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Magnesium carbonate
Magnesium carbonate
Magnesium carbonate
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Magnesium carbonate
Names
Other names
Magnesite
Barringtonite (dihydrate)
Nesequehonite (trihydrate)
Lansfordite (pentahydrate)
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.008.106 Edit this at Wikidata
E number E504(i) (acidity regulators, ...)
RTECS number
  • OM2470000
UNII
  • InChI=1S/CH2O3.Mg/c2-1(3)4;/h(H2,2,3,4);/q;+2/p-2 checkY
    Key: ZLNQQNXFFQJAID-UHFFFAOYSA-L checkY
  • InChI=1/CH2O3.Mg/c2-1(3)4;/h(H2,2,3,4);/q;+2/p-2
    Key: ZLNQQNXFFQJAID-NUQVWONBAU
  • [Mg+2].[O-]C([O-])=O
Properties
MgCO3
Molar mass 84.3139 g/mol (anhydrous)
Appearance Colourless crystals or white solid
Hygroscopic
Odor Odorless
Density 2.958 g/cm3 (anhydrous)
2.825 g/cm3 (dihydrate)
1.837 g/cm3 (trihydrate)
1.73 g/cm3 (pentahydrate)
Melting point 350 °C (662 °F; 623 K)
decomposes (anhydrous)
165 °C (329 °F; 438 K)
(trihydrate)
Anhydrous:
0.0139 g/100 ml (25 °C)
0.0063 g/100 ml (100 °C)[1]
10−7.8[2]
Solubility Soluble in acid, aqueous CO2
Insoluble in acetone, ammonia
−32.4·10−6 cm3/mol
1.717 (anhydrous)
1.458 (dihydrate)
1.412 (trihydrate)
Structure
Trigonal
R3c, No. 167[3]
Thermochemistry
75.6 J/mol·K[1]
65.7 J/mol·K[1][4]
−1113 kJ/mol[4]
−1029.3 kJ/mol[1]
Pharmacology
A02AA01 (WHO) A06AD01 (WHO)
Hazards
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 1: Exposure would cause irritation but only minor residual injury. E.g. turpentineFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
1
0
0
Flash point Non-flammable
NIOSH (US health exposure limits):
PEL (Permissible)
  • TWA 15 mg/m3 (total)
  • TWA 5 mg/m3 (resp)[5]
Safety data sheet (SDS) ICSC 0969
Related compounds
Other anions
 Magnesium bicarbonate
Other cations
 Beryllium carbonate
Calcium carbonate
Strontium carbonate
Barium carbonate
Radium carbonate
Related compounds
 Artinite
Hydromagnesite
Dypingite
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 carbonate, MgCO3 (archaic name magnesia alba), is an inorganic salt that is a colourless or white solid. Several hydrated and basic forms of magnesium carbonate also exist as minerals.

Forms

[edit]

The most common magnesium carbonate forms are the anhydrous salt called magnesite (MgCO3), and the di, tri, and pentahydrates known as barringtonite (MgCO3·2H2O), nesquehonite (MgCO3·3H2O), and lansfordite (MgCO3·5H2O), respectively.[6] Some basic forms such as artinite (Mg2CO3(OH)2·3H2O), hydromagnesite (Mg5(CO3)4(OH)2·4H2O), and dypingite (Mg5(CO3)4(OH)2·5H2O) also occur as minerals. All of those minerals are colourless or white.

Magnesite consists of colourless or white trigonal crystals. The anhydrous salt is practically insoluble in water, acetone, and ammonia. All forms of magnesium carbonate react with acids. Magnesite crystallizes in the calcite structure wherein Mg2+ is surrounded by six oxygen atoms.[3]

Crystal structure of magnesium carbonate
Carbonate coordination Magnesium coordination Unit cell

The dihydrate has a triclinic structure, while the trihydrate has a monoclinic structure.

References to "light" and "heavy" magnesium carbonates actually refer to the magnesium hydroxy carbonates hydromagnesite and dypingite, respectively.[7] The "light" form is precipitated from magnesium solutions using alkali carbonate at "normal temperatures" while the "heavy" may be produced from boiling concentrated solutions followed by precipitation to dryness, washing of the precipitate, and drying at 100 C. [8]

Preparation

[edit]

Magnesium carbonate is ordinarily obtained by mining the mineral magnesite. Seventy percent of the world's supply is mined and prepared in China.[9]

Magnesium carbonate can be prepared in laboratory by reaction between any soluble magnesium salt and sodium bicarbonate:

MgCl2(aq) + 2 NaHCO3(aq) → MgCO3(s) + 2 NaCl(aq) + H2O(l) + CO2(g)

If magnesium chloride (or sulfate) is treated with aqueous sodium carbonate, a precipitate of basic magnesium carbonate – a hydrated complex of magnesium carbonate and magnesium hydroxide – rather than magnesium carbonate itself is formed:

5 MgCl2(aq) + 5 Na2CO3(aq) + 5 H2O(l) → Mg4(CO3)3(OH)2·3H2O(s) + Mg(HCO3)2(aq) + 10 NaCl(aq)

High purity industrial routes include a path through magnesium bicarbonate, which can be formed by combining a slurry of magnesium hydroxide and carbon dioxide at high pressure and moderate temperature.[6] The bicarbonate is then vacuum dried, causing it to lose carbon dioxide and a molecule of water:

Mg(OH)2 + 2 CO2 → Mg(HCO3)2
Mg(HCO3)2 → MgCO3 + CO2 + H2O

Chemical properties

[edit]

With acids

[edit]

Like many common group 2 metal carbonates, magnesium carbonate reacts with aqueous acids to release carbon dioxide and water:

MgCO3 + 2 HCl → MgCl2 + CO2 + H2O
MgCO3 + H2SO4 → MgSO4 + CO2 + H2O

Decomposition

[edit]

At high temperatures MgCO3 decomposes to magnesium oxide and carbon dioxide. This process is important in the production of magnesium oxide.[6] This process is called calcining:

MgCO3 → MgO + CO2 (ΔH = +118 kJ/mol)

The decomposition temperature is given as 350 °C (662 °F).[10][11] However, calcination to the oxide is generally not considered complete below 900 °C due to interfering readsorption of liberated carbon dioxide.

The hydrates of the salts lose water at different temperatures during decomposition.[12] For example, in the trihydrate MgCO3·3H2O, which molecular formula may be written as Mg(HCO3)(OH)·2H2O, the dehydration steps occur at 157 °C and 179 °C as follows:[12]

Mg(HCO3)(OH)·2(H2O) → Mg(HCO3)(OH)·(H2O) + H2O at 157 °C
Mg(HCO3)(OH)·(H2O) → Mg(HCO3)(OH) + H2O at 179 °C

Uses

[edit]

The primary use of magnesium carbonate is the production of magnesium oxide by calcining. Magnesite and dolomite minerals are used to produce refractory bricks.[6] MgCO3 is also used in flooring, fireproofing, fire extinguishing compositions, cosmetics, dusting powder, and toothpaste. Other applications are as filler material, smoke suppressant in plastics, a reinforcing agent in neoprene rubber, a drying agent, and colour retention in foods.

Because of its low solubility in water and hygroscopic properties, MgCO3 was first added to table salt (NaCl) in 1911 to make it flow more freely. The Morton Salt company adopted the slogan "When it rains it pours", highlighting that its salt, which contained MgCO3, would not stick together in humid weather.[13]

Climber Jan Hojer blows surplus chalk from his hand. Boulder World Cup 2015

Powdered magnesium carbonate, known as climbing chalk or gym chalk is also used as a drying agent on athletes' hands in rock climbing, gymnastics, powerlifting, weightlifting and other sports in which a firm grip is necessary.[9] A variant is liquid chalk and another is mesoporous magnesium carbonate.

As a food additive, magnesium carbonate is known as E504. Its only known side effect is that it may work as a laxative in high concentrations.[14]

Magnesium carbonate is used in taxidermy for whitening skulls. It can be mixed with hydrogen peroxide to create a paste, which is spread on the skull to give it a white finish.

Magnesium carbonate is used as a matte white coating for projection screens.[15]

Medical use

[edit]

It is a laxative to loosen the bowels.

In addition, high purity magnesium carbonate is used as an antacid and as an additive in table salt to keep it free flowing. Magnesium carbonate can do this because it does not dissolve in water, only in acid, where it will effervesce (bubble).[16]

Compendial status

[edit]

See also

[edit]

Notes and references

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

Magnesium carbonate

View on Grokipedia
from Grokipedia
Magnesium carbonate is an with the MgCO₃, consisting of magnesium ions (Mg²⁺) and ions (CO₃²⁻), and it occurs naturally as the mineral , a white, odorless, crystalline powder that is practically insoluble in but soluble in dilute acids. Its molecular weight is 84.31 g/mol, and it has a density ranging from 2.96 to 3.1 g/cm³, decomposing upon heating above approximately 350 °C to form and . This compound exists in both and hydrated forms, with the trihydrate (MgCO₃·3H₂O) and pentahydrate being common, and it is primarily obtained by mining the or through synthetic production methods. Industrially, magnesium carbonate is valued for its thermal stability and is used in the production of refractory materials that withstand high temperatures, as well as in the manufacture of for various applications. It also serves as a source of when reacted with acids, contributing to processes like fire extinguishers and effervescent formulations. In medical and pharmaceutical contexts, magnesium carbonate functions as an to neutralize stomach acid and relieve , often combined with in over-the-counter remedies. As a (E504), it acts as an anti-caking agent in products like table salt and , and as a aid in powdered foods to enhance flowability. Additionally, it finds applications in for its mattifying effects on without irritation, in as a recycled coagulant for removing color and impurities from industrial wastes, and in paints, inks, and fertilizers to provide magnesium nutrients. Despite its low toxicity, it can irritate the eyes, , and upon prolonged exposure.

Structure and Forms

Anhydrous Magnesium Carbonate

Anhydrous magnesium carbonate, with the MgCO₃, has a of 84.3139 g/mol. It exhibits a rhombohedral , belonging to the trigonal system with R3c (No. 167), and parameters of a ≈ 4.63 and c ≈ 15.01 . In nature, anhydrous occurs primarily as the mineral , which forms through the carbonation of ultramafic rocks such as or serpentinized during metamorphic processes in belts of regionally metamorphosed terrains. Commercial forms of anhydrous are typically produced with high purity levels ranging from 98% to 100%, often available in grades up to 99.999% for specialized applications.

Hydrated Forms

Hydrated forms of magnesium carbonate incorporate water molecules into their crystal lattices, distinguishing them from the by enabling formation in aqueous environments and altering their structural and stability profiles. Common hydrated minerals include barringtonite (MgCO₃·2H₂O), nesquehonite (MgCO₃·3H₂O), lansfordite (MgCO₃·5H₂O), and artinite (Mg₂CO₃(OH)₂·3H₂O). These minerals exhibit diverse crystal structures that integrate hydration bonds. For instance, nesquehonite crystallizes in the monoclinic (space group P2₁/m), featuring infinite chains of edge-sharing MgO₆ octahedra along the b-axis, with groups bridging the chains and molecules facilitating between the octahedra and CO₃ units. Barringtonite adopts a triclinic structure, while lansfordite and artinite are both monoclinic, with lansfordite in space group P2₁/a and artinite in C2/m; in artinite, the structure involves ribbons of MgO₆ octahedra linked by OH⁻ and CO₃²⁻ groups, coordinated by molecules. These hydrated lattices contrast with the rhombohedral form by incorporating H₂O as ligands or interstitial molecules, which influences interlayer spacing and flexibility. In natural settings, these hydrates typically occur as secondary minerals in deposits, low-temperature hydrothermal veins, or as alteration products of primary magnesium-bearing rocks like or . Barringtonite forms encrustations on weathered in areas such as Barrington Tops, , often associated with nesquehonite. Nesquehonite appears in mine efflorescences, deposits, and fractures within serpentinites, such as in Lavrion, , or Pennsylvania, . Lansfordite is rare and found in mines under cool, humid conditions, exemplified by occurrences in Nesquehoning, . Artinite develops as veinlets or crusts in serpentinized ultramafic rocks, including sites in , , and . The stability of these hydrated forms is limited compared to anhydrous magnesite, as they preferentially precipitate under ambient aqueous conditions but dehydrate upon heating or drying, reverting toward more stable phases. Nesquehonite, for example, remains stable at low temperatures and pressures but undergoes phase transitions above 2.4 GPa or decomposes above 50–100°C in air. Lansfordite, as a low-temperature polymorph, is even less stable at room conditions, slowly losing over months. Overall, hydration enhances and reactivity in water-saturated environments, facilitating their role as transient phases in the magnesium carbonate system.

Physical Properties

Appearance and Density

Magnesium carbonate is typically observed as a white, odorless powder or crystalline solid. The form, , presents as colorless to white crystals, often with a vitreous luster, though massive varieties may appear dull, earthy, or chalky. Hydrated forms, such as nesquehonite, can display fibrous or needle-like morphologies, contributing to their distinct visual characteristics compared to the anhydrous variant. The density of anhydrous magnesium carbonate is approximately 3.0 g/cm³, with measured values ranging from 2.98 to 3.02 g/cm³ for . Hydrated forms exhibit lower densities; for example, the trihydrate nesquehonite has a density of 1.82 to 1.85 g/cm³. These variations arise from the incorporation of water molecules in the . Magnesite crystals register a hardness of 3.5 to 4.5 on the , indicating moderate scratch resistance suitable for industrial handling. In commercial applications, magnesium carbonate powders are commonly produced with particle sizes ranging from 10 to 50 μm, facilitating uniform dispersion in formulations.

Solubility and Stability

Magnesium carbonate exhibits low in , approximately 0.0139 g per 100 mL at 25 °C for the form, rendering it slightly soluble overall. It is practically insoluble in and acetone, which limits its dissolution in organic solvents commonly used in chemical . The solubility product constant (Ksp) for its dissociation into Mg²⁺ and CO₃²⁻ ions is approximately 3.5 × 10⁻⁸ at 25 °C, reflecting its sparingly soluble nature in aqueous environments. Upon limited dissolution, magnesium carbonate forms basic solutions due to the partial of the , which reacts with to produce and ions, typically resulting in a of 9–10 for suspensions. This dependence influences its behavior in aqueous systems, where the equilibrium shifts toward hydrolysis products under neutral or acidic conditions. Under standard ambient conditions, magnesium carbonate remains stable, showing no significant or phase changes. However, thermal stability is limited, with to and occurring above approximately 350 °C, as evidenced by thermogravimetric analyses indicating onset around 396 °C under controlled atmospheres. The form displays particular sensitivity to , adsorbing at relative humidities above 70% and converting to hydrated phases such as the trihydrate, which affects its handling and storage in moist environments.

Chemical Properties

Reactions with Acids

Magnesium undergoes an acid-base reaction with acids, acting as a base to neutralize the acid while evolving gas, which results in . The general reaction can be represented as: \ceMgCO3+2H+−>Mg2++H2O+CO2\ce{MgCO3 + 2H+ -> Mg^2+ + H2O + CO2} This process converts the insoluble carbonate into a soluble magnesium salt, water, and gaseous CO₂, making it a key demonstration of carbonate reactivity. A representative example is the reaction with hydrochloric acid, a strong acid commonly used in laboratory settings: \ceMgCO3(s)+2HCl(aq)−>MgCl2(aq)+H2O(l)+CO2(g)\ce{MgCO3(s) + 2HCl(aq) -> MgCl2(aq) + H2O(l) + CO2(g)} This reaction proceeds vigorously, producing as the soluble product. Similarly, with , magnesium carbonate yields : \ceMgCO3(s)+H2SO4(aq)−>MgSO4(aq)+H2O(l)+CO2(g)\ce{MgCO3(s) + H2SO4(aq) -> MgSO4(aq) + H2O(l) + CO2(g)} This is employed in the preparation of magnesium salts through neutralization. The rate of these reactions is influenced by the strength and concentration of the acid; strong acids like HCl react more rapidly than weaker ones, with the reaction speed increasing as acid concentration rises due to higher availability of H⁺ ions for collision with carbonate particles. also accelerates the process by enhancing molecular . In qualitative chemical analysis, the effervescent evolution of CO₂ upon addition of dilute acid to a sample is a standard confirmatory test for the presence of ions, including in magnesium carbonate; the gas can be verified by passing it through limewater, which turns milky due to formation. This test distinguishes carbonates from other anions like sulfates or phosphates that do not produce gas under similar conditions.

Thermal Decomposition

Magnesium carbonate decomposes thermally via an endothermic reaction, yielding and gas: \ceMgCO3(s)−>MgO(s)+CO2(g)\ce{MgCO3(s) -> MgO(s) + CO2(g)} with a standard change of ΔH=+101\Delta H = +101 kJ/mol. The onset of this typically occurs at around 350 °C under standard atmospheric conditions. The process exhibits a characteristic profile, featuring gradual mass loss primarily between 300 and 540 °C due to progressive CO₂ release, with complete conversion to MgO requiring temperatures up to 800 °C. This stepwise progression reflects the solid-state nature of the reaction, where initial surface accelerates internal breakdown as heat penetrates the particles. In industrial , controlled heating within this range ensures high-purity magnesia production while minimizing . Kinetically, the decomposition follows a first-order mechanism, where the rate depends on the concentration of undecomposed MgCO₃. Factors such as particle size significantly influence the process, with smaller particles exhibiting faster decomposition due to increased surface area exposure to heat. Additionally, the surrounding atmosphere affects the rate; for instance, a CO₂-rich environment slows decomposition by shifting the equilibrium toward the reactant, necessitating higher temperatures for completion. The primary byproduct, magnesia (MgO), is a highly reactive form of valued for its applications in refractories, chemicals, and environmental processes; this material is routinely obtained through the of magnesium carbonate ores or synthetic forms.

Synthesis and Preparation

Natural Extraction

Magnesium carbonate occurs naturally primarily as the mineral (MgCO₃), which is extracted from deposits formed through sedimentary, metamorphic, or hydrothermal processes. The world's major magnesite deposits are concentrated in regions with suitable geological conditions, such as formations and basins. Global production of magnesite reached approximately 22 million metric tons in 2023, with estimates for 2024 remaining stable at around 20-22 million tons annually. China dominates global magnesite production, accounting for about 60% of the total output, primarily from deposits in Province, which holds roughly 85% of the country's reserves and has a mining capacity exceeding 22 million tons per year. Other significant producers include , , and , but China's share underscores its role as the primary supplier. Extraction typically begins with for near-surface deposits, which is the most common method due to the mineral's frequent occurrence in shallow layers; underground is used less frequently for deeper ores. Following extraction, the undergoes crushing to reduce , typically in multi-stage processes using and crushers to achieve sizes suitable for further beneficiation. Flotation is then employed to separate from minerals like dolomite and silica, utilizing collectors such as fatty acids to enhance hydrophobicity of magnesite particles; this yields concentrates with purity levels up to 95% MgCO₃ and recovery rates often exceeding 85%. Screening, washing, and may supplement flotation to remove impurities. Optional calcination of the beneficiated magnesite at temperatures around 700-800°C can produce magnesia (MgO), which may then be used to form basic magnesium carbonate through recarbonation, though this step is not always required for direct use of the carbonate concentrate. Synthetic production methods serve as alternatives when higher purity levels beyond 95% are needed for specialized applications.

Synthetic Production Methods

Magnesium carbonate can be synthesized in laboratory settings through reactions involving soluble magnesium salts and carbonate sources. A common method entails the reaction of (MgCl₂) with (Na₂CO₃) in , yielding magnesium carbonate precipitate according to the equation MgCl₂ + Na₂CO₃ → MgCO₃ + 2NaCl, often resulting in hydrated forms such as nesquehonite (MgCO₃·3H₂O) depending on temperature and concentration conditions. For basic magnesium carbonate variants, (NaHCO₃) is employed as the carbonate source, where controlled addition to MgCl₂ solution produces hydromagnesite-like structures (e.g., 4MgCO₃·Mg(OH)₂·4H₂O) through partial and . These processes are typically conducted at ambient temperatures (20–80 °C) under stirring to ensure uniform particle size and minimize impurities like co-precipitated sodium salts. Solvothermal methods offer an alternative route for producing hydrated magnesium carbonates with enhanced crystallinity and control over morphology, particularly for pharmaceutical or material applications requiring specific particle sizes. These involve reacting magnesium precursors, such as MgCl₂, with carbonate ions in mixed solvent systems like water-ethanol at elevated temperatures of 100–150 °C under sealed conditions, promoting the formation of dense phases like lansfordite (MgCO₃·5H₂O) or other trihydrates through solvothermal and recrystallization. The ethanol component aids in reducing and improving , leading to higher yields (up to 95%) of nanoscale or microcrystalline hydrates compared to conventional aqueous . This approach, often performed in autoclaves for 4–12 hours, allows tailoring of hydration levels by adjusting solvent ratios and , avoiding the need for high-energy steps. On an industrial scale, magnesium carbonate is predominantly produced from or sources via the of (Mg(OH)₂). The process begins with of Mg(OH)₂ from magnesium-rich brines using (Ca(OH)₂), followed by sparging (CO₂) gas through the slurry in a tower, converting it to basic magnesium carbonate per the reaction Mg(OH)₂ + CO₂ → MgCO₃ + H₂O (with excess CO₂ yielding hydrated or basic forms). This method leverages abundant resources, achieving production rates of thousands of tons annually with efficiencies improved by CO₂ from industrial emissions, and results in products suitable for fillers or antacids after and . Natural serves occasionally as a supplementary for and re-carbonation in hybrid processes.

Applications

Industrial and Material Uses

Magnesium carbonate plays a significant role in the production of materials, particularly in processes where it is incorporated into bricks to enhance heat resistance. During high-temperature exposure, it decomposes to form and , providing structural stability and resistance to in furnaces and kilns. This application leverages its ability to withstand extreme conditions, making it essential for lining steel converters and ladles. As a filler , magnesium carbonate is widely used in rubber, plastics, and industries to improve key properties without compromising performance. In rubber production, it enhances elasticity, durability, and wear resistance while maintaining a lightweight profile, suitable for applications in automotive and components. For plastics, it serves as a reinforcing agent that boosts mechanical strength and cost efficiency. In the sector, it acts as an internal filler and pigment, increasing brightness, opacity, and printability by effectively and achieving whiteness levels exceeding 98% in high-purity forms. In sports applications, is the primary component of absorbent used by gymnasts, climbers, and weightlifters to improve grip by absorbing moisture from hands. High-purity , often pharmaceutical-grade without fillers or additives, is used to minimize skin irritation and provide a more consistent, skin-friendly performance during extended use. Magnesium carbonate is also used in cosmetics for its mattifying effects on skin without irritation, in water treatment as a recycled coagulant for removing color and impurities from industrial wastes, and in paints, inks, and fertilizers to provide magnesium nutrients. As a food additive designated E504, magnesium carbonate functions as an anti-caking agent in products like table salts and candies, preventing clumping by improving powder flowability and stability in humid conditions. This role ensures product quality in powdered seasonings and confections without altering taste or texture.

Medical and Pharmaceutical Uses

Magnesium carbonate serves as an in pharmaceutical applications, neutralizing excess in the to alleviate symptoms of , , and upset , and is classified under the Anatomical Therapeutic Chemical (ATC) code A02AA01. It reacts with to produce gas and soluble , providing rapid symptomatic relief without significantly altering gastric for extended periods. The typical oral dose for antacid use ranges from 0.5 to 1 gram, taken up to four times daily between meals or as needed, often in chewable tablet or suspension form. In addition to its antacid properties, magnesium carbonate exhibits effects through an osmotic mechanism, drawing water into the intestines to soften stool and promote bowel movements, and is designated under ATC code A06AD01. This osmotic action makes it suitable for short-term relief of occasional , particularly when higher doses are administered, and it is sometimes incorporated into combination products with other magnesium salts for enhanced gastrointestinal regulation, similar to variants of milk of magnesia formulations. Its mild profile helps minimize cramping compared to stimulant laxatives, though overuse can lead to . Magnesium carbonate meets stringent pharmacopeial standards outlined in the United States Pharmacopeia (USP) and National Formulary (NF) , which specifies that it must contain the equivalent of not less than 40.0% and not more than 43.5% (MgO) on a dried basis. The also mandates purity tests, including limits on (not exceeding 30 ppm (0.003%) using Method I <231>), soluble salts, acid-insoluble substances, and calcium content to ensure safety for oral ingestion. These requirements confirm its suitability as a pharmaceutical-grade and , free from contaminants that could pose health risks. Historically, has been employed in medical practice since the to treat and related dyspeptic conditions, building on earlier alchemical knowledge of magnesium compounds for digestive relief. In modern formulations, it continues to play a role in managing (GERD) symptoms, often as part of over-the-counter combinations that provide quick onset of action for acid-related discomfort.

Safety, Toxicology, and Environmental Impact

Health and Safety Considerations

Magnesium carbonate dust can cause irritation upon , potentially leading to coughing and discomfort, particularly in occupational settings where exposure is prolonged. The (OSHA) has established a (PEL) of 15 mg/m³ for total dust and 5 mg/m³ for the respirable fraction to mitigate these risks. Ingestion of magnesium carbonate exhibits low acute toxicity, with an oral LD50 greater than 2,000 mg/kg in rats, indicating it is not highly poisonous in small amounts. However, overdose may result in , characterized by symptoms such as , , and cardiac effects, especially in individuals with impaired function. Contact with magnesium carbonate can act as a mild irritant to the skin and eyes, causing redness or discomfort upon direct exposure, though it is not absorbed systemically through the skin. It is classified as non-carcinogenic by the International Agency for Research on Cancer (IARC), with no evidence of mutagenic or in standard assessments. For safe handling, magnesium carbonate should be used in well-ventilated areas to minimize dust inhalation, with appropriate such as gloves, goggles, and respirators recommended during processing. The compound is chemically stable under normal conditions but should be stored away from strong acids to prevent and gas release; in case of exposure, first aid measures include moving to for inhalation incidents, rinsing with water for skin or , and seeking medical attention for . While utilized medicinally as an for its low toxicity profile, adherence to dosage guidelines is essential to avoid adverse effects.

Environmental Effects

Magnesite mining, the primary source of magnesium carbonate, often employs open-pit methods that lead to significant disruption by removing large volumes of and rock, resulting in , , and fragmentation. These operations also generate substantial dust , which settles on surrounding and water bodies, impairing air quality and contributing to in nearby areas. The production of magnesium carbonate derivatives, particularly through to yield magnesia, releases as a byproduct of , contributing to the material's and overall from the mineral industry. However, magnesium carbonate systems offer potential within a framework, where captured CO₂ can react with to reform carbonates, enabling repeated sequestration cycles that mitigate emissions over time. Brine extraction methods for magnesium carbonate, often sourced from seawater or salt lakes, require processing vast quantities of water—exceeding 800 tons per ton of magnesium produced—which strains local and exacerbates in arid regions. Effluents from these processes can alter receiving bodies by increasing through dissolved ions, potentially shifting levels and affecting aquatic ecosystems. Studies have shown concerns over climbing chalk, a powdered form of magnesium carbonate, where fine particles dispersed in outdoor environments can accumulate on rocks and soils, altering pH levels and disrupting microbial communities such as as well as and growth. These micro-particles from overuse in popular climbing areas contribute to long-term soil degradation, underscoring the need for sustainable alternatives to minimize ecological harm.

References

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