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.

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.

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: