|Molar mass||171.34 g/mol (anhydrous)
189.39 g/mol (monohydrate)
315.46 g/mol (octahydrate)
|Density||3.743 g/cm3 (monohydrate)
2.18 g/cm3 (octahydrate, 16 °C)
|Melting point||78 °C (172 °F; 351 K) (octahydrate)
300 °C (monohydrate)
407 °C (anhydrous)
|Boiling point||780 °C (1,440 °F; 1,050 K)|
|mass of BaO (not Ba(OH)2):
1.67 g/100 mL (0 °C)
3.89 g/100 mL (20 °C)
4.68 g/100 mL (25 °C)
5.59 g/100 mL (30 °C)
8.22 g/100 mL (40 °C)
11.7 g/100 mL (50 °C)
20.94 g/100 mL (60 °C)
101.4 g/100 mL (100 °C)
|Solubility in other solvents||low|
Refractive index (nD)
Std enthalpy of
|EU classification||Harmful (Xn)|
|Supplementary data page|
|Refractive index (n),
Dielectric constant (εr), etc.
|UV, IR, NMR, MS|
Except where noted otherwise, data is given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
|what is: / ?)(|
Barium hydroxide are chemical compounds with the formulae Ba(OH)2(H2O)x. The monohydrate (x =1) is known as baryta, or baryta-water, it is one of the principal compounds of barium. This white granular monohydrate is the usual commercial form.
Preparation and structure
Barium hydroxide can be prepared by dissolving barium oxide (BaO) in water:
- BaO + 9 H2O → Ba(OH)2·8H2O
It crystallises as the octahydrate, which converts to the monohydrate upon heating in air. At 100 °C in a vacuum, the monohydrate gives BaO. The monohydrate adopts a layered structure (see picture above). The Ba2+ centers adopt a square anti-prismatic geometry. Each Ba2+ center is bound by two water ligands and six hydroxide ligands, which are respectively doubly and triply bridging to neighboring Ba2+ centers. sites. In the octahydrate, the individual Ba2+ centers are again eight coordinate but do not share ligands.
Industrially, barium hydroxide is used as the precursor to other barium compounds. The monohydrate is used to dehyrate and remove sulfate from various products. This application exploits the very low solubility of barium sulfate. This industrial application is also applied to laboratory uses.
Barium hydroxide is used in analytical chemistry for the titration of weak acids, particularly organic acids. Its clear aqueous solution is guaranteed to be free of carbonate, unlike those of sodium hydroxide and potassium hydroxide, as barium carbonate is insoluble in water. This allows the use of indicators such as phenolphthalein or thymolphthalein (with alkaline colour changes) without the risk of titration errors due to the presence of carbonate ions, which are much less basic.
It has been used to hydrolyse one of the two equivalent ester groups in dimethyl hendecanedioate.
Barium hydroxide is used, as well, in the decarboxylation of amino acids liberating barium carbonate in the process. 
Barium hydroxide decomposes to barium oxide when heated to 800 °C. Reaction with carbon dioxide gives barium carbonate. Its aqueous solution, being highly alkaline, undergoes neutralization reactions with acids. Thus, it forms barium sulfate and barium phosphate with sulfuric and phosphoric acids, respectively. Reaction with hydrogen sulfide produces barium sulfide. Precipitation of many insoluble, or less soluble barium salts, may result from double replacement reaction when a barium hydroxide aqueous solution is mixed with many solutions of other metal salts. 
Reactions of barium hydroxide with ammonium salts are strongly endothermic. The reaction of barium hydroxide octahydrate with ammonium chloride  or ammonium thiocyanate is often used as a classroom chemistry demonstration, producing temperatures cold enough to freeze water and enough water to dissolve the resulting mixture.
Barium hydroxide presents the same hazards as other strong bases and as other water-soluble barium compounds: it is corrosive and toxic.
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