Sodium hypochlorite

• 1. Introduction

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  Sodium hypochlorite, written NaOCl, is the compound behind the pale greenish-yellow solution most people know as chlorine bleach. In British homes alone, an estimated 3,300 accidents a year need hospital treatment because of solutions like it, according to a 2002 figure from RoSPA. Yet the same chemical is used to disinfect hospital areas, treat water and even bathe eczema-prone skin. How can one substance be so useful and so hazardous? Why does its pure form explode while a diluted solution can be stored at all? And who first worked out how to make it? The answers begin in a Paris laboratory in the eighteenth century.

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• 2. Javel Water

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  In 1789, Claude Louis Berthollet was working in his laboratory on the Quai de Javel in Paris. He passed chlorine gas through a solution of potash lye. The liquid that resulted was a weak solution of potassium hypochlorite. It became known as "Eau de Javel", or Javel water, after the street where it was made.

  Antoine Labarraque then replaced the potash lye with cheaper soda lye. That swap produced sodium hypochlorite instead, and the new liquid was called Eau de Labarraque.

  Eusol was a hospital antiseptic sold after World War I. Its name abbreviates Edinburgh University Solution Of (chlorinated) Lime, a reference to the university's pathology department, where it was developed. It was made by reacting sodium carbonate, known as washing soda, with chlorinated lime. Chlorinated lime, also called bleaching powder, is a mixture of calcium hypochlorite, calcium chloride and calcium hydroxide.

  Near the end of the nineteenth century, E. S. Smith patented the chloralkali process. It used the electrolysis of brine to make sodium hydroxide and chlorine gas, which then mixed to form sodium hypochlorite. Electric power and brine were both cheap at the time. Enterprising marketers took advantage and sold bottled solutions under numerous trade names. Smith's method would eventually be improved into the only process used at large industrial scale today.

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• 3. Making It Today

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  The Hooker process is named for Hooker Chemicals, a company later acquired by Occidental Petroleum. It is the only large-scale industrial method of making sodium hypochlorite. Chlorine is passed into a cold, dilute solution of sodium hydroxide, forming sodium hypochlorite and sodium chloride. The chlorine itself comes from electrolysis, with minimal separation between the anode and the cathode.

  Cooling coils keep the solution below 40 °C, because warmer conditions produce unwanted sodium chlorate. Commercial solutions therefore always carry a significant amount of common salt, the main by-product. In the reaction, the chlorine is simultaneously reduced and oxidized, a process known as disproportionation.

  A 45-48% sodium hydroxide solution is the starting point for the crystalline pentahydrate used in industry and laboratories. Chlorine gas is added, and some of the sodium chloride precipitates and is filtered out. Cooling the filtrate to 12 °C then yields the pentahydrate.

  A 1966 patent describes another route, to a solid called the dihydrate. A chloride-free solution of hypochlorous acid is reacted with concentrated sodium hydroxide. In a typical preparation, 255 mL of a solution with 118 g/L of the acid is slowly stirred into a solution of 40 g of sodium hydroxide in water at 0 °C. After filtering out some sodium chloride, the solution is evaporated under vacuum at 40-50 °C until crystals form. Vacuum drying leaves a free-flowing crystalline powder.

  A 1993 patent applied the same principle to make slurries of the pentahydrate. A 35% solution of hypochlorous acid by weight is combined with sodium hydroxide at about or below 25 °C. The slurry holds about 35% NaClO, and it is relatively stable because chloride is scarce. Stability is a rare quality for this chemical.

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• 4. Fragile Solids

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  Anhydrous sodium hypochlorite is a white solid with an orthorhombic crystal structure. It decomposes explosively on heating or friction. Carbon dioxide at the roughly 4 parts per ten thousand found in Earth's atmosphere speeds that decay.

  The crystalline pentahydrate, NaOCl*5H2O, is the tamer relative. It is not explosive and is much more stable than the anhydrous form. Its transparent, light greenish-yellow crystals contain 44% NaOCl by weight and melt at 25-27 °C. The Cl-O bond length in the crystal is 1.686 Å. At room temperature the compound breaks down rapidly, so it must be refrigerated. Kept at 7 °C, it reportedly lost only 1% after 360 days.

  The 1966 US patent claims a still more stable solid, the dihydrate NaOCl*2H2O. The trick is to exclude chloride ions. They come out of common manufacturing processes and are said to catalyze the decay into chlorate and chloride. In one test, the dihydrate showed only 6% decomposition after 13.5 months at -25 °C. The patent also says that vacuum drying at about 50 °C turns it into the anhydrous form. That solid showed no decomposition after 64 hours at -25 °C.

  Measuring how much hypochlorite a sample really holds is its own craft. Chemists often add a measured sample to an excess of acidified potassium iodide. They then titrate the iodine that is released with a standard solution of sodium thiosulfate or phenylarsine oxide. Starch serves as the indicator, and the endpoint is the moment the blue color disappears.

  Another US patent proposes tracking stability through the infrared absorption of the O-Cl bond. The characteristic wavelength is 140.25 μm for water solutions, 140.05 μm for the solid dihydrate and 139.08 μm for the anhydrous mixed salt Na2(OCl)(OH).

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• 5. Chlorine and pH

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  At 20 °C, a 5% solution of sodium hypochlorite has a density of 1.093 g/mL, and a 14% solution has a density of 1.21 g/mL. At ordinary temperatures the compound is more stable in dilute solution, where it exists as dissolved sodium ions and hypochlorite ions. The hypochlorite ion is a weak base, so plain solutions are fairly alkaline, with a pH of 11 or higher.

  At pH 2 and below, most of the chlorine in the solution is dissolved elemental chlorine. Around pH 4, the amount of undissociated hypochlorous acid is at its highest. Above pH 7.4 most of the chlorine is hypochlorite, and above pH 8 it is practically all hypochlorite. That is why adding an acid such as hydrochloric acid releases chlorine gas, especially strongly below pH 2. Adding a base such as sodium hydroxide pushes the balance back the other way.

  Even at pH 11-12, where solutions are fairly stable, they slowly fade. One report found that a conventional 13.6% reagent solution lost 17% of its strength after 360 days at 7 °C. Where that is not good enough, some applications use calcium hypochlorite or trichloroisocyanuric acid instead. Anhydrous sodium hypochlorite also dissolves in methanol, giving stable solutions.

  At high temperatures, the hypochlorite ion can disproportionate into chloride and chlorate, giving sodium chloride and sodium chlorate. Industry exploits this reaction to make sodium chlorate. A rival pathway yields sodium chloride and oxygen gas instead. Both reactions are maximized around pH 6.

  In one example at 80 °C, with 80 mM of both NaOCl and NaCl, decomposition was fastest at pH 6.5. Chlorate formed with about 95% efficiency. Above pH 11, decomposition is much slower, and chlorate efficiency is near 90%. Light and metal ions such as copper, nickel, cobalt and iridium also affect the rate. Industry can add sodium dichromate or sodium molybdate to reduce the oxygen pathway, though one report says only the molybdate is effective.

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• 6. Disinfection and Medicine

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  A 0.5% hypochlorite solution, holding approximately 5000 ppm of free chlorine, is called strong chlorine solution in healthcare settings. It is used on areas contaminated with body fluids, including large blood spills, after the area is first cleaned with detergent. It is normally made by mixing 1 part household bleach with 9 parts water. Such solutions have been shown to inactivate both C. difficile and HPV.

  A weak chlorine solution of 0.05% is used for washing hands, though it is normally prepared with calcium hypochlorite granules. Dakin's Solution holds an even lower concentration of sodium hypochlorite, along with some boric acid or sodium bicarbonate to stabilize the pH. It works with NaOCl concentrations as low as 0.025%. In October 2008, the UK's National Institute for Health and Care Excellence recommended that it not be used in routine wound care.

  Neutrophils in the human immune system run a small version of this chemistry. They produce small amounts of hypochlorite inside phagosomes, where it helps digest bacteria and viruses. In disinfectant solutions, undissociated hypochlorous acid is believed to react with and inactivate bacterial and viral enzymes.

  In root canal treatment, concentrations from 0.5% to 5.25% are used because the compound fights pathogens and digests pulp. At low concentrations it dissolves mainly dead tissue. At higher concentrations it also dissolves living tissue and more species of bacteria. In one study, Enterococcus faecalis was still present in the dentin after 40 minutes at 1.3% and 2.5%. Forty minutes at 5.25% removed it. Longer exposure and warming the solution to 60 °C also improve results.

  A concentration of 2% is common because it lowers the risk of a hypochlorite incident. In that event the solution escapes the tooth into the tissue around the root tip. The result is immediate severe pain, followed by swelling, a collection of blood and bruising. It may follow excessive pressure on the syringe or an unusually large opening at the root tip.

  Dilute bleach baths have been used for decades to treat moderate to severe eczema, though why they work was long unclear. Eczema often leads to secondary infections, especially from Staphylococcus aureus, and bleach baths lower the risk of staph infection. Researchers at the Stanford University School of Medicine reported in November 2013 that a very dilute 0.005% solution treated inflammatory skin damage in laboratory mice. The damage came from radiation therapy, excess sun exposure or aging.

  Mice with radiation dermatitis were given daily 30-minute baths in bleach solution. They had less severe skin damage, better healing and more hair regrowth than animals bathed in water. In elderly mice, blocking the inflammation regulator NF-κB with the baths made old, fragile skin look thicker and younger. The effect faded after the baths stopped, so regular exposure was needed to keep the skin thick.

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• 7. Cleaning and Industry

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  Household bleach for laundering is a 3-8% solution of sodium hypochlorite at the time of manufacture. Its strength varies between formulations and gradually falls with long storage. Small amounts of sodium hydroxide, 0.01-0.05% by weight, are added to slow the decay into sodium chloride and sodium chlorate.

  Tannins in tea leave stains on crockery, and bleach removes them. It also lifts mold stains and dental stains caused by fluorosis. Its bleaching, cleaning, deodorizing and caustic effects come from oxidation and hydrolysis, known as saponification. Organic dirt that meets hypochlorite becomes water-soluble and non-volatile. That cuts its odor and makes it easier to rinse away.

  Patio blackspot removers for home use are about 10% solutions of sodium hypochlorite. Waterworks widely use a 12% solution to chlorinate water, while wastewater plants more commonly use 15% for disinfection. For point-of-use disinfection of drinking water, 0.2-2 mg of sodium hypochlorite per liter is enough. Disinfecting sprays and wipes for hard surfaces run from 50 ppm to 1.5%. Some toilet cleaners are made viscous so they do not drain quickly from vertical surfaces such as the inside of a toilet bowl.

  In the United States, regulations allow food processing equipment and food contact surfaces to be sanitized with bleach solutions of no more than 200 ppm available chlorine. That equals one tablespoon of typical 5.25% household bleach per gallon of water. The solution must drain adequately before food touches the surface. Higher concentrations require a rinse with potable water afterward.

  A similar concentration in warm water sanitizes surfaces before brewing beer or wine. These surfaces must be rinsed with boiled water to avoid imparting flavors to the brew. The chlorinated byproducts of sanitizing are also harmful.

  Solutions above 500 ppm of available chlorine are corrosive to some metals, alloys and many thermoplastics, such as acetal resin. They need to be thoroughly removed afterward. For that reason, bleach disinfection is sometimes followed by an ethanol disinfection.

  Electroplating wastes contain dilute cyanide wastewater, and sodium hypochlorite has been used to treat it. In batch operations it has also handled more concentrated waste, such as silver cyanide plating solutions. The toxic cyanide is oxidized to cyanate, which is not toxic. Power plants and pulp and paper mills use 10-15% solutions by weight as a biocide to control slime and bacteria in their water systems.

  Oxidation of starch by sodium hypochlorite adds carbonyl and carboxyl groups, which is relevant to making modified starch products. With a phase-transfer catalyst, it oxidizes alcohols to aldehydes or ketones. It can also turn organic sulfides into sulfoxides or sulfones, and it can de-aromatize phenols. Its reactions with metal coordination complexes are exploited in the Jacobsen epoxidation.

  At nerve agent destruction facilities throughout the United States, 0.5-2.5% sodium hypochlorite removes all traces of nerve or blister agent from protective equipment. This happens after personnel enter toxic areas. The same strength neutralizes any accidental release of nerve agent inside those areas. Lesser concentrations in the Pollution Abatement System help ensure that no nerve agent reaches the furnace flue gas.

  Sodium thiosulfate is an effective chlorine neutralizer. Rinsing with a 5 mg/L solution, followed by washing with soap and water, removes chlorine odor from the hands.

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• 8. Hazards and Footprint

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  Vinegar and hydrochloric acid share a dangerous trait. Each releases toxic chlorine gas when mixed with liquid bleach, as do the acids found in limescale-removing products.

  A common misconception names ammonia as a source of chlorine gas. In reality, bleach and ammonia react to produce chloramines, such as nitrogen trichloride. With excess ammonia and sodium hydroxide, hydrazine may be generated. Other amines behave similarly, including ammonium salts, urea and biological materials such as urine. These gaseous products can cause acute lung injury. Chronic exposure, for example from the air at swimming pools where chlorine is the disinfectant, can lead to atopic asthma.

  A 2008 study found that sodium hypochlorite reacts with organic chemicals, such as surfactants and fragrances, in several household cleaning products. The reaction generates chlorinated organic compounds. While these products were in use, indoor air concentrations rose 8-52 times above baseline for chloroform and 1-1170 times for carbon tetrachloride.

  Hydrogen peroxide is another poor partner. Bleach can react violently with it 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.

  The slippery feel of bleach on skin comes from the lye used to stabilize it. This caustic soda defats and saponifies skin oils and destroys tissue, so it can irritate or burn on its own. Dilute household bleach mainly irritates the skin and respiratory tract, and short contact can dry the skin. Concentrated solutions burn skin and cause eye damage. Although it is a strong oxidizer, the NFPA does not classify the solutions as oxidizers. They do not readily yield oxygen or other oxidizing gases, and they do not promote combustion.

  Stainless steel and aluminium are corroded by hypochlorite solutions, and contact with metals may release flammable hydrogen gas. Titanium and tantalum are among the few compatible metals, though titanium is not compatible with dry chlorine. Glass is safe, as are polyethylene, polypropylene, PVC, PTFE, PVDF, ethylene propylene rubber and Viton. Containers must vent the oxygen produced by slow decomposition, or they may burst. Heated containers may explode from the release of chlorine gas.

  The hypochlorite ion itself has limited environmental impact, since it degrades rapidly before living beings can absorb it. The concern is persistent chlorinated organic compounds, including known carcinogens, that can enter the food chain. When household bleach and wastewater were mixed, 1-2% of the available chlorine was observed forming organic compounds. Chloroform and carbon tetrachloride are among those identified, and exposure from use is estimated to fall within occupational exposure limits. As of 1994, not all of the byproducts had been identified.

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