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2Xray.jpg) | |
| Names | |
|---|---|
| IUPAC name Magnesium hydroxide | |
Other names
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| Identifiers | |
3D model (JSmol) | |
| ChEBI | |
| ChEMBL | |
| ChemSpider |
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| DrugBank | |
| ECHA InfoCard | 100.013.792 |
| EC Number |
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| E number | E528(acidity regulators, ...) |
| 485572 | |
| KEGG | |
| RTECS number |
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| UNII | |
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| Properties | |
| Mg(OH)2 | |
| Molar mass | 58.3197 g/mol |
| Appearance | White solid |
| Odor | Odorless |
| Density | 2.3446 g/cm3 |
| Melting point | 350 °C (662 °F; 623 K) decomposes |
| |
Solubility product (Ksp) | 5.61×10−12 |
| −22.1×10−6 cm3/mol | |
Refractive index (nD) | 1.559[1] |
| Structure | |
| Hexagonal,hP3[2] | |
| P3m1 No. 164 | |
a = 0.312 nm,c = 0.473 nm | |
| Thermochemistry | |
| 77.03 J/mol·K | |
Std molar entropy(S⦵298) | 64 J·mol−1·K−1[3] |
Std enthalpy of formation(ΔfH⦵298) | −924.7 kJ·mol−1[3] |
Gibbs free energy(ΔfG⦵) | −833.7 kJ/mol |
| Pharmacology | |
| A02AA04 (WHO) G04BX01 (WHO) | |
| Hazards | |
| GHS labelling: | |
 [4] | |
| Warning[4] | |
| H315,H319,H335[4] | |
| P261,P280,P304+P340,P305+P351+P338,P405,P501[4] | |
| NFPA 704 (fire diamond) | |
| Flash point | Non-flammable |
| Lethal dose or concentration (LD, LC): | |
LD50 (median dose) | 8500 mg/kg (rat, oral) |
| Safety data sheet (SDS) | External MSDS |
| Related compounds | |
Otheranions | Magnesium oxide |
Othercations | |
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |
Magnesium hydroxide is aninorganic compound with the chemical formula Mg(OH)2. It occurs in nature as the mineralbrucite. It is a white solid with low solubility in water (Ksp =5.61×10−12).[5] Magnesium hydroxide is a common component ofantacids, such asmilk of magnesia.
Treating the solution of different soluble magnesium salts withalkaline water induces theprecipitation of the solid hydroxide Mg(OH)2:
AsMg2+
is the second most abundant cation present inseawater afterNa+
, it can be economically extracted directly from seawater byalkalinisation as described above. On an industrial scale, Mg(OH)2 is produced by treating seawater withlime (Ca(OH)2). A volume of 600 m3 (160,000 US gal) of seawater gives about 1 tonne (2,200 lb) of Mg(OH)2. Ca(OH)2(Ksp =5.02×10−6)[6] is far more soluble than Mg(OH)2(Ksp =5.61×10−12) and dramatically increases thepH value of seawater from 8.2 to 12.5. The less solubleMg(OH)
2 precipitates because of thecommon ion effect due to theOH−
added by the dissolution ofCa(OH)
2:[7]
For seawater brines, precipitating agents other thanCa(OH)2 can be utilized, each with their own nuances:
It has been demonstrated that sodium hydroxide,NaOH, is the better precipitating agent compared toCa(OH)2 andNH4OH due to higher recovery and purity rates, and the settling and filtration time can be improved at low temperatures and higher concentration of precipitates. Methods involving the use of precipitating agents are typically batch processes.[8]
It is also possible to obtainMg(OH)2 from seawater using electrolysis chambers separated with acation exchange membrane. This process is continuous, lower-cost, and produces oxygen gas, hydrogen gas, sulfuric acid (ifNa2SO4 is used;NaCl can alternatively be used to yieldHCl), andMg(OH)2 of 98% or higher purity. It is crucial todeaerate the seawater to mitigate co-precipitation of calcium precipitates.[9]
Most Mg(OH)2 that is produced industrially, as well as the small amount that is mined, is converted to fusedmagnesia (MgO). Magnesia is valuable because it is both a poor electrical conductor and an excellent thermal conductor.[7]
Only a small amount of the magnesium from magnesium hydroxide is usually absorbed by the intestine (unless one is deficient in magnesium). However, magnesium is mainly excreted by the kidneys; so long-term, daily consumption of milk of magnesia by someone suffering from kidney failure could lead in theory tohypermagnesemia. Unabsorbed magnesium is excreted in feces; absorbed magnesium is rapidly excreted in urine.[10]
As an antacid, magnesium hydroxide is dosed at approximately 0.5–1.5 g in adults and works by simpleneutralization, in which thehydroxideions from the Mg(OH)2 combine withacidic H+ions (orhydronium ions) produced in the form of hydrochloric acid byparietal cells in thestomach, to produce water.
As a laxative, magnesium hydroxide is dosed at 5–10 grams (0.18–0.35 oz), and works in a number of ways. First, Mg2+ is poorly absorbed from the intestinal tract, so it draws water from the surrounding tissue byosmosis. Not only does this increase in water content to soften the feces, it also increases the volume of feces in the intestine (intraluminal volume) which naturally stimulates intestinalmotility. Furthermore, Mg2+ ions cause the release ofcholecystokinin (CCK), which results in intraluminal accumulation of water and electrolytes, and increased intestinal motility. Some sources claim that the hydroxide ions themselves do not play a significant role in the laxative effects of milk of magnesia, as alkaline solutions (i.e., solutions of hydroxide ions) are not strongly laxative, and non-alkaline Mg2+ solutions, likeMgSO4, are equally strong laxatives,mole for mole.[11]
On May 4, 1818, American inventor Koen Burrows received a patent (No. X2952) for magnesium hydroxide.[12] In 1829,Sir James Murray used a "condensed solution of fluid magnesia" preparation of his own design[13] to treat theLord Lieutenant of Ireland, theMarquess of Anglesey, for stomach pain. This was so successful (advertised in Australia and approved by the Royal College of Surgeons in 1838)[14] that he was appointed resident physician to Anglesey and two subsequent Lords Lieutenant, and knighted. His fluid magnesia product was patented two years after his death, in 1873.[15]
The termmilk of magnesia was first used byCharles Henry Phillips in 1872 for asuspension of magnesium hydroxide formulated at about 8% w/v.[16] It was sold under the brand namePhillips' Milk of Magnesia for medicinal usage.
USPTO registrations show that the terms "Milk of Magnesia"[17] and "Phillips' Milk of Magnesia"[18] have both been assigned toBayer since 1995. In the UK, the non-brand (generic) name of "Milk of Magnesia" and "Phillips' Milk of Magnesia" is "Cream of Magnesia" (Magnesium Hydroxide Mixture,BP).
It is added directly to human food, and is affirmed asgenerally recognized as safe by theFDA.[19] It is known asE numberE528.
Magnesium hydroxide is marketed for medical use in the form of chewable tablets, capsules, powder, and as liquidsuspensions, which are sometimes flavored. These products are sold asantacids to neutralize stomachacid and relieveindigestion andheartburn.
It is also a laxative used to alleviateconstipation. As a laxative, theosmotic force of the magnesia acts to draw fluids from the body. High doses can lead todiarrhea and can deplete the body's supply ofpotassium, sometimes leading tomuscle cramps.[20] Some magnesium hydroxide products sold for antacid use (such asMaalox) are formulated to minimize unwanted laxative effects through the inclusion ofaluminum hydroxide, which inhibits the contractions ofsmooth muscle cells in the gastrointestinal tract,[21] thereby counterbalancing the contractions induced by the osmotic effects of the magnesium hydroxide.
Magnesium hydroxide is also a component ofantiperspirant.[22]
Magnesium hydroxide powder is used industrially to neutralize acidic wastewaters.[23] It is also a component of theBiorock method of buildingartificial reefs. The main advantage ofMg(OH)
2 overCa(OH)
2, is to impose a lower pH better compatible with that of seawater and sea life: pH 10.5 forMg(OH)
2 in place of pH 12.5 withCa(OH)
2.
Natural magnesium hydroxide (brucite) is used commercially as a fire retardant. Most industrially used magnesium hydroxide is produced synthetically.[24] Like aluminum hydroxide, solid magnesium hydroxide has smoke suppressing andflame retardant properties. This property is attributable to theendothermic decomposition it undergoes at 332 °C (630 °F):
The heat absorbed by the reaction retards the fire by delaying ignition of the associated substance. The water released dilutes combustible gases. Common uses of magnesium hydroxide as a flame retardant include additives to cable insulation, insulation plastics, roofing, and various flame retardant coatings.[25][26][27][28][29]
Brucite, the mineral form of Mg(OH)2 commonly found in nature also occurs in the 1:2:1clay minerals amongst others, inchlorite, in which it occupies the interlayer position usually filled by monovalent and divalentcations such as Na+, K+, Mg2+ and Ca2+. As a consequence, chlorite interlayers are cemented by brucite and cannot swell or shrink.
Brucite, in which some of the Mg2+ cations have been substituted by Al3+ cations, becomes positively charged and constitutes the main basis oflayered double hydroxide (LDH). LDH minerals ashydrotalcite are powerful anion sorbents but are relatively rare in nature.
Brucite may also crystallize incement andconcrete in contact withseawater. Indeed, theMg2+ cation is the second-most-abundant cation in seawater, just behindNa+ and beforeCa2+.
Whencement orconcrete are exposed toMg2+ andSO2−4 ions simultaneously present inseawater, theprecipitation of the poorlysoluble brucite contributes to enhance the formation ofgypsum in the sulfate attack :
The precipitation of insolubleMg(OH)2 helps to considerably drive thechemical equilibrium of the reaction to the right. It exacerbates thesulfate attack resulting in the formation ofgypsum andettringite (anexpansive phase) responsible for themechanical stress in the hardened cement paste. However, brucite, a phase with a smallmolar volume (24.63 cm3/mol),[30] may contribute to clogging the porous network in the hardened cement paste, hindering the diffusion of these harmful reactive species in the cement matrix. This can delay the decalcification of theC-S-H phase (the "glue" phase in the hardened cement paste responsible for the cohesion in concrete) and its transformation into an M-S-H phase.
Prolonged contact between seawater or Mg-richbrines and concrete may induce durability issues for regularly immersed concrete components or structures.
The exact mechanism of brucite degradation of hardened cement paste remains a matter of debate.[31] If brucite had a high molar volume, it could bede facto considered a swelling phase (likeettringite, or highly hydrated minerals), but this does not appear to be the case. It is unclear if it causes expansion or not, and how. If it replaces another phase locally (topo chemical replacement), in cases where its molar volume is smaller than that of the phase it replaces, no expansion is expected; rather, a decrease inporosity is anticipated. However, if it crystallizes in a large number of tiny crystals growing between existing ones, even with a small molar volume, it could exert a considerable crystallization pressure in the cement matrix, resulting in tensile stress, expansion and cracking.
For the same reason,dolomite cannot be used asconstruction aggregate for making concrete. The reaction ofmagnesium carbonate with the free alkalihydroxides present in the cement porewater also leads to the formation of brucite, a mineral phase with a low molar volume, but often accompanied by other expansive reaction products (with a higher molar volume than brucite compensating for its shrinkage).
This reaction, one of the two prominentalkali–aggregate reaction (AAR), is also known asalkali–carbonate reaction.
