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Sodium sulfate

-->{{chembox | Verifiedfields = changed | Watchedfields = changed | verifiedrevid = 477315271 | Name = Sodium sulfate | ImageFileL1 = Sodium sulfate.jpg | ImageFileR1 = Sodium sulfate.png | ImageSize = 150px | ImageName = Sodium sulfate | OtherNames = Thenardite (mineral)Glauber's salt (decahydrate)Sal mirabilis (decahydrate) Mirabilite (decahydrate)Disodium sulfate |Section1={{Chembox Identifiers | UNII_Ref = | UNII = 36KCS0R750 | 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 = | CASNo2_Ref = | CASNo2 = 7727-73-3 | CASNo2_Comment = (decahydrate) | ChEMBL_Ref = | ChEMBL = 233406 | PubChem = 24436 | RTECS = WE1650000 | ChemSpiderID_Ref = | ChemSpiderID = 22844 | ChEBI_Ref = | ChEBI = 32149 | StdInChIKey_Ref = | StdInChIKey = PMZURENOXWZQFD-UHFFFAOYSA-L | SMILES = Na+. Na+. O-S( O-)(=O)=O | StdInChI_Ref = | 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)322.20 g/mol (decahydrate) | Appearance = white crystalline solid hygroscopic | Odor = odorless | Density = 2.664 g/cm3 (anhydrous)1.464 g/cm3 (decahydrate) | Solubility = anhydrous: 4.76 g/100 mL (0 °C)13.9 g/100 mL (20 °C)42.7 g/100 mL (100 °C) heptahydrate: 19.5 g/100 mL (0 °C) 44 g/100 mL (20 °C) | SolubleOther = insoluble in ethanol soluble in glycerol, water and hydrogen iodide | MeltingPtC = 884 | MeltingPt_notes = (anhydrous) 32.38 °C (decahydrate) | BoilingPtC = 1429 | BoilingPt_notes = (anhydrous) | RefractIndex = 1.468 (anhydrous) 1.394 (decahydrate) | MagSus = −52.0·10−6 cm3/mol }} |Section3={{Chembox Structure | Coordination = | CrystalStruct = orthorhombic or hexagonal (anhydrous) monoclinic (decahydrate) }} |Section6={{Chembox Pharmacology | ATCCode_prefix = A06 | ATCCode_suffix = AD13 | ATC_Supplemental = }} |Section7={{Chembox Hazards | ExternalSDS = ICSC 0952 | MainHazards = Irritant | NFPA-H = 1 | NFPA-F = 0 | NFPA-R = 0 | FlashPt = Non-flammable }} |Section8={{Chembox Related | OtherAnions = Sodium selenate Sodium tellurate | OtherCations = Lithium sulfate Potassium sulfate Rubidium sulfate Caesium sulfate | OtherCompounds = Sodium bisulfate Sodium sulfite Sodium persulfate }} }} Sodium sulfate, also known as sulfate of soda, is the inorganic compound with formula Na2SO4 as well as several related hydrates. All forms are white solids that are highly soluble in water. With an annual production of 6 million tonnes, the decahydrate is a major commodity chemical product. It is mainly used for the manufacture of detergents and in the kraft process of paper pulping.Helmold Plessen "Sodium Sulfates" in Ullmann's Encyclopedia Of Industrial Chemistry Wiley-VCH, 2000, Weinheim.

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 Dutch/ German chemist and apothecary Johann Rudolf Glauber (1604–1670), who discovered it in 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 the supply of sodium sulfate had to increase in line. Therefore, in the nineteenth century, the large scale 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 very stable, being unreactive toward most oxidizing or reducing agents at normal temperatures. At high temperatures, it can be converted to sodium sulfide by carbothermal reduction: Na2SO4 + 2 C → Na2S + 2 CO2

Acid-base

Sodium sulfate is a neutral salt: its aqueous solutions exhibit a pH of 7. The neutrality of such solutions reflects the fact that sulfate is derived, formally, from the strong acid sulfuric acid. Furthermore, the Na+ ion, with only a single positive charge, only weakly polarizes its water ligands provided there are metal ions in solution. Sodium sulfate reacts with sulfuric acid to give the acid salt sodium bisulfate: Na2SO4 + H2SO4 ⇌ 2 NaHSO4 The equilibrium constant for this process depends on concentration and temperature.

Solution and ion exchange

Sodium sulfate has unusual solubility characteristics in water. Its solubility in water rises more than tenfold between 0 °C to 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 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 is a typical electrostatically bonded ionic sulfate, containing Na+ ions and SO42− ions. The existence of sulfate in solution is indicated by the easy formation of insoluble sulfates when these solutions are treated with Ba2+ or Pb2+ salts: Na2SO4 + BaCl2 → 2 NaCl + BaSO4 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 Na2SO4·3K2SO4 which occurs naturally as the mineral glaserite. Formation of glaserite by reaction of sodium sulfate with potassium chloride has been used as the basis of a method for producing potassium sulfate, a fertiliser. Other double salts include 3Na2SO4·CaSO4, 3Na2SO4·MgSO4 ( vanthoffite) and NaF·Na2SO4.

Structure

Crystals consist of Na(OH2)6+ ions with octahedral molecular geometry. These octahedral share edges such that eight of the 10 water molecules are bound to sodium and two others are interstitial, being hydrogen bonded to sulfate. These cations are linked to the sulfate anions via hydrogen bonds. The Na-O distances are 240 pm.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. Crystalline sodium sulfate decahydrate is also unusual among hydrated salts in having a measureable residual entropy (entropy at absolute zero) of 6.32 J·K−1·mol−1. This is ascribed to its ability to distribute water much more rapidly compared to most hydrates.

Production

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. 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, United States and Canada around 350,000 tonnes each. Natural resources are estimated at over 1 billion tonnes. 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 (Sichuan, China), Hongze Yinzhu Chemical Group (Jiangsu, China), Nafine Chemical Industry Group (Shanxi, China), Sichuan Province Chuanmei Mirabilite (Sichuan, China), 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 is 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 sodium hydroxide is neutralised by sulfuric acid, 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) In the laboratory it can also be synthesized from the reaction between sodium bicarbonate and magnesium sulfate. 2NaHCO3 + MgSO4 → Na2SO4 + Mg(OH)2 + 2CO2 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 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).

Applications

Commodity industries

With US pricing at $30 per tonne in 1970,6 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 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 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 . Hence, the use of sodium sulfate in the US 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 US 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 °C (90 °F) 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/kghttp://www.eng.mie-u.ac.jp/research/activities/29/29_31.pdf), 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 18 °C (64 °F). The heat of fusion of NaCl·Na2SO4·10H2O, is actually increased slightly to 286 kJ/kg.http://www.eng.mie-u.ac.jp/research/activities/29/29_31.pdf p.8

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 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, was historically used as a laxative. It is effective for the removal of certain drugs such as paracetamol (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, 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.

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 applies.

References

External links

"green air" © 2007 - Ingo Malchow, Webdesign Neustrelitz
This article based upon the http://en.wikipedia.org/wiki/Sodium_sulfate, the free encyclopaedia Wikipedia and is licensed under the GNU Free Documentation License.
Further informations available on the list of authors and history: http://en.wikipedia.org/w/index.php?title=Sodium_sulfate&action=history
presented by: Ingo Malchow, Mirower Bogen 22, 17235 Neustrelitz, Germany