Wikipedia:WikiProject Chemicals/Chembox validation/VerifiedDataSandbox and Sodium sulfate: Difference between pages

{{Short description|Chemical compound with formula Na₂SO₄}}

<!-- Spelling: this article is in EN-GB spelling. Note that, chemically, sulfate is with an 'f', even in EN-GB -->

{{Chembox

| Verifiedfields =

| Watchedfields = changed

| verifiedrevid = 477315271

| Name = Sodium sulfate

| ImageFileR1 = Sodium sulfate.jpg

| ImageFileL1 = Sodium sulfate.svg

| ImageSizeL1 = 150px

| ImageSizeR1 = 120px

| ImageName = Sodium sulphate

| IUPACName = Sodium sulfate

| OtherNames = Sodium sulphate<br>Disodium sulfate<br>Sulfate of sodium<br>[[Thénardite]] (anhydrous mineral)<br>Glauber's salt (decahydrate)<br>''Sal mirabilis'' (decahydrate)<br>[[Mirabilite]] (decahydrate mineral)

| Section1 = {{Chembox Identifiers

|UNII_Ref = {{fdacite|correct|FDA}}

|UNII = 36KCS0R750

|UNII2_Ref = {{fdacite|correct|FDA}}

|UNII2 = 0YPR65R21J

|UNII2_Comment = (decahydrate

|InChI = 1S/2Na.H2O4S/c;;1-5(2,3)4/h;;(H2,1,2,3,4)/q2*+1;/p-2

|InChIKey1 = PMZURENOXWZQFD-UHFFFAOYSA-L

|InChI1 = 1S/2Na.H2O4S/c;;1-5(2,3)4/h;;(H2,1,2,3,4)/q2*+1;/p-2

|CASNo = 7757-82-6

|CASNo_Ref = {{cascite|correct|CAS}}

|CASNo2_Ref = {{cascite|correct|CAS}}

|CASNo2 = 7727-73-3

|CASNo2_Comment = (decahydrate)

|ChEMBL_Ref = {{ebicite|correct|EBI}}

|ChEMBL = 233406

|PubChem = 24436

|RTECS = WE1650000

|ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}

|ChemSpiderID = 22844

|ChEBI_Ref = {{ebicite|correct|EBI}}

|ChEBI = 32149

|StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}

|StdInChIKey = PMZURENOXWZQFD-UHFFFAOYSA-L

|SMILES = [Na+].[Na+].[O-]S([O-])(=O)=O

|StdInChI_Ref = {{stdinchicite|correct|chemspider}}

|StdInChI = 1S/2Na.H2O4S/c;;1-5(2,3)4/h;;(H2,1,2,3,4)/q2*+1;/p-2

| Section2 = {{Chembox Properties

|Formula = Na₂SO₄

|MolarMass = 142.04 g/mol (anhydrous)<br>322.20 g/mol (decahydrate)

|Appearance = white crystalline solid<br>[[hygroscopic]]

|Odor = odorless

|Density = 2.664 g/cm³ (anhydrous)<br>1.464&nbsp;g/cm³ (decahydrate)

|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&nbsp;°C) <hr> ''heptahydrate:''<br>19.5 g/100 mL (0 °C)<br>44 g/100 mL (20 °C)

|SolubleOther = insoluble in [[ethanol]]<br>soluble in [[glycerol]], [[water]], and [[hydrogen iodide]]

|MeltingPtC = 884

|MeltingPt_notes = (anhydrous)<br>32.38 °C (decahydrate)

|BoilingPtC = 1429

|BoilingPt_notes = (anhydrous)

|RefractIndex = 1.468 (anhydrous)<br>1.394 (decahydrate)

|MagSus = &minus;52.0·10⁻⁶ cm³/mol

| Section3 = {{Chembox Structure

|CrystalStruct = [[orthorhombic]] (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=[[Akademische Verlagsgesellschaft]] |year=1932 |volume=81 |issue=1–6 |pages=92–101 |location=Wiesbaden |s2cid=102107891 |doi=10.1524/zkri.1932.81.1.92}}</ref><br>[[monoclinic]] (decahydrate)

| Section4 = {{Chembox Pharmacology

|ATCCode_prefix = A06

|ATCCode_suffix = AD13

|ATC_Supplemental = {{ATC|A12|CA02}}

| Section5 = {{Chembox Hazards

|ExternalSDS = [http://www.ilo.org/public/english/protection/safework/cis/products/icsc/dtasht/_icsc09/icsc0952.htm ICSC 0952]

|MainHazards = Irritant

|NFPA-H = 1

|NFPA-F = 0

|NFPA-R = 0

|FlashPt = Non-flammable

| Section6 = {{Chembox Related

|OtherAnions = [[Sodium selenate]]<br>[[Sodium tellurate]]

|OtherCations = [[Lithium sulfate]]<br>[[Potassium sulfate]]<br>[[Rubidium sulfate]]<br>[[Caesium sulfate]]

|OtherCompounds = [[Sodium bisulfate]]<br>[[Sodium sulfite]]<br>[[Sodium persulfate]]<br>[[Sodium pyrosulfate]]

}}

'''Sodium sulfate''' (also known as '''sodium sulphate''' or '''sulfate of soda''') is the [[inorganic compound]] with formula Na₂SO₄ as well as several related [[hydrate]]s. All forms are white solids that are highly soluble in water. With an annual production of 6 million [[tonne]]s, the decahydrate is a major [[commodity]] chemical product. It is mainly used as a filler in the manufacture of powdered home [[laundry detergent]]s and in the [[Kraft process]] of paper [[pulping]] for making highly alkaline [[sulfide]]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>

==Forms==

*Anhydrous sodium sulfate, known as the rare mineral [[thenardite]], used as a drying agent in [[organic synthesis]].

*Heptahydrate sodium sulfate, a very rare form.

*Decahydrate sodium sulfate, known as the mineral [[mirabilite]], widely used by [[chemical industry]]. It is also known as Glauber's salt.

==History==

The decahydrate of sodium sulfate is known as Glauber's salt after the [[Netherlands|Dutch]]–[[Germany|German]] chemist and [[apothecary]] [[Johann Rudolf Glauber]] (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 [[laxative]], 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, [[Johann von Löwenstern-Kunckel|J. Kunckel]] 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>

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]]). Demand for soda ash increased, and the supply of sodium sulfate had to increase in line. Therefore, in the 19th century, the large-scale [[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=11–16 |isbn=978-0-8122-1297-6}}</ref>

==Chemical properties==

Sodium sulfate is a typical electrostatically bonded [[ion]]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 [[barium|Ba²⁺]] or [[lead|Pb²⁺]] salts:

:{{Chem2 | Na2SO4 + BaCl2 -> 2 NaCl + BaSO4 }}

Sodium sulfate is unreactive toward most [[redox|oxidizing or reducing agents]]. At high temperatures, it can be converted to [[sodium sulfide]] by [[carbothermal reduction]] (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=[[CRC Press]] |location=Ann Arbor, Michigan |year=1990 |isbn=9780849304712}}</ref>

:{{Chem2 | Na2SO4 + 2 C -> Na2S + 2 CO2 }}

This reaction was employed in the [[Leblanc process]], a defunct industrial route to [[sodium carbonate]].

Sodium sulfate reacts with sulfuric acid to give the [[acid salt]] [[sodium bisulfate]]:<ref name=merck>{{cite book |title=The Merck Index |edition=7th |publisher=[[Merck & Co.]] |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=[[McGraw-Hill]] |location=New York |year=1968}}</ref>

:{{Chem2 | Na2SO4 + H2SO4 <-> 2 NaHSO4 }}

Sodium sulfate displays a moderate tendency to form [[double salt]]s. The only [[alum]]s formed with common trivalent metals are [[sodium alum|NaAl(SO₄)₂]] (unstable above 39&nbsp;°C) and NaCr(SO₄)₂, in contrast to [[potassium sulfate]] and [[ammonium sulfate]] 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=[[Proceedings of the Royal Society A]] |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 Na₂SO₄·3K₂SO₄ which occurs naturally as the mineral [[aphthitalite]]. 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=978-0-12-276151-5}}</ref> Other double salts include 3Na₂SO₄·CaSO₄, 3Na₂SO₄·MgSO₄ ([[vanthoffite]]) and NaF·Na₂SO₄.<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>

==Physical properties==

Sodium sulfate has unusual solubility characteristics in water.<ref>{{cite book |first=W.&nbsp;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&nbsp;°C and 32.384&nbsp;°C, where it reaches a maximum of 49.7&nbsp;g/100&nbsp;mL. At this point the solubility curve changes slope, and the solubility becomes almost independent of temperature. This temperature of 32.384&nbsp;°C, corresponding to the release of crystal water and melting of the hydrated salt, serves as an accurate temperature reference for thermometer [[calibration]].

[[File:Na2SO4 solubility.svg|thumb|Temperature dependence of Na₂SO₄ solubility in water]]

==Structure==

Crystals of the decahydrate consist of [Na(OH₂)₆]⁺ ions with [[octahedral molecular geometry]]. 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 [[hydrogen bond]]s. The Na–O distances are about 240&nbsp;[[picometer|pm]].<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 [[residual entropy]] (entropy at [[absolute zero]]) of 6.32&nbsp;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.&nbsp;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 |issue=9 |pages=2042–2044 |year=1958 |doi=10.1021/ja01542a003}}</ref>

==Production==

The world production of sodium sulfate, almost exclusively 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 |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=[[US Geological Survey]], 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.

===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]], [[United States]], and [[Canada]] around 350,000&nbsp;tonnes each.<ref name=usgs/> Natural resources are estimated at over 1 billion tonnes.<ref name=ceh/><ref name=usgs/>

Major producers of 200,000 to 1,500,000 tonnes/year in 2006 included [[Searles Valley Minerals]] (California, US), Airborne Industrial Minerals (Saskatchewan, Canada), [[Química del Rey]] (Coahuila, Mexico), Minera de Santa Marta and Criaderos Minerales Y Derivados, also known as [[Grupo Crimidesa]] (Burgos, Spain), Minera de Santa Marta (Toledo, Spain), Sulquisa (Madrid, Spain), Chengdu Sanlian Tianquan Chemical ([[Tianquan County]], Sichuan, China), Hongze Yinzhu Chemical Group ([[Hongze District]], 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}}, [[Dongpo District]], [[Meishan]], Sichuan, China), 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.{{citation needed|date=February 2017}}

===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 [[James Hargreaves (chemist)|Hargreaves process]].<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'''''.

:Mannheim: {{Chem2 | 2 NaCl + H2SO4 -> 2 HCl + Na2SO4 }}

:Hargreaves: {{Chem2 | 4 NaCl + 2 SO2 + O2 + 2 H2O -> 4 HCl + 2 Na2SO4 }}

The second major production of sodium sulfate are the processes where surplus [[sodium hydroxide]] is [[neutralization (chemistry)|neutralised]] by sulfuric acid to obtain [[sulfate]] ({{Chem2|SO4(2-)}}) by using [[copper(II) sulfate|copper sulfate]] (CuSO₄) (as historically applied on a large scale in the production of [[rayon]] by using [[copper(II) hydroxide]]). This method is also a regularly applied and convenient laboratory preparation.

:{{Chem2 | 2 NaOH(aq) + H2SO4(aq) -> Na2SO4(aq) + 2 H2O(l) }}&nbsp;&nbsp;&nbsp;&nbsp;ΔH = -112.5 kJ (highly exothermic)

In the laboratory it can also be synthesized from the reaction between [[sodium bicarbonate]] and [[magnesium sulfate]], by precipitating [[magnesium carbonate]].

:{{Chem2 | 2 NaHCO3 + MgSO4 -> Na2SO4 + MgCO3 + CO2 + H2O }}

However, as commercial sources are readily available, laboratory synthesis is not practised often.

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 50–80 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/>

==Applications==

[[File:Sulfate clump.ogv|thumb|Sodium sulfate used to dry an organic liquid. Here clumps form, indicating the presence of water in the organic liquid.]]

[[File:Sulfate noclump.ogg|thumb|By further application of sodium sulfate the liquid may be brought to dryness, indicated here by the absence of clumping.]]

===Commodity industries===

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 [[filler (materials)|filler]] in powdered home [[laundry detergent]]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/>

===Papermaking===

Another formerly major use for sodium sulfate, notably in the US 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, 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&nbsp;tonnes in 2006.<ref name=ceh/>

===Glassmaking===

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/>

===Textiles===

Sodium sulfate is important in the manufacture of [[textile]]s, particularly in Japan, where this is the largest application. Sodium sulfate is added to increase the [[ionic strength]] 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 [[diffuse double layer]] (DDL) elaborated by [[double layer (surface science)#Gouy–Chapman|Gouy and Chapman]]). Unlike the alternative [[sodium chloride]], it does not corrode the [[stainless steel]] vessels used in dyeing. This application in Japan and US consumed in 2006 approximately 100,000&nbsp;tonnes.<ref name=ceh/>

===Food industry===

Sodium sulfate is used as a diluent for food colours.<ref name=WHO2000/> It is known as [[E number]] additive '''E514'''.

===Heat storage===

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.

For cooling applications, a mixture with common [[sodium chloride]] salt (NaCl) lowers the melting point to {{cvt|18|C}}. The heat of fusion of NaCl·Na₂SO₄·10H₂O, 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>

===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. |author2=B.V. Smith |author3=N.M. Waldron |edition=3rd |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 {{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.

Glauber's salt, the decahydrate, is used as a [[laxative]]. It is effective for the removal of certain drugs, such as [[paracetamol]] (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>

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 |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>

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

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>

==Safety==

Although sodium sulfate is generally regarded as non-toxic,<ref name=WHO2000>{{cite web |title=Sodium sulfate (WHO Food Additives Series 44) |publisher=[[World Health Organization]] |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 [[list of R-phrases|Risk Phrase]] or [[list of S-phrases|Safety Phrase]] 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>

==References==

{{Reflist|30em}}

==External links==

*Calculators: [http://www.aim.env.uea.ac.uk/aim/surftens/surftens.php surface tensions], and [http://www.aim.env.uea.ac.uk/aim/density/density_electrolyte.php densities, molarities, and molalities] of aqueous sodium sulfate

{{Sodium compounds}}

{{Sulfates}}

{{Authority control}}

[[Category:Sodium compounds]]

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