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Magnesium sulfate
Magnesium sulfate
Magnesium sulfate
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Magnesium sulfate
Magnesium sulfate hexahydrate
Anhydrous magnesium sulfate
Epsomite (Magnesium sulfate heptahydrate)
Names
IUPAC name
Magnesium sulfate
Other names
  • Epsom salt (Magnesium sulfate heptahydrate)
  • English salt
  • Bitter salts
  • Bath salt
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
DrugBank
ECHA InfoCard 100.028.453 Edit this at Wikidata
EC Number
  • 231-298-2
KEGG
RTECS number
  • OM4500000
UNII
  • InChI=1S/Mg.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2 checkY
    Key: CSNNHWWHGAXBCP-UHFFFAOYSA-L checkY
  • InChI=1/Mg.8H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2
    Key: CSNNHWWHGAXBCP-NUQVWONBAQ
  • [Mg+2].[O-]S([O-])(=O)=O
Properties
MgSO4
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 (undecahydrate)
Melting point
  • decomposes
  • 1,124 °C (2,055 °F) (anhydrous)
  • 200 °C (392 °F) (monohydrate)
  • 150 °C (302 °F) (heptahydrate)
  • 2 °C (36 °F) (undecahydrate)
  • anhydrous
  • 26.9 g/100 mL @ 0 °C (32 °F)
  • 35.1 g/100 mL @ 20 °C (68 °F)
  • 50.2 g/100 mL @ 100 °C (212 °F)
  • heptahydrate
  • 113 g/100 mL @ 20 °C (68 °F)
738 (502 g/L)
Solubility in diethyl ether 1.16 g/100 mL @ 18 °C (64 °F)
Solubility in ethanol slight
Solubility in glycerol slight
−50×10−6 cm3/mol
  • 1.523 (monohydrate)
  • 1.433 (heptahydrate)
Structure
monoclinic (hydrate)
Pharmacology
A06AD04 (WHO) A12CC02 (WHO) B05XA05 (WHO) D11AX05 (WHO) V04CC02 (WHO)
Related compounds
Other cations
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)

Magnesium sulfate or magnesium sulphate is a chemical compound, a salt with the formula MgSO4, consisting of magnesium cations Mg2+ (20.19% by mass) and sulfate anions SO2−4. It is a white crystalline solid, soluble in water.

Magnesium sulfate is usually encountered in the form of a hydrate MgSO4·nH2O, for various values of n between 1 and 11. The most common is the heptahydrate MgSO4·7H2O,[1] known as Epsom salt, which is a household chemical with many traditional uses, including bath salts.[2]

The main use of magnesium sulfate is in agriculture, to correct soils deficient in magnesium (an essential plant nutrient because of the role of magnesium in chlorophyll and photosynthesis). The monohydrate is favored for this use; by the mid 1970s, its production was 2.3 million tons[clarification needed] per year.[3][page needed] The anhydrous form and several hydrates occur in nature as minerals, and the salt is a significant component of the water from some springs.

Hydrates

[edit]

Magnesium sulfate can crystallize as several hydrates, including:

As of 2017, the existence of the decahydrate apparently has not been confirmed.[9]

All the hydrates lose water upon heating. Above 320 °C (608 °F), only the anhydrous form is stable. It decomposes without melting at 1,124 °C (2,055 °F) into magnesium oxide (MgO) and sulfur trioxide (SO3).

Heptahydrate

[edit]

Main articles: Epsomite and Magnesium sulfate (medical use)

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 the impervious London clay.

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

It is a natural source of both magnesium and sulphur. Epsom salts are commonly used in bath salts, exfoliants, muscle relaxers and pain relievers.

Monohydrate

[edit]
Main article: Kieserite

Magnesium sulfate monohydrate, or kieserite, can be prepared by heating the heptahydrate to 120 °C (248 °F).[10] Further heating to 250 °C (482 °F) gives anhydrous magnesium sulfate.[10] Kieserite exhibits monoclinic symmetry at pressures lower than 2.7 gigapascals (27,000 atm) after which it transforms to phase of triclinic symmetry.[6]

Undecahydrate

[edit]

The undecahydrate MgSO4·11H2O, meridianiite, is stable at atmospheric pressure only below 2 °C (36 °F). Above that temperature, it liquefies into a mix of solid heptahydrate and a saturated solution. It has a eutectic point with water at −3.9 °C (25.0 °F) and 17.3% (mass) of MgSO4.[7] Large crystals can be obtained from solutions of the proper concentration kept at 0 °C (32 °F) for a few days.[7]

At pressures of about 0.9 gigapascals (8,900 atm) and at 240 K (−33 °C; −28 °F), meridianiite decomposes into a mixture of ice VI and the enneahydrate MgSO4·9H2O.[9]

Enneahydrate

[edit]

The enneahydrate MgSO4·9H2O was identified and characterized only recently, even though it seems easy to produce (by cooling a solution of MgSO4 and sodium sulfate (Na2SO4) in suitable proportions).

The structure is monoclinic, with unit-cell parameters at 250 K (−23 °C; −10 °F): a=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 MgSeO4·9H2O, but with a different crystal structure.[9]

Natural occurrence

[edit]

As Mg2+ and SO2−4 ions are respectively the second most abundant cation and anion present in seawater after Na+ and Cl−, 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.[citation needed]

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.[11]

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 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.[citation needed]

Preparation

[edit]

Magnesium sulfate is usually obtained directly from dry lake beds and other natural sources. It can also be prepared by reacting magnesite (magnesium carbonate, MgCO3) or magnesia (oxide, MgO) with sulfuric acid (H2SO4):

H2SO4 + MgCO3 → MgSO4 + H2O + CO2

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.

Also, magnesium sulfate heptahydrate (epsomite, MgSO4·7H2O) can be manufactured by dissolution of magnesium sulfate monohydrate (kieserite, MgSO4·H2O) in water and subsequent crystallization of the heptahydrate.

Physical properties

[edit]

Magnesium sulfate relaxation is the primary mechanism that causes the absorption of sound in seawater at frequencies above 10 kHz[12] (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.[13]

Uses

[edit]

Medical

[edit]
Main article: Magnesium sulfate (medical use)

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.[14] Health effects of magnesium sulfate that have been proposed include improvement of treatment resistant depression[15] and as an analgesic for migraine and chronic pain.[16]

Magnesium sulfate is usually the main component of the concentrated salt solution used in isolation tanks to increase its specific gravity to approximately 1.25–1.26. This high density allows an individual to float effortlessly on the surface of water in the closed tank, eliminating stimulation of as many of the external senses as possible.[17]

In the UK, a medication containing magnesium sulfate and phenol, called "drawing paste", is useful for small boils or localized infections[18] and removing splinters.[19]

Internally, magnesium sulfate may be administered by oral, respiratory, or intravenous routes. Internal uses include replacement therapy for magnesium deficiency, [20] [21] treatment of acute and severe arrhythmias,[21] as a bronchodilator in the treatment of asthma,[22] preventing eclampsia[23] and cerebral palsy,[24][25] a tocolytic agent,[26] and as an anticonvulsant.[26] The effectiveness and safety of magnesium sulfate for treating acute bronchiolitis in children under the age of 2 years old is not well understood.[27]

It also may be used as laxative.[28]

Agriculture

[edit]

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 the slightly 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.[26] Contrary to the popular belief that magnesium sulfate is able to control pests and slugs, helps seeds germination, produce more flowers, improve nutrient uptake, and is environmentally friendly, it does none of the purported claims except for correcting magnesium deficiency in soils. Magnesium sulfate can even pollute water if used in excessive amounts.[29]

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.[30] 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 had inadvertently ingested lead.[31][32]

Food preparation

[edit]

Magnesium sulfate is used as:

[34]

Chemistry

[edit]

Anhydrous magnesium sulfate is commonly used as a desiccant in organic synthesis owing to its affinity for water and compatibility with most organic compounds.[37] During work-up, an organic phase is treated with anhydrous magnesium sulfate. The hydrated solid is then removed by filtration, decantation, or by 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.

Construction

[edit]

Magnesium sulfate is used to prepare specific cements by the reaction between magnesium oxide and magnesium sulfate solution.[38] This cement is mainly used in the production of lightweight insulation panels, although its poor 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 is also used to test the resistance of concrete to external sulfate attack (ESA).[39]

Aquaria

[edit]

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 difficult to control because not enough magnesium is present to stabilize these ions in the saltwater and prevent their spontaneous precipitation into calcium carbonate.[40]

Double salts

[edit]

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 found to occur in mine tailings and was given the mineral name alpersite.[41]

See also

[edit]

References

[edit]
[edit]
Revisions and contributorsEdit on WikipediaRead on Wikipedia

Magnesium sulfate

View on Grokipedia
from Grokipedia
Magnesium sulfate is an inorganic salt with the MgSO₄ and a molecular weight of 120.37 g/mol, appearing as a white, odorless, crystalline solid with a saline, bitter . It commonly exists in hydrated forms, most notably the heptahydrate (MgSO₄·7H₂O), known as Epsom salt, which is highly soluble in and used for various applications. In , magnesium sulfate serves as an , , and replenisher, particularly for treating and by inhibiting myometrial muscle action potentials and blocking calcium influx to prevent seizures. It is also indicated for hypomagnesemia, acute in children, uterine , and , with administration routes including intravenous, intramuscular, oral, or topical forms. Common side effects include , , and respiratory depression, requiring careful monitoring. Beyond healthcare, magnesium sulfate functions as a providing essential magnesium and nutrients for crops, helping to correct deficiencies and improve growth in . Industrially, it acts as a in , a component in textiles, processing, explosives, and ceramics, while also serving as a bath soak for muscle relief and a in . Safety considerations include mild irritation to eyes and , with recommendations to avoid or .

Chemical identity

Formula and molecular structure

Magnesium sulfate is an ionic compound with the \ceMgSO4\ce{MgSO4}, composed of the magnesium dication \ceMg2+\ce{Mg^2+} and the dianion \ceSO42−\ce{SO4^2-}. This formula represents the form, where the magnesium achieves a +2 through the loss of its two valence electrons, balanced by the -2 charge of the tetrahedral . The molecular weight of anhydrous magnesium sulfate is 120.366 g/mol, calculated from the atomic masses of magnesium (24.305 g/mol), sulfur (32.065 g/mol), and four oxygen atoms (15.999 g/mol each). As an ionic salt, it exhibits strong electrostatic attractions between the oppositely charged ions, resulting in a stable crystal lattice rather than discrete molecules. This ionic bonding nature distinguishes it from covalent compounds, contributing to its high melting point and solubility characteristics, though the latter are addressed elsewhere. In the anhydrous state, magnesium sulfate crystallizes in an orthorhombic lattice with space group Cmcm (No. 63) and four formula units per unit cell (Z = 4). The atomic arrangement features \ceMg2+\ce{Mg^2+} cations octahedrally coordinated by six oxygen atoms from surrounding \ceSO42−\ce{SO4^2-} anions, forming distorted \ceMgO6\ce{MgO6} octahedra. These octahedra share corners with adjacent sulfate tetrahedra, creating a three-dimensional framework that stabilizes the ionic structure through a network of \ceMg−O\ce{Mg-O} and \ceS−O\ce{S-O} bonds, with \ceMg−O\ce{Mg-O} distances typically around 2.05–2.15 Å at low temperatures. This coordination geometry reflects the preference of \ceMg2+\ce{Mg^2+} for octahedral environments due to its ionic radius and charge density.

Hydrated forms

Magnesium sulfate exists in several hydrated forms, distinguished by the number of water molecules incorporated into their lattices, which influences their stability and applications. The most common hydrates include the monohydrate (MgSO₄·H₂O, known as ) and the heptahydrate (MgSO₄·7H₂O, known as or Epsom salt). Less frequently encountered are higher hydrates such as the enneahydrate (MgSO₄·9H₂O) and the undecahydrate (MgSO₄·11H₂O, known as meridianiite), which form under specific environmental conditions like high humidity or low temperatures. The heptahydrate adopts an orthorhombic crystal structure with space group P2₁2₁2₁ and lattice parameters a = 11.86 Å, b = 11.99 Å, and c = 6.86 Å, featuring coordinated water molecules that stabilize the sulfate ions within the lattice. In contrast, the monohydrate exhibits a monoclinic structure, making it more compact and less hygroscopic than the heptahydrate. The enneahydrate and undecahydrate possess more complex monoclinic and triclinic structures, respectively, with extended hydrogen-bonding networks that accommodate additional water layers, though these forms are metastable under standard atmospheric conditions. Phase transitions among these hydrates occur through processes driven by and . For instance, the heptahydrate transitions to the hexahydrate at approximately 48.5°C via incongruent melting at the with the , with further to lower hydrates like the monohydrate requiring temperatures above 75°C under controlled vapor pressures. Stability fields shift with environmental conditions; lower s and higher relative favor more hydrated forms, while the monohydrate remains stable up to 150°C before decomposing to the state. These transitions are reversible under appropriate , enabling hydration- cycles in applications like . Commercially, the heptahydrate is widely used as Epsom salt in bath products for its soothing properties and in medical applications as a , owing to its high and mild osmotic effects. The monohydrate, or , serves primarily as a to supply magnesium and to crops, particularly in magnesium-deficient soils, due to its stability and ease of handling in agricultural formulations. Higher hydrates like the enneahydrate and undecahydrate have limited commercial roles but are studied for potential in thermochemical storage systems exploiting their hydration reversibility.

History

Discovery and early uses

The medicinal properties of what would later be identified as magnesium sulfate were first observed in 1618 near , , during a period of . A local cowherd named Henry Wicker noticed his cattle avoiding a spring due to its bitter taste but, upon drinking from it himself, experienced improvements in his and skin condition, prompting others to use the water for treating sores, rashes, and digestive issues. This discovery quickly drew visitors seeking relief from various ailments, transforming the site into one of 's earliest health spas by the mid-17th century, where the mineral-rich waters were bathed in or ingested for their purported curative effects. The compound gained scientific recognition in 1695 through the work of English physician and botanist , who conducted a detailed chemical of the Epsom spring water and successfully isolated the solid crystalline form. Grew documented his findings in the treatise A Treatise of the Nature and Use of the Bitter Purging Salt Contain'd in Epsom and Such Other Wells, naming it "Epsom salts" after its origin and highlighting its composition as a of an unknown earth (later determined to be magnesium). His efforts marked the first systematic study of the substance, confirming its presence in similar bitter springs across and establishing its potential for broader medicinal application beyond natural bathing. Early adoption of salts focused on pre-industrial medicinal uses throughout the 17th and 18th centuries, primarily as an oral purgative to alleviate , cleanse the digestive system, and treat conditions like or urinary stones. Externally, it was dissolved in baths to soothe inflamed , heal wounds, and ease muscle aches, capitalizing on its osmotic properties to draw out impurities. Commercialization accelerated around when Grew sought a royal patent for its production and sale, sparking a legal dispute with apothecary brothers Francis and George Moult, who claimed prior independent and began mass-producing the salts at a factory near to meet growing demand from physicians and the public. This rivalry underscored the substance's rapid transition from folk remedy to a standardized pharmaceutical product.

Industrial development

In the , chemical advancements facilitated the isolation of pure through the reaction of with , enabling higher-purity production for emerging industrial needs. This method, developed amid growing chemical manufacturing capabilities, supported the compound's application in various processes. Concurrently, during the , 's use as a expanded significantly to address magnesium deficiencies in soils, enhancing crop yields and aligning with the era's push for intensified agriculture. The marked key milestones in magnesium sulfate production, with mass-scale extraction from and brines commencing after the , leveraging bitterns from salt as a source. The first reported commercial production of natural magnesium sulfate occurred in 1923, primarily from deposits in regions like Washington and , scaling up to meet industrial demands. In , the saw standardization of magnesium sulfate for treatment, with intravenous administration popularized by physician Edmond M. at General Hospital starting in 1924, reducing maternal mortality from seizures. By 2024, global production of magnesium sulfate surpassed 2.5 million tons annually, fueled by rising pharmaceutical applications for conditions like and strong agricultural demand for soil amendment in magnesium-deficient regions. This growth reflects ongoing innovations in extraction efficiency and market expansion, particularly in where agricultural use dominates.

Natural occurrence

Mineral forms

Magnesium sulfate occurs naturally as several hydrated forms, primarily (MgSO₄·7H₂O), the heptahydrate, which typically appears as white, powdery, fibrous, or botryoidal crusts formed through efflorescence on magnesium-rich rocks or around mineral springs. Another key is (MgSO₄·H₂O), the monohydrate, which forms coarse- to fine-grained masses in deposits and is noted for its slow solubility in water. (Na₂SO₄·10H₂O), a , often co-occurs with these magnesium sulfate forms in mixed settings, such as efflorescent crusts alongside and . Major deposits of these minerals include the Stassfurt evaporite basin in , a primary source of since the 19th century, where it intergrows with , , and in marine salt layers. Other notable deposits include the Carlsbad potash district in , , where is commercially mined. In volcanic regions, and appear rarely as fumarolic encrustations from gas exhalations. These minerals form primarily through the of sulfate-rich waters in arid environments, where dissolved precipitates as hydrates in basins, salt lakes, or near-surface settings like mineral springs and fumaroles. This process is common in closed basins with high rates, leading to sequential of salts.

In natural waters

Magnesium sulfate occurs naturally in various water bodies, primarily as dissolved ions of magnesium (Mg²⁺) and (SO₄²⁻), contributing significantly to overall . In typical , the concentration of Mg²⁺ is approximately 1.29 g/L, while SO₄²⁻ is about 2.71 g/L at a salinity of 35 practical salinity units (psu), making these ions major components that together account for roughly 10% of the total dissolved salts. These concentrations vary slightly with geographic location and depth but remain relatively conservative due to the long residence times of these elements in the , influencing the and circulation patterns of marine waters. High-concentration sources of magnesium sulfate are found in hypersaline brines, such as those in the Dead Sea and . In the Dead Sea, magnesium concentrations exceed 40 g/L, predominantly as MgCl₂ but with sulfate present at levels around 0.3–0.4 g/L, in this extremely saline environment (total salinity ~340 g/L). Similarly, brines in the north arm of the reach magnesium levels over 80 g/L in concentrated solar evaporation ponds, accompanied by sulfate concentrations up to 50 g/L, forming a magnesium sulfate-rich subtype that supports mineral extraction industries. Geochemically, magnesium sulfate plays a key role in marine cycles by facilitating interactions between the magnesium, , and carbon cycles through processes like authigenic formation and oceanic . In marine sediments, Mg²⁺ and SO₄²⁻ contribute to diagenetic reactions that influence global and chemistry over geological timescales. Additionally, concentrated magnesium sulfate appears as a byproduct in processes, where or thermal methods enrich brines, yielding MgSO₄ suitable for recovery without calcium impurities.

Preparation

Extraction from natural sources

Magnesium sulfate is primarily extracted from natural sources through underground mining of kieserite deposits and solar evaporation of brines containing dissolved magnesium sulfate. These methods leverage geological formations where the compound has concentrated over time through evaporation processes in ancient seas or lakes. Kieserite (MgSO₄·H₂O), the monohydrate form of magnesium sulfate, is obtained via conventional underground mining from evaporite deposits in salt domes, particularly in Germany where such formations are abundant. Major producers like K+S extract kieserite as a byproduct during potash mining from sylvinite ores; the raw mineral is mined using room-and-pillar or longwall techniques, then ground and processed using the ESTA® electrostatic separation to remove impurities such as halite and clay, yielding a high-purity product. This process supplies a significant portion of the global magnesium sulfate market, with Germany's deposits being uniquely rich in kieserite due to specific depositional conditions during the Permian period. The heptahydrate form, epsomite (MgSO₄·7H₂O), is recovered through solar evaporation of natural brines from seawater, salt lakes, or subsurface waters rich in sulfates. In this technique, brine is pumped into shallow evaporation ponds where solar heat and wind concentrate the solution, causing less soluble salts like calcium sulfate to precipitate first, followed by the crystallization of epsomite as the magnesium concentration increases. Historical and ongoing operations, such as those at ancient deposits in Washington state, utilize this method to harvest epsomite crystals, which are then separated by flotation or filtration, dissolved if needed for purification, and recrystallized to achieve pharmaceutical or industrial grades. From high-sulfate s, evaporation processes produce with high purity after processing, though yields can vary based on initial brine composition and climatic conditions.

Synthetic production

Magnesium sulfate is synthetically produced primarily through the neutralization of with , following the reaction: MgO+H2SO4→MgSO4+H2O\text{MgO} + \text{H}_2\text{SO}_4 \rightarrow \text{MgSO}_4 + \text{H}_2\text{O} This is conducted in under controlled conditions to form the soluble magnesium sulfate, which is subsequently evaporated and crystallized to yield hydrated forms such as the heptahydrate (MgSO₄·7H₂O). The precursor is typically obtained from calcined or precipitated , ensuring high purity in the final product through recrystallization steps. Alternative synthetic routes include the reaction of with , which proceeds as a double displacement: MgCl2+H2SO4→MgSO4+2HCl\text{MgCl}_2 + \text{H}_2\text{SO}_4 \rightarrow \text{MgSO}_4 + 2\text{HCl} This method recovers magnesium sulfate from solutions derived from ores or , while generating as a valuable byproduct. Additionally, magnesium sulfate can be obtained from byproducts of electrolytic magnesium production, where excess is treated with to convert it into the sulfate form. China is the world's largest producer of magnesium sulfate, accounting for a significant portion of global supply through both natural and synthetic methods as of 2023. On an industrial scale, these syntheses are performed in batch reactors to produce either magnesium sulfate or its hydrates, with the choice depending on end-use requirements. The processes are energy-efficient compared to metal extraction but require precise control of and to minimize impurities. While natural extraction from minerals like remains a cost-effective alternative for large volumes, synthetic methods offer greater flexibility for high-purity applications.

Physical properties

Appearance and crystal structure

Magnesium sulfate exists in various solid forms, with the anhydrous variant appearing as a , hygroscopic powder composed of orthorhombic crystals in the Cmcm . This form readily absorbs moisture from the air due to its hygroscopic nature, often leading to clumping in storage. The most common hydrate, magnesium sulfate heptahydrate (MgSO₄·7H₂O), also known as , manifests as colorless, prismatic or needle-like crystals that are odorless and possess a cool, bitter taste. These crystals have a of 1.68 g/cm³ and adopt an . X-ray diffraction serves as a key method for identifying magnesium sulfate forms, particularly the heptahydrate, which displays characteristic peaks at 2θ values around 14.8°, 21.0°, and 23.6° under Cu Kα radiation. These diffraction patterns confirm the structural integrity and purity of the crystals in analytical contexts.

Solubility and density

Magnesium sulfate's solubility in water depends on its form and temperature, with the heptahydrate demonstrating higher mass solubility due to its water content. The heptahydrate (MgSO₄·7H₂O) dissolves at approximately 710 g/L at 20 °C, and solubility rises with increasing temperature, reaching higher concentrations as thermal energy facilitates dissociation. The form (MgSO₄) is less soluble on a basis, with a of about 35 g per 100 g of at 20 °C, though it also exhibits increasing with —for instance, up to 74 g per 100 g of at 100 °C. Densities differ markedly between forms: magnesium sulfate has a of 2.66 g/cm³, reflecting its compact ionic lattice, while hydrated forms have reduced densities due to incorporated molecules; the heptahydrate, for example, measures 1.68 g/cm³. Solubility shows a positive dependence, approximately linear over common ranges, underscoring the endothermic dissolution that drives higher saturation at elevated temperatures.

Chemical properties

Reactivity and stability

Magnesium sulfate is chemically stable under standard ambient conditions, including and neutral , without significant or hazardous reactions. Its and hydrated forms remain intact in dry storage and show no tendency to react with air or common inert materials. In aqueous solutions, magnesium sulfate exhibits slight hydrolysis due to the Mg²⁺ ion, resulting in a mildly acidic environment. A 0.1 M solution has a pH of approximately 6, ranging from 5.5 to 6.5 depending on concentration and preparation. This acidity arises from the equilibrium Mg(H₂O)₆²⁺ ⇌ Mg(H₂O)₅OH⁺ + H⁺, with a hydrolysis constant Kh ≈ 3.6 × 10−12 at 25°C. The minor production of H⁺ ions leads to limited formation of basic magnesium species, such as hydroxy complexes, but the overall solution remains stable without precipitation under dilute, ambient conditions. Magnesium sulfate demonstrates reactivity through precipitation reactions with anions like carbonate and hydroxide, which are key in analytical chemistry. It forms a white, insoluble magnesium carbonate precipitate upon addition of carbonate ions, as in MgSO₄ + CO₃²⁻ → MgCO₃ ↓ + SO₄²⁻./Qualitative_Analysis/Characteristic_Reactions_of_Select_Metal_Ions/Characteristic_Reactions_of_Magnesium_Ions_(Mg)) Similarly, hydroxide ions yield a gelatinous white magnesium hydroxide precipitate: MgSO₄ + 2OH⁻ → Mg(OH)₂ ↓ + SO₄²⁻, which is sparingly soluble in water (Ksp ≈ 5.6 × 10−12)./Qualitative_Analysis/Characteristic_Reactions_of_Select_Metal_Ions/Characteristic_Reactions_of_Magnesium_Ions_(Mg)) These specific precipitations are employed in qualitative analysis schemes to confirm the presence of magnesium ions in unknown samples, often after separation from interfering metals.

Thermal decomposition

Magnesium sulfate, commonly encountered as the heptahydrate (MgSO₄·7H₂O), undergoes stepwise upon heating, losing molecules to form lower hydrates and eventually the form. The initial step converts the heptahydrate to the hexahydrate (MgSO₄·6H₂O) at approximately 81 °C, followed by to the monohydrate (MgSO₄·H₂O) at around 110 °C, and finally to MgSO₄ at 258 °C. This multi-stage process, spanning roughly 80–280 °C, is endothermic, absorbing during each release and enabling potential applications in systems. No products form during these steps, preserving the structure. The anhydrous magnesium sulfate remains stable up to about 850 °C but decomposes at higher temperatures via the endothermic reaction: MgSO4→MgO+SO3\text{MgSO}_4 \rightarrow \text{MgO} + \text{SO}_3 This decomposition initiates around 875 °C and completes near 1044 °C, producing (MgO) and (SO₃) gas, with SO₃ potentially further breaking down to SO₂ and O₂ under certain conditions. The reaction's strong endothermicity requires significant energy input, influencing process design in high-temperature environments. This , particularly the of MgSO₄, serves as a method to produce magnesia (MgO), a key industrial material used in refractories and chemicals, often sourced from sulfate-rich feedstocks.

Uses

Medical applications

Magnesium sulfate is a cornerstone therapy in for the prevention and treatment of seizures associated with , a severe complication of in . The recommended regimen includes an initial intravenous of 4 to 6 g administered over 15 to 20 minutes, followed by a continuous maintenance infusion of 1 to 2 g per hour until delivery or resolution of symptoms, with monitoring of serum magnesium levels to avoid toxicity. This approach has been shown to reduce the risk of recurrent seizures by over 50% compared to other anticonvulsants like . In the management of hypomagnesemia, magnesium sulfate serves as the primary replenishment agent, particularly in cases of severe deficiency often seen in critically ill patients, alcoholics, or those with gastrointestinal losses. Intravenous administration of 1 to 2 g over 5 to 60 minutes is typical for acute correction, with subsequent dosing adjusted based on serum levels to achieve 1.5 to 2.5 mEq/L. Oral forms may be used for milder cases or maintenance, though absorption is variable. As an osmotic , oral magnesium sulfate is employed for rapid bowel evacuation, such as in for procedures or treatment of . A standard dose is 15 g dissolved in 8 ounces of , producing a effect within 30 minutes to 6 hours by drawing into the intestinal lumen. This non-absorbable salt is particularly useful when prompt colonic cleansing is required. The mechanism of magnesium sulfate in involves antagonism of N-methyl-D-aspartate (NMDA) receptors in the , which inhibits excitatory and stabilizes neuronal membranes to prevent propagation. For its laxative action, magnesium sulfate exerts an osmotic effect in the gut, where unabsorbed magnesium and sulfate ions retain fluid, increasing intraluminal volume and promoting . Recent developments include its investigation as an intravenous adjunct for acute exacerbations, particularly in children. A 2025 systematic review and of randomized controlled trials demonstrated that adding IV magnesium sulfate to standard therapies (inhaled beta-agonists and systemic corticosteroids) significantly lowers hospitalization rates (risk ratio 0.70, 95% CI 0.54-0.90), with benefits also seen in reduced need for . This bronchodilatory effect stems from magnesium's blockade, relaxing airway .

Agricultural applications

Magnesium sulfate serves as an essential fertilizer and soil amendment in agriculture, providing plants with magnesium (Mg) and sulfur (S), two critical secondary nutrients often deficient in intensively cropped soils. It is particularly effective in correcting Mg deficiencies that limit plant growth and yield, as Mg is a key component of chlorophyll and various enzymes involved in photosynthesis and metabolism. The compound is commonly applied in forms such as kieserite (MgSO₄·H₂O), which contains 20-27% MgO, allowing for targeted supplementation in soils low in these elements. Global agricultural consumption of magnesium sulfate exceeds 1.5 million metric tons annually, with significant use in crops such as and that are prone to Mg and S deficiencies. In production, it helps maintain sulfur levels essential for protein synthesis and oil formation, while in , applications have been shown to increase leaf yield and nutrient concentrations. This widespread adoption underscores its role in enhancing overall crop productivity in regions with marginal soils. One primary benefit of magnesium sulfate is its contribution to chlorophyll formation, which promotes vigorous and greener foliage, ultimately boosting and harvestable yield. For rapid correction of deficiencies, it is often used in foliar sprays at concentrations of 1-2%, enabling quick absorption through leaf surfaces during critical growth stages. This method is especially valuable for high-value crops, where timely nutrient delivery can prevent and support optimal development.

Industrial applications

In , magnesium sulfate serves as an effective drying agent due to its strong hygroscopic properties, absorbing from solvents and reaction mixtures to facilitate purification without reacting with most organic compounds. It forms hydrates upon uptake, with a theoretical capacity equivalent to seven molecules of per , enabling efficient removal of residual moisture in and industrial processes. In the , magnesium sulfate, designated as additive E345, functions as a firming agent in canned , helping to maintain texture and structural during and storage. It is typically used in controlled amounts to prevent softening without altering flavor or nutritional profile significantly. Magnesium sulfate contributes to construction-related applications, particularly in , where it is incorporated as a for fabrics, reducing flammability by releasing upon heating and forming a protective char layer. In , especially systems, it is added to adjust magnesium levels, targeting 1250–1350 ppm to support health, stabilize calcium and balance, and mimic natural conditions. Additionally, in leather tanning, magnesium sulfate enhances suppleness by preventing drying during processing and aiding the penetration and binding of tanning agents to fibers. In production, it supplies magnesium ions that strengthen pulp fibers, improve sheet formation, and act as a protector against degradation during bleaching and delignification stages.

Double salts

Common double salts

Double salts of are compounds formed by the combination of with another salt, typically involving metals such as sodium or , resulting in crystalline structures with shared anions and of hydration. These salts arise through co-crystallization processes when solutions containing and the corresponding are evaporated or cooled under controlled conditions. One principal example is astrakhanite, also known as blödite, with the \ceNa2Mg(SO4)2⋅4H2O\ce{Na2Mg(SO4)2 \cdot 4H2O}. This sodium-magnesium sulfate forms via the slow evaporation of mixed aqueous solutions of and magnesium sulfate, leading to its precipitation as colorless to white monoclinic crystals. Astrakhanite occurs naturally in deposits, such as those in saline lakes like Lake Bai Shagyr in , where it precipitates during the evaporation of sulfate-rich brines. Another key double salt is schönite, synonymous with picromerite, having the formula \ceK2Mg(SO4)2⋅6H2O\ce{K2Mg(SO4)2 \cdot 6H2O}. It is produced synthetically by co-crystallization from solutions of and magnesium sulfate through methods like isothermal evaporation, yielding orthorhombic prismatic crystals. Naturally, schönite is found in formations in alkaline lakes and deposits, where sequential evaporation of mixed sulfate waters favors its formation over individual salts.

Properties and uses

Double salts of magnesium sulfate demonstrate enhanced stability relative to individual single salts within mixed aqueous systems, where they form persistent crystalline phases that resist dissociation under varying temperature and concentration conditions. This stability arises from the integrated lattice structures incorporating both magnesium and counter-cation sulfates, as observed in phase equilibrium studies of sulfate systems. For instance, schönite (K₂SO₄·MgSO₄·6H₂O), a prominent potassium-magnesium , exhibits a of approximately 111 g/L at 20°C, balancing adequate dissolution with controlled release in practical applications. These double salts find primary use as fertilizers, delivering potassium-magnesium blends that supply essential nutrients— for fruit development, for synthesis, and for protein formation—promoting balanced crop nutrition and higher yields. They are particularly valuable for chlorine-intolerant crops like , potatoes, and , where traditional fertilizers pose risks, offering a chlorine-free alternative that supports vigorous growth and stress resistance. Ammonium-magnesium double salts, such as (NH₄)₂Mg(SO₄)₂·6H₂O, are also utilized in industrial fertilizer formulations derived from processes, enhancing economic viability through recovery. A key advantage of these double salts lies in their superior , which enables efficient nutrient delivery via foliar sprays or systems, ensuring rapid uptake by and minimizing losses from fixation compared to less soluble single-salt formulations. This property not only improves nutrient efficiency in but also supports eco-friendly practices by reducing overall application rates.

Safety and toxicology

Health effects

Magnesium sulfate exhibits low via oral exposure, with an LD50 exceeding 2,000 mg/kg body weight in rats according to Test Guideline 425. Acute poisoning primarily manifests as when significant amounts are ingested or administered intravenously, leading to mild symptoms such as , , flushing, , and at serum magnesium levels above approximately 2 mmol/L. As serum concentrations rise beyond 4-5 mmol/L, more severe effects emerge, including due to and neuromuscular blockade, electrocardiographic changes like prolonged PR intervals, and loss of deep tendon reflexes. At levels of 5-7.5 mmol/L, respiratory depression and can occur, potentially progressing to and if untreated. Prolonged misuse of magnesium sulfate, such as repeated overuse as a , can result in imbalances, including , , and disturbances in and sodium levels, particularly in individuals with impaired renal function. These imbalances may contribute to cardiac arrhythmias, , and gastrointestinal disturbances over time. Within its therapeutic window, magnesium sulfate is generally safe when administered at controlled medical doses, such as intravenous infusions maintaining serum levels of 2-3.5 mmol/L for conditions like . However, intravenous overdose poses significant risks, including rapid onset of respiratory depression due to relaxation, which can lead to apnea and require in severe cases. Management includes discontinuing magnesium, administering IV to counteract effects, hydration, and in cases of renal impairment. Monitoring of serum levels, reflexes, and respiratory function is essential to prevent such complications.

Handling and environmental impact

Magnesium sulfate should be stored in a cool, dry environment, ideally between 68°F and 110°F with relative of 54% to 87%, to minimize caking caused by absorption. During handling, generation must be minimized through appropriate and , such as respirators, gloves, and , to prevent or skin contact. The (OSHA) (PEL) for respirable from magnesium sulfate, classified as particulates not otherwise regulated, is 5 mg/m³ as an 8-hour time-weighted average. Environmentally, magnesium sulfate exhibits low toxicity to aquatic organisms, with LC50 values exceeding 1000 mg/L for fish such as Gambusia affinis (15,500 mg/L at 96 hours) and invertebrates like ( of 1700 mg/L at 24 hours). However, agricultural runoff containing magnesium sulfate can elevate concentrations in surface waters, contributing to increased and osmotic stress on sensitive aquatic species in soft freshwaters, where chronic no-effect levels are as low as 39–65 mg SO₄/L. In soils, such runoff promotes salinization, which can alter balance and impair growth by disrupting uptake and . Under the EU REACH framework, magnesium sulfate is not subject to harmonized hazard classification and is generally regarded as non-hazardous, though some registrations note potential for allergic skin reactions due to impurities. The EU Fertilising Products Regulation (EU 2019/1009), as amended to include provisions for digital labelling of nutrient content as of 2024, supports mitigation of from use through improved transparency and compliance.

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