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Sodium sulfate
History
The hydrate of sodium sulfate is known as Glauber's Salt after the Dutch/German chemist and apothecary Johann Rudolf Glauber (16041670), who discovered it 1625 in Austrian spring water. He 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.
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). Demand 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.
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. 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:
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. Its solubility rises more than tenfold between 0 C to 32.384 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. This temperature at 32.384 C, corresponding to the release of crystal water and melting of the hydrated salt, serves as an accurate temperature reference for thermometer calibration.
Sodium sulfate decahydrate is also unusual among hydrated salts in having a measureable residual entropy (entropy at absolute zero) of 6.32 JK-1mol-1. This is ascribed to its ability to distribute water much more rapidly compared to most hydrates.
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. Double salts with some other alkali metal sulfates are known, including Na2SO43K2SO4 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. Other double salts include 3Na2SO4CaSO4, 3Na2SO4MgSO4 (vanthoffite) and NaFNa2SO4.
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. 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. Natural resources are estimated as over 1 billion tonnes.
Major producers of 200,0001,500,000 tonnes/a in 2006 include Searles Valley Minerals (California, USA), Airborne Industrial Minerals (Saskatchewan, Canada), Qumica 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).
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 sulfate 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. 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.
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 5080 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).
Applications
Sodium sulfate used to dry an organic liquid. Here clumps form, indicating the presence of water in the organic liquid.
By further application of sodium sulfate the liquid may be brought to dryness, indicated here by the absence of clumping.
Commodity industries
With USA pricing at $30 per tonne in 1970, in 2006 up to $90 per 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.
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.
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.
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.
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 applications 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 solareated 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. 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.
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).
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]
At least one company makes a laptop computer chill mat using sodium sulfate decahydrate inside a quilted plastic pad. The material slowly turns to liquid as the heat from the laptop is transferred.[citation needed]
Safety
Although sodium sulfate is generally regarded as non-toxic, 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.
References
^ Szydlo, Zbigniew (1994). Water which does not wet hands: The Alchemy of Michael Sendivogius. London-Warsaw: Polish Academy of Sciences.
^ Westfall, Richard S. (1995). "Glauber, Johann Rudolf". The Galileo Project. http://galileo.rice.edu/Catalog/NewFiles/glauber.html.
^ Aftalion, Fred (1991). A History of the International Chemical Industry. Philadelphia: University of Pennsylvania Press. pp. 1116. ISBN 0-8122-1297-5.
^ Handbook of Chemistry and Physics (71st ed.). Ann Arbor, Michigan: CRC Press. 1990.
^ The Merck Index (7th ed.). Rahway, New Jersey, USA: Merck & Co.. 1960.
^ Nechamkin, Howard (1968). The Chemistry of the Elements. New York: McGraw-Hill.
^ Linke, W.F.; A. Seidell (1965). Solubilities of Inorganic and Metal Organic Compounds (4th ed.). Van Nostrand.
^ 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 (ACS) 80: pp. 20422044. doi:10.1021/ja01542a003.
^ 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): 66480. doi:10.1098/rspa.1935.0040.
^ Garrett, Donald E. (2001). Sodium sulfate : handbook of deposits, processing, properties, and use. San Diego: Academic Press. ISBN 9780122761515.
^ Mellor, Joseph William (1961). Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry. Volume II (new impression ed.). London: Longmans. pp. 656673.
^ a b c d e f g h i Suresh, Bala; Kazuteru Yokose (May 2006). Sodium sulfate. Zurich: Chemical Economic Handbook SRI Consulting. pp. 771.1000A771.1002J. http://www.sriconsulting.com/CEH/Public/Reports/771.1000/?Abstract.html.
^ a b c "Statistical compendium Sodium sulfate". Reston, Virginia: US Geological Survey, Minerals Information. 1997. http://minerals.usgs.gov/minerals/pubs/commodity/sodium_sulfate/stat. Retrieved 2007-04-22.
^ a b The economics of sodium sulphate (Eighth ed.). London: Roskill Information Services. 1999. pp. 195 pages and appendices.
^ The sodium sulphate business. London: Chem Systems International. November 1984.
^ Butts, D. (1997). Kirk-Othmer Encyclopedia of Chemical Technology. v22 (4th ed.).
^ Hargreaves, J. (1873). Chem. News 27: p. 183.
^ Vogel, Arthur I.; B.V. Smith, N.M. Waldron (1980). Vogel's Elementary Practical Organic Chemistry 1 Preparations (3rd ed.). London: Longman Scientific & Technical.
^ Cocchetto, D.M.; G. Levy (1981). "Absorption of orally administered sodium sulfate in humans". J Pharm Sci 70 (3): p. 3313. doi:10.1002/jps.2600700330. http://www.ncbi.nlm.nih.gov/pubmed/7264905?dopt=Citation. Retrieved 2007-06-06.
^ Prescott, L.F.; J.A.J.H. Critchley (1979). "The Treatment of Acetaminophen Poisoning". Annual Review of Pharmacology and Toxicology 23: pp. 87101. doi:10.1146/annurev.pa.23.040183.000511.
^ Telkes, Maria (1953). Improvements in or relating to a device and a composition of matter for the storage of heat. http://v3.espacenet.com/textdes?DB=EPODOC&IDX=GB694553&F=0&QPN=GB694553.
^ "Sodium sulfate (WHO Food Additives Series 44)". World Health Organization. 2000. http://www.inchem.org/documents/jecfa/jecmono/v44jec07.htm. Retrieved 2007-06-06.
^ "MSDS Sodium Sulfate Anhydrous". James T Baker. 2006. http://www.jtbaker.com/msds/englishhtml/S5022.htm. Retrieved 2007-04-21.
External links
Sodium sulfate information of Airborne Industrial Minerals
Sodium sulfate website of Elementis Chromium
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Sodium compounds
NaAlO2 NaBH3(CN) NaBH4 NaBr NaBrO4 NaCH3COO NaCN NaC6H5CO2 NaC6H4(OH)CO2 NaCl NaClO NaClO2 NaClO3 NaClO4 NaF NaH NaHCO3 NaHSO3 NaHSO4 NaI NaIO3 NaIO4 NaMnO4 NaNH2 NaNO2 NaNO3 NaN3 NaOH NaO2 NaPO2H2 NaReO4 NaSCN NaSH NaTcO4 NaVO3 Na2CO3 Na2C2O4 Na2CrO4 Na2Cr2O7 Na2MnO4 Na2MoO4 Na2O Na2O2 Na2O(UO3)2 Na2S Na2SO3 Na2SO4 Na2S2O3 Na2S2O4 Na2S2O5 Na2S2O6 Na2S2O7 Na2S2O8 Na2SeO3 Na2SeO4 Na2SiO3 Na2Te Na2TeO3 Na2Ti3O7 Na2U2O7 NaWO4 Na2Zn(OH)4 Na3N Na3P Na3VO4 Na4Fe(CN)6 Na5P3O10
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