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Sodium sulfate



Sodium sulfate
Other names Salt cake
Thenardite (mineral)
Glauber's salt (decahydrate)
Sal mirabilis (decahydrate)
Mirabilite (decahydrate)
Identifiers
CAS number 7757-82-6
RTECS number WE1650000 (anhydrous)
Properties
Molar mass 142.04 g/mol (anhydrous)
268.15 g/mol (heptahydrate)
322.20 g/mol (decahydrate)
Appearance White crystalline solid,
hygroscopic
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)
Structure
Crystal structure monoclinic, orthorhombic or
hexagonal
Hazards
MSDS External MSDS
MSDS External MSDS
Main hazards Irritant
NFPA 704
0
1
0
 
R/S statement None
Related Compounds
Other anions Sodium bisulfate
Sodium sulfite
Sodium bisulfite
Sodium persulfate
Other cations Lithium sulfate
Potassium sulfate
Magnesium sulfate
Supplementary data page
Structure and
properties
n, εr, etc.
Thermodynamic
data
Phase behaviour
Solid, liquid, gas
Spectral data UV, IR, NMR, MS
Except where noted otherwise, data are given for
materials in their standard state
(at 25 °C, 100 kPa)

Infobox disclaimer and references

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.

Additional recommended knowledge

Contents

History

Glauber's salt is named after Johann Rudolf Glauber, who discovered it in the 17th century in Hungarian spring water, and named it sal mirabilis (miraculous salt), because of its medicinal properties. The white or colourless crystals of Glauber's salt were used as a laxative until more sophisticated alternatives came about in 1900s.[1]

In the eighteenth century, Glauber's salt began to be used as a raw material for the production of soda ash, by reaction with potash. In the nineteenth century the Leblanc process became the principal method of soda production, and one of the world's major chemical processes, using synthetic sodium sulfate as a key intermediate.[2]

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.[3] 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[4][5]:

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)

Sodium sulfate has unusual solubility characteristics in water.[6] 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.[7]

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.[8] 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.[9] Other double salts include 3Na2SO4.CaSO4, 3Na2SO4.MgSO4 (vanthoffite) and NaF.Na2SO4.[10]

Production

The world production of sodium sulfate, mostly in the form of the decahydrate amounts to approximately 5.5 to 6 million tonnes annually (Mt/a). In 1985, production was 4.5 Mt/a, half from natural sources, and half from chemical production. After 2000, at a stable level until 2006, natural production had increased to 4 Mt/a, and chemical production decreased to 1.5 to 2 Mt/a, with a total of 5.5 to 6 Mt/a.[11][12][13][14] For all applications, naturally produced and chemically produced sodium sulfate are practically interchangeable.

Natural sources

Two thirds of the world's production of the decahydrate (Glauber's salt) is from the natural mineral form mirabilite, for example as found in lake beds in southern Saskatchewan. In 1990, Mexico and Spain were the world's main producers of natural sodium sulfate (each around 500,000 tonnes), with Russia, USA and Canada around 350,000 tonnes each.[12] Estimatedly, natural resources amount to over 1 billion tonnes.[11][12]

Major producers of 200–1500 Mt/a in 2006 include Searles Valley Minerals (California, USA), Airborne Industrial Minerals (Saskatchewan, Canada), Química del Rey (Coahuila, Mexico), Criaderos Minerales Y Derivados and Minera de Santa Marta, also known as Grupo Crimidesa (Burgos, Spain), FMC Foret (Toledo, Spain), Sulquisa (Madrid, Spain), and in China Chengdu Sanlian Tianquan Chemical (Sichuan), Hongze Yinzhu Chemical Group (Jiangsu), Nafine Chemical Industry Group (Shanxi), and Sichuan Province Chuanmei Mirabilite (Sichuan), and Kuchuksulphat JSC (Altai Krai, Siberia, Russia).[11][13]

Anhydrous sodium sulfate occurs in arid environments as the mineral thenardite. It slowly turns to mirabilite in damp air. Sodium sulfate is also found as glauberite, a calcium sodium sulfate mineral. Both minerals are less common than mirabilite.

Chemical industry

About one third of the world's sodium sulfate is produced as by-product of other processes in chemical industry. Most of this production is chemically inherent to the primary process, and only marginally economical. By effort of the industry, therefore, sodium sulfate production as by-product is declining.

The most important chemical sodium sulphate production is during hydrochloric acid production, either from sodium chloride (salt) and sulfuric acid, in the Mannheim process, or from sulfur dioxide in the Hargreaves process.[15][16] The resulting sodium sulfate from these processes are known as salt cake.

Mannheim: 2 NaCl + H2SO4 → 2 HCl + Na2SO4
Hargreaves: 4 NaCl + 2 SO2 + O2 + 2 H2O → 4 HCl + 2 Na2SO4

The second major production of sodium sulfate are the processes where surplus sulfuric acid is neutralised by sodium hydroxide, as applied on a large scale in the production of rayon. This method is also a regularly applied and convenient laboratory preparation.

2 NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2 H2O(l)

Formerly, sodium sulfate was also a by-product of the manufacture of sodium dichromate, where sulfuric acid is added to sodium chromate solution forming sodium dichromate, or subsequently chromic acid. Alternatively, sodium sulfate is or was formed in the production of lithium carbonate, chelating agents, resorcinol, ascorbic acid, silica pigments, nitric acid, and phenol.[11]

Bulk sodium sulfate is usually purified via the decahydrate form, since the anhydrous form tends to attract iron compounds and organic compounds. The anhydrous form is easily produced from the hydrated form by gentle warming.

Major sodium sulfate by-product producers of 50–80 Mt/a in 2006 include Elementis Chromium (chromium industry, Castle Hayne, NC, USA), Lenzing AG (200 Mt/a, rayon industry, Lenzing, Austria), Addiseo (formerly Rhodia, methionine industry, Les Roches-Roussillon, France), Elementis (chromium industry, Stockton-on-Tees, UK), Shikoku Chemicals (Tokushima, Japan) and Visko-R (rayon industry, Russia).[11]

Applications

Sodium sulfate as a drying agent(1)

Drying a wet organic phase using sodium sulfate, which clumps, indicating that more sodium sulfate is needed.

details and copyright info
In case of problems, see media help.
Sodium sulfate as a drying agent(2)

Drying a fairly dry organic phase using sodium sulfate, which does not clump, indicating that the solution is dry.

details and copyright info
In case of problems, see media help.

Commodity industries

With USA pricing at $30 per metric tonne in 1970, in 2006 up to $90 per metric tonne for salt cake quality and $130 for better grades, sodium sulfate is a very cheap material. The largest use is as filler in powdered home laundry detergents, consuming approx. 50% of world production. This use is waning as domestic consumers are increasingly switching to compact or liquid detergents that do not include sodium sulfate.[11]

Another formerly major use for sodium sulfate, notably in the USA and Canada, is in the Kraft process for the manufacture of wood pulp. Organics present in the "black liquor" from this process are burnt to produce heat, needed to drive the reduction of sodium sulfate to sodium sulfide. However, this process is being replaced by newer processes; use of sodium sulfate in the USA and Canadian pulp industry declined from 1.4 Mt/a in 1970 to only approx. 150,000 tonnes in 2006.[11]

The glass industry provides another significant application for sodium sulfate, as second largest application in Europe. Sodium sulfate is used as a fining agent, to help remove small air bubbles from molten glass. It fluxes the glass, and prevents scum formation of the glass melt during refining. The glass industry in Europe has been consuming from 1970 to 2006 a stable 110,000 tonnes annually.[11]

Sodium sulfate is important in the manufacture of textiles, particularly in Japan, where it is the largest application. Sodium sulfate helps in "levelling", reducing negative charges on fibres so that dyes can penetrate evenly. Unlike the alternative sodium chloride, it does not corrode the stainless steel vessels used in dyeing. This application in Japan and USA consumed in 2006 approximately 100,000 tonnes.[11]

Thermal storage

The high heat storage capacity in the phase change from solid to liquid, and the advantageous phase change temperature of 32 degrees Celsius (90 degrees Fahrenheit) makes this material especially appropriate for storing low grade solar heat for later release in space heating applications. In some application the material is incorporated into thermal tiles that are placed in an attic space while in other applications the salt is incorporated into cells surrounded by solar–heated water. The phase change allows a substantial reduction in the mass of the material required for effective heat storage (83 calories per gram stored across the phase change, versus one calorie per gram per degree Celsius using only water), with the further advantage of a consistency of temperature as long as sufficient material in the appropriate phase is available.

Small scale applications

In the laboratory, anhydrous sodium sulfate is widely used as an inert drying agent, for removing traces of water from organic solutions.[17] It is more efficient, but slower-acting, than the similar agent magnesium sulfate. It is only effective below about 30 °C, but it can used with a variety of materials since it is chemically fairly inert. Sodium sulfate is added to the solution until the crystals no longer clump together; the two video clips (see above) demonstrate how the crystals clump when still wet, but some crystals flow freely once a sample is dry.

Glauber's salt, the decahydrate, was historically used as a laxative. It is effective for the removal of certain drugs such as acetaminophen from the body, for example, after an overdose.[18][19]

In 1953, sodium sulfate was proposed for heat storage in passive solar heating systems. This takes advantage of its unusual solubility properties, and the high heat of crystallisation (78.2 kJ/mol).[20]

Other uses for sodium sulfate include de-frosting windows, in carpet fresheners, starch manufacture, and as an additive to cattle feed.

Lately, sodium sulfate has been found effective in dissolving very finely electroplated micrometre gold that is found in gold electroplated hardware on electronic products such as pins, and other connectors and switches. It is safer and cheaper than other reagents used for gold recovery, with little concern for adverse reactions or health effects.[citation needed]

Safety

Although sodium sulfate is generally regarded as non-toxic,[21] 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.[22]

References

  1. ^ Szydlo, Zbigniew (1994). Water which does not wet hands: The Alchemy of Michael Sendivogius. London-Warsaw: Polish Academy of Sciences. 
  2. ^ Aftalion, Fred (1991). A History of the International Chemical Industry. Philadelphia: University of Pennsylvania Press, pp. 11–16. ISBN 0-8122-1297-5. 
  3. ^ (1990) Handbook of Chemistry and Physics, 71st edition, Ann Arbor, Michigan: CRC Press. 
  4. ^ (1960) The Merck Index, 7th edition, Rahway, New Jersey, USA: Merck & Co.. 
  5. ^ Nechamkin, Howard (1968). The Chemistry of the Elements. New York: McGraw-Hill. 
  6. ^ Linke, W.F.; A. Seidell (1965). Solubilities of Inorganic and Metal Organic Compounds, 4th edition, Van Nostrand. 
  7. ^ Brodale, G.; W.F. Giauque (1958). "The Heat of Hydration of Sodium Sulfate. Low Temperature Heat Capacity and Entropy of Sodium Sulfate Decahydrate". Journal of the American Chemical Society 80: pp. 2042–2044. ACS.
  8. ^ Lipson, Henry; C.A. Beevers (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.
  9. ^ Garrett, Donald E. (2001). Sodium sulfate : handbook of deposits, processing, properties, and use. San Diego: Academic Press. ISBN 9780122761515. 
  10. ^ Mellor, Joseph William (1961). Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry, new impression, London: Longmans, pp. 656–673. 
  11. ^ a b c d e f g h i Suresh, Bala; Kazuteru Yokose (May 2006). Sodium sulfate. Zurich: Chemical Economic Handbook SRI Consulting, pp. 771.1000A–771.1002J. 
  12. ^ a b c Statistical compendium Sodium sulfate. US Geological Survey, Minerals Information (1997). Retrieved on 2007-04-22.
  13. ^ a b (1999) The economics of sodium sulphate, Eighth edition, London: Roskill Information Services, 195 pages and appendices. 
  14. ^ (Nov 1984) The sodium sulphate business. London: Chem Systems International. 
  15. ^ Butts, D. (1997). Kirk-Othmer Encyclopedia of Chemical Technology, 4th edition. 
  16. ^ Hargreaves, J. (1873). "". Chem. News 27: p. 183.
  17. ^ Vogel, Arthur I.; B.V. Smith, N.M. Waldron (1980). Vogel's Elementary Practical Organic Chemistry 1 Preparations, 3rd Edition, London: Longman Scientific & Technical. 
  18. ^ Cocchetto, D.M.; G. Levy (1981). "Absorption of orally administered sodium sulfate in humans.". J Pharm Sci 70 (3): p. 331–3. doi:10.1002/jps.2600700330. Retrieved on 2007-06-06.
  19. ^ Prescott, L.F.; J.A.J.H. Critchley (1979). "The Treatment of Acetaminophen Poisoning". Annual Review of Pharmacology and Toxicology 23: pp. 87–101. doi:10.1146/annurev.pa.23.040183.000511.
  20. ^ Telkes, Maria (1953). Improvements in or relating to a device and a composition of matter for the storage of heat. 
  21. ^ Sodium sulfate (WHO Food Additives Series 44). World Health Organization (2000). Retrieved on 2007-06-06.
  22. ^ MSDS Sodium Sulfate Anhydrous. James T Baker (2006). Retrieved on 2007-04-21.
 
This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Sodium_sulfate". A list of authors is available in Wikipedia.
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