products.
In particular, mixing hypochlorite bleaches with amines (for example, cleaning products that contain or releaseammonia,ammonium salts,urea, or related compounds and biological materials such asurine) produces chloramines.[62][12] These gaseous products can cause acute lung injury. Chronic exposure, for example, from the air at swimming pools where chlorine is used as the disinfectant, can lead to the development of atopic asthma.[63]
Bleach can react violently withhydrogen peroxide and produce oxygen gas:
Explosive reactions or byproducts can also occur in industrial and laboratory settings when sodium hypochlorite is mixed with diverse organic compounds.[12]
The UK'sNational Institute for Health and Care Excellence in October 2008 recommended thatDakin's solution should not be used in routine wound care.[64]
In spite of its strong biocidal action, sodium hypochloriteper se has limited environmental impact, since the hypochlorite ion rapidly degrades before it can be absorbed by living beings.[65]
However, one major concern arising from sodium hypochlorite use is that it tends to form persistentchlorinated organic compounds, including knowncarcinogens, that can be absorbed by organisms and enter thefood chain. These compounds may be formed during household storage and use as well as during industrial use.[41] For example, when household bleach and wastewater were mixed, 1–2% of the available chlorine was observed to form organic compounds.[41] As of 1994, not all the byproducts had been identified, but identified compounds includechloroform andcarbon tetrachloride.[41][needs update] The exposure to these chemicals from use is estimated to be within occupational exposure limits.[41]
This Support Dossier deals with information on the environmental and human safety evaluation of hypochlorite, and on its benefits as a disinfecting, deodorising, and stain removing agent.
