Sodium Sulfate Acid Or Base
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| Other names Sodium sulphate | |||
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| ECHA InfoCard | 100.028.928 | ||
| E number | E514(i) (acerbity regulators, ...) | ||
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| CompTox Dashboard (EPA) |
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| Properties | |||
| Chemical formula | NatwoAnd sofour | ||
| Molar mass | 142.04 thou/mol (anhydrous) 322.20 g/mol (decahydrate) | ||
| Advent | white crystalline solid hygroscopic | ||
| Odour | odorless | ||
| Density | two.664 one thousand/cm3 (anhydrous) 1.464 g/cm3 (decahydrate) | ||
| Melting bespeak | 884 °C (1,623 °F; 1,157 K) (anhydrous) 32.38 °C (decahydrate) | ||
| Boiling point | 1,429 °C (2,604 °F; 1,702 Grand) (anhydrous) | ||
| Solubility in h2o | anhydrous: 4.76 g/100 mL (0 °C) 28.1 g/100 mL (25 °C)[one] 42.7 thou/100 mL (100 °C) heptahydrate: 19.five thou/100 mL (0 °C) 44 k/100 mL (xx °C) | ||
| Solubility | insoluble in ethanol soluble in glycerol, water and hydrogen iodide | ||
| Magnetic susceptibility (χ) | −52.0·10−vi cmiii/mol | ||
| Refractive index (northward D) | one.468 (anhydrous) 1.394 (decahydrate) | ||
| Construction | |||
| Crystal structure | orthorhombic (anhydrous)[2] monoclinic (decahydrate) | ||
| Pharmacology | |||
| ATC code | A06AD13 (WHO) A12CA02 (WHO) | ||
| Hazards | |||
| Occupational condom and health (OHS/OSH): | |||
| Master hazards | Irritant | ||
| NFPA 704 (fire diamond) | i 0 0 | ||
| Flash point | Non-flammable | ||
| Safety information 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 information page | |||
| Sodium sulfate (information folio) | |||
| Except where otherwise noted, data are given for materials in their standard land (at 25 °C [77 °F], 100 kPa). Infobox references | |||
Sodium sulfate (besides known as 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 6 million tonnes, the decahydrate is a major commodity chemical product. It is mainly used as a filler in the manufacture of powdered home laundry detergents and in the Kraft process of paper pulping for making highly alkali metal sulfides.[3]
Forms [edit]
- Anhydrous sodium sulfate, known every bit the rare mineral thenardite, used every bit 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. Information technology is also known equally Glauber's table salt.
History [edit]
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 Austrian spring water in 1625. He named information technology sal mirabilis (miraculous salt), considering of its medicinal properties: the crystals were used as a general-purpose laxative, until more sophisticated alternatives came well-nigh in the 1900s.[4] [v]
In the 18th century, Glauber's table salt began to be used as a raw material 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 synthetic sodium sulfate every bit a key intermediate, became the principal method of soda-ash production.[six]
Chemical properties [edit]
Sodium sulfate is a typical electrostatically bonded ionic sulfate. The existence of gratuitous sulfate ions in solution is indicated by the easy formation of insoluble sulfates when these solutions are treated with Ba2+ or Pb2+ salts:
- Na2Then4 + BaCl2 → 2 NaCl + BaSO4
Sodium sulfate is unreactive toward about oxidizing or reducing agents. At high temperatures, it can exist converted to sodium sulfide by carbothermal reduction (aka thermo-chemical sulfate reduction (TSR), high temperature heating with charcoal, etc.):[7]
- NaiiThen4 + two C → NaiiS + 2 CO2
This reaction was employed in the Leblanc process, a defunct industrial route to sodium carbonate.
Sodium sulfate reacts with sulfuric acrid to give the acrid salt sodium bisulfate:[8] [ix]
- Na2SO4 + HiiSOfour ⇌ 2 NaHSOfour
Sodium sulfate displays a moderate trend to form double salts. The only alums formed with common trivalent metals are NaAl(SOiv)ii (unstable to a higher place 39 °C) and NaCr(SO4)2, in contrast to potassium sulfate and ammonium sulfate which form many stable alums.[10] Double salts with some other brine metal sulfates are known, including Na2Then4·3K2So4 which occurs naturally equally the mineral aphthitalite. Germination of glaserite past 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 water rises more than 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 bend changes slope, and the solubility becomes nigh 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 authentic temperature reference for thermometer scale.
Temperature dependence of NatwoAnd sofour solubility in water
Structure [edit]
Crystals of the decahydrate consist of [Na(OH2)6]+ ions with octahedral molecular geometry. These octahedra share edges such that 8 of the x h2o molecules are jump to sodium and two others are interstitial, being hydrogen-bonded to sulfate. These cations are linked to the sulfate anions by hydrogen bonds. The Na–O distances are nearly 240 pm.[xiv] Crystalline sodium sulfate decahydrate is as well unusual amid hydrated salts in having a measurable residual entropy (entropy at absolute zero) of vi.32 J/(K·mol). This is ascribed to its ability to distribute water much more quickly compared to most hydrates.[15]
Production [edit]
The world product of sodium sulfate, almost exclusively in the form of the decahydrate, amounts to approximately 5.five to vi million tonnes annually (Mt/a). In 1985, production was 4.5 Mt/a, half from natural sources, and half from chemical product. After 2000, at a stable level until 2006, natural product had increased to four Mt/a, and chemic production decreased to 1.five to two Mt/a, with a full of 5.5 to 6 Mt/a.[16] [17] [xviii] [19] For all applications, naturally produced and chemically produced sodium sulfate are practically interchangeable.
Natural sources [edit]
Two thirds of the globe's production of the decahydrate (Glauber's salt) is from the natural mineral form mirabilite, for case as plant in lake beds in southern Saskatchewan. In 1990, Mexico and Spain were the world'south main producers of natural sodium sulfate (each effectually 500,000 tonnes), with Russia, Us and Canada around 350,000 tonnes each.[17] Natural resource are estimated at over one billion tonnes.[16] [17]
Major producers of 200,000 to 1,500,000 tonnes/year in 2006 included Searles Valley Minerals (California, U.s.a.), Airborne Industrial Minerals (Saskatchewan, Canada), Química del Rey (Coahuila, United mexican states), Minera de Santa Marta and Criaderos Minerales Y Derivados, likewise known as Grupo Crimidesa (Burgos, Kingdom of spain), Minera de Santa Marta (Toledo, Spain), Sulquisa (Madrid, Spain), Chengdu Sanlian Tianquan Chemical (Tianquan County, Sichuan, China), Hongze Yinzhu Chemical Grouping (Hongze District, Jiangsu, China), Nafine Chemical Manufacture Group (Shanxi, China), Sichuan Province Chuanmei Mirabilite (万胜镇, Dongpo Commune, Meishan, Sichuan, China), and Kuchuksulphat JSC (Altai Krai, Siberia, Russian federation).[sixteen] [18]
Anhydrous sodium sulfate occurs in arid environments as the mineral thenardite. Information technology slowly turns to mirabilite in clammy air. Sodium sulfate is also constitute as glauberite, a calcium sodium sulfate mineral. Both minerals are less common than mirabilite.[ commendation needed ]
Chemical manufacture [edit]
Near i third of the world'south sodium sulfate is produced every bit past-production of other processes in chemic industry. Near of this production is chemically inherent to the primary procedure, and only marginally economical. By endeavor of the manufacture, therefore, sodium sulfate production every bit past-product is declining.
The nearly important chemic sodium sulfate production is during muriatic acid product, either from sodium chloride (salt) and sulfuric acrid, in the Mannheim procedure, or from sulfur dioxide in the Hargreaves process.[xx] The resulting sodium sulfate from these processes is known as salt cake .
- Mannheim: 2 NaCl + HiiSO4 → 2 HCl + Na2SO4
- Hargreaves: 4 NaCl + 2 And thenii + O2 + 2 H2O → iv HCl + ii NatwoSO4
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 user-friendly laboratory preparation.
- 2 NaOH(aq) + HtwoAnd theniv(aq) → Na2SO4(aq) + 2 H2O(50) ΔH = -112.five kJ (highly exothermic)
In the laboratory it can also exist synthesized from the reaction between sodium bicarbonate and magnesium sulfate.
- 2 NaHCO3 + MgSO4 → Na2And so4 + Mg(OH)2 + 2 CO2
However, as commercial sources are readily available, laboratory synthesis is non practised oft. Formerly, sodium sulfate was also a by-production of the manufacture of sodium dichromate, where sulfuric acid 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 acrid, and phenol.[16]
Bulk sodium sulfate is usually purified via the decahydrate form, since the anhydrous class tends to attract fe compounds and organic compounds. The anhydrous form is hands produced from the hydrated course past gentle warming.
Major sodium sulfate past-product producers of l–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, French republic), Elementis (chromium manufacture, Stockton-on-Tees, United kingdom), Shikoku Chemicals (Tokushima, Japan) and Visko-R (rayon manufacture, Russia).[xvi]
Applications [edit]
Sodium sulfate used to dry an organic liquid. Hither clumps form, indicating the presence of water in the organic liquid.
By further awarding of sodium sulfate the liquid may be brought to dryness, indicated here by the absence of clumping.
Commodity industries [edit]
With US pricing at $thirty per tonne in 1970, up to $ninety per tonne for common salt cake quality, and $130 for better grades, sodium sulphate is a very cheap fabric. The largest utilize is as filler in powdered home laundry detergents, consuming approx. 50% of world production. This apply is waning as domestic consumers are increasingly switching to compact or liquid detergents that do not include sodium sulfate.[xvi]
Another formerly major use for sodium sulfate, notably in the US and Canada, is in the Kraft procedure for the industry of wood pulp. Organics present in the "blackness liquor" from this procedure 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.[21] Hence, the use of sodium sulfate in the United states and Canadian lurid industry declined from one,400,000 tonnes per twelvemonth in 1970 to only approx. 150,000 tonnes in 2006.[16]
The drinking glass industry provides another significant awarding for sodium sulfate, as 2nd largest application in Europe. Sodium sulfate is used as a fining amanuensis, to help remove minor 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.[16]
Sodium sulfate is of import in the manufacture of textiles, particularly in Japan, where it is the largest application. Sodium sulfate is added to increment the ionic strength of the solution and and so helps in "levelling", reducing negative electrical charges on textile fibres and so that dyes can penetrate evenly (see the theory of the diffuse double layer (DDL) elaborated past Gouy and Chapman). Dissimilar 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.[sixteen]
Food industry [edit]
Sodium sulfate is used equally a diluent for food colours.[22] Information technology is known equally Due east number additive E514.
Thermal storage [edit]
The high heat storage capacity in the phase change from solid to liquid, and the advantageous stage change temperature of 32 °C (ninety °F) makes this material especially advisable for storing low grade solar heat 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 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 effective heat storage (the estrus of fusion of sodium sulfate decahydrate is 82 kJ/mol or 252 kJ/kg[23]), with the farther reward of a consistency of temperature equally long equally sufficient material in the advisable phase is available.
For cooling applications, a mixture with mutual sodium chloride salt (NaCl) lowers the melting indicate to 18 °C (64 °F). The heat of fusion of NaCl·NaiiSO4·10H2O, is actually increased slightly to 286 kJ/kg.[24]
Small-scale 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 than efficient, but slower-acting, than the similar agent magnesium sulfate. Information technology is only effective below virtually 30 °C, but it tin be used with a diversity of materials since it is chemically fairly inert. Sodium sulfate is added to the solution until the crystals no longer clump together; the 2 video clips (see higher up) demonstrate how the crystals dodder when still wet, but some crystals menstruum freely in one case a sample is dry.
Glauber's common salt, the decahydrate, is used equally a laxative. Information technology is effective for the removal of sure drugs such as paracetamol (acetaminophen) from the body, for case, after an overdose.[26] [27]
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 rut of crystallisation (78.2 kJ/mol).[28]
Other uses for sodium sulfate include de-frosting windows, starch manufacture, as an additive in carpet fresheners, and every bit an additive to cattle feed.
At least one company, Thermaltake, makes a laptop figurer chill mat (iXoft Notebook Libation) using sodium sulfate decahydrate inside a quilted plastic pad. The cloth slowly turns to liquid and recirculates, equalizing laptop temperature and acting equally an insulation.[29]
Safety [edit]
Although sodium sulfate is by and large regarded every bit non-toxic,[22] information technology should be handled with care. The dust tin cause temporary asthma or eye irritation; this gamble can exist prevented by using eye protection and a paper mask. Ship is not limited, and no Risk Phrase or Safety Phrase applies.[30]
References [edit]
- ^ National Center for Biotechnology Information. PubChem Compound Summary for CID 24436, Sodium sulfate. https://pubchem.ncbi.nlm.nih.gov/compound/Sodium-sulfate. Accessed Nov. ii, 2020.
- ^ Zachariasen, West. H.; Ziegler, Chiliad. 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 periodical}}: CS1 maint: multiple names: authors list (link) - ^ Helmold Plessen (2000). "Sodium Sulfates". Ullmann'south Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:x.1002/14356007.a24_355. ISBN978-3527306732.
- ^ Szydlo, Zbigniew (1994). Water which does non wet hands: The Alchemy of Michael Sendivogius. London–Warsaw: Smooth Academy of Sciences.
- ^ Westfall, Richard S. (1995). "Glauber, Johann Rudolf". The Galileo Project. Archived from the original on 2011-11-xviii.
- ^ Aftalion, Fred (1991). A History of the International Chemical Industry. Philadelphia: University of Pennsylvania Printing. pp. 11–16. ISBN978-0-8122-1297-vi.
- ^ Handbook of Chemistry and Physics (71st ed.). Ann Arbor, Michigan: CRC Press. 1990. ISBN9780849304712.
- ^ The Merck Index (seventh ed.). Rahway, New Jersey, U.s.: Merck & Co. 1960.
- ^ Nechamkin, Howard (1968). The Chemistry of the Elements . New York: McGraw-Hill.
- ^ Lipson, Henry; Beevers, C. A. (1935). "The Crystal Structure of the Alums". Proceedings of the Royal Society A. 148 (865): 664–80. Bibcode:1935RSPSA.148..664L. doi:10.1098/rspa.1935.0040.
- ^ Garrett, Donald Eastward. (2001). Sodium sulfate : handbook of deposits, processing, properties, and use. San Diego: Academic Press. ISBN978-0-12-276151-five.
- ^ 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 Metallic 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 Heat of Hydration of Sodium Sulfate. Low Temperature Heat Chapters and Entropy of Sodium Sulfate Decahydrate". Journal of the American Chemical Lodge. 80 (9): 2042–2044. doi:ten.1021/ja01542a003.
- ^ a b c d due east f 1000 h i Suresh, Bala; Kazuteru Yokose (May 2006). Sodium sulfate. CEH Marketing Enquiry 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: 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 (8th ed.). London: Roskill Data Services. 1999.
- ^ The sodium sulphate business. London: Chem Systems International. Nov 1984.
- ^ Butts, D. (1997). Kirk-Othmer Encyclopedia of Chemic Engineering. Vol. v22 (quaternary 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 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 spider web}}: CS1 maint: archived copy as title (link) - ^ "Archived copy" (PDF). Archived (PDF) from the original on 2015-09-24. Retrieved 2014-06-19 .
{{cite spider web}}: CS1 maint: archived re-create as championship (link) p.8 - ^ Vogel, Arthur I.; B.V. Smith; N.M. Waldron (1980). Vogel's Elementary Practical Organic Chemistry 1 Preparations (tertiary ed.). London: Longman Scientific & Technical.
- ^ Cocchetto, D.Thousand.; 1000. Levy (1981). "Absorption of orally administered sodium sulfate in humans". J Pharm Sci. seventy (iii): 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 composition of matter for the storage of rut. 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 Bakery. 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 Acid Or Base,
Source: https://en.wikipedia.org/wiki/Sodium_sulfate
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