Magnesium sulfate (magnesium sulphate in British English) is a chemical compound, a salt with the formula MgSO
4, consisting of magnesium cations Mg2+
(20.19% by mass) and sulfate anions SO2−
4. It is a white crystalline solid, soluble in water but not in ethanol.
Anhydrous magnesium sulfate
Epsom salt (heptahydrate)
3D model (JSmol)
|E number||E518 (acidity regulators, ...)|
CompTox Dashboard (EPA)
|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|
|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
Refractive index (nD)
|A06AD04 (WHO) A12CC02 (WHO) B05XA05 (WHO) D11AX05 (WHO) V04CC02 (WHO)|
|Safety data sheet||External MSDS|
|NFPA 704 (fire diamond)|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
Magnesium sulfate is usually encountered in the form of a hydrate MgSO
2O, for various values of n between 1 and 11. The most common is the heptahydrate MgSO
2O, known as Epsom salt, which is a household chemical with many traditional uses, including bath salts.
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. 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:
- Anhydrous, MgSO
4; unstable in nature, hydrates to form epsomite.
- Monohydrate, MgSO
2O; kieserite, monoclinic.
2O or 8MgSO
- Dihydrate, MgSO
2O or 2MgSO
- Trihydrate, MgSO
- Tetrahydrate, MgSO
2O; starkeyite, monoclinic.
- Pentahydrate, MgSO
2O; pentahydrite, triclinic.
- Hexahydrate, MgSO
2O; hexahydrite, monoclinic.
- Heptahydrate, MgSO
2O ("Epsom salt"); epsomite, orthorhombic.
- Enneahydrate, MgSO
- Decahydrate, MgSO
- Undecahydrate, MgSO
2O; meridianiite, triclinic.
As of 2017, the existence of the decahydrate apparently has not been confirmed.
Heptahydrate (Epsom salt)Edit
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 nonporous London clay.
The heptahydrate readily loses one equivalent of water to form the hexahydrate.
It is a natural and organic source of both magnesium and sulfur. Epsom salts are commonly used in bath salts, exfoliants, muscle relaxers and pain relievers. However, these are different from Epsom salts that are used for gardening, as they contain aromas and perfumes not suitable for plants.
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 MgSO
2O, 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. Large crystals can be obtained from solutions of the proper concentration kept at 0 °C for a few days.
The enneahydrate MgSO
2O was identified and characterized only recently, even though it seems easy to produce (by cooling a solution of MgSO
4 and sodium sulfate Na
4 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, β = 95.1°, V = 1.177 nm3 with Z = 4. The most probable space group is P21/c. Magnesium selenate also forms an enneahydrate MgSeO
2O, but with a different crystal structure.
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.
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, MgCO
3) or magnesia (oxide, MgO) 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 relaxation is the primary mechanism that causes the absorption of sound in seawater at frequencies above 10 kHz  (acoustic energy is converted to thermal energy). Lower frequencies are less absorbed by the salt, so that low frequency sound travels farther in the ocean. Boric acid and magnesium carbonate also contribute to absorption.
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 from 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. It is also the usual component of the solution used in isolation tanks.
Internally, magnesium sulfate may be administered by oral, respiratory, or intravenous routes. Internal uses include replacement therapy for magnesium deficiency, treatment of acute and severe arrhythmias, as a bronchodilator in the treatment of asthma, and preventing eclampsia.
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.
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. 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 that inadvertently ingested lead.
Anhydrous magnesium sulfate is commonly used as a desiccant in organic synthesis owing 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.
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.
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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.
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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.
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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.
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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.
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