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{{Short description|Chemical compound with formula Na<sub>2</sub>SO<sub>4</sub>}} |
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{{ambox | text = This page contains a copy of the infobox ({{tl|chembox}}) taken from revid of page ] with values updated to verified values.}} |
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<!-- Spelling: this article is in EN-GB spelling. Note that, chemically, sulfate is with an 'f', even in EN-GB --> |
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{{chembox |
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{{Chembox |
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| verifiedrevid = 476992246 |
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| Verifiedfields = |
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| Name = Sodium sulfate |
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| Watchedfields = changed |
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| ImageFileL1 = Sodium sulfate.jpg |
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| verifiedrevid = 477315271 |
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| ImageFileR1 = Sodium sulfate.png |
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| Name = Sodium sulfate |
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| ImageSize = 150px |
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| ImageName = Sodium sulfate |
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| ImageFileR1 = Sodium sulfate.jpg |
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| ImageFileL1 = Sodium sulfate.svg |
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| OtherNames = ] (mineral)<br/>Glauber's salt (decahydrate)<br/>Sal mirabilis (decahydrate)<br/>] (decahydrate) |
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| ImageClassL1 = skin-invert |
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| Section1 = {{Chembox Identifiers |
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| ImageSizeL1 = 150px |
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| UNII_Ref = {{fdacite|correct|FDA}} |
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| ImageSizeR1 = 120px |
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| UNII = 36KCS0R750 |
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| ImageName = Sodium sulphate |
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| InChI = 1S/2Na.H2O4S/c;;1-5(2,3)4/h;;(H2,1,2,3,4)/q2*+1;/p-2 |
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| IUPACName = Sodium sulfate |
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| InChIKey1 = PMZURENOXWZQFD-UHFFFAOYSA-L |
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| OtherNames = Sodium sulphate<br>Disodium sulfate<br>Sulfate of sodium<br>] (anhydrous mineral)<br>Glauber's salt (decahydrate)<br>''Sal mirabilis'' (decahydrate)<br>] (decahydrate mineral) |
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| InChI1 = 1S/2Na.H2O4S/c;;1-5(2,3)4/h;;(H2,1,2,3,4)/q2*+1;/p-2 |
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| Section1 = {{Chembox Identifiers |
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| CASNo = 7757-82-6 |
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| CASNo_Ref = {{cascite|correct|CAS}} |
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|UNII_Ref = {{fdacite|correct|FDA}} |
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|UNII = 36KCS0R750 |
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| CASOther = <br/>{{CAS|7727-73-3}} (decahydrate) |
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| ChEMBL_Ref = {{ebicite|correct|EBI}} |
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|UNII2_Ref = {{fdacite|correct|FDA}} |
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|UNII2 = 0YPR65R21J |
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| ChEMBL = 233406 |
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|UNII2_Comment = (decahydrate |
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| PubChem = 24436 |
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|InChI = 1S/2Na.H2O4S/c;;1-5(2,3)4/h;;(H2,1,2,3,4)/q2*+1;/p-2 |
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| RTECS = WE1650000 |
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|InChIKey1 = PMZURENOXWZQFD-UHFFFAOYSA-L |
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| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}} |
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|InChI1 = 1S/2Na.H2O4S/c;;1-5(2,3)4/h;;(H2,1,2,3,4)/q2*+1;/p-2 |
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| ChemSpiderID = 22844 |
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|CASNo = 7757-82-6 |
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| ChEBI_Ref = {{ebicite|correct|EBI}} |
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|CASNo_Ref = {{cascite|correct|CAS}} |
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| ChEBI = 32149 |
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| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}} |
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|CASNo2_Ref = {{cascite|correct|CAS}} |
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|CASNo2 = 7727-73-3 |
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| StdInChIKey = PMZURENOXWZQFD-UHFFFAOYSA-L |
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|CASNo2_Comment = (decahydrate) |
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| SMILES = ..S()(=O)=O |
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| StdInChI_Ref = {{stdinchicite|correct|chemspider}} |
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|ChEMBL_Ref = {{ebicite|correct|EBI}} |
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|ChEMBL = 233406 |
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| StdInChI = StdInChI_Ref = {{stdinchicite|correct|chemspider}} |
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|PubChem = 24436 |
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| StdInChI=1S/2Na.H2O4S/c;;1-5(2,3)4/h;;(H2,1,2,3,4)/q2*+1;/p-2 |
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|RTECS = WE1650000 |
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| ATCCode_prefix = A06 |
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|ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}} |
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| ATCCode_suffix = AD13 |
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|ChemSpiderID = 22844 |
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| ATC_Supplemental = {{ATC|A12|CA02}} |
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|ChEBI_Ref = {{ebicite|correct|EBI}} |
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|ChEBI = 32149 |
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|StdInChIKey_Ref = {{stdinchicite|correct|chemspider}} |
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|StdInChIKey = PMZURENOXWZQFD-UHFFFAOYSA-L |
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|SMILES = ..S()(=O)=O |
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|StdInChI_Ref = {{stdinchicite|correct|chemspider}} |
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|StdInChI = 1S/2Na.H2O4S/c;;1-5(2,3)4/h;;(H2,1,2,3,4)/q2*+1;/p-2 |
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}} |
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}} |
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| Section2 = {{Chembox Properties |
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| Section2 = {{Chembox Properties |
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| Formula = Na<sub>2</sub>SO<sub>4</sub> |
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|Formula = Na<sub>2</sub>SO<sub>4</sub> |
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| MolarMass = 142.04 g/mol (anhydrous)<br/>322.20 g/mol (decahydrate) |
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|MolarMass = 142.04 g/mol (anhydrous)<br>322.20 g/mol (decahydrate) |
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| Appearance = white crystalline solid <br> ] |
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|Appearance = white crystalline solid<br>] |
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|Odor = odorless |
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| Density = 2.664 g/cm<sup>3</sup> (anhydrous)<br/>1.464 g/cm<sup>3</sup> (decahydrate) |
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| Solubility = 47.6 g/L (0 °C)<br />427 g/L (100 °C) |
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|Density = 2.664 g/cm<sup>3</sup> (anhydrous)<br>1.464 g/cm<sup>3</sup> (decahydrate) |
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|Solubility = ''anhydrous:''<br>4.76 g/100 mL (0 °C)<br>28.1 g/100 mL (25 °C)<ref>National Center for Biotechnology Information. PubChem Compound Summary for CID 24436, Sodium sulfate. https://pubchem.ncbi.nlm.nih.gov/compound/Sodium-sulfate. Accessed Nov. 2, 2020.</ref><br>42.7 g/100 mL (100 °C) <hr> ''heptahydrate:''<br>19.5 g/100 mL (0 °C)<br>44 g/100 mL (20 °C) |
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| SolubleOther = insoluble in ] |
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|SolubleOther = insoluble in ]<br>soluble in ], ], and ] |
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| MeltingPt = 884 °C (anhydrous) <br> 32.4 °C (decahydrate) |
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|MeltingPtC = 884 |
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| BoilingPt = 1429 °C (anhydrous) |
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| RefractIndex = 1.468 (anhydrous) <br> 1.394 (decahydrate) |
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|MeltingPt_notes = (anhydrous)<br>32.38 °C (decahydrate) |
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|BoilingPtC = 1429 |
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|BoilingPt_notes = (anhydrous) |
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|RefractIndex = 1.468 (anhydrous)<br>1.394 (decahydrate) |
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|MagSus = −52.0·10<sup>−6</sup> cm<sup>3</sup>/mol |
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}} |
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}} |
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| Section3 = {{Chembox Structure |
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| Section3 = {{Chembox Structure |
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|CrystalStruct = ] (anhydrous)<ref>{{cite journal |vauthors=Zachariasen WH, Ziegler GE |title=The crystal structure of anhydrous sodium sulfate Na2SO4 |journal=Zeitschrift für Kristallographie, Kristallgeometrie, Kristallphysik, Kristallchemie |publisher=] |year=1932 |volume=81 |issue=1–6 |pages=92–101 |location=Wiesbaden |s2cid=102107891 |doi=10.1524/zkri.1932.81.1.92}}</ref><br>] (decahydrate) |
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| Coordination = |
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| CrystalStruct = ] or ] (anhydrous) <br> ] (decahydrate) |
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}} |
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}} |
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| Section7 = {{Chembox Hazards |
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| Section4 = {{Chembox Pharmacology |
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|ATCCode_prefix = A06 |
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| ExternalMSDS = |
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|ATCCode_suffix = AD13 |
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| MainHazards = Irritant |
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|ATC_Supplemental = {{ATC|A12|CA02}} |
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| EUIndex = Not listed |
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| NFPA-H = 1 |
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| NFPA-F = 0 |
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| NFPA-R = 0 |
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| FlashPt = Non-flammable |
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}} |
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}} |
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| Section8 = {{Chembox Related |
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| Section5 = {{Chembox Hazards |
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|ExternalSDS = |
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| OtherAnions = ]<br/>] |
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|MainHazards = Irritant |
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| OtherCations = ]<br/>]<br/>]<br/>] |
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|NFPA-H = 1 |
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| OtherCpds = ]<br/>]<br/>] |
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|NFPA-F = 0 |
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|NFPA-R = 0 |
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|FlashPt = Non-flammable |
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}} |
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}} |
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| Section6 = {{Chembox Related |
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|OtherAnions = ]<br>] |
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|OtherCations = ]<br>]<br>]<br>] |
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|OtherCompounds = ]<br>]<br>]<br>] |
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'''Sodium sulfate''' (also known as '''sodium sulphate''' or '''sulfate of soda''') is the ] with formula Na<sub>2</sub>SO<sub>4</sub> as well as several related ]s. All forms are white solids that are highly soluble in water. With an annual production of 6 million ]s, the decahydrate is a major ] chemical product. It is mainly used as a filler in the manufacture of powdered home ]s and in the ] of paper ] for making highly alkaline ]s.<ref>{{cite journal |author=Helmold Plessen |title=Sodium Sulfates |journal=Ullmann's Encyclopedia of Industrial Chemistry |publisher=Wiley-VCH |year=2000 |location=Weinheim |doi=10.1002/14356007.a24_355 |isbn=978-3527306732}}</ref> |
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==Forms== |
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*Anhydrous sodium sulfate, known as the rare mineral ], used as a drying agent in ]. |
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*Heptahydrate sodium sulfate, a very rare form. |
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*Decahydrate sodium sulfate, known as the mineral ], widely used by ]. It is also known as Glauber's salt. |
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==History== |
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The decahydrate of sodium sulfate is known as Glauber's salt after the ]–] chemist and ] ] (1604–1670), who discovered it in Austrian spring water in 1625. He named it {{lang|la|sal mirabilis}} (miraculous salt), because of its medicinal properties: the crystals were used as a general-purpose ], until more sophisticated alternatives came about in the 1900s.<ref name=szydlo>{{cite book |first=Zbigniew |last=Szydlo |author-link=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 |url-status=live |archive-url=https://web.archive.org/web/20111118122205/http://galileo.rice.edu/Catalog/NewFiles/glauber.html |archive-date=2011-11-18}}</ref> However, ] later alleged that it was known as a secret medicine in Saxony already in the mid-16th century.<ref>{{1911 Encyclopædia Britannica|no-prescript=1|wstitle=Glauber's Salt}}</ref> |
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In the 18th century, Glauber's salt began to be used as a raw material for the ] production of soda ash (]), by reaction with potash (]). Demand for soda ash increased, and the supply of sodium sulfate had to increase in line. Therefore, in the 19th century, the large-scale ], 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=11–16 |isbn=978-0-8122-1297-6}}</ref> |
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==Chemical properties== |
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Sodium sulfate is a typical electrostatically bonded ]ic sulfate. The existence of free sulfate ions in solution is indicated by the easy formation of insoluble sulfates when these solutions are treated with ] or ] salts: |
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:{{Chem2 | Na2SO4 + BaCl2 -> 2 NaCl + BaSO4 }} |
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Sodium sulfate is unreactive toward most ]. At high temperatures, it can be converted to ] by ] (aka thermo-chemical sulfate reduction (TSR), high temperature heating with charcoal, etc.):<ref name=crc>{{cite book |title=Handbook of Chemistry and Physics |url=https://archive.org/details/crchandbookofche00lide |url-access=registration |edition=71st |publisher=] |location=Ann Arbor, Michigan |year=1990 |isbn=9780849304712}}</ref> |
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:{{Chem2 | Na2SO4 + 2 C -> Na2S + 2 CO2 }} |
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This reaction was employed in the ], a defunct industrial route to ]. |
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Sodium sulfate reacts with sulfuric acid to give the ] ]:<ref name=merck>{{cite book |title=The Merck Index |edition=7th |publisher=] |location=Rahway, New Jersey, US |year=1960 |title-link=Merck Index}}</ref><ref>{{cite book |first=Howard |last=Nechamkin |title=The Chemistry of the Elements |url=https://archive.org/details/chemistryofeleme00nech |url-access=registration |publisher=] |location=New York |year=1968}}</ref> |
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:{{Chem2 | Na2SO4 + H2SO4 <-> 2 NaHSO4 }} |
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Sodium sulfate displays a moderate tendency to form ]s. The only ]s formed with common trivalent metals are ] (unstable above 39 °C) and NaCr(SO<sub>4</sub>)<sub>2</sub>, in contrast to ] and ] which form many stable alums.<ref name=Lipson1935>{{cite journal |last1=Lipson |first1=Henry |author-link1=Henry Lipson |first2=C. A. |last2=Beevers |author-link2=C. Arnold Beevers |year=1935 |title=The Crystal Structure of the Alums |journal=] |volume=148 |issue=865 |pages=664–80 |doi=10.1098/rspa.1935.0040 |bibcode=1935RSPSA.148..664L |doi-access=free}}</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 ]. Formation of ] by reaction of sodium sulfate with ] has been used as the basis of a method for producing ], a ].<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=978-0-12-276151-5}}</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> (]) 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=II |publisher=Longmans |year=1961 |edition=new impression |location=London |pages=656–673 |isbn=978-0-582-46277-9}}</ref> |
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==Physical properties== |
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Sodium sulfate has unusual solubility characteristics in water.<ref>{{cite book |first=W. F. |last=Linke |author2=A. Seidell |title=Solubilities of Inorganic and Metal Organic Compounds |edition=4th |publisher=Van Nostrand |year=1965 |isbn=978-0-8412-0097-5}}</ref> Its solubility in water rises more than tenfold between 0 °C and 32.384 °C, where it reaches a maximum of 49.7 g/100 mL. At this point the solubility curve changes slope, and the solubility becomes almost independent of temperature. This temperature of 32.384 °C, corresponding to the release of crystal water and melting of the hydrated salt, serves as an accurate temperature reference for thermometer ]. |
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] |
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==Structure== |
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Crystals of the decahydrate consist of <sup>+</sup> ions with ]. These octahedra share edges such that 8 of the 10 water molecules are bound to sodium and 2 others are interstitial, being hydrogen-bonded to sulfate. These cations are linked to the sulfate anions by ]s. The Na–O distances are about 240 ].<ref>Helena W. Ruben, David H. Templeton, Robert D. Rosenstein, Ivar Olovsson, "Crystal Structure and Entropy of Sodium Sulfate Decahydrate", J. Am. Chem. Soc. 1961, volume 83, pp. 820–824. {{doi|10.1021/ja01465a019}}.</ref> Crystalline sodium sulfate decahydrate is also unusual among hydrated salts in having a measurable ] (entropy at ]) of 6.32 J/(K·mol). 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. |author2=W. F. Giauque |title=The Heat of Hydration of Sodium Sulfate. Low Temperature Heat Capacity and Entropy of Sodium Sulfate Decahydrate |journal=] |volume=80 |issue=9 |pages=2042–2044 |year=1958 |doi=10.1021/ja01542a003}}</ref> |
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==Production== |
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The world production of sodium sulfate, almost exclusively 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.<ref name=ceh>{{cite book |last=Suresh |first=Bala |author2=Kazuteru Yokose |title=Sodium sulfate |url=http://www.sriconsulting.com/CEH/Public/Reports/771.1000/?Abstract.html |location=Zurich |publisher=Chemical Economic Handbook SRI Consulting |date=May 2006 |pages=771.1000A–771.1002J |work=CEH Marketing Research Report |url-status=live |archive-url=https://web.archive.org/web/20070314084954/http://www.sriconsulting.com/CEH/Public/Reports/771.1000/?Abstract.html |archive-date=2007-03-14}}</ref><ref name=usgs>{{cite web |url=http://minerals.usgs.gov/minerals/pubs/commodity/sodium_sulfate/stat |title=Statistical compendium Sodium sulfate |publisher=], Minerals Information |location=Reston, Virginia |year=1997 |access-date=2007-04-22 |url-status=live |archive-url=https://web.archive.org/web/20070307171936/http://minerals.usgs.gov/minerals/pubs/commodity/sodium_sulfate/stat/ |archive-date=2007-03-07}}</ref><ref name=roskill>{{cite book |title=The economics of sodium sulphate |edition=Eighth |year=1999 |location=London |publisher=Roskill Information Services}}</ref><ref name=chemsys>{{cite book |title=The sodium sulphate business |date=November 1984 |location=London |publisher=Chem Systems International}}</ref> For all applications, naturally produced and chemically produced sodium sulfate are practically interchangeable. |
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===Natural sources=== |
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Two thirds of the world's production of the decahydrate (Glauber's salt) is from the natural mineral form ], for example as found in lake beds in southern ]. In 1990, ] and ] were the world's main producers of natural sodium sulfate (each around 500,000 ]s), with ], ], and ] around 350,000 tonnes each.<ref name=usgs/> Natural resources are estimated at over 1 billion tonnes.<ref name=ceh/><ref name=usgs/> |
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Major producers of 200,000 to 1,500,000 tonnes/year in 2006 included ] (California, US), Airborne Industrial Minerals (Saskatchewan, Canada), ] (Coahuila, Mexico), Minera de Santa Marta and Criaderos Minerales Y Derivados, also known as ] (Burgos, Spain), Minera de Santa Marta (Toledo, Spain), Sulquisa (Madrid, Spain), Chengdu Sanlian Tianquan Chemical (], Sichuan, China), Hongze Yinzhu Chemical Group (], Jiangsu, China), {{ill|Nafine Chemical Industry Group|lt=|zh|南风化工}} (Shanxi, China), Sichuan Province Chuanmei Mirabilite ({{ill|万胜镇|lt=|zh|万胜镇}}, ], ], Sichuan, China), and Kuchuksulphat JSC (Altai Krai, Siberia, Russia).<ref name=ceh/><ref name=roskill/> |
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Anhydrous sodium sulfate occurs in arid environments as the mineral ]. It slowly turns to mirabilite in damp air. Sodium sulfate is also found as ], a calcium sodium sulfate mineral. Both minerals are less common than mirabilite.{{citation needed|date=February 2017}} |
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===Chemical industry=== |
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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. |
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The most important chemical sodium sulfate production is during ] production, either from ] (salt) and ], in the ], or from ] in the ].<ref name=kirk-othmer>{{cite book |first=D. |last=Butts |title=Kirk-Othmer Encyclopedia of Chemical Technology |edition=4th |volume=v22 |pages=403–411 |year=1997}}</ref> The resulting sodium sulfate from these processes is known as '''''salt cake'''''. |
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:Mannheim: {{Chem2 | 2 NaCl + H2SO4 -> 2 HCl + Na2SO4 }} |
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:Hargreaves: {{Chem2 | 4 NaCl + 2 SO2 + O2 + 2 H2O -> 4 HCl + 2 Na2SO4 }} |
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The second major production of sodium sulfate are the processes where surplus ] is ] by sulfuric acid to obtain ] ({{Chem2|SO4(2-)}}) by using ] (CuSO<sub>4</sub>) (as historically applied on a large scale in the production of ] by using ]). This method is also a regularly applied and convenient laboratory preparation. |
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:{{Chem2 | 2 NaOH(aq) + H2SO4(aq) -> Na2SO4(aq) + 2 H2O(l) }} ΔH = -112.5 kJ (highly exothermic) |
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In the laboratory it can also be synthesized from the reaction between ] and ], by precipitating ]. |
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:{{Chem2 | 2 NaHCO3 + MgSO4 -> Na2SO4 + MgCO3 + CO2 + H2O }} |
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However, as commercial sources are readily available, laboratory synthesis is not practised often. |
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Formerly, sodium sulfate was also a by-product of the manufacture of ], 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 ], ]s, ], ], ] pigments, ], and ].<ref name=ceh/> |
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Bulk sodium sulfate is usually purified via the decahydrate form, since the anhydrous form tends to attract ] compounds and ]s. The anhydrous form is easily produced from the hydrated form by gentle warming. |
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Major sodium sulfate by-product producers of 50–80 Mt/a in 2006 include Elementis Chromium (chromium industry, Castle Hayne, NC, US), 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/> |
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==Applications== |
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] |
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] |
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===Commodity industries=== |
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With US pricing at $30 per tonne in 1970, up to $90 per tonne for salt cake quality, and $130 for better grades, sodium sulphate is a very cheap material. The largest use is as ] in powdered home ]s, consuming approximately 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/> |
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===Papermaking=== |
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Another formerly major use for sodium sulfate, notably in the US and Canada, is in the ] for the manufacture of ]. Organics present in the "black liquor" from this process are burnt to produce heat, needed to drive the ] of sodium sulfate to ]. However, due to advances in the thermal efficiency of the Kraft recovery process in the early 1960s, more efficient sulfur recovery was achieved and the need for sodium sulfate makeup was drastically reduced.<ref>{{cite book |last=Smook |first=Gary |title=Handbook for Pulp and Paper Technologists |url=http://imisrise.tappi.org/TAPPI/Products/02/SMO/0202SMOOK4.aspx |year=2002 |page=143 |url-status=live |archive-url=https://web.archive.org/web/20160807043026/http://imisrise.tappi.org/TAPPI/Products/02/SMO/0202SMOOK4.aspx |archive-date=2016-08-07}}</ref> Hence, the use of sodium sulfate in the US and Canadian pulp industry declined from 1,400,000 tonnes per year in 1970 to only approx. 150,000 tonnes in 2006.<ref name=ceh/> |
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===Glassmaking=== |
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The ] industry provides another significant application for sodium sulfate, as second largest application in Europe. Sodium sulfate is used as a ], 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.<ref name=ceh/> |
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===Textiles=== |
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Sodium sulfate is important in the manufacture of ]s, particularly in Japan, where this is the largest application. Sodium sulfate is added to increase the ] of the solution and so helps in "levelling", i.e. reducing negative electrical charges on textile fibres, so that dyes can penetrate evenly (see the theory of the ] (DDL) elaborated by ]). Unlike the alternative ], it does not corrode the ] vessels used in dyeing. This application in Japan and US consumed in 2006 approximately 100,000 tonnes.<ref name=ceh/> |
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===Food industry=== |
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Sodium sulfate is used as a diluent for food colours.<ref name=WHO2000/> It is known as ] additive '''E514'''. |
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===Heat storage=== |
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The high heat-storage capacity in the phase change from solid to liquid, and the advantageous phase change temperature of {{cvt|32|C}} 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 solar–heated water. The phase change allows a substantial reduction in the mass of the material required for effective heat storage (the heat of fusion of sodium sulfate decahydrate is 82 kJ/mol or 252 kJ/kg<ref>{{cite web |title=Phase-Change Materials for Low-Temperature Solar Thermal Applications |url=https://www.eng.mie-u.ac.jp/research/activities/29/29_31.pdf |access-date=2014-06-19 |url-status=live |archive-url=https://web.archive.org/web/20150924000749/http://www.eng.mie-u.ac.jp/research/activities/29/29_31.pdf |archive-date=2015-09-24}}</ref>), with the further advantage of a consistency of temperature as long as sufficient material in the appropriate phase is available. |
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For cooling applications, a mixture with common ] salt (NaCl) lowers the melting point to {{cvt|18|C}}. The heat of fusion of NaCl·Na<sub>2</sub>SO<sub>4</sub>·10H<sub>2</sub>O, is actually ''increased'' slightly to 286 kJ/kg.<ref>{{cite web |title=Phase-Change Materials for Low-Temperature Solar Thermal Applications |page=8 |url=https://www.eng.mie-u.ac.jp/research/activities/29/29_31.pdf |access-date=2014-06-19 |url-status=live |archive-url=https://web.archive.org/web/20150924000749/http://www.eng.mie-u.ac.jp/research/activities/29/29_31.pdf |archive-date=2015-09-24}}</ref> |
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===Small-scale applications=== |
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In the laboratory, anhydrous sodium sulfate is widely used as an inert ], for removing traces of water from organic solutions.<ref name=vogel>{{cite book |last=Vogel |first=Arthur I. |author2=B.V. Smith |author3=N.M. Waldron |edition=3rd |title=Vogel's Elementary Practical Organic Chemistry 1 Preparations |publisher=] Scientific & Technical |location=London |year=1980}}</ref> It is more efficient, but slower-acting, than the similar agent ]. It is only effective below about {{cvt|30|C}}, but it can be 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. |
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Glauber's salt, the decahydrate, is used as a ]. It is effective for the removal of certain drugs, such as ] (acetaminophen) from the body; thus it can be used after an overdose.<ref name=Cocchetto1981>{{cite journal |last=Cocchetto |first=D.M. |author2=G. Levy |year=1981 |title=Absorption of orally administered sodium sulfate in humans |journal=J Pharm Sci |volume=70 |issue=3 |pages=331–3 |doi=10.1002/jps.2600700330 |pmid=7264905}}</ref><ref name=Prescott1979>{{cite journal |last1=Prescott |first1=L. F. |first2=J. A. J. H. |last2=Critchley |year=1979 |title=The Treatment of Acetaminophen Poisoning |journal=Annual Review of Pharmacology and Toxicology |volume=23 |pages=87–101 |doi=10.1146/annurev.pa.23.040183.000511 |pmid=6347057}}</ref> |
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In 1953, sodium sulfate was proposed for heat storage in passive ] systems. This takes advantage of its unusual solubility properties, and the high heat of ] (78.2 kJ/mol).<ref>{{cite book |last=Telkes |first=Maria |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> |
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Other uses for sodium sulfate include de-frosting windows, ] manufacture, as an additive in carpet fresheners, and as an additive to cattle feed. |
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At least one company, Thermaltake, makes a laptop computer chill mat (iXoft Notebook Cooler) using sodium sulfate decahydrate inside a quilted plastic pad. The material slowly turns to liquid and recirculates, equalizing laptop temperature and acting as an insulation.<ref>{{cite web |url=http://www.thermaltake.com/products-model_Specification.aspx?id=C_00000712 |title=IXoft Specification |publisher=Thermaltake Technology Co., Ltd. |access-date=2015-08-15 |url-status=live |archive-url=https://web.archive.org/web/20160312234730/http://www.thermaltake.com/products-model_Specification.aspx?id=C_00000712 |archive-date=2016-03-12}}</ref> |
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==Safety== |
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Although sodium sulfate is generally regarded as non-toxic,<ref name=WHO2000>{{cite web |title=Sodium sulfate (WHO Food Additives Series 44) |publisher=] |year=2000 |url=http://www.inchem.org/documents/jecfa/jecmono/v44jec07.htm |access-date=2007-06-06 |url-status=live |archive-url=https://web.archive.org/web/20070904064342/http://www.inchem.org/documents/jecfa/jecmono/v44jec07.htm |archive-date=2007-09-04}}</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 ] or ] applies.<ref name=msds>{{cite web |title=MSDS Sodium Sulfate Anhydrous |publisher=James T Baker |year=2006 |access-date=2007-04-21 |url=http://hazard.com/msds/mf/baker/baker/files/s5022.htm |url-status=usurped |archive-url=https://web.archive.org/web/20030619125307/http://hazard.com/msds/mf/baker/baker/files/s5022.htm |archive-date=2003-06-19}}</ref> |
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==References== |
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{{Reflist|30em}} |
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==External links== |
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*Calculators: , and of aqueous sodium sulfate |
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{{Sodium compounds}} |
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{{Sulfates}} |
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{{Authority control}} |
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] |
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] |
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] |
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] |
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] |