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Magnesium and its Compounds

1. Discovery and Naming of Magnesium

In 1808 Sir Humphrey Davy reported the production of magnesium in the form of amalgam by electrolytic reduction of its oxide using a mercury cathode. In 1828 the French scientist A. Bussy fused magnesium chloride with metallic potassium and became the first to produce free magnesium. Michael Faraday in 1833 was the first to produce magnesium by electrolytic reduction from the chloride...

Source: Kirk-Othmer Encyclopedia Of Chemical Technology fifth Edition volume 15 page 320 (WSU Reference TP9.K54 2004 V.15)

The name magnesium goes back many centuries. It was selected in honor of a region in Greece known as Magnesia. The region contains large supplies of magnesium compounds.

Source: Chemical Elements (David E Newton) Pages 309-310 (WSU Reference QD 466.N464 1999V.2)

Three different substances were described during antiquity by terms related to our word "magnesia" (named from a district in Thessaly Greece). One, magnesius lapis, referred to magnetite, and another, magnesia niger, to pyrolusite. The third, magnesia, was probably applied to the soft white mineral steatite, also known as soapstone or talc, a hydrated magnesium silicate.

Source: Comprehensive Inorganic Chemistry (J.C. Bailar et al.) Page 593 (WSU Library QD 451.2 C64 V.1)

2. Occurrence and Extraction of Magnesium

Magnesium is the eighth most abundant element in the earth’s crust. It occurs naturally in a number of minerals such as dolomite CaCO3MgCO3 and magnesite MgCO3 and is the third most abundant element in seawater from which it is commonly extracted. The extraction from seawater relies on the fact that magnesium hydroxide is less soluble than calcium hydroxide, because the solubility of the salts of mononegative anions increases down the group. Either CaO (quick lime) or Ca(OH)2 (slaked lime) is added to the seawater and Mg(OH)2 precipitates. The hydroxide is converted to the chloride by treatment with hydrochloric acid.

The magnesium is then extracted by electrolysis of molten magnesium chloride.

Source: Inorganic Chemistry fourth edition (Shriver & Atkins) page 276 (Stoker Library)

Mg was formerly prepared by heating MgO and C to 2000 °C, at which temperature C reduces MgO. The gaseous mixture of Mg and CO was then cooled very rapidly to deposit the metal. This 'quenching' or 'shock-cooling' was necessary as the reaction is reversible, and if cooled slowly the reaction will come to equilibrium further to the left. MgO + C ↔ Mg + CO Magnesium is now produced by high temperature reduction, and by electrolysis.

Source: Concise Inorganic Chemistry fifth edition (J.D. Lee) page 328 (WSU Library QD 453.2.L44 1996)

3. Uses of Magnesium

Magnesium is the only Group 2 metal to be produced on a large scale. World production was 303000 tones in 1993. The largest producers were the USA 48%, the Soviet Union 12%, Canada 9% and Norway 8%. Mg is an extremely important lightweight structural metal because of its low density (1.74gcm-3 compared with steel 7.8gcm-3 or aluminum 2.7 g cm-3). Mg forms many binary alloys, often containing up to 9% AI, 3% Zn and 1 % Mn; traces of the lanthanides praseodymium Pr and neodymium Nd, and traces of thorium. The metal and its alloys can be cast, machined and welded quite easily. It is used to make the bodies of aircraft, aircraft parts and motor car engines. Up to 5% Mg is usually added to Al to improve its properties.

Source: Concise Inorganic Chemistry fifth edition page 328 (WSU Library QD 453.2.L44 1996)

Perhaps the best know magnesium compound is magnesium sulfate (MgSO4). It is popularly known as Epson salts.

One of the earliest stories about Epsom salts dates back to 1618. The town of Epsom, in Surrey, England, was suffering from a severe drought. A farmer named Henry Wicker brought his cattle to drink from a water hole on the town commons (central park). But the cattle would not drink the water. Wicker was surprised because he knew they were very thirsty. He tasted the water himself and found that it was very bitter. The bitterness was due to magnesium sulfate in the water. This compound became known as Epsom salts.

People soon learned that soaking in the natural waters that contained Epsom salts made them feel better. The salts seemed to have properties that soothed the body. Before long, soaking in these waters became very popular.

Source: Chemical Elements (David E Newton) Pages 309-310 (WSU Reference QD 466.N464 1999V.2)

4. General Properties of Magnesium

Magnesium is the easiest of all structural metals to machine (108). Because of this machine ability, it is sometimes used in applications where a large number of machining operations are required. The machine ability of magnesium alloys relative to that of other metals based on the lowest power required to remove 16 cm3 of metal and when magnesium is assigned a value of unity, is (108): Some of the advantages of excellent machine ability include reduced machining time, resulting in higher productivity for the machine tools and thus lower capital investment; greatly increased tool life; an excellent surface finish with a single large cut; well-broken chips which minimize handling costs; and less tool buildup.

If a coolant or cutting fluid must be used for a particular operation, mineral oil is usually preferred over water-based coolants. This is due to the fact that magnesium reacts with water to some degree, over time, producing hydrogen gas and magnesium hydroxide. As a result, wet chips should be treated with caution in order to prevent ignition of the evolved hydrogen, and to prevent partial drying of the wet mass, which may result in spontaneous ignition due to the heat evolved in the reaction in combination with poor heat-transfer characteristics of the partially dried mass. If wet chips are generated, the chips should be kept fully submerged in excess water.

Source: Kirk-Othmer Encyclopedia Of Chemical Technology fifth Edition volume 15 page 368 (WSU Reference TP9.K54 2004 V.15)

5. Magnesium Alloys

Aluminum alloys are the largest single consumer of magnesium. In 1992 over 50 percent of the magnesium shipped was consumed in this way. The aluminum beverage can contains about 4.5 percent magnesium in the lid and 1.1 percent in the can body. When combined with ferrosilicon, magnesium produces ductile iron. Automotive engines are a major consumer of ductile iron.

Source: MacMillian Encyclopedia of Chemistry (Lagowski) Page 879 (WSU Reference QD 4.M33 1997 V.3)

Magnesium has become important as a sttuctural material. Its great advantage is that it is very light, with a density of only 1.74 grams/cm3. For comparison's sake, water has a density of 1 gram/cm3. Magnesium has a density that is only about one-fifth that of iron and two-thirds that of aluminum. It is usually mixed with these metals to form an alloy. When alloyed with aluminum, magne- sium makes the latter metal stronger, lighter, and even more corrosion-resistant than it normally is. Many people, for example, have aluminum-magnesium alloy ladders in their homes. Its light weight also makes it ideal for fabricating automobile and aircraft parts, as well as power tools, lawn mower housings, and racing bikes.

Source: A Guide to the Elements second edition (Stwertka) Pages 57-58

6. Magnesium-Oxygen Compounds

In a close packing of oxide ions, the cations are inserted where, in accordance with their size, room is to be found. The inclusion of different cations thus changes the dimensions of the lattice of the oxide ions but little. The spaces in the latter are of two kinds, octahedral and tetrahedral. The octahedral spaces are the larger, so that the larger cations find room in them. The octahedral space provides room for the Mg2+ … ions. In order that such may enter a silicate lattice, larger spaces must be present between the closely packed oxide ions. These may be created either by an open arrangement in places, as in the feldspars or by layer structures.

Source: Structural Chemisry of inorganic compounds (Huckel) Page 772 (WSU Library QD 151 H833 V.2)

Magnesium Oxide, MgO, also known as magnesia, occurs in nature only infrequently as the mineral periclase, most commonly as a group of crystals in marble. Refractory magnesia represents the largest tonnage use of magnesium in compounds. The iron and steel industry is the largest consumer of these products in the United States and most other magnesia-consuming countries. Dead-burned magnesia from magnesite, seawater, or well and lake brines is used as a principal constituent in metallurgical furnace refractory products.

Source: Kirk-Othmer Encyclopedia of Chemical Technology fifth Edition volume 15 page 413 (WSU Reference TP9.K54 2004 V.15)

7. Magnesium-Halide Compounds

Chlorides Magnesium chloride is among the industrially most important of the magnesium salts. It is marketed in the anhydrous and hexahydrated forms. Both are deliquescent and form saturated solutions if left standing in a moist atmosphere.

Anhydrous magnesium chloride can be prepared by the direct chlorination of magnesium oxide, but it is usually prepared from a hydrate. The dehydration cannot be simply accomplished with heat because decomposition to the oxychloride occurs. After all the water has been removed, the compound is melted at 712°.

Approaching half a million tons of MgCl₂ is produced in the United States each year. Aside from its use as a source of the metal, it is used in oxychloride cements and wall plasters. It is also used in fire-proofing and dust-laying compositions.

Magnesium bromide is found in sea water, some mineral springs, the Dead Sea, the Utah Salt Lake and in deposits such as those at Strassfurt. At temperatures down to O°C, it crystallizes from aqueous solution as the hexahydrate which is very hygroscopic.

Magnesium iodide chemistry is very similar to that of the bromide.

Source: Comprehensive Inorganic Chemistry (J.C. Bailar et al.) Page 629-635 (WSU Library QD 451.2 C64 V.1)

Dicyclopentadienylmagnesium (magnesium cyclopentadienide), (C₅H₅)₂Mg, is a pyrophoric white fluffy solid that can be made directly from magnesium metal and cyclopentadiene at elevated temperatures: 500°C

Mg + 2C₅H₆ ↔ (C₅H₅)₂Mg + H₂

Dicyclopentadienylmagnesium is a useful reagent for the preparation of cycloIpentadienyl transition metal derivatives.

Source: Inorganic Chemistry of Main Group Elements (R. Bruce King) Page 271 (WSU Library QD 151.2.K5 1995 c2)

8. Magnesium in the Synthesis of Organic Compounds

[Grignard reagents] a very large and important group of alkyl and aryl magnesium halides together with the magnesium dialkyls and diaryls. There are two binary carbides of magnesium, which cannot be got pure enough for analysis, but from their behavior must have the compositions MgC2 and Mg2Ca. Neither can be made by heating magnesium or its oxide with carbon to a high temperature, and both are thermodynamically unstable above 800°.

In most cases the original magnesium compound is not isolated, but the solution immediately used for the synthesis, the second organic component, usually dissolved in ether or benzene, being added as soon as the magnesium has gone into solution. The reactivity essentially depends on the fact that magnesium has a much weaker affinity for carbon than for oxygen, nitro- gen, or the halogens, so that it very readily exchanges its hydrocarbon radical for one of the other elements. The number of reactions which the Grignard reagents can undergo is enormous.

Source: Chemical Elements and Their Compounds (Sidgwick) pages 224-225 (WSU Library QD 466 S53 v.1)

9. Magnesium in Life

An adult body contains about 25 g of magnesium, 60% of which is present in the skeleton while most of the remainder resides inside cells where it is the next most important cation after potassium. The daily requirement in the human diet is not known with certainty but has been estimated to be 200-300 mg.

All enzymes that utilize ATP in phosphate transfer, and many which aid either transfer of other groups or hydrolysis, require magnesium as the cofactor. For example, in phosphorylation reactions Mg²⁺ ion is thought to complex with oxygen on the various phosphate groups. As it is the smallest, doubly-charged cation which is biologically available, magnesium will normally form the strongest complexes of this type.

Source: Main Group Chemistry (A.G. Massey) page 142 (WSU Library QD 151.2.M37 2000)

Magnesium also plays a crucial role in the makeup of the green chlorophylls present in all green plant cells. The ability of the chlorophyll to capture solar energy and convert it by photosynthesis to energy is the ultimate source of all biological energy. The structure of the chlorophyll molecule, a ring of complex atomic structures around the magnesium atom, endows the molecule with its deep color and its ability to absorb light.

Source: A Guide to the Elements second edition (Stwertka) Pages 58

10. Magnesium and the Treatment of Waste Water

The largest use for magnesium hydroxide in the United States is for environmental applications. This portion of the market, which includes industrial water treatment, heavy-metals removal, and fluegas desulfurization, accounts for about more than half of total U.S. consumption.

For water treatment, magnesium hydroxide is supplied as a suspension containing about 58% solids, and it is used primarily to lower the pH of acidic solutions. In this market it competes with other acid-neutralizing compounds; the most common are lime and caustic soda. Magnesium hydroxide has advantages and disadvantages when compared to the other materials in this use. One of the advantages is that it is a pH buffer, and wastewater treated with it will not exceed a pH of 9.5 even if excess magnesium hydroxide is added. In contrast, excess addition of lime can raise the pH to 12,

Source: Kirk-Othmer Encyclopedia of Chemical Technology fifth Edition volume 15 page 406 (WSU Reference TP9.K54 2004 V.15)

Silica may be adsorbed on solid magnesium oxide or hydroxide, or on crushed dolomite. This is not a stoichiometric reaction involving the whole solid phase, but occurs to a variable extent on the surface only. The extent of adsorption is adequately expressed by the Freundlich adsorytion isotherm*: mass SiO2 adsorbed. This procedure can be used in conjunction with lime-soda softening since this results in a precipitate of Mg(OH)2. Effective adsorption of silica onto the Mg(OH)2 requires a temperature of at least 50 °C and preferably 100-110°C (the hot lime-soda process), and it may be necessary to add MgSO4 to get enough Mg(OH)2.

Source: Inorganic Chemistry and Industrial and Environmental Perspective (T.W. Swaddle) pages 274-275 (WSU Library QD 151.5.S93 1997)