View source for Magnesium hydroxide

{{Short description|Inorganic compound of formula Mg(OH)2}}
{{Chembox
| Verifiedfields = changed
| Watchedfields  = changed
| verifiedrevid  = 476992492
| ImageFile      = Hydroxid hořečnatý.PNG
| ImageSize      = 250px
| ImageName      = Magnesium hydroxide
| ImageFile1     = Mg(OH)2Xray.jpg
| ImageSize1     = 250px
| ImageName1     = Magnesium hydroxide
| IUPACName      = Magnesium hydroxide
| OtherNames     = {{Unbulleted list|Magnesium dihydroxide|Milk of magnesia}}
| pronounce      = 
| Section1       = {{Chembox Identifiers
| CASNo = 1309-42-8
| CASNo_Ref = {{cascite|correct|CAS}}
| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
| ChemSpiderID = 14107
| ChEBI_Ref = {{ebicite|correct|EBI}}
| ChEBI = 6637
| ChEMBL_Ref = {{ebicite|correct|EBI}}
| ChEMBL = 1200718
| DrugBank = DB09104
| EINECS = 215-170-3
| Gmelin = 485572
| KEGG = C07876
| PubChem = 14791
| RTECS = OM3570000
| UNII_Ref = {{fdacite|correct|FDA}}
| UNII = NBZ3QY004S
| InChI = 1/Mg.2H2O/h;2*1H2/q+2;;/p-2
| InChIKey = VTHJTEIRLNZDEV-NUQVWONBAW
| SMILES = [Mg+2].[OH-].[OH-]
| StdInChI_Ref = {{stdinchicite|correct|chemspider}}
| StdInChI = 1S/Mg.2H2O/h;2*1H2/q+2;;/p-2
| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
| StdInChIKey = VTHJTEIRLNZDEV-UHFFFAOYSA-L
 }}
| Section2       = {{Chembox Properties
| Formula = Mg(OH)<sub>2</sub>
| MolarMass = 58.3197 g/mol
| Appearance = White solid
| Odor = Odorless
| Density = 2.3446{{nbsp}}g/cm<sup>3</sup>
| MeltingPtC = 350
| MeltingPt_notes = decomposes
| Solubility = {{ubl
 | 0.00064{{nbsp}}g/100{{nbsp}}mL (25{{nbsp}}°C)
 | 0.004{{nbsp}}g/100{{nbsp}}mL (100{{nbsp}}°C)
 }}
| SolubilityProduct =  {{val|5.61e-12}}
| RefractIndex = 1.559<ref>{{Cite book |last=Patnaik |first=Pradyot |title=Handbook of inorganic chemicals |date=2003 |publisher=McGraw-Hill |isbn=0-07-049439-8 |location=New York |oclc=50252041}}</ref>
| MagSus = {{val|-22.1e-6|u=cm<sup>3</sup>/mol}}
}}
| Section3       = {{Chembox Structure
| CrystalStruct = Hexagonal, [[Pearson symbol|hP3]]<ref>{{cite journal|doi=10.1143/JPSJ.39.317|title=Magnetic Behaviours of a Random Magnet, Ni<sub>''p''</sub>Mg<sub>(1&minus;''p'')</sub>(OH)<sub>2</sub>|last1=Toshiaki Enoki |last2=Ikuji Tsujikawa |name-list-style=and |journal=J. Phys. Soc. Jpn. |volume=39|year=1975|pages=317–323|issue=2|bibcode=1975JPSJ...39..317E}}</ref>
| SpaceGroup = P{{overline|3}}m1 No. 164
| Coordination =
| LattConst_a = 0.312{{nbsp}}nm
| LattConst_b =
| LattConst_c = 0.473{{nbsp}}nm
| LattConst_alpha =
| LattConst_beta =
| LattConst_gamma =
| MolShape =
| OrbitalHybridisation =
| Dipole = 
 }}
| Section5       = {{Chembox Thermochemistry
| DeltaHf = −924.7&nbsp;kJ·mol<sup>−1</sup><ref name="b1">{{Cite book |last=Zumdahl |first=Steven S. |title=Chemical Principles |publisher=Houghton Mifflin Company |year=2009 |isbn=978-0-618-94690-7 |edition=6th |page=A22}}</ref>
| DeltaGf = −833.7{{nbsp}}kJ/mol
| Entropy = 64&nbsp;J·mol<sup>−1</sup>·K<sup>−1</sup><ref name=b1/>
| HeatCapacity = 77.03{{nbsp}}J/mol·K
 }}
| Section6       = {{Chembox Pharmacology
| ATCCode_prefix = A02
| ATCCode_suffix = AA04
| ATC_Supplemental = {{ATC|G04|BX01}}
}}
| Section7       = {{Chembox Hazards
| ExternalSDS = [https://fscimage.fishersci.com/msds/13405.htm External MSDS]
| GHSPictograms = {{GHS07}}<ref name=sds>{{cite web |title=Magnesium Hydroxide |url=https://www.americanelements.com/magnesium-hydroxide-1309-42-8 |publisher=[[American Elements]] |access-date=May 9, 2019}}</ref>
| GHSSignalWord = Warning<ref name=sds/>
| HPhrases = {{H-phrases|315|319|335}}<ref name=sds/>
| PPhrases = {{P-phrases|261|280|305+351+338|304+340|405|501}}<ref name=sds/>
| MainHazards =
| NFPA-H = 1
| NFPA-F = 0
| NFPA-R = 0
| NFPA-S =
| FlashPt = Non-flammable
| AutoignitionPt =
| LD50 = 8500{{nbsp}}mg/kg (rat, oral)
 }}
| Section8       = {{Chembox Related
| OtherAnions = [[Magnesium oxide]]
| OtherCations = {{ubl
 | [[Beryllium hydroxide]]
 | [[Calcium hydroxide]]
 | [[Strontium hydroxide]]
 | [[Barium hydroxide]]
  }}
 }}
}}

'''Magnesium hydroxide''' is an [[inorganic compound]] with the chemical formula Mg(OH)<sub>2</sub>. It occurs in nature as the mineral [[brucite]]. It is a white solid with low solubility in water ({{nowrap|[[solubility product|''K''<sub>sp</sub>]] {{=}} {{val|5.61e-12}}}}).<ref>{{cite book|title=Handbook of Chemistry and Physics|date=12 March 1996|publisher=CRC Press|isbn=0-8493-0596-9|edition=76th}}</ref> Magnesium hydroxide is a common component of [[antacid]]s, such as '''milk of magnesia'''.

==Preparation==
Treating the solution of different soluble magnesium salts with [[alkaline]] water induces the [[Precipitation (chemistry)|precipitation]] of the solid hydroxide Mg(OH)<sub>2</sub>:
{{block indent|Mg<sup>2+</sup> + 2&nbsp;OH<sup>−</sup> → Mg(OH)<sub>2</sub>}}

As {{chem|Mg|2+}} is the second most abundant cation present in [[seawater]] after {{chem|Na|+}}, it can be economically extracted directly from seawater by [[alkalinisation]] as described above. On an industrial scale, Mg(OH)<sub>2</sub> is produced by treating seawater with [[calcium hydroxide|lime]] (Ca(OH)<sub>2</sub>). A volume of {{cvt|600|m3|USgal}} of seawater gives about {{convert|1|t|lb}} of Mg(OH)<sub>2</sub>. Ca(OH)<sub>2</sub> {{nowrap|([[Solubility product|''K''<sub>sp</sub>]] {{=}} {{val|5.02e-6}}}})<ref name="CRC">{{cite book |first=John |last=Rumble |title=CRC Handbook of Chemistry and Physics |date=June 18, 2018 |publisher=CRC Press |isbn=978-1-138-56163-2 |pages=5-188<!-- hyphen not range -->|edition=99th |language=English}}</ref> is far more soluble than Mg(OH)<sub>2</sub> {{nowrap|([[solubility product|''K''<sub>sp</sub>]] {{=}} {{val|5.61e-12}}}}) and dramatically increases the [[pH]] value of seawater from 8.2 to 12.5. The less soluble {{chem|Mg|(OH)|2}} precipitates because of the [[common ion effect]] due to the {{chem|OH|-}} added by the dissolution of {{chem|Ca|(OH)|2}}:<ref name="ullmann">{{Ullmann|title=Magnesium Compounds|doi=10.1002/14356007.a15_595.pub2|last1=Seeger|first1=Margarete|last2=Otto|first2=Walter|last3=Flick|first3=Wilhelm|last4=Bickelhaupt|first4=Friedrich|last5=Akkerman|first5=Otto S.}}</ref>
{{block indent|{{chem2|Mg(2+) + Ca(OH)2 -> Mg(OH)2 + Ca(2+)}}}}

For seawater brines, precipitating agents other than {{chem2|Ca(OH)2}} can be utilized, each with their own nuances:
* Use of {{chem2|Ca(OH)2}} can yield [[CaSO4|CaSO<sub>4</sub>]] or [[CaCO3|CaCO<sub>3</sub>]], which reduces the final purity of {{chem2|Mg(OH)2}}.
* [[NH4OH|NH<sub>4</sub>OH]] can produce explosive [[nitrogen trichloride]] when the brine is used for chlorine production.
* [[NaOH]] as the precipitating agent has longer settling times and is difficult to filter.

It has been demonstrated that sodium hydroxide, {{chem2|NaOH}}, is the better precipitating agent compared to [[Ca(OH)2|Ca(OH)<sub>2</sub>]] and {{chem2|NH4OH}} 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.<ref>{{cite journal |last1=Fontana |first1=Danilo |last2=Forte |first2=Federica |last3=Pietrantonio |first3=Massimiliana |last4=Pucciarmati |first4=Stefano |last5=Marcoaldi |first5=Caterina |title=Magnesium recovery from seawater desalination brines: a technical review |journal=Environment, Development and Sustainability |date=2023-12-01 |volume=25 |issue=12 |pages=13733–13754 |doi=10.1007/s10668-022-02663-2 |ref=fontana2023|doi-access=free |bibcode=2023EDSus..2513733F }}</ref>

It is also possible to obtain {{chem2|Mg(OH)2}} from seawater using electrolysis chambers separated with a [[cation exchange membrane]]. This process is continuous, lower-cost, and produces oxygen gas, hydrogen gas, sulfuric acid (if {{chem2|Na2SO4}} is used; {{chem2|NaCl}} can alternatively be used to yield {{chem2|HCl}}), and {{chem2|Mg(OH)2}} of 98% or higher purity. It is crucial to [[Degasification|deaerate]] the seawater to mitigate co-precipitation of calcium precipitates.<ref name="sano2018">{{cite journal |last1=Sano |first1=Yoshihiko |last2=Hao |first2=YiJia |last3=Kuwahara |first3=Fujio |title=Development of an electrolysis based system to continuously recover magnesium from seawater |journal=Heliyon |date=2018-11-01 |volume=4 |issue=11 |article-number=e00923 |doi=10.1016/j.heliyon.2018.e00923 |doi-access=free |pmid=30839823 |pmc=6249789 |bibcode=2018Heliy...400923S }}</ref>

==Uses==
===Precursor to MgO===
Most Mg(OH)<sub>2</sub> that is produced industrially, as well as the small amount that is mined, is converted to fused [[Magnesium oxide|magnesia]] (MgO). Magnesia is valuable because it is both a poor electrical conductor and an excellent thermal conductor.<ref name=ullmann/>

===Medical===
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 to [[hypermagnesemia]]. Unabsorbed magnesium is excreted in feces; absorbed magnesium is rapidly excreted in urine.<ref>{{Cite web |title=magnesium hydroxide |url=https://www.glowm.com/resources/glowm/cd/pages/drugs/m001.html |archive-url=https://web.archive.org/web/20180114020205/https://www.glowm.com/resources/glowm/cd/pages/drugs/m001.html |archive-date=14 January 2018 |access-date=2023-03-14 |publisher=[[Global Library of Women's Medicine]] |language=en}}</ref>
[[File:Santo Domingo - Museo de Ámbar 0669.JPG|thumb|left|Bottle used for ''Phillips' Leche de Magnesia'' (Milk of Magnesia) in the Amber Museum, [[Santo Domingo]], [[Dominican Republic]]]]

====Applications====

=====Antacid=====
As an antacid, magnesium hydroxide is dosed at approximately 0.5–1.5{{nbsp}}g in adults and works by simple [[Neutralization (chemistry)|neutralization]], in which the [[hydroxide]] [[ion]]s from the Mg(OH)<sub>2</sub> combine with [[acid]]ic H<sup>+</sup> [[ion]]s (or [[hydronium]] ions) produced in the form of hydrochloric acid by [[parietal cell]]s in the [[stomach]], to produce water.

=====Laxative=====
As a laxative, magnesium hydroxide is dosed at {{convert|5|–|10|g}}, and works in a number of ways. First, Mg<sup>2+</sup> is poorly absorbed from the intestinal tract, so it draws water from the surrounding tissue by [[osmosis]]. 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 intestinal [[motility]]. Furthermore, Mg<sup>2+</sup> ions cause the release of [[cholecystokinin]] (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 Mg<sup>2+</sup> solutions, like [[Magnesium sulfate|MgSO<sub>4</sub>]], are equally strong laxatives, [[Mole (unit)|mole]] for mole.<ref>{{Cite journal |last1=Tedesco |first1=Frances J. |last2=DiPiro |first2=Joseph T. |year=1985 |title=Laxative use in constipation |journal=[[The American Journal of Gastroenterology]] |volume=80 |issue=4 |pages=303–309 |pmid=2984923}}</ref>

====History of milk of magnesia====
[[File:Earle's Mint-O-Mag.jpeg|thumb|Vintage bottle for milk of magnesia]]
On May 4, 1818, American inventor Koen Burrows received a patent (No. X2952) for magnesium hydroxide.<ref>{{US patent|X2952|Patent USX2952 - Magnesia, medicated, liquid - Google Patents}}</ref> In 1829, [[James Murray (physician)|Sir James Murray]] used a "condensed solution of fluid magnesia" preparation of his own design<ref>[[Michael Hordern]]. ''A World Elsewhere'' (1993), p. 2.</ref> to treat the [[Lord Lieutenant of Ireland]], the [[Henry Paget, 1st Marquess of Anglesey|Marquess of Anglesey]], for stomach pain. This was so successful (advertised in Australia and approved by the Royal College of Surgeons in 1838)<ref>{{Cite news |url=https://news.google.com/newspapers?nid=1301&dat=18461007&id=ERkRAAAAIBAJ&pg=6490,1153513 |title=Sir James Murray's condensed solution of fluid magnesia |work=[[The Sydney Morning Herald]] |date=October 7, 1846 |volume=21 |issue=2928 |page=1, column 4}}</ref> 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.<ref>[http://www.ulsterhistory.co.uk/jamesmurray.htm Ulster History. Sir James Murray – Inventor of Milk of Magnesia. 1788 to 1871] {{webarchive|url=https://web.archive.org/web/20110605061937/http://www.ulsterhistory.co.uk/jamesmurray.htm |date=2011-06-05 }}, 24 February 2005</ref>

The term ''milk of magnesia'' was first used by [[Charles Henry Phillips]] in 1872 for a [[Suspension (chemistry)|suspension]] of magnesium hydroxide formulated at about 8[[Mass concentration (chemistry)|% w/v]].<ref>[http://www.phillipsrelief.com/phillips-frequently-asked-questions/#quest19 When was Phillips' Milk of Magnesia introduced?] {{Webarchive|url=https://web.archive.org/web/20170622220811/https://www.phillipsdigestive.com/frequently-asked-question/#quest19 |date=2017-06-22 }} FAQ, phillipsrelief.com, accessed 4 July 2016</ref> It was sold under the brand name ''Phillips' Milk of Magnesia'' for medicinal usage.

[[USPTO]] registrations show that the terms "Milk of Magnesia"<ref>[http://tarr.uspto.gov/servlet/tarr?regser=serial&entry=70024049 results from the TARR web server: Milk of Magnesia]</ref> and "Phillips' Milk of Magnesia"<ref>[http://tarr.uspto.gov/servlet/tarr?regser=serial&entry=71016576 results from the TARR web server: Phillips' Milk of Magnesia]</ref> have both been assigned to [[Bayer]] 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, [[British Pharmacopoeia|BP]]).

===As food additive===

It is added directly to human food, and is affirmed as [[generally recognized as safe]] by the [[FDA]].<ref>{{cite web |title=Compound Summary for CID 14791 - Magnesium Hydroxide |url=https://pubchem.ncbi.nlm.nih.gov/compound/14791 |publisher=PubChem}}</ref> It is known as [[E number]] '''E528'''.

Magnesium hydroxide is marketed for medical use in the form of chewable tablets, capsules, powder, and as liquid [[suspension (chemistry)|suspensions]], which are sometimes flavored. These products are sold as [[antacid]]s to neutralize stomach [[acid]] and relieve [[indigestion]] and [[heartburn]]. 

It is also a laxative used to alleviate [[constipation]]. As a laxative, the [[osmotic]] force of the magnesia acts to draw fluids from the body. High doses can lead to [[diarrhea]] and can deplete the body's supply of [[potassium]], sometimes leading to [[muscle cramps]].<ref>[http://www.revolutionhealth.com/articles/magnesium-hydroxide/hn-drug_magnesium_hydroxide Magnesium Hydroxide – Revolution Health]</ref> Some magnesium hydroxide products sold for antacid use (such as [[Maalox]]) are formulated to minimize unwanted laxative effects through the inclusion of [[aluminum hydroxide]], which inhibits the contractions of [[smooth muscle]] cells in the gastrointestinal tract,<ref name="Washington1991">{{cite book|last1=Washington|first1=Neena|title=Antacids and Anti Reflux Agents|date=2 August 1991|publisher=CRC Press|location=Boca Raton, FL|isbn=0-8493-5444-7|page=10}}</ref> thereby counterbalancing the contractions induced by the osmotic effects of the magnesium hydroxide.

===Other niche uses===

Magnesium hydroxide is also a component of [[antiperspirant]].<ref>[needs citation]</ref>

====Waste water treatment====
Magnesium hydroxide powder is used industrially to neutralize acidic wastewaters.<ref>Aileen Gibson and Michael Maniocha. [http://www.wateronline.com/article.mvc/The-Use-Of-Magnesium-Hydroxide-Slurry-For-Bio-0001?VNETCOOKIE=NO White Paper: The Use Of Magnesium Hydroxide Slurry For Biological Treatment Of Municipal and Industrial Wastewater], August 12, 2004</ref> It is also a component of the [[Biorock]] method of building [[artificial reef]]s. The main advantage of {{chem|Mg|(OH)|2}} over {{chem|Ca|(OH)|2}}, is to impose a lower pH better compatible with that of seawater and sea life: pH 10.5 for {{chem|Mg|(OH)|2}} in place of pH 12.5 with {{chem|Ca|(OH)|2}}.

====Fire retardant====
Natural magnesium hydroxide ([[brucite]]) is used commercially as a fire retardant. Most industrially used magnesium hydroxide is produced synthetically.<ref>{{cite book|last=Rothon|first=RN|title=Particulate Filled Polymer Composites|year=2003|publisher=Rapra Technology|location=Shrewsbury, UK|pages=53–100}}</ref> Like aluminum hydroxide, solid magnesium hydroxide has smoke suppressing and [[flame retardant]] properties. This property is attributable to the [[Thermal decomposition|endothermic decomposition]] it undergoes at {{cvt|332|C}}:

{{block indent|Mg(OH)<sub>2</sub> → MgO + H<sub>2</sub>O}}

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.<ref name="Rev1">{{cite journal |last1=Hollingbery |first1=LA |last2=Hull |first2=TR |year=2010 |title=The Thermal Decomposition of Huntite and Hydromagnesite - A Review |url=http://clok.uclan.ac.uk/1139 |journal=Thermochimica Acta |volume=509 |issue=1–2 |pages=1–11 |doi=10.1016/j.tca.2010.06.012}}</ref><ref name="Rev2">{{cite journal |last1=Hollingbery |first1=LA |last2=Hull |first2=TR |year=2010 |title=The Fire Retardant Behaviour of Huntite and Hydromagnesite - A Review |url=http://clok.uclan.ac.uk/1432 |journal=Polymer Degradation and Stability |volume=95 |issue=12 |pages=2213–2225 |doi=10.1016/j.polymdegradstab.2010.08.019}}</ref><ref name="Fire1">{{cite journal |last1=Hollingbery |first1=LA |last2=Hull |first2=TR |year=2012 |title=The Fire Retardant Effects of Huntite in Natural Mixtures with Hydromagnesite |url=http://clok.uclan.ac.uk/3420/ |journal=Polymer Degradation and Stability |volume=97 |issue=4 |pages=504–512 |doi=10.1016/j.polymdegradstab.2012.01.024}}</ref><ref name="Therm1">{{cite journal |last1=Hollingbery |first1=LA |last2=Hull |first2=TR |year=2012 |title=The Thermal Decomposition of Natural Mixtures of Huntite and Hydromagnesite |url=http://clok.uclan.ac.uk/3414 |journal=Thermochimica Acta |volume=528 |pages=45–52 |doi=10.1016/j.tca.2011.11.002|bibcode=2012TcAc..528...45H }}</ref><ref name="Fire2">{{cite journal |last1=Hull |first1=TR |last2=Witkowski |first2=A |last3=Hollingbery |first3=LA |year=2011 |title=Fire Retardant Action of Mineral Fillers |url=http://clok.uclan.ac.uk/2963 |journal=Polymer Degradation and Stability |volume=96 |issue=8 |pages=1462–1469 |doi=10.1016/j.polymdegradstab.2011.05.006 |s2cid=96208830}}</ref>

==Mineralogy==
[[File:Brucite-169935.jpg|thumb|[[Brucite]] crystals (mineral form of Mg(OH)<sub>2</sub>) from the Sverdlovsk Region, Urals, Russia (size: {{cvt|10.5|×|7.8|×|7.4|cm|disp=or}})]]
[[Brucite]], the mineral form of Mg(OH)<sub>2</sub> commonly found in nature also occurs in the 1:2:1 [[clay mineral]]s amongst others, in [[Chlorite group|chlorite]], in which it occupies the interlayer position usually filled by monovalent and divalent [[cation]]s such as Na<sup>+</sup>, K<sup>+</sup>, Mg<sup>2+</sup> and Ca<sup>2+</sup>. As a consequence, chlorite interlayers are cemented by brucite and cannot swell or shrink.

Brucite, in which some of the Mg<sup>2+</sup> cations have been substituted by Al<sup>3+</sup> cations, becomes positively charged and constitutes the main basis of [[layered double hydroxide]] (LDH). LDH minerals as [[hydrotalcite]] are powerful anion sorbents but are relatively rare in nature.

== Concrete degradation ==
Brucite may also crystallize in [[cement]] and [[concrete]] in contact with [[seawater]]. Indeed, the {{Chem2|Mg(2+)}} cation is the second-most-abundant cation in seawater, just behind {{Chem2|Na+}} and before {{Chem2|Ca(2+)}}. 

When [[cement]] or [[concrete]] are exposed to {{Chem2|Mg(2+)}} and {{Chem2|SO4(2-)}} ions simultaneously present in [[seawater]], the [[Precipitation (chemistry)|precipitation]] of the poorly [[Solubility|soluble]] brucite contributes to enhance the formation of [[gypsum]] in the sulfate attack : 

{{block indent|{{Chem2|MgSO4 + Ca(OH)2 + 2 H2O -> Mg(OH)2 + CaSO4 . 2H2O}}}}

The precipitation of insoluble {{Chem2|Mg(OH)2}} helps to considerably drive the [[chemical equilibrium]] of the reaction to the right. It exacerbates the [[Sulfate attack in concrete and mortar|sulfate attack]] resulting in the formation of [[gypsum]] and [[ettringite]] (an [[Tension (physics)|expansive]] phase) responsible for the [[Stress (mechanics)|mechanical stress]] in the hardened cement paste. However, brucite, a phase with a small [[molar volume]] ({{Nowrap|24.63 cm<sup>3</sup>/mol}}),<ref name="Molar_Volume">{{cite web | title=Data consultation – Minerals – Molar volume | website=BRGM, Thermoddem | url=https://thermoddem.brgm.fr/data/minerals?title=brucite&field_mformulechimique_value= | access-date=2025-07-22}}</ref> 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 the [[Calcium silicate hydrate|C-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-rich [[brine]]s 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.<ref name="Lee2002">{{Cite journal| last=Lee| first=Hyomin|author2=Robert D. Cody|author3=Anita M. Cody|author4=Paul G. Spry| year=2002| title=Observations on brucite formation and the role of brucite in Iowa highway concrete deterioration| journal=Environmental and Engineering Geoscience| volume=8| issue=2| pages=137–145| access-date=2009-09-10| doi=10.2113/gseegeosci.8.2.137| bibcode=2002EEGeo...8..137L| url=http://eeg.geoscienceworld.org/cgi/content/abstract/8/2/137| url-access=subscription}}</ref> If brucite had a high molar volume, it could be ''de facto'' considered a swelling phase (like [[ettringite]], 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 in [[porosity]] 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 (rock)|dolomite]] cannot be used as [[construction aggregate]] for making concrete. The reaction of [[magnesium carbonate]] with the free alkali [[hydroxide]]s 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).

{{block indent|MgCO<sub>3</sub> + 2 NaOH → Mg(OH)<sub>2</sub> + Na<sub>2</sub>CO<sub>3</sub>}}

This reaction, one of the two prominent [[alkali–aggregate reaction]] (AAR), is also known as [[alkali–carbonate reaction]].

==See also==
* [[Portlandite]] – calcium hydroxide: {{chem|Ca|(OH)|2}}

==References==
{{reflist|30em}}

{{Magnesium compounds}}
{{Hydroxides}}
{{Urologicals}}
{{Antacids}}
{{Laxatives}}

{{DEFAULTSORT:Magnesium Hydroxide}}
[[Category:Antacids]]
[[Category:Laxatives]]
[[Category:Food acidity regulators]]
[[Category:Hydroxides]]
[[Category:Magnesium compounds]]
[[Category:E-number additives]]
[[Category:Over-the-counter drugs in the United States]]