Sodium Sulfate Ionic Or Covalent
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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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| CompTox Dashboard (EPA) |
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| Chemical formula | Na2Sofour | ||
| Molar mass | 142.04 chiliad/mol (anhydrous) 322.xx g/mol (decahydrate) | ||
| Advent | white crystalline solid hygroscopic | ||
| Smell | odorless | ||
| Density | 2.664 g/cm3 (anhydrous) 1.464 one thousand/cm3 (decahydrate) | ||
| Melting point | 884 °C (1,623 °F; one,157 One thousand) (anhydrous) 32.38 °C (decahydrate) | ||
| Boiling betoken | i,429 °C (2,604 °F; 1,702 K) (anhydrous) | ||
| Solubility in water | anhydrous: iv.76 g/100 mL (0 °C) 28.1 one thousand/100 mL (25 °C)[i] 42.vii thousand/100 mL (100 °C) heptahydrate: 19.v m/100 mL (0 °C) 44 g/100 mL (twenty °C) | ||
| Solubility | insoluble in ethanol soluble in glycerol, h2o and hydrogen iodide | ||
| Magnetic susceptibility (χ) | −52.0·ten−6 cm3/mol | ||
| Refractive index (northward D) | 1.468 (anhydrous) 1.394 (decahydrate) | ||
| Structure | |||
| Crystal construction | orthorhombic (anhydrous)[2] monoclinic (decahydrate) | ||
| Pharmacology | |||
| ATC lawmaking | A06AD13 (WHO) A12CA02 (WHO) | ||
| Hazards | |||
| Occupational safety and health (OHS/OSH): | |||
| Primary hazards | Irritant | ||
| NFPA 704 (fire diamond) | ane 0 0 | ||
| Flash indicate | Non-combustible | ||
| Safety information sail (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 folio) | |||
| Except where otherwise noted, information are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). Infobox references | |||
Sodium sulfate (as well known as sodium sulphate or sulfate of soda) is the inorganic compound with formula NaiiAnd then4 every bit well as several related hydrates. All forms are white solids that are highly soluble in h2o. With an annual production of 6 meg tonnes, the decahydrate is a major commodity chemical product. It is mainly used equally a filler in the manufacture of powdered home laundry detergents and in the Kraft process of paper pulping for making highly alkaline metal sulfides.[3]
Forms [edit]
- Anhydrous sodium sulfate, known every bit the rare mineral thenardite, used as a drying agent in organic synthesis.
- Heptahydrate sodium sulfate, a very rare grade.
- Decahydrate sodium sulfate, known equally the mineral mirabilite, widely used by chemical industry. It is also known as Glauber's common salt.
History [edit]
The decahydrate of sodium sulfate is known as Glauber's salt after the Dutch/German language chemist and apothecary Johann Rudolf Glauber (1604–1670), who discovered information technology in Austrian spring water in 1625. He named it sal mirabilis (miraculous common salt), because of its medicinal properties: the crystals were used equally a general-purpose laxative, until more than sophisticated alternatives came virtually in the 1900s.[4] [5]
In the 18th century, Glauber's table salt began to be used every bit a raw material for the industrial production of soda ash (sodium carbonate), past 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 synthetic sodium sulfate every bit a fundamental intermediate, became the master method of soda-ash product.[6]
Chemic backdrop [edit]
Sodium sulfate is a typical electrostatically bonded ionic sulfate. The existence of complimentary sulfate ions in solution is indicated by the easy germination of insoluble sulfates when these solutions are treated with Ba2+ or Atomic number 82two+ salts:
- NatwoAnd so4 + BaCltwo → 2 NaCl + BaSO4
Sodium sulfate is unreactive toward most oxidizing or reducing agents. At high temperatures, it tin be converted to sodium sulfide past carbothermal reduction (aka thermo-chemical sulfate reduction (TSR), high temperature heating with charcoal, etc.):[7]
- NaiiSO4 + 2 C → Na2Southward + 2 CO2
This reaction was employed in the Leblanc procedure, a defunct industrial route to sodium carbonate.
Sodium sulfate reacts with sulfuric acid to give the acid table salt sodium bisulfate:[8] [9]
- Na2SO4 + H2So4 ⇌ 2 NaHSO4
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(Then4)2, in contrast to potassium sulfate and ammonium sulfate which course many stable alums.[10] Double salts with some other alkaline sulfates are known, including NaiiSO4·3KtwoThen4 which occurs naturally equally 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 3Na2SOfour·CaSO4, 3Na2SO4·MgSO4 (vanthoffite) and NaF·Na2SO4.[12]
Concrete properties [edit]
Sodium sulfate has unusual solubility characteristics in water.[13] Its solubility in h2o rises more than tenfold between 0 °C and 32.384 °C, where it reaches a maximum of 49.vii one thousand/100 mL. At this point the solubility curve changes slope, and the solubility becomes almost contained of temperature. This temperature of 32.384 °C, corresponding to the release of crystal h2o and melting of the hydrated salt, serves as an accurate temperature reference for thermometer calibration.
Temperature dependence of NaiiAnd so4 solubility in water
Structure [edit]
Crystals of the decahydrate consist of [Na(OH2)six]+ ions with octahedral molecular geometry. These octahedra share edges such that 8 of the ten 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 bonds. The Na–O distances are virtually 240 pm.[14] Crystalline sodium sulfate decahydrate is as well unusual among hydrated salts in having a measurable residue entropy (entropy at accented zero) of 6.32 J/(K·mol). This is ascribed to its power to distribute water much more quickly compared to most hydrates.[fifteen]
Product [edit]
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 four.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 i.5 to ii Mt/a, with a full of 5.5 to six Mt/a.[sixteen] [17] [18] [xix] For all applications, naturally produced and chemically produced sodium sulfate are practically interchangeable.
Natural sources [edit]
Ii thirds of the world's production of the decahydrate (Glauber'due south 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 chief producers of natural sodium sulfate (each around 500,000 tonnes), with Russia, Us and Canada around 350,000 tonnes each.[17] Natural resources are estimated at over ane billion tonnes.[xvi] [17]
Major producers of 200,000 to 1,500,000 tonnes/year in 2006 included Searles Valley Minerals (California, United states), Airborne Industrial Minerals (Saskatchewan, Canada), Química del Rey (Coahuila, United mexican states), Minera de Santa Marta and Criaderos Minerales Y Derivados, also known as Grupo Crimidesa (Burgos, Spain), Minera de Santa Marta (Toledo, Kingdom of spain), Sulquisa (Madrid, Spain), Chengdu Sanlian Tianquan Chemical (Tianquan County, Sichuan, People's republic of china), Hongze Yinzhu Chemic Group (Hongze District, Jiangsu, China), Nafine Chemic Industry Group (Shanxi, People's republic of china), Sichuan Province Chuanmei Mirabilite (万胜镇, Dongpo District, Meishan, Sichuan, China), and Kuchuksulphat JSC (Altai Krai, Siberia, Russia).[sixteen] [18]
Anhydrous sodium sulfate occurs in arid environments as the mineral thenardite. It slowly turns to mirabilite in clammy air. Sodium sulfate is also found every bit glauberite, a calcium sodium sulfate mineral. Both minerals are less common than mirabilite.[ citation needed ]
Chemical industry [edit]
About 1 3rd of the globe'southward sodium sulfate is produced equally past-product of other processes in chemical industry. Most of this production is chemically inherent to the main process, and only marginally economical. By effort of the manufacture, therefore, sodium sulfate production as by-product is declining.
The most important chemic sodium sulfate product 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.[20] The resulting sodium sulfate from these processes is known as common salt cake .
- Mannheim: 2 NaCl + H2And then4 → 2 HCl + Na2Then4
- Hargreaves: 4 NaCl + 2 So2 + Otwo + 2 HtwoO → iv HCl + 2 Na2SO4
The second major production of sodium sulfate are the processes where surplus sodium hydroxide is neutralised by sulfuric acid, equally applied on a large scale in the production of rayon. This method is also a regularly applied and convenient laboratory preparation.
- 2 NaOH(aq) + H2And soiv(aq) → NatwoSOiv(aq) + two H2O(l) ΔH = -112.five kJ (highly exothermic)
In the laboratory information technology tin also be synthesized from the reaction between sodium bicarbonate and magnesium sulfate.
- ii NaHCOthree + MgSO4 → Na2SO4 + Mg(OH)two + 2 COii
Still, as commercial sources are readily bachelor, laboratory synthesis is non practised oft. Formerly, sodium sulfate was also a by-product of the manufacture of sodium dichromate, where sulfuric acrid is added to sodium chromate solution forming sodium dichromate, or afterward 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.[16]
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 hands produced from the hydrated class by gentle warming.
Major sodium sulfate by-product producers of 50–eighty 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, Britain), Shikoku Chemicals (Tokushima, Nippon) and Visko-R (rayon industry, Russia).[16]
Applications [edit]
Sodium sulfate used to dry an organic liquid. Hither clumps course, indicating the presence of water in the organic liquid.
By further application of sodium sulfate the liquid may be brought to dryness, indicated here by the absenteeism of clumping.
Commodity industries [edit]
With US pricing at $30 per tonne in 1970, up to $90 per tonne for table salt cake quality, and $130 for better grades, sodium sulphate is a very cheap material. The largest utilise is equally filler in powdered dwelling house laundry detergents, consuming approx. 50% of earth production. This utilize is waning as domestic consumers are increasingly switching to compact or liquid detergents that do non include sodium sulfate.[sixteen]
Some other formerly major utilize for sodium sulfate, notably in the United states and Canada, is in the Kraft process for the manufacture of woods pulp. Organics present in the "black liquor" from this process are burnt to produce oestrus, 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.[21] Hence, the employ of sodium sulfate in the US and Canadian lurid manufacture declined from ane,400,000 tonnes per year in 1970 to only approx. 150,000 tonnes in 2006.[xvi]
The glass industry provides another significant application for sodium sulfate, as second largest application in Europe. Sodium sulfate is used equally a fining amanuensis, to help remove small air bubbles from molten drinking glass. Information technology 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.[xvi]
Sodium sulfate is important in the manufacture of textiles, particularly in Japan, where it is the largest application. Sodium sulfate is added to increase the ionic force of the solution and so helps in "levelling", reducing negative electric charges on textile fibres so that dyes tin penetrate evenly (run across the theory of the diffuse double layer (DDL) elaborated by Gouy and Chapman). Unlike the alternative sodium chloride, it does not corrode the stainless steel vessels used in dyeing. This application in Japan and United states consumed in 2006 approximately 100,000 tonnes.[16]
Food industry [edit]
Sodium sulfate is used as a diluent for food colours.[22] It is known equally E number additive E514.
Thermal storage [edit]
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 peculiarly appropriate for storing low grade solar rut for afterwards release in space heating applications. In some applications the material is incorporated into thermal tiles that are placed in an cranium infinite while in other applications the salt is incorporated into cells surrounded by solar–heated h2o. The phase modify allows a substantial reduction in the mass of the textile 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 equally long every bit sufficient material in the advisable phase is bachelor.
For cooling applications, a mixture with common sodium chloride salt (NaCl) lowers the melting betoken to 18 °C (64 °F). The oestrus of fusion of NaCl·NaiiSOiv·10HtwoO, is actually increased slightly to 286 kJ/kg.[24]
Small-calibration applications [edit]
In the laboratory, anhydrous sodium sulfate is widely used equally an inert drying agent, for removing traces of water from organic solutions.[25] It is more than efficient, only slower-acting, than the like agent magnesium sulfate. It is only effective beneath about 30 °C, but it can be used with a diverseness of materials since it is chemically fairly inert. Sodium sulfate is added to the solution until the crystals no longer clump together; the ii video clips (come across above) demonstrate how the crystals clump when still moisture, just some crystals flow freely once a sample is dry.
Glauber's table salt, the decahydrate, is used as a laxative. Information technology is effective for the removal of sure drugs such as paracetamol (acetaminophen) from the body, for example, subsequently an overdose.[26] [27]
In 1953, sodium sulfate was proposed for oestrus storage in passive solar heating systems. This takes advantage of its unusual solubility properties, and the high heat of crystallisation (78.2 kJ/mol).[28]
Other uses for sodium sulfate include de-frosting windows, starch industry, as an additive in rug fresheners, and as an condiment to cattle feed.
At least ane company, Thermaltake, makes a laptop figurer arctic mat (iXoft Notebook Libation) using sodium sulfate decahydrate inside a quilted plastic pad. The material slowly turns to liquid and recirculates, equalizing laptop temperature and interim every bit an insulation.[29]
Safety [edit]
Although sodium sulfate is generally regarded as non-toxic,[22] it should be handled with care. The dust tin crusade temporary asthma or eye irritation; this risk can be prevented by using eye protection and a paper mask. Send is not limited, and no Hazard Phrase or Safety 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 November. two, 2020.
- ^ Zachariasen, W. H.; Ziegler, Yard. Due east. (1932). "The crystal structure 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 periodical}}: CS1 maint: multiple names: authors list (link) - ^ Helmold Plessen (2000). "Sodium Sulfates". Ullmann'southward Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:ten.1002/14356007.a24_355. ISBN978-3527306732.
- ^ Szydlo, Zbigniew (1994). H2o which does not wet hands: The Alchemy of Michael Sendivogius. London–Warsaw: Polish Academy of Sciences.
- ^ Westfall, Richard S. (1995). "Glauber, Johann Rudolf". The Galileo Projection. Archived from the original on 2011-11-18.
- ^ Aftalion, Fred (1991). A History of the International Chemical Industry. Philadelphia: University of Pennsylvania Printing. pp. xi–sixteen. 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 Jersey, The states: Merck & Co. 1960.
- ^ Nechamkin, Howard (1968). The Chemistry of the Elements . New York: McGraw-Colina.
- ^ Lipson, Henry; Beevers, C. A. (1935). "The Crystal Construction of the Alums". Proceedings of the Royal Society A. 148 (865): 664–lxxx. Bibcode:1935RSPSA.148..664L. doi:10.1098/rspa.1935.0040.
- ^ Garrett, Donald E. (2001). Sodium sulfate : handbook of deposits, processing, properties, and use. San Diego: Bookish Press. ISBN978-0-12-276151-5.
- ^ Mellor, Joseph William (1961). Mellor'southward Comprehensive Treatise on Inorganic and Theoretical Chemical science. Vol. II (new impression ed.). London: Longmans. pp. 656–673. ISBN978-0-582-46277-9.
- ^ Linke, Due west. F.; A. Seidell (1965). Solubilities of Inorganic and Metal Organic Compounds (quaternary ed.). Van Nostrand. ISBN978-0-8412-0097-5.
- ^ Helena Due west. 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, K.; W. F. Giauque (1958). "The Heat of Hydration of Sodium Sulfate. Low Temperature Rut Capacity and Entropy of Sodium Sulfate Decahydrate". Journal of the American Chemical Social club. fourscore (9): 2042–2044. doi:10.1021/ja01542a003.
- ^ a b c d e f g h i Suresh, Bala; Kazuteru Yokose (May 2006). Sodium sulfate. CEH Marketing Research Report. Zurich: Chemic 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: US Geological Survey, Minerals Information. 1997. Archived from the original on 2007-03-07. Retrieved 2007-04-22 .
- ^ a b The economic science 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 Chemic Applied science. Vol. v22 (4th 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 Food Additives Series 44)". Earth Health 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 web}}: CS1 maint: archived copy as championship (link) - ^ "Archived re-create" (PDF). Archived (PDF) from the original on 2015-09-24. Retrieved 2014-06-19 .
{{cite spider web}}: CS1 maint: archived re-create as title (link) p.8 - ^ Vogel, Arthur I.; B.V. Smith; Due north.M. Waldron (1980). Vogel'southward Simple Practical Organic Chemistry one Preparations (3rd ed.). London: Longman Scientific & Technical.
- ^ Cocchetto, D.1000.; Grand. Levy (1981). "Absorption of orally administered sodium sulfate in humans". J Pharm Sci. 70 (3): 331–3. doi:10.1002/jps.2600700330. PMID 7264905.
- ^ Prescott, L. F.; Critchley, J. A. J. H. (1979). "The Treatment of Acetaminophen Poisoning". Almanac 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 composition of matter for the storage of heat. British Patent No. GB694553.
- ^ "IXoft Specification". Thermaltake Technology Co., Ltd. Archived from the original on 2016-03-12. Retrieved 2015-08-15 .
- ^ "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
Sodium Sulfate Ionic Or Covalent,
Source: https://en.wikipedia.org/wiki/Sodium_sulfate#:~:text=Sodium%20sulfate%20is%20a%20typical%20electrostatically%20bonded%20ionic%20sulfate.
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