Stability of the solid Anhydrous<unk>  can be prepared but, like many<unk> s, it is highly unstable and decomposes explosively on heating or friction. The decomposition is accelerated by carbon dioxide at atmospheric levels. It is a white solid with the orthorhombic crystal structure.<unk>  can also be obtained as a crystalline pentahydrate ·5, which is not explosive and is much more stable than the anhydrous compound. The formula is sometimes given as 2·10.. The Cl–O bond length in the pentahydrate is 1.686 Å. The transparent, light greenish-yellow, orthorhombic crystals contain 44%<unk>  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<unk>  dihydrate ·2 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<unk>  into chlorate () and chloride. In one test, the dihydrate was claimed to<unk> ow only 6% decomposition after 13.5 months 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<unk> owed no decomposition after 64 hours at −25 °C. Equilibria and stability of solutions At typical ambient temperatures,<unk>  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. Stoichiometric solutions are fairly alkaline, with pH 11 or higher since hypochlorous acid is a weak acid: : + HOCl + The following species and equilibria are present in solutions of : : (aq) ⇌ + : (aq) + + ⇌ (aq) + : (aq) + ⇌ : (aq) ⇌ (g) The second equilibrium equation above will be<unk> ifted 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<unk>  . The equilibrium can be<unk> ifted by adding acids (such as hydrochloric acid) or bases (such as<unk>  hydroxide) to the solution: : (aq) + 2 HCl (aq) → (g) + (aq) + (aq) : (g) + 2 → (aq) + (aq) + (aq) 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: : + ⇌ HClO<unk>  solutions combined with acid evolve chlorine gas, particularly strongly at pH 8, the chlorine is practically all in the form of<unk>  anions (). The solutions are fairly stable at pH 11–12. Even so, one report claims that a conventional 13.6%<unk>  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<unk>  or trichloroisocyanuric acid . Anhydrous<unk>  is soluble in methanol, and solutions are stable. Decomposition to chlorate or oxygen In solution, under certain conditions, the<unk>  anion may also disproportionate (autoxidize) to chloride and chlorate: :3 + → + 2 In particular, this reaction occurs in<unk>  solutions at high temperatures, forming<unk>  chlorate and<unk>  chloride: :3 (aq) → 2 (aq) + (aq) This reaction is exploited in the industrial production of<unk>  chlorate. An alternative decomposition of<unk>  produces oxygen instead: :2 → 2 + In hot<unk>  solutions, this reaction competes with chlorate formation, yielding<unk>  chloride and oxygen gas: :2 (aq) → 2 (aq) + (g) These two decomposition reactions of solutions are maximized at pH around 6. The chlorate-producing reaction predominates at pH above 6, while the oxygen one becomes significant below that. For example, at 80 °C, with<unk>  and<unk> concentrations of 80 mM, and pH 6–6.5, the chlorate is produced with ∼95% efficiency. The oxygen pathway predominates at pH 10. This decomposition is affected by light and metal ion catalysts such as copper, nickel, cobalt, and iridium. Catalysts like<unk>  dichromate and<unk>  molybdate may be added industrially to reduce the oxygen pathway, but a report claims that only the latter is effective. Titration Titration of<unk>  solutions is often done by adding a measured sample to an excess amount of acidified solution of potassium iodide () and then titrating the liberated iodine () with a standard solution of<unk>  thiosulfate or phenyl arsine oxide, using starch as indicator, until the blue color disappears. According to one US patent, the stability of<unk>  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<unk> ·2, and 139.08 μm for the anhydrous mixed salt . Oxidation of organic compounds Oxidation of starch by<unk> , that adds carbonyl and carboxyl groups, is relevant to the production of modified starch products. In the presence of a phase-transfer catalyst, alcohols are oxidized to the corresponding carbonyl compound (aldehyde or ketone).<unk>  can also oxidize organic sulfides to sulfoxides or sulfones, disulfides or thiols to sulfonyl chlorides or bromides, imines to oxaziridines. It can also de-aromatize phenols. Oxidation of metals and complexes Heterogeneous reactions of<unk>  and metals such as zinc proceed slowly to give the metal oxide or hydroxide:<unk>  + Zn → ZnO +<unk> 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,<unk>  reacts with carbon dioxide to form<unk>  carbonate: :2<unk>  + + → + 2 HOCl<unk>  reacts with most nitrogen compounds to form volatile monochloramine, dichloramines, and nitrogen trichloride: : +<unk>  → + NaOH : +<unk>  → + NaOH : +<unk>  → + NaOH Neutralization<unk>  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.
Sodium hypochlorite
Amylmetacresol
Sodium hydroxide
Povidone-iodine
Tetrachlorodecaoxide
Chloroethane
Denatured alcohol
Silver nitrate