Is Naso4 Soluble In Water
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| Names | |||
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| Other names Sodium sulphate | |||
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| ECHA InfoCard | 100.028.928 | ||
| E number | E514(i) (acidity regulators, ...) | ||
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| Properties | |||
| Chemical formula | Na2SO4 | ||
| Tooth mass | 142.04 g/mol (anhydrous) 322.20 yard/mol (decahydrate) | ||
| Appearance | white crystalline solid hygroscopic | ||
| Odor | odorless | ||
| Density | 2.664 g/cm3 (anhydrous) ane.464 one thousand/cm3 (decahydrate) | ||
| Melting point | 884 °C (1,623 °F; 1,157 K) (anhydrous) 32.38 °C (decahydrate) | ||
| Humid signal | 1,429 °C (2,604 °F; 1,702 Yard) (anhydrous) | ||
| Solubility in water | anhydrous: 4.76 k/100 mL (0 °C) 28.1 g/100 mL (25 °C)[1] 42.7 g/100 mL (100 °C) heptahydrate: 19.5 1000/100 mL (0 °C) 44 k/100 mL (20 °C) | ||
| Solubility | insoluble in ethanol soluble in glycerol, h2o and hydrogen iodide | ||
| Magnetic susceptibility (χ) | −52.0·10−vi cm3/mol | ||
| Refractive index (due north D) | 1.468 (anhydrous) ane.394 (decahydrate) | ||
| Construction | |||
| Crystal structure | orthorhombic (anhydrous)[two] monoclinic (decahydrate) | ||
| Pharmacology | |||
| ATC lawmaking | A06AD13 (WHO) A12CA02 (WHO) | ||
| Hazards | |||
| Occupational prophylactic and health (OHS/OSH): | |||
| Main hazards | Irritant | ||
| NFPA 704 (fire diamond) | 1 0 0 | ||
| Flash indicate | Non-flammable | ||
| Safety data sheet (SDS) | ICSC 0952 | ||
| Related compounds | |||
| Other anions | Sodium selenate Sodium tellurate | ||
| Other cations | Lithium sulfate Potassium sulfate Rubidium sulfate Caesium sulfate | ||
| Related compounds | Sodium bisulfate Sodium sulfite Sodium persulfate | ||
| Supplementary data page | |||
| Sodium sulfate (data page) | |||
| Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). Infobox references | |||
Sodium sulfate (also known every bit sodium sulphate or sulfate of soda) is the inorganic compound with formula Na2SO4 besides as several related hydrates. All forms are white solids that are highly soluble in water. With an almanac production of six million tonnes, the decahydrate is a major commodity chemic production. Information technology is mainly used as a filler in the manufacture of powdered abode laundry detergents and in the Kraft process of paper pulping for making highly alkaline sulfides.[3]
Forms [edit]
- Anhydrous sodium sulfate, known equally the rare mineral thenardite, used equally a drying agent in organic synthesis.
- Heptahydrate sodium sulfate, a very rare form.
- Decahydrate sodium sulfate, known as the mineral mirabilite, widely used past chemical manufacture. It is also known as Glauber's salt.
History [edit]
The decahydrate of sodium sulfate is known equally Glauber's table salt afterwards the Dutch/German pharmacist and apothecary Johann Rudolf Glauber (1604–1670), who discovered it in Austrian bound water in 1625. He named it sal mirabilis (miraculous salt), because of its medicinal backdrop: the crystals were used as a full general-purpose laxative, until more than sophisticated alternatives came about in the 1900s.[4] [5]
In the 18th century, Glauber'south salt began to exist used as a raw cloth for the industrial product 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 procedure, producing constructed sodium sulfate as a key intermediate, became the principal method of soda-ash product.[6]
Chemical properties [edit]
Sodium sulfate is a typical electrostatically bonded ionic sulfate. The existence of gratis sulfate ions in solution is indicated past the piece of cake germination of insoluble sulfates when these solutions are treated with Ba2+ or Pb2+ salts:
- Na2So4 + BaCl2 → ii NaCl + BaSO4
Sodium sulfate is unreactive toward near oxidizing or reducing agents. At high temperatures, it tin can exist converted to sodium sulfide past carbothermal reduction (aka thermo-chemical sulfate reduction (TSR), loftier temperature heating with charcoal, etc.):[seven]
- Na2Then4 + 2 C → Na2S + two 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 common salt sodium bisulfate:[8] [9]
- Na2And then4 + H2SO4 ⇌ 2 NaHSO4
Sodium sulfate displays a moderate tendency to course double salts. The only alums formed with common trivalent metals are NaAl(SOiv)ii (unstable above 39 °C) and NaCr(Theniv)2, in dissimilarity to potassium sulfate and ammonium sulfate which form many stable alums.[x] Double salts with some other alkaline sulfates are known, including Na2Then4·3K2SO4 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 fertiliser.[11] Other double salts include 3NaiiSO4·CaSOiv, 3Na2Then4·MgSO4 (vanthoffite) and NaF·NaiiSO4.[12]
Concrete properties [edit]
Sodium sulfate has unusual solubility characteristics in water.[13] Its solubility in water rises more than than tenfold between 0 °C and 32.384 °C, where it reaches a maximum of 49.7 one thousand/100 mL. At this point the solubility curve changes slope, and the solubility becomes about 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 scale.
Temperature dependence of NaiiSO4 solubility in water
Structure [edit]
Crystals of the decahydrate consist of [Na(OHtwo)6]+ ions with octahedral molecular geometry. These octahedra share edges such that 8 of the x water molecules are bound to sodium and 2 others are interstitial, being hydrogen-bonded to sulfate. These cations are linked to the sulfate anions past hydrogen bonds. The Na–O distances are near 240 pm.[14] Crystalline sodium sulfate decahydrate is as well unusual among hydrated salts in having a measurable residual entropy (entropy at absolute zippo) of 6.32 J/(K·mol). This is ascribed to its ability to distribute water much more speedily compared to near hydrates.[15]
Production [edit]
The world production of sodium sulfate, almost exclusively in the form of the decahydrate, amounts to approximately 5.5 to six million tonnes annually (Mt/a). In 1985, production was iv.v Mt/a, half from natural sources, and one-half from chemical production. After 2000, at a stable level until 2006, natural production had increased to 4 Mt/a, and chemical product decreased to one.5 to 2 Mt/a, with a total of 5.v to 6 Mt/a.[16] [17] [eighteen] [19] For all applications, naturally produced and chemically produced sodium sulfate are practically interchangeable.
Natural sources [edit]
2 thirds of the globe's production of the decahydrate (Glauber's common salt) is from the natural mineral grade mirabilite, for example as found in lake beds in southern Saskatchewan. In 1990, Mexico and Spain were the earth's master producers of natural sodium sulfate (each around 500,000 tonnes), with Russia, The states and Canada around 350,000 tonnes each.[17] Natural resources are estimated at over 1 billion tonnes.[16] [17]
Major producers of 200,000 to 1,500,000 tonnes/twelvemonth in 2006 included Searles Valley Minerals (California, US), Airborne Industrial Minerals (Saskatchewan, Canada), Química del Rey (Coahuila, United mexican states), Minera de Santa Marta and Criaderos Minerales Y Derivados, also known every bit 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, People's republic of china), Nafine Chemical Manufacture Group (Shanxi, Cathay), Sichuan Province Chuanmei Mirabilite (万胜镇, Dongpo District, Meishan, Sichuan, Red china), and Kuchuksulphat JSC (Altai Krai, Siberia, Russia).[xvi] [eighteen]
Anhydrous sodium sulfate occurs in barren environments every bit the mineral thenardite. It slowly turns to mirabilite in damp air. Sodium sulfate is also establish equally glauberite, a calcium sodium sulfate mineral. Both minerals are less common than mirabilite.[ citation needed ]
Chemical industry [edit]
Nigh one 3rd of the world'south sodium sulfate is produced as by-product of other processes in chemical manufacture. Most of this product is chemically inherent to the primary process, and only marginally economical. By try of the manufacture, therefore, sodium sulfate production equally by-product is failing.
The about important chemical sodium sulfate product is during muriatic acid production, either from sodium chloride (common salt) and sulfuric acid, in the Mannheim process, or from sulfur dioxide in the Hargreaves process.[20] The resulting sodium sulfate from these processes is known as table salt block .
- Mannheim: 2 NaCl + H2Then4 → 2 HCl + NatwoSO4
- Hargreaves: 4 NaCl + two SOii + O2 + 2 H2O → 4 HCl + 2 NaiiThen4
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 too a regularly practical and convenient laboratory preparation.
- ii NaOH(aq) + HiiAnd then4(aq) → Na2SO4(aq) + two H2O(l) ΔH = -112.5 kJ (highly exothermic)
In the laboratory it can also exist synthesized from the reaction between sodium bicarbonate and magnesium sulfate.
- 2 NaHCO3 + MgSOfour → Na2SO4 + Mg(OH)2 + two COii
However, as commercial sources are readily available, laboratory synthesis is not practised oftentimes. Formerly, sodium sulfate was as well 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 product of lithium carbonate, chelating agents, resorcinol, ascorbic acrid, silica pigments, nitric acid, and phenol.[16]
Bulk sodium sulfate is usually purified via the decahydrate form, since the anhydrous form tends to attract fe compounds and organic compounds. The anhydrous form is easily produced from the hydrated form by gentle warming.
Major sodium sulfate by-product producers of l–80 Mt/a in 2006 include Elementis Chromium (chromium industry, Castle Hayne, NC, United states), Lenzing AG (200 Mt/a, rayon industry, Lenzing, Austria), Addiseo (formerly Rhodia, methionine industry, Les Roches-Roussillon, French republic), Elementis (chromium manufacture, Stockton-on-Tees, United kingdom), Shikoku Chemicals (Tokushima, Japan) and Visko-R (rayon industry, Russian federation).[16]
Applications [edit]
Sodium sulfate used to dry out an organic liquid. Here clumps form, indicating the presence of water in the organic liquid.
By farther application of sodium sulfate the liquid may exist brought to dryness, indicated hither past the absence of clumping.
Commodity industries [edit]
With US pricing at $thirty per tonne in 1970, up to $xc per tonne for table salt cake quality, and $130 for improve grades, sodium sulphate is a very inexpensive material. The largest utilise is every bit filler in powdered dwelling house 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.[16]
Another formerly major use for sodium sulfate, notably in the US and Canada, is in the Kraft process for the manufacture of wood lurid. Organics present in the "blackness 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 procedure in the early 1960s, more than efficient sulfur recovery was achieved and the need for sodium sulfate makeup was drastically reduced.[21] Hence, the employ of sodium sulfate in the United states and Canadian pulp industry declined from 1,400,000 tonnes per year in 1970 to merely approx. 150,000 tonnes in 2006.[sixteen]
The glass industry provides another significant application for sodium sulfate, every bit 2d largest awarding in Europe. Sodium sulfate is used as a fining agent, to assistance remove small air bubbles from molten glass. It fluxes the glass, and prevents scum formation of the drinking glass melt during refining. The glass industry in Europe has been consuming from 1970 to 2006 a stable 110,000 tonnes annually.[sixteen]
Sodium sulfate is of import in the manufacture of textiles, especially in Japan, where it is the largest awarding. Sodium sulfate is added to increment the ionic strength of the solution and so helps in "levelling", reducing negative electrical charges on textile fibres so that dyes can penetrate evenly (come across the theory of the lengthened double layer (DDL) elaborated by Gouy and Chapman). Unlike the alternative sodium chloride, information technology does not corrode the stainless steel vessels used in dyeing. This awarding in Japan and U.s.a. consumed in 2006 approximately 100,000 tonnes.[sixteen]
Food industry [edit]
Sodium sulfate is used as a diluent for food colours.[22] It is known as E number additive E514.
Thermal storage [edit]
The loftier heat storage chapters 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 depression class solar oestrus for subsequently 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 common salt is incorporated into cells surrounded by solar–heated water. The stage change allows a substantial reduction in the mass of the material required for constructive heat storage (the heat of fusion of sodium sulfate decahydrate is 82 kJ/mol or 252 kJ/kg[23]), with the further advantage of a consistency of temperature as long every bit sufficient fabric in the appropriate phase is bachelor.
For cooling applications, a mixture with mutual sodium chloride salt (NaCl) lowers the melting point to 18 °C (64 °F). The oestrus of fusion of NaCl·Na2And so4·10H2O, is actually increased slightly to 286 kJ/kg.[24]
Modest applications [edit]
In the laboratory, anhydrous sodium sulfate is widely used as an inert drying agent, for removing traces of water from organic solutions.[25] It is more efficient, but slower-interim, than the like agent magnesium sulfate. Information technology is simply effective beneath most xxx °C, just information technology tin 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 (encounter above) demonstrate how the crystals dodder when still wet, but some crystals flow freely one time a sample is dry out.
Glauber'south salt, the decahydrate, is used as a laxative. It is effective for the removal of certain drugs such equally paracetamol (acetaminophen) from the torso, for example, after an overdose.[26] [27]
In 1953, sodium sulfate was proposed for rut storage in passive solar heating systems. This takes advantage of its unusual solubility properties, and the high estrus of crystallisation (78.2 kJ/mol).[28]
Other uses for sodium sulfate include de-frosting windows, starch industry, as an condiment in carpet fresheners, and as an additive to cattle feed.
At to the lowest degree one company, Thermaltake, makes a laptop calculator 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 interim as an insulation.[29]
Safety [edit]
Although sodium sulfate is by and large regarded every bit non-toxic,[22] it should be handled with intendance. The dust tin cause temporary asthma or middle irritation; this hazard can be prevented by using heart protection and a paper mask. Transport is not express, and no Hazard Phrase or Condom Phrase applies.[30]
References [edit]
- ^ National Center for Biotechnology Information. PubChem Chemical compound Summary for CID 24436, Sodium sulfate. https://pubchem.ncbi.nlm.nih.gov/compound/Sodium-sulfate. Accessed Nov. 2, 2020.
- ^ Zachariasen, W. H.; Ziegler, 1000. E. (1932). "The crystal construction of anhydrous sodium sulfate Na2SO4". Zeitschrift für Kristallographie, Kristallgeometrie, Kristallphysik, Kristallchemie. Wiesbaden: Akademische Verlagsgesellschaft. 81 (1–6): 92–101. doi:10.1524/zkri.1932.81.1.92. S2CID 102107891.
{{cite journal}}: CS1 maint: multiple names: authors listing (link) - ^ Helmold Plessen (2000). "Sodium Sulfates". Ullmann'due south Encyclopedia of Industrial Chemical science. Weinheim: Wiley-VCH. doi:ten.1002/14356007.a24_355. ISBN978-3527306732.
- ^ Szydlo, Zbigniew (1994). Water which does non moisture hands: The Abracadabra of Michael Sendivogius. London–Warsaw: Polish Academy of Sciences.
- ^ Westfall, Richard South. (1995). "Glauber, Johann Rudolf". The Galileo Project. Archived from the original on 2011-11-18.
- ^ Aftalion, Fred (1991). A History of the International Chemical Manufacture. Philadelphia: University of Pennsylvania Press. pp. 11–16. ISBN978-0-8122-1297-6.
- ^ Handbook of Chemistry and Physics (71st ed.). Ann Arbor, Michigan: CRC Press. 1990. ISBN9780849304712.
- ^ The Merck Index (7th ed.). Rahway, New Bailiwick of jersey, U.s.: Merck & Co. 1960.
- ^ Nechamkin, Howard (1968). The Chemical science of the Elements . New York: McGraw-Colina.
- ^ Lipson, Henry; Beevers, C. A. (1935). "The Crystal Structure of the Alums". Proceedings of the Royal Gild A. 148 (865): 664–80. Bibcode:1935RSPSA.148..664L. doi:10.1098/rspa.1935.0040.
- ^ Garrett, Donald E. (2001). Sodium sulfate : handbook of deposits, processing, backdrop, and use. San Diego: Academic Printing. ISBN978-0-12-276151-5.
- ^ Mellor, Joseph William (1961). Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry. Vol. II (new impression ed.). London: Longmans. pp. 656–673. ISBN978-0-582-46277-9.
- ^ Linke, W. F.; A. Seidell (1965). Solubilities of Inorganic and Metal Organic Compounds (quaternary ed.). Van Nostrand. ISBN978-0-8412-0097-5.
- ^ 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.
- ^ Brodale, G.; W. F. Giauque (1958). "The Rut of Hydration of Sodium Sulfate. Low Temperature Heat Chapters and Entropy of Sodium Sulfate Decahydrate". Journal of the American Chemical Order. 80 (9): 2042–2044. doi:10.1021/ja01542a003.
- ^ a b c d due east f g h i Suresh, Bala; Kazuteru Yokose (May 2006). Sodium sulfate. CEH Marketing Inquiry Report. Zurich: Chemical Economic Handbook SRI Consulting. pp. 771.1000A–771.1002J. Archived from the original on 2007-03-14.
- ^ a b c "Statistical compendium Sodium sulfate". Reston, Virginia: U.s.a. Geological Survey, Minerals Information. 1997. Archived from the original on 2007-03-07. Retrieved 2007-04-22 .
- ^ a b The economics of sodium sulphate (Eighth ed.). London: Roskill Information Services. 1999.
- ^ The sodium sulphate business. London: Chem Systems International. November 1984.
- ^ Butts, D. (1997). Kirk-Othmer Encyclopedia of Chemical Technology. Vol. v22 (fourth ed.). pp. 403–411.
- ^ Smook, Gary (2002). Handbook for Pulp and Paper Technologists. p. 143. Archived from the original on 2016-08-07.
- ^ a b "Sodium sulfate (WHO Nutrient Additives Serial 44)". World Wellness Organization. 2000. Archived from the original on 2007-09-04. Retrieved 2007-06-06 .
- ^ "Archived copy" (PDF). Archived (PDF) from the original on 2015-09-24. Retrieved 2014-06-19 .
{{cite spider web}}: CS1 maint: archived copy as title (link) - ^ "Archived re-create" (PDF). Archived (PDF) from the original on 2015-09-24. Retrieved 2014-06-nineteen .
{{cite spider web}}: CS1 maint: archived copy as title (link) p.8 - ^ Vogel, Arthur I.; B.V. Smith; N.M. Waldron (1980). Vogel's Uncomplicated Applied Organic Chemical science 1 Preparations (tertiary ed.). London: Longman Scientific & Technical.
- ^ Cocchetto, D.M.; G. Levy (1981). "Assimilation of orally administered sodium sulfate in humans". J Pharm Sci. lxx (3): 331–3. doi:10.1002/jps.2600700330. PMID 7264905.
- ^ Prescott, L. F.; Critchley, J. A. J. H. (1979). "The Treatment of Acetaminophen Poisoning". Annual Review of Pharmacology and Toxicology. 23: 87–101. doi:10.1146/annurev.pa.23.040183.000511. PMID 6347057.
- ^ Telkes, Maria (1953). Improvements in or relating to a device and a limerick of matter for the storage of heat. British Patent No. GB694553.
- ^ "IXoft Specification". Thermaltake Engineering science Co., Ltd. Archived from the original on 2016-03-12. Retrieved 2015-08-fifteen .
- ^ "MSDS Sodium Sulfate Anhydrous". James T Baker. 2006. Archived from the original on 2003-06-19. Retrieved 2007-04-21 .
External links [edit]
- Calculators: surface tensions, and densities, molarities and molalities of aqueous sodium sulfate
Is Naso4 Soluble In Water,
Source: https://en.wikipedia.org/wiki/Sodium_sulfate
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