3D model (JSmol)
|Molar mass||74.442 g/mol|
|Appearance||greenish-yellow solid (pentahydrate)|
|Odor||chlorine-like and sweetish|
|Melting point||18 °C (64 °F; 291 K) pentahydrate|
|Boiling point||101 °C (214 °F; 374 K) (decomposes)|
|29.3 g/100mL (0 °C)|
Std enthalpy of
|Safety data sheet||ICSC 1119 (solution, >10% active chlorine)
ICSC 0482 (solution, <10% active chlorine)
Dangerous for the environment (N)
|R-phrases (outdated)||R31, R34, R50|
|S-phrases (outdated)||(S1/2), S28, S45, S50, S61|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
Sodium hypochlorite is a chemical compound with the formula NaClO. It is composed of a sodium cation (Na+
) and a hypochlorite anion (ClO−
); it may also be viewed as the sodium salt of hypochlorous acid. When dissolved in water it is commonly known as bleach or liquid bleach. Sodium hypochlorite is practically and chemically distinct from chlorine. Sodium hypochlorite is frequently used as a disinfectant or a bleaching agent. The mixture of sodium peroxide (Na2O2) and hydrochloric acid, which react to produce sodium hypochlorite is also termed as oxone.
- 1 Uses
- 2 Safety
- 3 Production
- 4 Packaging and sale
- 5 Reactions
- 6 Relationship to free chlorine
- 7 Neutralization
- 8 References
- 9 Bibliography
- 10 External links
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 into sodium chloride and sodium chlorate.
Sodium hypochlorite has destaining properties. Among other applications, it can be used to remove mold stains, dental stains caused by fluorosis, and stains on crockery, especially those caused by the tannins in tea. It has also been used in laundry detergents and as a surface cleaner.
Its bleaching, cleaning, deodorizing and caustic effects are due to oxidation and hydrolysis (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. 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). It may be made by diluting household bleach as appropriate (normally 1 part bleach to 9 parts water). Such solutions have been demonstrated to inactivate both C. difficile and HPV. "Weak chlorine solution" is a 0.05% solution of hypochlorite used for washing hands, but is normally prepared with calcium hypochlorite granules.
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). 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 prior to brewing of 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 are corrosive to some metals, alloys and many thermoplastics (such as acetal resin) and need to be thoroughly removed afterwards, so the bleach disinfection is sometimes followed by an ethanol disinfection. Liquids containing sodium hypochlorite as the main active component are also used for household cleaning and disinfection, for example toilet cleaners. Some cleaners are formulated to be thick so as not to drain quickly from vertical surfaces, such as the inside of a toilet bowl.
Sodium hypochlorite has deodorizing properties, which go hand in hand with its cleaning properties.
Sodium hypochlorite solutions have been used to treat dilute cyanide waste water, such as electroplating 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 to cyanate (OCN−) that is not toxic, idealized as follows:
- CN− + OCl− → OCN− + Cl−
Sodium hypochlorite is commonly used as a biocide 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 in endodontic 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 that Enterococcus 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 in E. faecalis removal. In addition to higher concentrations of sodium hypochlorite, longer time exposure and warming the solution also increases its effectiveness in removing soft tissue and bacteria within the root canal chamber. 2% is a common concentration as there is less risk of an iatrogenic hypochorite incident. A hypochlorite incident is an immediate reaction of severe pain, followed by edema, haematoma, and ecchymosis 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 large apical foramen.
Nerve agent neutralization
At the various nerve agent (chemical warfare nerve gas) destruction facilities throughout the United States, 50% 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. 50% sodium hypochlorite is also used to neutralize any accidental releases of nerve agent in the toxic areas. Lesser concentrations of sodium hypochlorite are used in similar fashion in the Pollution Abatement System to ensure that no nerve agent is released in furnace flue gas.
Reduction of skin damage
Dilute bleach baths have been used for decades to treat moderate to severe eczema in humans, but it has not been clear why they work. According to work published by researchers at the Stanford 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 an inflammatory component caused by radiation therapy, excess sun exposure or aging in laboratory mice. Mice with radiation 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. A molecule called nuclear 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 increased cell proliferation. The effect diminished after the baths were stopped, indicating that regular exposure was necessary to maintain skin thickness.
Sodium hypochlorite is a strong oxidizer. Oxidation reactions are corrosive. Solutions burn the skin and cause eye damage, especially when used in concentrated forms. However, as recognized by the NFPA, only solutions containing more than 40% sodium hypochlorite by weight are considered hazardous oxidizers. Solutions less than 40% are classified as a moderate oxidizing hazard (NFPA 430, 2000).
Mixing bleach with some household cleaners can be hazardous. For example, mixing an acid cleaner with sodium hypochlorite bleach generates toxic chlorine gas. Mixing bleach with amines (for example, cleaning products containing ammonia or related compounds and biological materials such as urine) produces nitrogen trichloride. This gaseous product 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.
Bleach can react violently with hydrogen peroxide and produce oxygen gas:
- H2O2(aq) + NaOCl(aq) → NaCl(aq) + H2O(l) + O2(g)
It is estimated that there are about 3300 accidents needing hospital treatment caused by sodium hypochlorite solutions each year in British homes (RoSPA, 2002).
Household bleach and pool chlorinator solutions are typically stabilized by a significant concentration of lye (caustic soda, NaOH) as part of the manufacturing reaction. Skin contact will produce caustic irritation or burns due to defatting and saponification of skin oils and destruction of tissue. The slippery feel of bleach on skin is due to this process.
A European study, published in 2008, indicated that sodium hypochlorite and organic chemicals (e.g., surfactants, fragrances) contained in several household cleaning products can react to generate chlorinated volatile organic compounds (VOCs). These chlorinated compounds are emitted during cleaning applications, some of which are toxic and probable human carcinogens. 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. The increase in chlorinated volatile organic compound concentrations was the lowest for plain bleach and the highest for the products in the form of "thick liquid and gel." The significant increases observed in indoor air concentrations of several chlorinated VOCs (especially carbon tetrachloride and chloroform) indicate that the bleach use may be a source that could be important in terms of inhalation exposure to these compounds. The authors suggested that using these cleaning products may significantly increase the cancer risk.
One major concern arising from sodium hypochlorite use is that it tends to form chlorinated organic compounds, some of which are carcinogenic. This can occur during household storage and use as well during industrial use. For example, when household bleach and wastewater were mixed, 1–2% of the available chlorine was observed to form organic compounds. As of 1994, not all the byproducts had been identified, but identified compounds include chloroform and carbon tetrachloride. The estimated exposure to these chemicals from use is estimated to be within occupational exposure limits.
Potassium hypochlorite was first produced in 1789 by Claude Louis Berthollet in his laboratory on the Quai de Javel in Paris, France, by passing chlorine gas through a solution of potash 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 cheaper soda lye, thus obtaining sodium hypochlorite (Eau de Labarraque). However, this process was not very efficient, and alternative production methods were sought. One such method involved the extraction of chlorinated lime (known as bleaching powder) with sodium carbonate to yield low levels of available chlorine. 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 – and "Dakin's Solution." The UK's National Institute for Health and Care Excellence in October 2008 recommended the preparation should not be used in routine wound care.
Near the end of the nineteenth century, E. S. Smith patented the chloralkali process: a method of producing sodium hypochlorite involving the electrolysis of brine to produce sodium hydroxide and chlorine gas, which then mixed to form sodium hypochlorite. Both electric power and brine solution 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.
Today, an improved version of this method, known as the Hooker process (named after Hooker Chemicals, acquired by Occidental Petroleum), is the only large scale industrial method of sodium hypochlorite production. In the process, sodium hypochlorite (NaClO) and sodium chloride (NaCl) are formed when chlorine is passed into cold dilute sodium hydroxide solution. It is prepared industrially by electrolysis with minimal separation between the anode and the cathode. The solution must be kept below 40 °C (by cooling coils) to prevent the undesired formation of sodium chlorate.
- Cl2(g) + 2 NaOH(aq) → NaCl(aq) + NaOCl(aq) + H2O(l)
Commercial solutions always contain significant amounts of sodium chloride (common salt) as the main by-product, as seen in the equation above.
Other production routes
Sodium hypochlorite can be easily produced for research purposes by reacting ozone with salt.
- NaCl + O3 → NaClO + O2
This reaction happens at room temperature and can be helpful for oxidizing alcohols.
Packaging and sale
Like many hypochlorites, anhydrous NaClO obtained by desiccation of the pentahydrate will decompose violently on heating or friction. However, it is more stable in cold dilute solutions.
Household bleach 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.
A 12% solution is widely used in waterworks for the chlorination of water, and a 15% solution is more commonly used for disinfection of waste water in treatment plants. Sodium hypochlorite can also be used for point-of-use disinfection of drinking water.
- NaOCl(aq) ⇌ Na⁺(aq) + OCl⁻(aq)
- OCl⁻(aq) + H₂O(l) ⇌ HOCl(aq) + OH⁻(aq)
- NaClO(aq) + 2 HCl(aq) → Cl2(g) + H2O(L) + NaCl(aq)
- NaClO + CH3COOH → HClO + CH3COONa
- 3 NaOCl(aq) → NaClO3(aq) + 2 NaCl(aq)
Sodium hypochlorite decomposes with increasing temperature and under the influence of light and such metals as copper, nickel, or cobalt:
- 2 OCl⁻(aq) → 2 Cl⁻(aq) + O₂(g)
- 2 NaOCl(aq) + CO₂(g) → Na₂CO₃(aq) + Cl₂(g)
Depending on the pH of solution the hypochlorous acid dissociates to form hypochlorite ion and hydrogen ion.
- HOCl ⇌ H+ + OCl−
The undissociated (nonionized) hypochlorous acid is believed to react with and inactivate bacterial and viral enzymes.
In NaClO solutions, the following species are thought to be present when the system is in equilibrium.
- HOCl ⇌ H+ + OCl−
- HOCl + Cl− + H+ ⇌ Cl2 + H2O
The ratio Cl2 : HOCl : OCl− is pH dependent. The amount of undissociated (nonionized) HOCl is highest at a pH of about 4. At pH < 2, chlorine gas (Cl₂) evolves. At pH > 7, only hypochlorite anions (OCl⁻) remain.
- NH3 + NaClO → NH2Cl + NaOH
- NH2Cl+ NaClO → NHCl2 + NaOH
- NHCl2 + NaClO → NCl3 + NaOH
- NaClO + Zn → ZnO + NaCl
Relationship to free chlorine
When used as a swimming pool disinfectant or for water treatment, the efficacy of the treated water is usually measured in the mass concentration of "free chlorine" or "available chlorine". The usual units are mg/L or, equivalently, ppm (parts per million). The unit of gpl (grams per liter) is also used. The NaClO content of household bleach is generally specified as a mass or weight percent.
The "free chlorine" is a measure of the amount of chlorine gas (Cl2) that would yield the same oxidizing power as the NaClO in solution. As one molecule of NaClO has the same oxidizing power as one molecule of Cl2, for a given mass of NaClO, the mass of "free chlorine" is 1/1.05 times that quantity (the ratio of the molecular weights of the two compounds).
To determine the mass concentration of NaClO or free chlorine in a diluted solution from the NaClO concentration, the density of the solutions must be considered. Since NaClO in solution is hydrophilic, water molecules are attracted to the hypochlorite ion, and the density of the solution is rather higher than might be expected from a simple calculation. Volumes are thus not strictly additive; the volume of a liter of NaClO solution and an equal volume of water will somewhat less than twice the original volume. The density of a solution (ρ(w)) will be a function of the mass fraction of NaClO (w) and tables are given in the OxyChem Handbook.
The masses of NaClO and water will be constant before and after dilution. If wb, ρ(wb) are the mass fraction and density of NaClO in the undiluted solution, and Vb its volume, and w and ρ(w) and V are likewise for the final solution, the masses of NaClO (Mb) and water (Mw) before and after dilution are conserved:
where ρw and Vw are the density and volume of the added water. Solving the above equations for w and V, the final NaClO mass fraction will be:
and the final NaClO mass concentration will be:
Dividing this by 1.05 will yield the mass concentration of free chlorine. For example, one ml of 5.25 wt% NaClO bleach added to ten liters of water, will yield a NaClO concentration of about 5.76 mg/L NaClO and 5.48 mg/L of free chlorine. Sodium hydroxide (NaOH or lye) is usually added in small amounts to household bleach to slow down the decomposition of NaClO. This will require a small correction to the above calculations.
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