Sodium sulfate 794439 223101657 2008-07-02T16:03:51Z DOI bot 6652755 Citation maintenance. Formatted: title. You can [[WP:DOI|use this bot]] yourself! Please [[User:DOI_bot/bugs|report any bugs]]. <!-- Spelling: this article is in EN-GB spelling. Note that, chemically, sulfate is with an 'f', even in EN-GB -->{{Chembox new <!-- infobox --> | Name = Sodium sulfate | ImageFile = Sodium sulfate.jpg | ImageSize = 150px | ImageName = Sodium sulfate | OtherNames = Sodium sulphate<br>Salt cake<br />[[Thenardite]] (mineral)<br />Glauber's salt (decahydrate)<br />Sal mirabilis (decahydrate)<br />[[Mirabilite]] (decahydrate) | Section1 = {{Chembox Identifiers | CASNo = 7757-82-6 | RTECS = WE1650000 (anhydrous) }} | Section2 = {{Chembox Properties | MolarMass = 142.04&nbsp;g/mol ([[anhydrous]])<br />268.15&nbsp;g/mol (heptahydrate)<br />322.20&nbsp;g/mol (decahydrate) | Appearance = White crystalline solid,<br />[[hygroscopic]] | Density = 2.68&nbsp;g/cm³, anhydrous<br />(orthorhombic form)<br />1.464&nbsp;g/cm³, decahydrate | Solubility = 4.76 g/100&nbsp;ml (0&nbsp;°C)<br />42.7&nbsp;g/100&nbsp;ml (100&nbsp;°C) | MeltingPt = 884&nbsp;°C (1157&nbsp;K) anhydrous<br />32.4&nbsp;°C decahydrate }} | Section3 = {{Chembox Structure | Coordination = | CrystalStruct = [[monoclinic]], [[orthorhombic]] or<br />[[hexagonal]] }} | Section7 = {{Chembox Hazards | ExternalMSDS = [[Sodium sulfate chemdata supplement#Material Safety Data Sheet|External MSDS]] | MainHazards = Irritant | RSPhrases = None | NFPA-H = 1 | NFPA-F = | NFPA-R = }} | Section8 = {{Chembox Related | OtherAnions = [[Sodium bisulfate]]<br />[[Sodium sulfite]]<br />[[Sodium bisulfite]]<br />[[Sodium persulfate]] | OtherCations = [[Lithium sulfate]]<br />[[Potassium sulfate]]<br />[[Magnesium sulfate]]}} }} '''Sodium sulfate''' is the [[sodium]] salt of [[sulfuric acid]]. [[Anhydrous]], it is a white crystalline solid of formula Na<sub>2</sub>SO<sub>4</sub>; the [[hydrate|decahydrate]] Na<sub>2</sub>SO<sub>4</sub>·10H<sub>2</sub>O has been known as '''Glauber's salt''' or, historically, '''sal mirabilis''' since the 17th century. With an annual production of 6 million [[tonne]]s, it is one of the world's major [[commodity]] chemicals. Sodium sulfate is mainly used for the manufacture of [[detergent]]s 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 industry|chemical processes]] such as [[hydrochloric acid]] production. ==History== The hydrate of sodium sulfate is known as Glauber's Salt after the [[Netherlands|Dutch]]/[[Germany|German]] [[apothecary]] [[Johann Rudolf Glauber]] (1604&ndash;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.<ref name='szydlo'>{{cite book | first = Zbigniew | last = Szydlo | authorlink = Zbigniew Szydlo | title = Water which does not wet hands: The Alchemy of Michael Sendivogius | location = London-Warsaw | publisher = Polish Academy of Sciences | year = 1994}}</ref><ref name='galileo'>{{cite web | url = http://galileo.rice.edu/Catalog/NewFiles/glauber.html | title = Glauber, Johann Rudolf | first = Richard S. | last = Westfall | publisher = The Galileo Project | year = 1995}}</ref> In the 18th century, Glauber's salt began to be used as a raw material for the [[chemical industry|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. <ref name='Aftalion'>{{cite book | first = Fred | last = Aftalion | title = A History of the International Chemical Industry | location = Philadelphia | publisher = University of Pennsylvania Press | year = 1991 | pages = pp.&nbsp;11&ndash;16 | isbn = 0-8122-1297-5}}</ref> ==Physical and chemical properties== Sodium sulfate is chemically very stable, being unreactive toward most [[Redox|oxidising or reducing agents]] at normal temperatures. At high temperatures, it can be reduced to [[sodium sulfide]].<ref name='crc'>{{cite book | title = Handbook of Chemistry and Physics | edition = 71st edition | publisher = [[CRC Press]] | location = [[Ann Arbor]], [[Michigan]] | year = 1990}}</ref> It is a [[Acidity|neutral]] salt, which forms aqueous solutions with pH of 7. The neutrality of such solutions reflects the fact that Na<sub>2</sub>SO<sub>4</sub> 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 [[Chemical equilibrium|equilibrium]] [[concentration]] of the [[acid salt]] [[sodium bisulfate]]<ref name='merck'>{{cite book | title = [[Merck Index|The Merck Index]] | edition = 7th edition | publisher = [[Merck & Co.]] | location = Rahway, New Jersey, USA | year = 1960}}</ref><ref>{{cite book | first = Howard |last = Nechamkin | title = The Chemistry of the Elements | publisher = [[McGraw-Hill]] | location = New York | year = 1968}}</ref>: : Na<sub>2</sub>SO<sub>4</sub>([[Aqueous|aq]]) + H<sub>2</sub>SO<sub>4</sub>(aq) {{unicode|⇌}} 2 [[sodium bisulfate|NaHSO<sub>4</sub>]](aq) In fact, the equilibrium is very complex, depending on [[concentration]] and temperature, with other acid salts being present. Sodium sulfate is a typical [[ion]]ic sulfate, containing Na<sup>+</sup> ions and [[sulfate|SO<sub>4</sub><sup>2&minus;</sup>]] ions. Aqueous solutions can produce precipitates when combined with salts of [[barium|Ba<sup>2+</sup>]] or [[lead|Pb<sup>2+</sup>]], which form insoluble sulfates : Na<sub>2</sub>SO<sub>4</sub>(aq) + [[Barium chloride|BaCl<sub>2</sub>]](aq) → 2 [[Sodium chloride|NaCl]](aq) + [[Barium sulfate|BaSO<sub>4</sub>]]([[Solid|s]]) [[Image:Na2SO4 solubility.png|left|350px|Graph showing solubility of Na<sub>2</sub>SO<sub>4</sub> vs. temperature]] Sodium sulfate has unusual solubility characteristics in water.<ref>{{cite book | first = W.F. | last = Linke | coauthors = A. Seidell | title = Solubilities of Inorganic and Metal Organic Compounds | edition = 4th edition | publisher = Van Nostrand | year = 1965}}</ref> Its solubility rises more than tenfold between 0&nbsp;°C to 32.4&nbsp;°C, where it reaches a maximum of 49.7&nbsp;g Na<sub>2</sub>SO<sub>4</sub> per 100&nbsp;g water. At this point the solubility curve changes slope, and the solubility becomes almost independent of temperature. In the presence of [[sodium chloride|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&nbsp;°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&nbsp;J·K<sup>-1</sup>·mol<sup>-1</sup>. This is ascribed to its ability to distribute water much more rapidly compared to most hydrates.<ref name="Brodale1957">{{cite journal | last = Brodale | first = G. | coauthors = W.F. Giauque | title = The Heat of Hydration of Sodium Sulfate. Low Temperature Heat Capacity and Entropy of Sodium Sulfate Decahydrate | journal = [[Journal of the American Chemical Society]] | volume = 80 | pages = pp.&nbsp;2042&ndash;2044 | publisher = [[American Chemical Society|ACS]] | year = 1958 | doi = 10.1021/ja01542a003 }}</ref> Sodium sulfate displays a moderate tendency to form [[double salt]]s. The only [[alum]]s formed with common trivalent metals are NaAl(SO<sub>4</sub>)<sub>2</sub> (unstable above 39&nbsp;°C) and NaCr(SO<sub>4</sub>)<sub>2</sub>, in contrast to [[potassium sulfate]] and [[ammonium sulfate]] which form many stable alums.<ref name=Lipson1935>{{cite journal | last = Lipson | first = Henry | authorlink = Henry Lipson | coauthors = C.A. Beevers | year = 1935 | title = The Crystal Structure of the Alums | journal = Proceedings of the [[Royal Society of London]]. Series A, Mathematical and Physical Sciences | volume = 148 | issue = 865 | pages = pp.&nbsp;664&ndash;80 | doi = 10.1098/rspa.1935.0040}}</ref> Double salts with some other alkali metal sulfates are known, including Na<sub>2</sub>SO<sub>4</sub>.3K<sub>2</sub>SO<sub>4</sub> 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 [[fertilizer|fertiliser]].<ref name=Garrett2001>{{cite book | last = Garrett | first = Donald E. | title = Sodium sulfate : handbook of deposits, processing, properties, and use | publisher = Academic Press | year = 2001 | location = San Diego | isbn = 9780122761515}}</ref> Other double salts include 3Na<sub>2</sub>SO<sub>4</sub>.CaSO<sub>4</sub>, 3Na<sub>2</sub>SO<sub>4</sub>.MgSO<sub>4</sub> ([[vanthoffite]]) and NaF.Na<sub>2</sub>SO<sub>4</sub>.<ref name=Mellor1961>{{cite book | last = Mellor | first = Joseph William | title = Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry | volume = Volume II | publisher = Longmans | year = 1961 | edition = new impression | location = London | pages = pp.&nbsp;656&ndash;673}}</ref> ==Production== The world production of sodium sulfate, mostly in the form of the decahydrate amounts to approximately 5.5 to 6&nbsp;million&nbsp;tonnes annually (Mt/a). In 1985, production was 4.5&nbsp;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&nbsp;Mt/a, and chemical production decreased to 1.5 to 2&nbsp;Mt/a, with a total of 5.5 to 6&nbsp;Mt/a.<ref name='ceh'>{{cite book | first = Bala | last = Suresh | coauthors = Kazuteru Yokose | title = Sodium sulfate | url = http://www.sriconsulting.com/CEH/Public/Reports/771.1000/?Abstract.html | works = CEH Marketing Research Report | location = Zurich | publisher = Chemical Economic Handbook SRI Consulting | year = May 2006 | pages = pp.&nbsp;771.1000A&ndash;771.1002J}}</ref><ref name ='usgs'>{{cite web | url = http://minerals.usgs.gov/minerals/pubs/commodity/sodium_sulfate/stat | title = Statistical compendium Sodium sulfate | publisher = [[US Geological Survey]], Minerals Information | location = Reston, Virginia | year = 1997 | accessdate = 2007-04-22}}</ref><ref name='roskill'>{{cite book | title = The economics of sodium sulphate | edition = Eighth edition | year = 1999 | location = London | publisher = Roskill Information Services | pages = 195 pages and appendices}}</ref><ref name='chemsys'>{{cite book | title = The sodium sulphate business | month = Nov | year = 1984 | location = London | publisher = Chem Systems International}}</ref> 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&nbsp;[[tonne]]s), with [[Russia]], [[USA]] and [[Canada]] around 350,000&nbsp;tonnes each.<ref name='usgs'/> Estimatedly, natural resources amount to over 1 billion tonnes.<ref name='ceh'/><ref name='usgs'/> Major producers of 200&ndash;1500 Mt/a in 2006 include [[Trona, California|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).<ref name='ceh'/><ref name='roskill'/> 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]].<ref name='kirk-othmer'>{{cite book | first = D. | last = Butts | title = Kirk-Othmer Encyclopedia of Chemical Technology | edition = 4th edition | volume = v22 | paged = p.&nbsp;403&ndash;411 | year = 1997}}</ref><ref name='hargreaves'>{{cite journal | first = J. | last = Hargreaves | title = | journal = Chem. News | volume = 27 | pages = p.&nbsp;183 | year = 1873}}</ref> The resulting sodium sulfate from these processes are known as ''salt cake''. :Mannheim: 2 [[Sodium chloride|NaCl]] + [[Sulfuric acid|H<sub>2</sub>SO<sub>4</sub>]] → 2 [[Hydrogen chloride|HCl]] + Na<sub>2</sub>SO<sub>4</sub> :Hargreaves: 4 NaCl + 2 [[Sulfur dioxide|SO<sub>2</sub>]] + [[Oxygen|O<sub>2</sub>]] + 2 [[Water (molecule)|H<sub>2</sub>O]] → 4 [[Hydrogen chloride|HCl]] + 2 Na<sub>2</sub>SO<sub>4</sub> The second major production of sodium sulfate are the processes where surplus sulfuric acid is [[neutralisation|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 [[Sodium hydroxide|NaOH]]([[Aqueous|aq]]) + H<sub>2</sub>SO<sub>4</sub>(aq) → Na<sub>2</sub>SO<sub>4</sub>(aq) + 2 H<sub>2</sub>O([[Liquid|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 agent]]s, [[resorcinol]], [[ascorbic acid]], [[silica]] pigments, [[nitric acid]], and [[phenol]].<ref name='ceh'/> Bulk sodium sulfate is usually purified via the decahydrate form, since the anhydrous form tends to attract [[iron]] compounds and [[organic compound]]s. The anhydrous form is easily produced from the hydrated form by gentle warming. Major sodium sulfate by-product producers of 50&ndash;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).<ref name='ceh'/> ==Applications== {| width=350px | align = right |{{video float | filename = Sulfate clump.ogg | title = Sodium sulfate as a drying agent(1) | description = Drying a wet organic phase using sodium sulfate, which clumps, indicating that more sodium sulfate is needed.}} |- |{{video float | filename = Sulfate noclump.ogg | title = Sodium sulfate as a drying agent(2) | description = Drying a fairly dry organic phase using sodium sulfate, which does not clump, indicating that the solution is dry.}} |} ===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 [[detergent]]s, 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.<ref name='ceh'/> 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 [[Redox|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&nbsp;tonnes in 2006.<ref name='ceh'/> 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&nbsp;tonnes annually.<ref name='ceh'/> 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&nbsp;tonnes.<ref name='ceh'/> ===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&ndash;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 [[desiccant|drying agent]], for removing traces of water from organic solutions.<ref name=vogel>{{cite book | last = Vogel | first = Arthur I. | coauthors = B.V. Smith, N.M. Waldron | edition = 3rd Edition | title = Vogel's Elementary Practical Organic Chemistry 1 Preparations | publisher = [[Longman]] Scientific & Technical | location = London | year = 1980}}</ref> It is more efficient, but slower-acting, than the similar agent [[magnesium sulfate]]. It is only effective below about 30&nbsp;°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.<ref name=Cocchetto1981>{{cite journal | last = Cocchetto | first = D.M. | coauthors = G. Levy | year = 1981 | title = Absorption of orally administered sodium sulfate in humans | journal = J Pharm Sci | volume = 70 | issue = 3 | pages = p.&nbsp;331&ndash;3| doi = 10.1002/jps.2600700330 | url = http://www.ncbi.nlm.nih.gov/pubmed/7264905?dopt=Citation | accessdate = 2007-06-06 }}</ref><ref name=Prescott1979>{{cite journal | last = Prescott | first = L.F. | coauthors = J.A.J.H. Critchley | year = 1979 | title = The Treatment of Acetaminophen Poisoning | journal = Annual Review of Pharmacology and Toxicology | volume = 23 | pages = pp.&nbsp;87&ndash;101 | doi = 10.1146/annurev.pa.23.040183.000511 }}</ref> 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&nbsp;kJ/mol).<ref>{{cite book | first = Maria | last = Telkes | title = Improvements in or relating to a device and a composition of matter for the storage of heat | url = http://v3.espacenet.com/textdes?DB=EPODOC&IDX=GB694553&F=0&QPN=GB694553 | work = British Patent No. GB694553 | year = 1953}}</ref> 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.{{Fact|date=July 2007}} 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. ==Safety== Although sodium sulfate is generally regarded as non-toxic,<ref name="WHO2000">{{cite web | title = Sodium sulfate (WHO Food Additives Series 44) | work = | publisher = [[World Health Organization]] | date = 2000 | url = http://www.inchem.org/documents/jecfa/jecmono/v44jec07.htm | accessdate = 2007-06-06}}</ref> 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.<ref name='msds'>{{cite web | url = http://www.jtbaker.com/msds/englishhtml/S5022.htm | title = MSDS Sodium Sulfate Anhydrous | publisher = James T Baker | accessdate = 2007-04-21 | year = 2006}}</ref> ==References== {{reflist|2}} ==External links== *[http://www.saskatchewanminerals.com/images/product%20info%20sheet-Dec-05.jpg Sodium sulfate information of Airborne Industrial Minerals] *[http://www.elementischromium.com/products/sodiumsulphate.htm Sodium sulfate website of Elementis Chromium] [[Category:Sulfates]] [[Category:Sodium compounds]] [[Category:Desiccants]] [[Category:Alchemical substances]] [[ar:كبريتات صوديوم]] [[bs:Natrijum sulfat]] [[bg:Глауберова сол]] [[da:Natriumsulfat]] [[de:Natriumsulfat]] [[es:Sulfato de sodio]] [[fr:Sulfate de sodium]] [[it:Solfato di sodio]] [[he:נתרן גופרתי]] [[la:Natrii sulfas]] [[lv:Nātrija sulfāts]] [[hu:Nátrium-szulfát]] [[nl:Natriumsulfaat]] [[ja:硫酸ナトリウム]] [[no:Natriumsulfat]] [[pl:Siarczan(VI) sodu]] [[pt:Sulfato de sódio]] [[ru:Сульфат натрия]] [[sv:Natriumsulfat]] [[zh:硫酸钠]]