Sodium sulfate

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Sodium sulfate
Sodium sulfate.jpg
Other names Salt cake
Thenardite (mineral)
Glauber's salt (decahydrate)
Sal mirabilis (decahydrate)
Mirabilite (decahydrate)
CAS number 7757-82-6
RTECS number WE1650000 (anhydrous)
Molar mass 142.04 g/mol (anhydrous)
268.15 g/mol (heptahydrate)
322.20 g/mol (decahydrate)
Appearance White crystalline solid,
Density 2.68 g/cm³, anhydrous
(orthorhombic form)
1.464 g/cm³, decahydrate
Melting point

884 °C (1157 K) anhydrous
32.4 °C decahydrate

Solubility in water 4.76 g/100 ml (0 °C)
42.7 g/100 ml (100 °C)
Crystal structure monoclinic, orthorhombic or
MSDS External MSDS
Main hazards Irritant
NFPA 704

NFPA 704.svg

R/S statement None
Related Compounds
Other anions Sodium bisulfate
Sodium sulfite
Sodium bisulfite
Sodium persulfate
Other cations Lithium sulfate
Potassium sulfate
Magnesium sulfate
Except where noted otherwise, data are given for
materials in their standard state
(at 25 °C, 100 kPa)

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Sodium sulfate is the sodium salt of sulfuric acid. With an annual production of 6 million tonnes, it is one of the world's major commodity chemicals. Anhydrous, it is a white crystalline solid of formula Na2SO4; the decahydrate Na2SO4·10H2O has been known as Glauber's salt or, historically, sal mirabilis since the 17th century.

Sodium sulfate is mainly used for the manufacture of detergents and in the Kraft process of paper pulping. About two thirds of the world's production is from mirabilite, the natural mineral form of the decahydrate, and the remainder from by-products of chemical processes such as hydrochloric acid production.


The hydrate of sodium sulfate is known as Glauber's Salt after the Dutch/German apothecary Johann Rudolf Glauber (1604–1670), who discovered it in Hungarian spring water. He himself named it sal mirabilis (miraculous salt), because of its medicinal properties: the crystals were used as a general purpose laxative, until more sophisticated alternatives came about in the 1900s.[1][2]

In the 18th century, Glauber's salt began to be used as a raw material for the industrial production of soda ash (sodium carbonate), by reaction with potash (potassium carbonate). Requirement for soda ash increased and supply of sodium sulfate had to increase in line. Therefore, in the nineteenth century, the Leblanc process, producing synthetic sodium sulfate as a key intermediate, became the principal method of soda ash production. [3]

Physical and chemical properties

Sodium sulfate is chemically very stable, being unreactive toward most oxidising or reducing agents at normal temperatures. At high temperatures, it can be reduced to sodium sulfide.[4] It is a neutral salt, which forms aqueous solutions with pH of 7. The neutrality of such solutions reflects the fact that Na2SO4 is derived, formally speaking, from the strong acid sulfuric acid and a strong base sodium hydroxide. Sodium sulfate reacts with an equivalent amount of sulfuric acid to give an equilibrium concentration of the acid salt sodium bisulfate[5][6]:

Na2SO4(aq) + H2SO4(aq) 2 NaHSO4(aq)

In fact, the equilibrium is very complex, depending on concentration and temperature, with other acid salts being present.

Sodium sulfate is a typical ionic sulfate, containing Na+ ions and SO42− ions. Aqueous solutions can produce precipitates when combined with salts of Ba2+ or Pb2+, which form insoluble sulfates

Na2SO4(aq) + BaCl2(aq) → 2 NaCl(aq) + BaSO4(s)
Graph showing solubility of Na2SO4 vs. temperature

Sodium sulfate has unusual solubility characteristics in water.[7] Its solubility rises more than tenfold between 0 °C to 32.4 °C, where it reaches a maximum of 49.7 g Na2SO4 per 100 g water. At this point the solubility curve changes slope, and the solubility becomes almost independent of temperature. In the presence of NaCl, the solubility of sodium sulfate is markedly diminished. Such changes provide the basis for the use of sodium sulfate in passive solar heating systems, as well is in the preparation and purification of sodium sulfate. This nonconformity can be explained in terms of hydration, since 32.4 °C corresponds with the temperature at which the crystalline decahydrate (Glauber's salt) changes to give a sulfate liquid phase and an anhydrous solid phase.

Sodium sulfate decahydrate is also unusual among hydrated salts in having a measureable residual entropy (entropy at absolute zero) of 6.32 J·K-1·mol-1. This is ascribed to its ability to distribute water much more rapidly compared to most hydrates.[8]

Sodium sulfate displays a moderate tendency to form double salts. The only alums formed with common trivalent metals are NaAl(SO4)2 (unstable above 39 °C) and NaCr(SO4)2, in contrast to potassium sulfate and ammonium sulfate which form many stable alums.[9] Double salts with some other alkali metal sulfates are known, including Na2SO4.3K2SO4 which occurs naturally as the mineral glaserite. Formation of glaserite by reaction of sodium sulfate with potassium chloride has been used as the basis of a method for producing potassium sulfate, a fertiliser.[10] Other double salts include 3Na2SO4.CaSO4, 3Na2SO4.MgSO4 (vanthoffite) and NaF.Na2SO4.[11]


Although sodium sulfate is generally regarded as non-toxic,[12] it should be handled with care. The dust can cause temporary asthma or eye irritation; this risk can be prevented by using eye protection and a paper mask. Transport is not limited, and no Risk Phrase or Safety Phrase apply.[13]


  1. Szydlo, Zbigniew (1994). Water which does not wet hands: The Alchemy of Michael Sendivogius. London-Warsaw: Polish Academy of Sciences. 
  2. Westfall, Richard S. (1995). "Glauber, Johann Rudolf". The Galileo Project. 
  3. Aftalion, Fred (1991). A History of the International Chemical Industry. Philadelphia: University of Pennsylvania Press. pp. pp. 11–16. ISBN 0-8122-1297-5. 
  4. Handbook of Chemistry and Physics (71st edition ed.). Ann Arbor, Michigan: CRC Press. 1990. 
  5. The Merck Index (7th edition ed.). Rahway, New Jersey, USA: Merck & Co. 1960. 
  6. Nechamkin, Howard (1968). The Chemistry of the Elements. New York: McGraw-Hill. 
  7. Linke, W.F. (1965). Solubilities of Inorganic and Metal Organic Compounds (4th edition ed.). Van Nostrand. 
  8. Brodale, G. (1958). "The Heat of Hydration of Sodium Sulfate. Low Temperature Heat Capacity and Entropy of Sodium Sulfate Decahydrate". Journal of the American Chemical Society. ACS. 80: pp. 2042–2044. 
  9. Lipson, Henry (1935). "The Crystal Structure of the Alums". Proceedings of the Royal Society of London. Series A, Mathematical and Physical Sciences. 148 (865): pp. 664–80. 
  10. Garrett, Donald E. (2001). Sodium sulfate : handbook of deposits, processing, properties, and use. San Diego: Academic Press. ISBN 9780122761515. 
  11. Mellor, Joseph William (1961). Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry. Volume II (new impression ed.). London: Longmans. pp. pp. 656–673. 
  12. "Sodium sulfate (WHO Food Additives Series 44)". World Health Organization. 2000. Retrieved 2007-06-06. 
  13. "MSDS Sodium Sulfate Anhydrous". James T Baker. 2006. Retrieved 2007-04-21. 

External links

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