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| | Magnesium Chemical Properties |
| Melting point | 648 °C (lit.) | | Boiling point | 1090 °C (lit.) | | density | 1.74 g/mL at 25 °C (lit.) | | vapor density | 6 (vs air) | | vapor pressure | 1 mm Hg ( 621 °C) | | Fp | −26 °F | | storage temp. | Store below +30°C. | | solubility | H2O: 1 M at 20 °C, clear, colorless | | form | turnings | | color | White | | Specific Gravity | 1.74 | | resistivity | 4.46 μΩ-cm, 20°C | | Water Solubility | REACTS | | Sensitive | Hygroscopic | | Crystal Structure | HCP, Space Group P63/mmc | | Merck | 14,5674 | | BRN | 4948473 | | Exposure limits | ACGIH: TWA 2 ppm; STEL 4 ppm
OSHA: TWA 2 ppm(5 mg/m3)
NIOSH: IDLH 25 ppm; TWA 2 ppm(5 mg/m3); STEL 4 ppm(10 mg/m3) | | LogP | -0.57 at 20℃ | | CAS DataBase Reference | 7439-95-4(CAS DataBase Reference) | | NIST Chemistry Reference | Magnesium(7439-95-4) | | EPA Substance Registry System | Magnesium (7439-95-4) |
| | Magnesium Usage And Synthesis |
| Description | In Nature:
Magnesium is a chemical element with symbol Mg and atomic number 12. It is a shiny gray solid which bears a close physical resemblance to the other five elements in the second column (Group 2, or alkaline earth metals) of the periodic table: all Group 2 elements have the same electron configuration in the outer electron shell and a similar crystal structure.
Elemental magnesium is a gray-white lightweight metal, two-thirds the density of aluminium. It tarnishes slightly when exposed to air, although, unlike the other alkaline earth metals, an oxygen-free environment is unnecessary for storage because magnesium is protected by a thin layer of oxide that is fairly impermeable and difficult to remove. Magnesium has the lowest melting (923 K (1,202 °F)) and the lowest boiling point 1,363 K (1,994 °F) of all the alkaline earth metals.
Magnesium is probably one of the most common metals distributed in nature, constituting about 2.4% of the earth’s crust. The metal, however, does not occur in nature in elemental form. The principal minerals are dolomite [CaMg(CO3)2], magnesite MgCO3; carnallite KCl•MgCl2•6H2O, and silicate materials, such as talc Mg3(Si4O10)(OH)2 and asbestos H4Mg3Si2O9. Magnesium also is found in seawater, natural underground brines and salt deposits. Its concentration in sea water is 1,350 mg/L. Magnesium also occurs in all plants. Its porphyrin complex, chlorophyll, is essential for photosynthesis.
In Human body:
It is an essential nutrient element for humans. The dietary requirement for adults is about 300 mg per day. Magnesium plays an important role in over 300 enzymatic reactions within the body including the metabolism of food, synthesis of fatty acids and proteins, and the transmission of nerve impulses. It is one of the seven essential macrominerals; these are minerals that need to be consumed in relatively large amounts-at least 100 milligrams per day.
| | Uses | Magnesium metal and its alloys have numerous uses in chemical, electrochemical, metallurgy, and electronic industries. Its thermal and electrical properties, lightness, and ease of fabrication into useful shapes make it an attractive choice in industrial applications. The metal is alloyed with aluminum for various structural uses. Its alloys with zinc, copper, nickel, lead, zirconium and other metals have many uses too.
Magnesium alloys are used in automobile parts, aircraft, missiles, space vehicles, ship hulls, underground pipelines, memory discs, machine tools, furniture, lawn mowers, ladders, toys, and sporting goods. It also is used in making small and lightweight dry cell batteries. Chemical applications of magnesium include its use as a reducing agent, to prepare Grignard reagent for organic syntheses, and to purify gases.
Magnesium also is used in blasting compositions, explosive sensitizers, incendiaries, signal flares, and pyrotechnics. Magnesium salts have numerous uses. They are discussed individually.
| | Production Methods | Although many commercial processes have been developed since the first electrolytic isolation of Mg metal by Davy and Faraday, and Bussy, by chemical reduction, the principles of the manufacturing processes have not changed. At present, the metal is most commonly manufactured by electrolytic reduction of molten magnesium chloride, in which chlorine is produced as a by-product. In chemical reduction processes, the metal is obtained by reduction of magnesium oxide, hydroxide, or chloride at elevated temperatures.
All the magnesium produced in the world currently is derived from its minerals dolomite and carnallite, as well as from the underground brines and seawaters. In most processes, magnesium is recovered from its mineral or brine either as magnesium chloride or converted to the latter for electrolytic production.
Many subterranean brines are very rich in magnesium chloride, often containing about 11% MgCl2. Sodium and calcium chlorides are the other two major components (c.12% NaCl and 2% CaCl2) in such brines. Solar evaporation of the brine solution and repeated heating increases the MgCl2 concentration in the brine to above 25% at which the solubility of NaCl significantly decreases and it can be filtered out. Repeated spray drying and purification by chlorination yields anhydrous magnesium chloride.
Magnesium chloride produced from dolomite for electrolysis involves a series of steps that include calcinations of the mineral to oxide and then conversion to magnesium hydroxide, neutralization of the hydroxide with hydrochloric acid to form hydrated chloride, addition of sulfuric acid to separate out calcium as its insoluble sulfate, and dehydration of the hydrated salt to yield anhydrous MgCl2. Similar steps are also followed to obtain the metal from seawater. The average concentration of magnesium ion in seawater is about 1,200 mg/L, thus making ocean water an enormous source of magnesium. Magnesium is precipitated as hydroxide by treatment with lime in an agitated flocculator:
MgCl2 + Ca(OH)2 → Mg(OH)2 + CaCl2
The insoluble Mg(OH)2 is filtered off and the seawater containing calcium chloride is returned to the sea. The hydroxide is then neutralized with hydrochloric acid. Evaporation of the solution yields hexahydrate, MgCl2•6H2O. The hexahydrate is either fully dehydrated to anhydrous MgCl2 by heating in dryers or partially dehydrated to monohydrate for electrolytic 512 MAGNESIUMproduction of metal. Magnesium hydroxide produced from seawater alternatively may be calcined to magnesium oxide, MgO. The latter is reduced with carbon and converted to magnesium chloride by heating in an electric furnace in the presence of chlorine gas:
MgO + C + Cl2 → MgCl2 + CO
MgO + CO + Cl2 → MgCl2 + CO2
Manufacturing processes, based on thermal reduction of magnesium oxide employ ferrosilicon or carbon as a reducing agent and use dolomite as the starting material. In these processes, the mineral is first calcined to produce oxides of magnesium and calcium, MgO•CaO. In one such batch process, known as the Pidgeon process, calcined dolomite is mixed with pulverized ferrosilicon powder, briquetted, and charged into an electrically-heated retort made of nickel-chrome-steel alloy and operated under vacuum (0.1 to 0.2 mm Hg). The reaction is carried out at about 1,150°C for several hours (8 hours). Silicon reduces magnesium oxide to metallic magnesium produced as vapor. The vapors condense into crystals in the cooler zone of the retort (500°C). The reactions are as follows:
2(MgO•CaO) + Si(Fe) → 2 Mg + 2CaO•SiO2(Fe)
The ferrosilicon alloy required in the above process is produced by thermal reduction of silica with carbon in the presence of iron:
SiO2 + 2C + Fe → Si(Fe) + 2CO
In the Pidgeon process discussed above, a secondary side reaction occurs between the CaO and SiO2 forming dicalcium silicate:
2CaO + SiO2 → Ca2SiO4
In a modified method known as Magnetherm process, sufficient aluminum oxide is added to melt this Ca2SiO4 slag. This allows the products to be removed in the molten state and, in addition, heats the reactor by the electrical resistance of the slag.
Magnesium also is produced by thermal reduction of its oxide by carbon:
MgO + C → Mg + CO
The above reaction is reversible above 1,850°C. The metal produced as vapor must be cooled rapidly to prevent any reversible reactions. Rapid cooling (shock cooling) can quench the reaction giving finely divided pyrophoric dust of the metal. The separation, however, is difficult. This makes the carbon reduction process less attractive than the other two thermal reduction processes, namely Pidgeon and Magnetherm processes.
| | Category | Water flammable items | | Explosive hazardous characteristics | It is easily explosive after reacting with water and producing hydrogen. | | Storage characteristics | Ventilated warehouse, low temperature, dry; separated storage with oxidants and acid | | Extinguishing agent | Graphite powder, dry sand. | | Description | Magnesium is a Group 2 element (Group IIA in older labeling schemes). This element has the symbol Mg, atomic number 12, atomic weight of 24.305 g/mol and common oxidation number ＋2. It is the eighth most abundant element in the earth s crust by mass, although ninth in the Universe as a whole. This preponderance of magnesium in the Universe is related to the fact that it is easily built up in supernova stars from a sequential addition of three helium nuclei to carbon (which in turn is made from three helium nuclei). Magnesium constitutes about 2% of the Earth s crust by mass, which makes it the eighth most abundant element in the crust. Magnesium ion’s high solubility in water helps to
ensure that it is the third most abundant element dissolved
in seawater. | | Chemical Properties | Silvery, moderately hard, alkaline-earth
metal; readily fabricated by all standard methods.
Lightest of the structural metals; strong reducing
agent; electrical conductivity similar to aluminum.
Soluble in acids; insoluble in water. | | Chemical Properties | Magnesium is a light, silvery-white metal in
various forms, and is a fire hazard. | | Physical properties | Magnesium is a lightweight, silvery-white, malleable alkali earth metal that is flammable.It has a weak electronegativity (–1.31), which means it is highly reactive as it combines withsome nonmetals. As with other alkali earth metals, magnesium is a good conductor of heatand electricity. Its melting point is 648.8°C, its boiling point is 1090°C, and its density is1.74 g/cm3, making it about one-fifth the density of iron and only two-thirds as dense asaluminum. | | Isotopes | There are 15 isotopes of magnesium, ranging from Mg-20 to Mg-34. Threeof these isotopes are stable: Mg-24 makes up 78.99% of all magnesium found in theEarth’s crust. Mg-25 makes up 10%, and Mg-26 constitutes most of the rest at 11%.The other 12 isotopes are radioactive and are produced artificially with half-lives rangingfrom microseconds to a few hours. | | Isotopes | Magnesium has three stable isotopes: 24Mg, 25Mg and
26Mg. All are present in significant amounts. About 79% of Mg is 24Mg. The isotope 28Mg is radioactive and in the 1950s to 1970s was made commercially
by several nuclear power plants for use in scientific
experiments. This isotope has a relatively short
half-life (21 h) and so its use was limited by shipping
times. 26Mg has found application in isotopic geology,
similar to that of aluminum. 26Mg is a radiogenic
daughter product of 26Al, which has a half-life of
717,000 years. Large enrichments of stable 26Mg have
been observed in the Ca–Al-rich inclusions of some
carbonaceous chrondrite meteorites. The anomalous
abundance of 26Mg is attributed to the decay of its parent
26Al in the inclusions. Therefore, the meteorite must
have formed in the solar nebula before the 26Al had
decayed. Hence, these fragments are among the oldest
objects in the solar system and have preserved information
about its early history.
It is conventional to plot 26Mg/24Mg against an Al/
Mg ratio. In an isochronic dating plot, the Al/Mg ratio
plotted is 27Al/24Mg. The slope of the isochron has no
age significance, but indicates the initial 26Al/27Al ratio in the sample at the time when the systems were separated
from a common reservoir. | | Origin of Name | Magnesium is named after Magnesia, an ancient region of Thessaly,
Greece, where it was mined. Magnesium is often confused with another element, manganese. One way to eliminate the confusion is to think of magnesium (Mg) as “12” and
manganese (Mn) as “25” and to use the mental trick of remembering that “g” comes
before “n” in the alphabet, so magnesium is the one with lower atomic number. | | Occurrence | Magnesium is the eighth most abundant of the elements found in the entire universe, andthe seventh most abundant found in the Earth’s crust. Its oxide (MgO) is second in abundance to oxide of silicon (SiO2), which is the most abundant oxide found in the Earth’s crust.Magnesium is found in great quantities in seawater and brines, which provide an endless supply. Each cubic mile of seawater contains about 12 billion pounds of magnesium. Althoughmagnesium metal cannot be extracted from seawater directly, it can be extracted by severalchemical processes through which magnesium chloride (MgCl2) is produced. Electrolysis isthen used with the magnesium chloride as the electrolyte at 714°C to produce metallic magnesium and chlorine gas. Another method of securing magnesium is known as the Pigeonprocess. This procedure uses the magnesium minerals dolomite or ferrosilicon. Dolomite(CaCO3), which also contains MgCO3, is crushed and then heated to produce oxides of Caand Mg. The oxides are heated to about 1200°C along with the ferrosilicon (an alloy of ironand silicon), and the silicon reduces the magnesium, producing a vapor of metallic magnesiumthat, as it cools, condenses to pure magnesium metal. | | History | The name originates from the Greek word for a district in Thessaly called Magnesia . It is related to the terms magnetite and manganese , which also originated from this area, and required differentiation as separate substances. Magnesium is the seventh most abundant element in the Earth s crust by mass and eighth by molarity. It is found in large deposits of Magnesite, Dolomite and other minerals, and in mineral waters, where the magnesium ion is soluble. In 1618 a farmer at Epsom in England attempted to give his cows water from a well. They refused to drink because of the water s bitter taste. However the farmer noticed that the water seemed to heal scratches and rashes. The fame of Epsom Salts spread. Eventually the compound was recognized to be hydrated magnesium sulfate, MgSO4. The first person to propose that magnesium was an element was Joseph Black of Edinburgh in 1755. In 1792, an impure form of metallic magnesium was produced by Anton Rupprecht who heated magnesia (magnesium oxide, MgO) with charcoal. He named the element Austrium after his native Austria. In 1808, a small sample of the pure metal was isolated by Humphry Davy by the electrolysis of moist MgO. He proposed the name magnium based on the mineral Magnesite (MgCO3) that came from Magnesia in Greece. Neither name survived and eventually the metal was called magnesium. The metal itself was first produced in quantity in England by Davy in 1808 using then the new method of electrolysis of a mixture of molten magnesia and mercuric oxide. Antoine Bussy prepared it in a consistent form in 1831. | | Characteristics | While in a thin solid form, magnesium ignites at 650°C, and it is more easily ignited ina fine powder form. Burning magnesium produces a brilliant white light. It is also used asan oxidizer to displace several other metals from their compound minerals, salts, and ores. Itis alloyed with other metals to make them lighter and more machinable, so that they can berolled, pounded, formed into wires, and worked on a lathe.The ground water in many regions of the United States contains relatively high percentagesof magnesium, as well as some other minerals. A small amount improves the taste of water,but larger amounts result in “hard” water, which interferes with the chemical and physicalaction of soaps and detergents. The result is a scum-like precipitate that interferes with thecleansing action. The solution is the use of water softeners that treat hard water with eithersodium chloride or potassium chloride, which displace the magnesium—making the water“soft,” resulting in a more effective cleansing action. | | History | Compounds of magnesium have long been known. Black recognized magnesium as an element in 1755. It was isolated by Davy in 1808, and prepared in coherent form by Bussy in 1831. Magnesium is the eighth most abundant element in the Earth’s crust. It does not occur uncombined, but is found in large deposits in the form of magnesite, dolomite, and other minerals. The metal is now principally obtained in the U.S. by electrolysis of fused magnesium chloride derived from brines, wells, and sea water. Magnesium is a light, silvery-white, and fairly tough metal. It tarnishes slightly in air, and finely divided magnesium readily ignites upon heating in air and burns with a dazzling white flame. It is used in flashlight photography, flares, and pyrotechnics, including incendiary bombs. It is one third lighter than aluminum, and in alloys is essential for airplane and missileconstruction. The metal improves the mechanical, fabrication, and welding characteristics of aluminum when used as an alloying agent. Magnesium is used in producing nodular graphite in cast iron, and is used as an additive to conventional propellants. It is also used as a reducing agent in the production of pure uranium and other metals from their salts. The hydroxide (milk of magnesia), chloride, sulfate (Epsom salts), and citrate are used in medicine. Dead-burned magnesite is employed for refractory purposes such as brick and liners in furnaces and converters. Calcined magnesia is also used for water treatment and in the manufacture of rubber, paper, etc. Organic magnesium compounds (Grignard’s reagents) are important. Magnesium is an important element in both plant and animal life. Chlorophylls are magnesiumcentered porphyrins. The adult daily requirement of magnesium is about 300 mg/day, but this is affected by various factors. Great care should be taken in handling magnesium metal, especially in the finely divided state, as serious fires can occur. Water should not be used on burning magnesium or on magnesium fires. Natural magnesium contains three isotopes. Twelve other isotopes are recognized. Magnesium metal costs about $100/kg (99.8%). | | Uses | magnesium plays an important role in various processes within the skin, including amino acid synthesis and protein synthesis (e.g., collagen), and in the metabolism of calcium, sodium, and phosphorus. | | Uses | Solid state synthesis with Ca and Sn resulted in a new phase, Ca6.2Mg3.8Sn7, which has an unprecedented type of tin chain composed of square-planar tin units.1 | | Uses | In alloys to produce light weight structural metals. In aluminum alloys to improve mechanical properties; in Grignard reagents; in dry cell batteries; in pyrotechnics. For hot metal desulfurization, especially. molten iron; production of ductile iron; metal reduction to produce elemental boron, titanium, zirconium; corrosion protection of steel structures; sacrificial anodes for corrosion protection. | | Uses | Small particles of powdered magnesium metal burn with a bright white flame that makesthe magnesium ideal for aerial flares dropped from airplanes that will light up ground areas. Itis has also been used in aerial firebombs during wars to devastate a city by fire because waterwill not extinguish the flames—sand must be used. In the past decades, thin magnesium wireor foil was placed inside glass bulbs containing pure oxygen to form flash bulbs for photographic purposes. When an electric charge ignites the magnesium, a brilliant light is produced.Today most flash cameras use a strobe light instead of flash bulbs.Pure magnesium metal is lighter in weight than aluminum and, thus, would make anexcellent construction metal were it not for its high reactivity and flammability at a rather lowtemperature when compared to other metals. It is an excellent metal to alloy with other metalsfor use in the aircraft, space, and automobile industries.It is used for the production (thermal reduction) of other metals, such as zinc, iron, titanium, zirconium, and nickel. For instance, because of its strong electropositive nature, magnesium can “desulfurize” molten iron when it combines with the sulfur impurities in the ironto produce high-grade metallic iron plus MgS.Milk of Magnesia is an alkaline (basic) water suspension and “creamy-like” suspended formof magnesium hydroxide, Mg(OH)2. It is used as an antacid to neutralize excess stomach acid.Magnesium can also be used in the form of Epsom salts as a treatment for rashes and as alaxative. A more important commercial use of Epsom salts is in the tanning of leather, as wellas in the dyeing of fabrics.Magnesium is essential for proper nutrition in humans as well as other living organisms.It plays an important role in the process of photosynthesis in plant chlorophyll and is thusessential to green plants, which are, in turn, essential for most living organisms. Magnesiumis also used as a dietary supplement for both humans and animals for maintaining properenzyme levels.Magnesium is an important element that acts as a catalyst in many life processes. In addition to photosynthesis, it is also required for the oxidation in animal cells that produce energyand for the production of healthy red blood cells. Humans cannot live without magnesium—which we acquire mainly from various foods. | | Uses | Magnesium is a metallic element that is involved in certain bodily
functions. sources of magnesium include magnesium chloride and
magnesium oxide. it functions as a nutrient and dietary supplement. | | Uses | Magnesium is used in the manufacture ofalloys, optical mirrors, and precision instruments;in pyrotechnics; as a deoxidizing anddesulfurizing agent in metallurgy; in signallights, flash bulbs, and dry batteries; and inGrignard reagent. | | Uses | Magnesium powder is used in the manufacture of
fireworks and marine flares where a brilliant white light
is required. Flame temperatures of magnesium and
magnesium alloys can reach 1371°C (2500 F), although
flame height above the burning metal is usually less
than 300 mm (12 in). Magnesium may be used as an
ignition source for “thermite”, or otherwise difficult to
ignite mixture of aluminum and iron oxide powder.
Magnesium compounds are typically white crystals.
Most are soluble in water, providing the sour-tasting
magnesium ion, Mg2+. Small amounts of dissolved
magnesium ion contribute to the tartness and taste of
natural waters. Magnesium ion in large amounts is an
ionic laxative, and magnesium sulfate (known as
“Epsom Salts”) is sometimes used for this purpose. Socalled
“milk of magnesia” is a water suspension of one
of the few insoluble magnesium compounds,
Mg(OH)2. The undissolved particles give rise to its
appearance and name. Milk of magnesia is a mild base
commonly used as an antacid.
Commercially, the chief use for the metal is as an alloying agent to make Al Mg alloys, sometimes called magnalium or magnelium . Since magnesium is less dense than aluminum, these alloys are valued for their relative lightness and strength. Magnesium is an important element for plant and animal life. The adult human daily requirement of magnesium is about 0.3 g/day. Magnesium is the 11th most abundant element by mass in the human body. Its ions are essential to all living cells, where they play a major role in manipulating important biological polyphosphate compounds like ATP, DNA and RNA. Hundreds of enzymes thus require magnesium ions in order to function. Magnesium, being the metallic ion at the center of chlorophyll, is thus a common additive to fertilizers. Magnesium compounds are used medicinally as common laxatives, antacids (i.e. Milk of Magnesia ), and in a number of situations where stabilization of abnormal nerve excitation and blood vessel spasm is required (i.e. to treat eclampsia). Magnesium ions are sour to the taste, and in low concentrations help to impart a natural tartness to fresh mineral waters.
Magnesium is also used:
? To remove sulfur from iron and steel.
? To refine titanium in the “Kroll” process.
? To photoengrave plates in the printing industry.
? To combine in alloys, where this metal is essential for
airplane and missile construction.
? In the form of turnings or ribbons, to prepare “Grignard
Reagents”, which are useful in organic synthesis.
? As an alloying agent, improving the mechanical,
fabrication and welding characteristics of aluminum.
? As an additive agent in conventional propellants and
the production of “nodular graphite” in cast iron.
? As a reducing agent for the production of uranium
and other metals from their salts.
? As a desiccant, since it easily reacts with water.
? As a sacrificial (galvanic) anode to protect underground
tanks, pipelines, buried structures, and water heaters. | | Definition | magnesium: Symbol Mg. A silverymetallic element belonging to group 2 (formerly IIA) of the periodic table; a.n. 12;r.a.m. 24.305; r.d. 1.74; m.p. 648.8°C;b.p. 1090°C. The element is found ina number of minerals, includingmagnesite (MgCO3), dolomite(MgCO3.CaCO3), and carnallite(MgCl2.KCl.6H2O). It is also present in sea water, and it is an essential element for living organisms. Extraction is by electrolysis of the fusedchloride. The element is used in a number of light alloys (e.g. for aircraft).Chemically, it is very reactive.In air it forms a protective oxide coating but when ignited it burns with an intense white flame. It also reacts with the halogens, sulphur, and nitrogen.Magnesium was first isolatedby Bussy in 1828. | | Definition | Metallic element of atomic
number 12, group IIA of the periodic table, aw
24.305, valence = 2; three isotopes. Magnesium
is the central element of the chlorophyll molecule;
it is also an important component of red blood corpuscles. | | General Description | A light silvery metal. The more finely divided material reacts with water to liberate hydrogen, a flammable gas, though this reaction is not as vigorous as that of sodium or lithium with water. In finely divided forms is easily ignited. Burns with an intense white flame. Can be wax coated to render magnesium as nonreactive. | | Air & Water Reactions | Pyrophoric in dust form [Bretherick 1979, p. 104]. Magnesium ribbon and fine magnesium shavings can be ignited at air temperatures of about 950°F and very finely divided powder has been ignited at air temperatures below 900°F. [Magnesium Standard 1967 p. 4]. The more finely divided material reacts with water to liberate hydrogen, a flammable gas, though this reaction is not as vigorous as that of sodium or lithium | | Reactivity Profile | Magnesium slowly oxidizes in moist air. Reacts very slowly with water at ordinary temperatures, less slowly at 100°C. Reacts with aqueous solutions of dilute acids with liberation of hydrogen [Merck 11th ed. 1989]. In the presence of carbon, the combination of chlorine trifluoride with aluminum, copper, lead, magnesium, silver, tin, or zinc results in a violent reaction [Mellor 2, Supp. 1. 1956]. A mixture of powdered magnesium with trichloroethylene or with carbon tetrachloride will flash or spark under heavy impact [ASESB Pot. Incid, 39. 1968]. Stannic oxide, heated with magnesium explodes [Mellor 7:401. 1946-47]. When carbon dioxide gas is passed over a mixture of powdered magnesium and sodium peroxide, the mixture exploded [Mellor 2:490. 1946-47]. Powdered magnesium plus potassium (or sodium) perchlorate is a friction- sensitive mixture [Safety Eng. Reports. 1947]. An explosion occurred during heating of a mixture of potassium chlorate and magnesium [Chem. Eng. News 14:451. 1936]. Powdered magnesium can decompose performic acid violently [Berichte 48:1139. 1915]. A mixture of finely divided magnesium and nitric acid is explosive [Pieters 1957. p. 28]. Magnesium exposed to moist fluorine or chlorine is spontaneously flammable [Mellor 4:267. 1946-47]. | | Hazard | (Solid metal) Combustible at 650C. (Powder, flakes, etc.) Flammable, dangerous fire hazard.
Use dry sand or talc to extinguish. | | Hazard | Magnesium metal, particularly in the form of powder or small particles, can be ignited atrelatively low temperatures. The resulting fires are difficult to extinguish, requiring dry sandor dirt. Water will just accelerate the fire as hydrogen that will intensify the fire is releasedfrom the water.
Some magnesium compounds, whose molecules contain several atoms of oxygen—Mg(ClO4), for example—are extremely explosive when in contact with moist organic substance,such as your hands.
Although traces of magnesium are required for good nutrition and health, some compoundsof magnesium are poisonous when ingested. | | Health Hazard | Inhalation of magnesium dust can produceirritation of the eyes and mucous membranes.Magnesium may react with waterin the bronchial passage to form magnesiumhydroxide, which is caustic and maycause adverse effects on lungs. The fumescan cause metal fever. | | Health Hazard | Dust irritates eyes in same way as any foreign material. Penetration of skin by fragments of metal is likely to produce local irritation, blisters, and ulcers which may become infected. | | Fire Hazard | Behavior in Fire: Forms dense white smoke. Flame is very bright. | | Agricultural Uses | Magnesium (Mg) is an essential element for plant and
animal growth. It belongs to Group 2, and has an atomic
weight of 24.32 and atomic number of 13.
Magnesium is the eighth most abundant element in the
earth's crust. It is made by electrolysis of fused
magnesium chloride taken from sea water. Magnesium is
a light, silvery-white, hard, reactive metal. It plays a
crucial role in the life of both plants and animals.
Magnesium and its compounds are also used in light
metal alloys, incendiary devices, flash bulbs, flares,
fertilizers and in medicine.
Magnesium is a constituent of chlorophyll,
protochlorophyll, pectin and Phyllis. While its role in
plant metabolism is not very clear, it seems to perform
many functions in plants. For example, as the only
metallic constituent of chlorophyll, Mg gives green color to leaves and has the structure of hemoglobin. It plays a
role in photosynthesis, forming hexose sugar from water
and carbon dioxide in the presence of sunlight.
Magnesium regulates the uptake of other materials by the
plant, and acts as a carrier of phosphorus to the seeds in
the plant. Mg promotes the formation of oils and fats. It
plays a role in the translocation of starch. Almost the
whole of magnesium dissolves in the cell sap of the plant
and becomes readily mobile in the plant. Many important
colloidal chemical functions are ascribed to this fraction
of magnesium.
Magnesium also participates in the production of
proteins, fats, vitamins and some catalytic reactions in
the enzyme system. It is mobile in plants and serves as a
structural component in the ribosome, playing an
important role in protein synthesis.
The above ground portion of most mature grain crops
and grasses contain about 0.1 to 0.4% of magnesium,
whereas that of cotton, soybean and alfalfa plants contain
0.3 to 0.6%. Plants absorb magnesium as a divalent
cation (Mg2+). Its absorption depends on many factors,
such as the amount of solution Mg2+, the soil pH and
type, the percentage of Mg saturation on the cation
exchange complex (CEC), and quantities of other
exchangeable ions. Many soils absorb magnesium in a
non-exchangeable form (MgCO3). Nitrate ions promote
its absorption, whereas the ions of ammonium, potassium
and calcium ions restrict it.
Plant species and varieties differ in their magnesium
requirement. For instance, corn, potato, oil palm, cotton,
citrus, tobacco, sugar beet and pastures respond to a high
magnesium content. Seasonal and environmental
conditions interact with plant varieties for magnesium
uptake and cause magnesium deficiency.
The non-availability of magnesium in soils having
less than one mole of the exchangeable magnesium per kg
of soil, or the presence of magnesium in amounts less
than 4% of the CEC, are indications of magnesium
deficiency. Magnesium deficiencies occur in soils with
high ratios of exchangeable Ca/Mg which should not
exceed 10: 1 or 15: 1, depending on specific conditions. A
high level of exchangeable potassium may interfere with
the uptake of Mg by crops. The recommended ratios of K
to Mg are less than 5 : 1 for field crops, 3: 1 for vegetables
and sugar beets, and 2: 1 for fruits and green house crops.
The symptoms of magnesium deficiency, which do
not occur too frequently, first appear on older leaves and
then spread to younger ones. The green chlorophyll disappears, leaving behind spots between the leaf veins.
The leaf margin then tums yellow (interveinal chlorosis
in older leaves).The leaves exhibit a stripy or spotty
appearance. However, unlike the deficiencies of K and
Cu, the Mg deficiency symptoms of necrosis seldom
occur, except for chlorotic discoloration. A large number
of leaves may fall as a result of magnesium deficiency,
especially in fruit and berry crops. Magnesium deficiency
causes significant injuries, particularly in fruit crops,
which may extend to the roots, and create phosphorus
deficiency in oil plants, such as palm and linseed.
Magnesium deficiency in cotton and grapes appears as
purplish red leaves with green veins. As the leaves
become older, they turn brown. The lower leaves are
affected first in corn, as whitish stripes appear along the
veins and a purplish color is seen on the underside. In
tobacco, it is known as sand down and appears as loss of
green color at the tips of the lower leaves. As the
deficiency worsens, the upper leaves become bleached
and turn white in color. The deficiency in animals shows
up as low blood-serum magnesium and muscle spasm,
finally leading to death.
Soil analysis is widely used to detect the Mg
deficiency and to estimate the Mg requirement of the
plant. The most effective material for correcting
magnesium deficiency and soil acidity is dolomitic
limestone or dolomite. Magnesium uptake is greater
from fine dolomite than from the coarse variety, while it
is less than that from magnesium sulphate. An addition of
16.8 to 33.6 kg/ha of dolomite significantly increases the
dry weight of corn. Similarly, the clover yield is higher
with soluble magnesium than with dolomite. For soil with
a pH more than 6.0, water-soluble magnesium sulphate is preferred to dolomite as a source of magnesium.
Other materials containing magnesium are magnesia,magnesium nitrate, magnesium silicate, serpentine,magnesium chloride solution, synthetic chelates and natural organic complexing substances. Magnesium sulphate (MgSO4), magnesium chloride (MgCl2),magnesium nitrate [Mg(NO3)2] and synthetic and natural chelates are well suited for clear-liquid foliar applications.
The double sulphate of potassium and magnesium is
the most widely used magnesium additive for
suspensions. Magnesium ammonium phosphate has nonburning
and non-leaching characteristics. These are
especially valuable when the fertilizer comes in contact
with seeds or roots.
Depending on factors like the magnesium content,
the rate of weathering, uptake by plants etc., magnesium
ions (Mg2+) an be leached from soils.
Forage crops, particularly forage grass with
magnesium concentration less than 2 g/kg, are dangerous
for the cattle which on consumption of such grass, may
get a disease called hypomagnesemia or grass tetany, in
which the blood magnesium level decreases abnormally. | | Safety Profile | Inhalation of dust and
fumes can cause metal fume fever. The
powdered metal igrutes readily on the skin
causing burns. Particles embedded in the
skin can produce gaseous blebs that heal
A dangerous fire hazard in the form of
dust or flakes when exposed to flame or
oxiduing agents. In solid form, magnesium
is difficult to ipte because heat is conducted rapidly away from the source of
ignition; it must be heated above its melting
point before it will burn. However, in finely
divided form, it may be ignited by a spark or
the flame of a match. Magnesium fires do
not flare up violently unless there is
moisture present. Therefore, it must be kept
away from water, moisture, etc. It may ignited spontaneously when the material is
finely divided and damp, particularly with
water-oil emulsion. Moderately explosive in
the form of dust when exposed to flame.
Also, magnesium reacts with moisture, acids,
etc., to evolve hydrogen, a highly dangerous
fire and explosion hazard.
Explosive reaction or ignition with
calcium carbonate + hydrogen + heat, gold
cyanide + heat, mercury cyanide + heat,
silver oxide + heat, fused nitrates, phosphates, or sulfates (e.g., ammonium nitrate,
metal nitrates), chloroformamidinium nitrate
+ water (when ignited with powder). The
powder may explode on contact with
halocarbons (e.g., chloromethane,
chloroform, or carbon tetrachloride), and
explodes when sparked in dichlorodifluoromethane. Hypergolic reaction with nitric
acid + 2-nitroanhe. Mixtures of
powdered magnesium and methanol are
more powerful than some mihtary
explosives. Mixtures of magnesium powder
+ water can be detonated. Reacts with
acetylenic compounds including traces of
acetylene found in ethylene gas to form
explosive magnesium acetylide.
chlorate salts, beryllium fluoride, boron
diiodophosphide, carbon tetrachloride +
methanol, 1,1,1 -trichloroethane, 1,2
dibromoethane, halogens or interhalogens
(e.g., fluorine, chlorine, bromine, iodine
vapor, chlorine trifluoride, iodine
heptafluoride), hydrogen iodide, metal
oxides + heat (e.g., berylhum oxide,
cadmium oxide, copper oxide, mercury
oxide, molybdenum oxide, tin oxide, zinc
oxide), nitrogen (when ipted), silicon
dioxide powder + heat,
polytetrafluoroethylene powder + heat, sulfur + heat, tellurium + heat, barium
peroxide, nitric acid vapor, hydrogen
peroxide, ammonium nitrate, sodium iodate
+ heat, sodium nitrate + heat, dinitrogen
tetraoxide (when ignited), lead dioxide.
Ignites in carbon dioxide at 780°C, molten
barium carbonate + water, fluorocarbon
polymers + heat, carbon tetrachloride or
trichloroethylene (on impact),
dichlorodifluoromethane + heat.
Incompatible with ethylene oxide, metal
oxosalts, oxidants, potassium carbonate, Al
+ KClO4, [Ba(NO3)2 + BaO2 + Zn],
bromobenzyl trifluoride, CaC, carbonates,
CHCb, LCuSO4 (anhydrous) + NH4NO3 +
KClO3 + H2O], CuSO4, (H2 + CaCO3),
CH3Cl, N02, liquid oxygen, metal cyanides
(e.g., cadmium cyanide, cobalt cyanide,
copper cyanide, lead cyanide, nickel cyanide,
zinc cyanide), performic acid, phosphates,
KClO3, KClO4, AgNO3, NaClO4, (Na2O2 +
CO2), sulfates, trichloroethylene, Na2O2.
To fight fire, operators and firefighters can
approach a magnesium fEe to within a few
feet if no moisture is present. Water and
ordinary extinguishers, such as CO2, carbon
tetrachloride, etc., should not be used on
magnesium fires. G-1 powder or powdered
talc should be used on open fires.
Dangerous when heated; burns violently in
air and emits fumes; will react with water or
steam to produce hydrogen. See also
MAGNESIUM COMPOUNDS. | | Potential Exposure | Magnesium alloyed with manganese,
aluminum, thorium, zinc, cerium, and zirconium, is used in
aircraft, ships, automobiles, hand tools, etc., because of its
lightness. Dow metal is the general name for a large group
of alloys containing over 85% magnesium. Magnesium
wire and ribbon are used for degassing valves in the radio
industry and in various heating appliances; as a deoxidizer
and desulfurizer in copper, brass, and nickel alloys;
in chemical reagents; as the powder in the manufacture
of flares, incendiary bombs, tracer bullets, and flashlight
powders; in the nuclear energy process; and in a cement
of magnesium oxide and magnesium chloride for floors.
Magnesium is an essential element in human and animal
nutrition and also in plants, where it is a component of altypes of chlorophyll. It is the most abundant intracellular
divalent cation in both plants and animals. It is an activator
of many mammalian enzymes | | Carcinogenicity | MgO is regarded as an
“experimental tumorigen”, although the only reference
in the literature that could be found relating to the carcinogenicity of MgO was an instillation study, in which
MgO dust instilled intratracheally for 30 weeks resulted in
induction of histiocytic lymphomas in hamsters. It was
also demonstrated that MgO enhanced the tumorigenesis of
benzo[a]pyrene and was an effective carrier agent for the
experimental induction of respiratory tract tumors. | | Environmental Fate | Aquatic fate: Because of magnesium ion’s high solubility in
water, it is the third most abundant element dissolved in
seawater, because rainwater falling on rocks can also increase
the level of magnesium in river and seawater. | | Shipping | UN1869 Magnesium pellets, turnings or ribbons,
Hazard Class: 4.1; Labels: 4.1-Flammable solid. UN1418
Magnesium, powder or Magnesium alloys, powder, Hazard
Class: 4.3; Labels: 4.3-Dangerous when wet material,
4.2-Spontaneously combustible material. UN2950 Magnesium
granules, coated, particle size not <149 μm, Hazard Class:
4.3; Labels: 4.3-Dangerous when wet material | | Purification Methods | It slowly oxidises in moist air and tarnishes. If dark in colour, do not use. The shiny solid should be degreased by washing with dry Et2O, dry it in vacuo and keep it in a N2 atmosphere. It can be activated by stirring it in Et2O containing a crystal of I2 then filtering it off, before drying and storing. [Gmelin’s Magnesium (8th edn) 27A 121 1937.] | | Toxicity evaluation | Magnesium is a vital cation involved in cell homeostasis. In
normal humans, magnesium increases cardiac output, dilates
the coronary arterioles, and decreases systolic arterial pressure
and systemic vascular resistance. However, magnesium levels
outside of the normal range alter cellular ion balances and
activity, especially Ca2+ activity, which directly affects neural
and muscular functions. Parenterally administered magnesium
decreases acetylcholine in motor nerve terminals and
acts on myocardium by slowing rate of SA node impulse
formation and prolonging conduction time. One study also
found magnesium in relatively high amounts in about half of
human colon cancers, but the relationship is unknown and
animal studies have found that magnesium actually reduces
sarcoma incidence in some nickel- and cadmium-induced
tumors. | | Incompatibilities | Dust may form explosive mixture with
air. Capable of self-ignition in moist air. The substance is
a strong reducing agent. Reacts violently with, oxidizers,
strong acids; acetylene, ammonium salts; arsenic, beryllium
fluoride, carbon tetrachloride, carbonates, chloroform,
cyanides, chlorinated hydrocarbons; ethylene oxide; hydrocarbons, metal oxides; methanol, phosphates, silver nitrate;
sodium peroxide; sulfates, trichloroethylene, and many
other substances, causing fire and explosion hazards. Finely
divided material, in powdered, chip or sheet form, reacts
with moisture or acids, evolving flammable hydrogen gas,
causing fire and explosion hazard. Finely divided form
is readily ignited by a spark or flame. It splatters and burns
at above 1260℃ |
| | Magnesium Preparation Products And Raw materials |
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