Reston Glass

Sodium sulfate
History
Hydrate of sodium sulfate is known as Glauber's salt after the Dutch / German chemist Johann Rudolf Glauber and apothecary (16041670), who discovered that 1625 in Austrian water source. He was named sal mirabilis (miraculous salt), because of its medicinal properties: the crystals were used as a laxative for general use until more sophisticated solutions have emerged in the 1900s.
In the 18th century, Glauber's salt has begun to be used as raw material for industrial production of soda ash (sodium carbonate) by reaction with potash (potassium carbonate). The demand for soda ash increased and the supply of sodium sulfate has increased apace. Therefore, in the nineteenth century the Leblanc process, sodium sulfate synthesized product as a key intermediate, has become the main method of producing sodium carbonate.
Physical and chemical
Sodium sulfate is chemically very stable, being responsive to most oxidizing agents or reducing to normal temperatures. At high temperature, it can be reduced by sodium sulfide. It is a neutral salt, which forms aqueous solutions with a pH of 7. The neutrality of such solutions reflects the fact that Na2SO4 is derived, formally speaking, the strong acid sulfuric acid and sodium hydroxide solid base. Sodium sulfate reacts with an equivalent amount of sulfuric acid to obtain an equilibrium concentration of sodium bisulfate salt of the acid:
Na2SO4 (aq) + H2SO4 (aq) 2 NaHSO4 (aq)
In fact, the balance is very complex, depending on concentration and temperature, with the presence of other acid salts.
Sodium sulfate is an ionic sulfate, containing Na + and SO42 ions. Aqueous solutions can produce precipitates when combined with salts of Ba2 + or Pb2 +, which form insoluble sulphates
Na2SO4 (aq) + BaCl2 (aq) 2 NaCl (aq) + BaSO4 (s)
Sodium sulfate has unusual solubility characteristics in water. Its solubility increases more than tenfold between 0 ° C to 32.384 C, where it reaches a maximum of 49.7 g Na2SO4 per 100 g of water. At this point, the solubility curve changes slope, and the solubility becomes almost independent of temperature. The temperature at 32.384 C, corresponding to the output of water and melting of the hydrated salt crystals, used as reference for precise measurement of the temperature of the thermometer.
Sodium sulfate decahydrate is also rare among hydrated salts have a measurable residual entropy (entropy at absolute zero) of 6.32 JK-1mol-1. This is attributed to its ability to distribute water much more rapidly compared to most hydrates.
sodium sulfate shows a moderate trend to form double salts. Alums only trained with base metals are trivalent Naal (SO4) 2 (unstable above 39 C) and nacre (SO4) 2, in contrast to potassium sulfate and ammonium sulfate, which forms many alums stable. Double salts with some other alkalis are known, including Na2SO43K2SO4 naturally occurring mineral that glaserite. Training glaserite by reaction of sodium sulfate with potassium chloride was used as the basis of a method of producing potassium sulfate, a fertilizer. Other double salts include 3Na2SO4CaSO4, 3Na2SO4MgSO4 (vanthoffite) and NaFNa2SO4.
Production
World production of sodium sulfate, mainly in the form of decahydrate amounts to about 5.5 to 6 million tons per year (Mt / a). In 1985, production was 4.5 Mt / a, half from natural sources, and half of chemical production. After 2000, a steady level until 2006, natural production was increased to 4 Mt / y, and chemical production fell from 1.5 to 2 Mt / y, with a total of 5.5 to 6 Mt / a. For all applications, naturally produced and chemically produced sodium sulfate are practically interchangeable.
Natural sources
Two thirds of production in the world of the decahydrate (Glauber's salt) is the natural mineral form mirabilite, for example.