Magnesium Sulfate
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Magnesium Sulfate

Magnesium sulphate
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Magnesium sulfate anhydrous.jpg
Anhydrous magnesium sulfate
Magnesium sulfate.JPG
Epsomite (heptahydrate)
IUPAC name
Magnesium sulfate
Other names
Epsom salt (heptahydrate)
English salt
Bitter salts
Bath salts
3D model (JSmol)
ECHA InfoCard 100.028.453 Edit this at Wikidata
E number E518 (acidity regulators, ...)
RTECS number
  • OM4500000
Molar mass 120.366 g/mol (anhydrous)
138.38 g/mol (monohydrate)
174.41 g/mol (trihydrate)
210.44 g/mol (pentahydrate)
228.46 g/mol (hexahydrate)
246.47 g/mol (heptahydrate)
Appearance white crystalline solid
Odor odorless
Density 2.66 g/cm3 (anhydrous)
2.445 g/cm3 (monohydrate)
1.68 g/cm3 (heptahydrate)
1.512 g/cm3 (11-hydrate)
Melting point anhydrous decomposes at 1,124 °C
monohydrate decomposes at 200 °C
heptahydrate decomposes at 150 °C
undecahydrate decomposes at 2 °C
26.9 g/100 mL (0 °C)
35.1 g/100 mL (20 °C)
50.2 g/100 mL (100 °C)
113 g/100 mL (20 °C)
Solubility 1.16 g/100 mL (18 °C, ether)
slightly soluble in alcohol, glycerol
insoluble in acetone
-50·10-6 cm3/mol
1.523 (monohydrate)
1.433 (heptahydrate)
monoclinic (hydrate)
A06AD04 (WHO) A12CC02 (WHO) B05XA05 (WHO) D11AX05 (WHO) V04CC02 (WHO)
Safety data sheet External MSDS
NFPA 704 (fire diamond)
Flammability code 0: Will not burn. E.g. waterHealth code 1: Exposure would cause irritation but only minor residual injury. E.g. turpentineReactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no codeNFPA 704 four-colored diamond
Related compounds
Other cations
Beryllium sulfate
Calcium sulfate
Strontium sulfate
Barium sulfate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

Magnesium sulfate (Magnesium sulphate in British English) is a chemical compound, a salt with the formula , consisting of magnesium cations (20.19% by mass) and sulfate anions . It is a white crystalline solid, soluble in water but not in ethanol.

Magnesium sulfate is usually encountered in the form of a hydrate , for various values of n between 1 and 11. The most common is the heptahydrate , known as Epsom salt, which is a household chemical with many traditional uses, including bath salts.[1]

The main use of magnesium sulfate is in agriculture, to correct soils deficient in magnesium (an essential plant nutrient). The monohydrate is favored for this use; by the mid-1970s, its production was 2.3 million tons per year.[2] The anhydrous form and several hydrates occur in nature as minerals, and the salt is a significant component of the water from some springs.


Magnesium sulfate can crystallize as several hydrates, including:

As of 2017, the existence of the decahydrate apparently has not been confirmed.[7]

All the hydrates lose water upon heating. Above 320 °C only the anhydrous form is stable. It decomposes without melting at 1124 °C into magnesium oxide (MgO) and sulfur trioxide (SO3).

Heptahydrate (Epsom salt)

The heptahydrate takes its common name "Epsom salt" from a bitter saline spring in Epsom in Surrey, England, where the salt was produced from the springs that arise where the porous chalk of the North Downs meets non-porous London clay.

The heptahydrate readily loses one equivalent of water to form the hexahydrate.


The monohydrate can be prepared by heating the hexahydrate to approximately 150 °C. Further heating to approximately 300-320 °C gives anhydrous magnesium sulfate.


The undecahydrate , meridianiite, is stable at atmospheric pressure only below 2 °C. Above that temperature, it liquefies into a mix of solid heptahydrate and a saturated solution. It has an eutectic point with water at -3.9 °C and 17.3% (mass) of MgSO4.[5] Large crystals can be obtained from solutions of the proper concentration kept at 0 °C for a few days.[5]

At pressures of about 0.9 GPa and at 240 K, meridianiite decomposes into a mixture of ice VI and the enneahydrate [7]


The enneahydrate was identified and characterized only recently, even though it seems easy to produce (by cooling a solution of and sodium sulfate in suitable proportions).

The structure is monoclinic, with unit-cell parameters at 250 Ka= 0.675 nm, b = 1.195 nm, c = 1.465 nm, ? = 95.1°, V = 1.177 nm3 with Z = 4. The most probable space group is P21/c. Magnesium selenate also forms an enneahydrate , but with a different crystal structure.[7]

Natural occurrence

Magnesium sulfates are common minerals in geological environments. Their occurrence is mostly connected with supergene processes. Some of them are also important constituents of evaporitic potassium-magnesium (K-Mg) salts deposits.

Bright spots observed by the Dawn Spacecraft in Occator Crater on the dwarf planet Ceres are most consistent with reflected light from magnesium sulfate hexahydrate.[8]

Almost all known mineralogical forms of MgSO4 are hydrates. Epsomite is the natural analogue of "Epsom salt". Meridianiite, MgSO4·11H2O, has been observed on the surface of frozen lakes and is thought to also occur on Mars. Hexahydrite is the next lower (6) hydrate. Three next lower hydrates--pentahydrite, starkeyite, and especially sanderite are rare. Kieserite is a monohydrate and is common among evaporitic deposits. Anhydrous magnesium sulfate was reported from some burning coal dumps.


Magnesium sulfate is usually obtained directly from dry lake beds and other natural sources. It can also be prepared by reacting magnesite (magnesium carbonate, ) or magnesia (oxide, ) with sulfuric acid.

Another possible method is to treat seawater or magnesium-containing industrial wastes so as to precipitate magnesium hydroxide, and react the precipitate with sulfuric acid.



Magnesium sulfate is used both externally (as Epsom salt) and internally.

The main external use is the formulation as bath salts, especially for foot baths to soothe sore feet. Such baths have been claimed to also soothe and hasten recovery of muscle pain, soreness, or injury. However, these claims have not been scientifically confirmed. The main benefit of the salt is cosmetic: it prevents the temporary skin wrinkling caused by prolonged immersion in plain water.[9][1] It is also the usual component of the solution used in isolation tanks.

In the UK, a medication containing magnesium sulphate and phenol, called "drawing paste", is claimed to be useful for small boils or localised infections,[10] and removing splinters.[11]

Internally, magnesium sulfate may be administered by oral, respiratory, or intravenous routes. Internal uses include: replacement therapy for magnesium deficiency,[12] treatment of acute and severe arrhythmias,[13] as a bronchodilator in the treatment of asthma,[14] and preventing eclampsia.[15]


In agriculture, magnesium sulfate is used to increase magnesium or sulfur content in soil. It is most commonly applied to potted plants, or to magnesium-hungry crops, such as potatoes, tomatoes, carrots, peppers, lemons, and roses. The advantage of magnesium sulfate over other magnesium soil amendments (such as dolomitic lime) is its high solubility, which also allows the option of foliar feeding. Solutions of magnesium sulfate are also nearly pH neutral, compared with alkaline salts of magnesium as found in limestone; therefore, the use of magnesium sulfate as a magnesium source for soil does not significantly change the soil pH.[16]

Magnesium sulfate was historically used as a treatment for lead poisoning prior to the development of chelation therapy, as it was hoped that any lead ingested would be precipitated out by the magnesium sulfate and subsequently purged from the digestive system.[17] This application saw particularly widespread use among veterinarians during the early-to-mid 20th century; Epsom salt was already available on many farms for agricultural use, and it was often prescribed in the treatment of farm animals which inadvertently ingested lead.[18][19]

Food preparation

Magnesium sulfate is used as a brewing salt in making beer.[20] It may also be used as a coagulant for making tofu.[21]


Anhydrous magnesium sulfate is commonly used as a desiccant in organic synthesis due to its affinity for water and compatibility with most organic compounds. During work-up, an organic phase is treated with anhydrous magnesium sulfate. The hydrated solid is then removed with filtration, decantation or distillation (if the boiling point is low enough). Other inorganic sulfate salts such as sodium sulfate and calcium sulfate may be used in the same way.


Magnesium sulfate is used to prepare specific cements by the reaction between magnesium oxide and magnesium sulfate solution, which are of good binding ability and more resistance than Portland cement. This cement is mainly adopted in the production of lightweight insulation panels. Weakness in water resistance limits its usage.

Magnesium (or sodium) sulfate is also used for testing aggregates for soundness in accordance with ASTM C88 standard, when there are no service records of the material exposed to actual weathering conditions. The test is accomplished by repeated immersion in saturated solutions followed by oven drying to dehydrate the salt precipitated in permeable pore spaces. The internal expansive force, derived from the rehydration of the salt upon re-immersion, simulates the expansion of water on freezing.


Magnesium sulfate heptahydrate is also used to maintain the magnesium concentration in marine aquaria which contain large amounts of stony corals, as it is slowly depleted in their calcification process. In a magnesium-deficient marine aquarium, calcium and alkalinity concentrations are very difficult to control because not enough magnesium is present to stabilize these ions in the saltwater and prevent their spontaneous precipitation into calcium carbonate.[22]

Double salts

Double salts containing magnesium sulfate exist. There are several known as sodium magnesium sulfates and potassium magnesium sulfates. A mixed copper-magnesium sulfate heptahydrate (Mg,Cu)SO4·7H2O was recently found to occur in mine tailings,and has been given the mineral name alpersite.[23]

See also


  1. ^ a b "Quick Cures/Quack Cures: Is Epsom Worth Its Salt?". The Wall Street Journal. 9 April 2012. Archived from the original on 12 April 2012. Retrieved 2019.
  2. ^ Industrial Inorganic Chemistry, Karl Heinz Büchel, Hans-Heinrich Moretto, Dietmar Werner, John Wiley & Sons, 2d edition, 2000, ISBN 978-3-527-61333-5
  3. ^
  4. ^ a b c Odochian, Lucia (1995). "Study of the nature of the crystallization water in some magnesium hydrates by thermal methods". Journal of Thermal Analysis and Calorimetry. 45 (6): 1437-1448. doi:10.1007/BF02547437. Archived from the original on 26 August 2011. Retrieved 2010.
  5. ^ a b c d e A. Dominic Fortes, Frank Browning, and Ian G. Wood (2012): "Cation substitution in synthetic meridianiite (MgSO4·11H2O) I: X-ray powder diffraction analysis of quenched polycrystalline aggregates". Physics and Chemistry of Minerals, volume 39, issue , pages 419-441. doi:10.1007/s00269-012-0497-9
  6. ^ a b c R. C. Peterson, W. Nelson, B. Madu, and H. F. Shurvell (2007): "Meridianiite: A new mineral species observed on Earth and predicted to exist on Mars". American Mineralogist, volume 92, issue 10, pages 1756-1759. doi:10.2138/am.2007.2668
  7. ^ a b c d A. Dominic Fortes, Kevin S. Knight, and Ian G. Wood (2017): "Structure, thermal expansion and incompressibility of MgSO4·9H2O, its relationship to meridianiite (MgSO4·11H2O) and possible natural occurrences". Acta Crystallographica Section B: Structureal Science, Crystal Engineering and Materials, volume 73, part 1, pages 47-64. doi:10.1107/S2052520616018266
  8. ^ M. C. De Sanctis; E. Ammannito; A. Raponi; S. Marchi; T. B. McCord; H. Y. McSween; F. Capaccioni; M. T. Capria; F. G. Carrozzo; M. Ciarniello; A. Longobardo; F. Tosi; S. Fonte; M. Formisano; A. Frigeri; M. Giardino; G. Magni; E. Palomba; D. Turrini; F. Zambon; J.-P. Combe; W. Feldman; R. Jaumann; L. A. McFadden; C. M. Pieters (2015). "Ammoniated phyllosilicates with a likely outer Solar System origin on (1) Ceres" (PDF). Nature. 528 (7581): 241-244. doi:10.1038/nature16172. PMID 26659184.
  9. ^ Ingraham, Paul. "Does Epsom Salt Work? The science of Epsom salt bathing for recovery from muscle pain, soreness, or injury". Pain Science. Archived from the original on 10 September 2016. Retrieved 2016.
  10. ^ "Boots Magnesium Sulfate Paste B.P. - Patient Information Leaflet (PIL) - (eMC)". Retrieved 2018.
  11. ^ "Removing a splinter with Magnesium Sulphate".
  12. ^ "Pharmaceutical Information - Magnesium Sulfate". RxMed. Archived from the original on 3 April 2009. Retrieved 2009.
  13. ^ "CPR and First Aid: Antiarrhythmic Drugs During and Immediately After Cardiac Arrest (section)". American Heart Association. Retrieved 2016. Previous ACLS guidelines addressed the use of magnesium in cardiac arrest with polymorphic ventricular tachycardia (ie, torsades de pointes) or suspected hypomagnesemia, and this has not been reevaluated in the 2015 Guidelines Update. These previous guidelines recommended defibrillation for termination of polymorphic VT (ie, torsades de pointes), followed by consideration of intravenous magnesium sulfate when secondary to a long QT interval.
  14. ^ Blitz M, Blitz S, Hughes R, Diner B, Beasley R, Knopp J, Rowe BH (2005). "Aerosolized magnesium sulfate for acute asthma: a systematic review". Chest. 128 (1): 337-344. doi:10.1378/chest.128.1.337. PMID 16002955..
  15. ^ Duley, L; Gülmezoglu, AM; Henderson-Smart, DJ; Chou, D (10 November 2010). "Magnesium sulphate and other anticonvulsants for women with pre-eclampsia". The Cochrane Database of Systematic Reviews (11): CD000025. doi:10.1002/14651858.CD000025.pub2. PMC 7061250. PMID 21069663.
  16. ^ "Pubchem: magnesium sulfate". Archived from the original on 18 October 2016.
  17. ^ Wood, H. C. (1877). A Treatise on Therapeutics, Comprising Materia Medica and Toxicology, with Especial Reference to the Application of the Physiological Action of Drugs to Clinical Medicine. Philadelphia: J. B. Lippincott & Co. p. 34. The treatment of acute lead-poisoning consists in the evacuation of the stomach, if necessary, the exhibition of the sulphate of sodium or of magnesium, and the meeting of the indications as they arrive. The Epsom and Glauber's salts act as chemical antidotes, by precipitating the insoluble sulphate of lead, and also, if in excess, empty the bowel of the compound formed.
  18. ^ Barker, C. A. V. (January 1945). "Experience with Lead Poisoning". Canadian Journal of Comparative Medicine and Veterinary Science. 9 (1): 6-8. PMC 1660962. PMID 17648099. Udall (1) suggests sodium citrate as of some value together with Epsom salts which will bring about a precipitation of the lead in the form of an insoluble compound. Nelson (3) reported a case that survived following the use of a 20% magnesium sulphate solution intravenously, subcutaneously and orally. McIntosh (5) has suggested that purgative doses of Epsom salts may be effective in combining with the lead and overcoming the toxicity.
  19. ^ Herriot, James (1972). All Creatures Great and Small. New York: St. Martin's Press. p. 157. ISBN 0-312-08498-6. The specific antidotes to metal poisoning had not been discovered and the only thing which sometimes did a bit of good was magnesium sulphate which caused the precipitation of insoluble lead sulphate. The homely term for magnesium sulphate is, of course, epsom salts.
  20. ^ "Magnesium Sulphate". National Home Brew. Archived from the original on 1 August 2016. Retrieved 2019.
  21. ^ US The present invention relates to a novel process for producing packed tofu, particularly a process for producing long-life packed tofu from sterilized soybean milk. 6042851, Matsuura, Masaru; Masaoki Sasaki & Jun Sasakib et al., "Process for producing packed tofu" 
  22. ^ "Do-It-Yourself Magnesium Supplements for the Reef Aquarium". Reefkeeping. 2006. Archived from the original on 22 March 2008. Retrieved 2008.
  23. ^ Peterson, Ronald C.; Hammarstrom, Jane M.; Seal, II, Robert R (February 2006). "Alpersite (Mg,Cu)SO4·7H2O, a new mineral of the melanterite group, and cuprian pentahydrite: Their occurrence within mine waste". American Mineralogist. 91 (2-3): 261-269. doi:10.2138/am.2006.1911.

External links

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