Sodium hypochlorite is an alkaline inorganic chemical compound with the formula NaOCl (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 (Na+) and hypochlorite anions (−OCl, also written as OCl− and ClO−). The anhydrous compound is unstable and may decompose explosively. It can be crystallized as a pentahydrate NaOCl·5H2O, 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 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 NaOCl·5H2O, which is not explosive and is much more stable than the anhydrous compound. The formula is sometimes given in its hydrous crystalline form as 2NaOCl·10H2O. 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 NaOCl·2H2O can be obtained by carefully excluding chloride ions (Cl−), which are present in the output of common manufacturing processes and are said to catalyze the decomposition of hypochlorite into chlorate (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.
Equilibria and stability of solutions At typical ambient temperatures, sodium hypochlorite is more stable in dilute solutions that contain solvated Na+ and OCl− 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 the hypochlorite ion is a weak base:
OCl− + H2O ⇌ HOCl + OH− The following species and equilibria are present in NaOCl/NaCl solutions:
HOCl(aq) ⇌ H+ + OCl− HOCl(aq) + Cl− + H+ ⇌ Cl2(aq) + H2O Cl2(aq) + Cl− ⇌ Cl−3 Cl2(aq) ⇌ Cl2(g) The second equilibrium equation above will be shifted to the right if the chlorine Cl2 is allowed to escape as gas. The ratios of Cl2, HOCl, and OCl− 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 Cl2. At pH greater than 7.4, the majority is in the form of hypochlorite ClO−. The equilibrium can be shifted by adding acids (such as hydrochloric acid) or bases (such as sodium hydroxide) to the solution:
ClO−(aq) + 2 HCl(aq) → Cl2(g) + H2O + Cl−(aq) Cl2(g) + 2 OH−(aq) → ClO−(aq) + Cl−(aq) + H2O 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:
−OCl + H+ ⇌ HOCl Sodium hypochlorite solutions combined with acid evolve chlorine gas, particularly strongly at pH < 2, by the reactions:
HOCl(aq) + Cl− + H+ ⇌ Cl2(aq) + H2O Cl2(aq) ⇌ Cl2(g) 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. For this reason, in some applications one may use more stable chlorine-releasing compounds, such as calcium hypochlorite Ca(ClO)2 or trichloroisocyanuric acid (CNClO)3. 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:
3 ClO− + H+ → HClO3 + 2 Cl− In particular, this reaction occurs in sodium hypochlorite solutions at high temperatures, forming sodium chlorate and sodium chloride:
3 NaOCl(aq) → 2 NaCl(aq) + NaClO3(aq) This reaction is exploited in the industrial production of sodium chlorate. An alternative decomposition of hypochlorite produces oxygen instead:
2 OCl− → 2 Cl− + O2 In hot sodium hypochlorite solutions, this reaction competes with chlorate formation, yielding sodium chloride and oxygen gas:
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