Sodium hypochlorite is an alkaline inorganic chemical compound with the formula (also written as NaClO). It is commonly known in a dilute aqueous solution as bleach or chlorine bleach. It is the sodium salt of hypochlorous acid, consisting of sodium cations () and hypochlorite anions (, also written as and ).
The anhydrous compound is unstable and may decompose explosively. It can be crystallized as a pentahydrate , a pale greenish-yellow solid which is not explosive and is stable if kept refrigerated.
Sodium hypochlorite is most often encountered as a pale greenish-yellow dilute solution referred to as chlorine bleach, which is a household chemical widely used (since the 18th century) as a disinfectant and bleaching agent. In solution, the compound is unstable and easily decomposes, liberating chlorine, which is the active principle of such products. Sodium hypochlorite is still the most important chlorine-based bleach.
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 in limescale-removing products, will release toxic chlorine gas. A common misconception is that mixing bleach with ammonia also releases chlorine, but in reality they react to produce still dangerously toxic chloramines such as nitrogen trichloride. With excess ammonia and sodium hydroxide, hydrazine may be generated.
Chemistry
Stability of the solid Anhydrous sodium hypochlorite can be prepared but, like many hypochlorites, it is highly unstable and decomposes explosively on heating or friction. The decomposition is accelerated by carbon dioxide at Earth's atmospheric levels - around 4 parts per ten thousand. It is a white solid with the orthorhombic crystal structure.
Sodium hypochlorite can also be obtained as a crystalline pentahydrate , which is not explosive and is much more stable than the anhydrous compound. The formula is sometimes given in its hydrous crystalline form as . The Cl–O bond length in the pentahydrate is 1.686 Å. The transparent, light greenish-yellow, orthorhombic" 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.
A 1966 US patent claims that stable solid sodium hypochlorite dihydrate can be obtained by carefully excluding chloride ions (), which are present in the output of common manufacturing processes and are said to catalyze the decomposition of hypochlorite into chlorate () 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.
Equilibria and stability of solutions At typical ambient temperatures, sodium hypochlorite is more stable in dilute solutions that contain solvated and ions. The density of the solution is 1.093 g/mL at 5% concentration, and 1.21 g/mL at 14%, 20 °C.Environment Canada (1985): "Tech Info for Problem Spills: Sodium Hypochlorite (Draft)". Stoichiometric solutions are fairly alkaline, with pH 11 or higher since the hypochlorite ion is a weak base:
The following species and equilibria are present in NaOCl/NaCl solutions:
The second equilibrium equation above will be shifted to the right if the chlorine is allowed to escape as gas. The ratios of , HOCl, and in solution are also pH dependent. At pH below 2, the majority of the chlorine in the solution is in the form of dissolved elemental . At pH greater than 7.4, the majority is in the form of hypochlorite . The equilibrium can be shifted by adding acids (such as hydrochloric acid) or bases (such as sodium hydroxide) to the solution:
At a pH of about 4, such as obtained by the addition of strong acids like hydrochloric 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 (). 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. For this reason, in some applications one may use more stable chlorine-releasing compounds, such as calcium hypochlorite or trichloroisocyanuric acid .
Anhydrous sodium hypochlorite is soluble in methanol, and solutions are stable.
Decomposition to chlorate or oxygen In solution, under certain conditions, the hypochlorite anion may also disproportionate (autoxidize) to chloride and chlorate:
In particular, this reaction occurs in sodium hypochlorite solutions at high temperatures, forming sodium chlorate and sodium chloride:
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:
These two decomposition reactions of solutions are maximized at pH around 6. For example, at 80 °C, with and concentrations of 80 mM, 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. Above pH 11, both reactions have rate proportional to , decomposition is much slower, and chlorate is produced with ~90% efficiency. This decomposition is affected by light and metal ion catalysts such as copper, nickel, cobalt, and iridium. Catalysts like sodium dichromate and sodium molybdate may be added industrially to reduce the oxygen pathway, but a report claims that only the latter is effective.
Titration Titration of hypochlorite solutions is often done by adding a measured sample to an excess amount of acidified solution of potassium iodide (KI) and then titrating the liberated iodine () with a standard solution of sodium thiosulfate or phenylarsine oxide, using starch as indicator, until the blue color disappears.
According to one US patent, the stability of sodium hypochlorite content of solids or solutions can be determined by monitoring the infrared absorption due to the O–Cl bond. The characteristic wavelength is given as 140.25 μm for water solutions, 140.05 μm for the solid dihydrate , and 139.08 μm for the anhydrous mixed salt .
Oxidation of organic compounds Oxidation of starch by sodium hypochlorite, which adds carbonyl and carboxyl groups, is relevant to the production of modified starch products.ASC – PT Asahimas Chemical (2009): "Sodium hypochlorite". Online product description. Accessed on 2018-06-14.
In the presence of a phase-transfer catalyst, alcohols are oxidized to the corresponding carbonyl compound (aldehyde or ketone). Sodium hypochlorite can also oxidize organic sulfides to sulfoxides or sulfones; disulfides or thiols to sulfonyl halides; and imines to oxaziridines. It can also de-aromatize phenols.
Oxidation of metals and complexes Heterogeneous reactions of sodium hypochlorite and metals such as zinc proceed slowly to give the metal oxide or hydroxide: NaOCl + Zn → ZnO + NaCl
Homogeneous reactions with metal coordination complexes proceed somewhat faster. This has been exploited in the Jacobsen epoxidation.
Other reactions If not properly stored in airtight containers, sodium hypochlorite reacts with carbon dioxide to form sodium carbonate:
Sodium hypochlorite reacts with most nitrogen compounds to form volatile monochloramines, dichloramines, and nitrogen trichloride:
Neutralization 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.
Production
Chlorination of soda 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).
Hence, chlorine is simultaneously reduced and oxidized; this process is known as disproportionation.
The process is also used to prepare the pentahydrate 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.
From calcium hypochlorite Another method involved the reaction of sodium carbonate ("washing soda") with chlorinated lime ("bleaching powder"), a mixture of calcium hypochlorite , calcium chloride , and calcium hydroxide :
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.
Electrolysis of brine 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. The key reactions are: (at the anode) (at the cathode) 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.
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 a cold dilute sodium hydroxide solution. The chlorine 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.
Commercial solutions always contain significant amounts of sodium chloride (common salt) as the main by-product, as seen in the equation above.
From hypochlorous acid and soda A 1966 patent describes the production of solid stable dihydrate 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 1–2 mmHg until the dihydrate crystallizes out. The crystals are vacuum-dried to produce a free-flowing crystalline powder.
The same principle was used in a 1993 patent to produce concentrated slurries of the pentahydrate . 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.
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