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Magnesium is a silver-shining light metal similar to aluminum, which quickly becomes coated with a thin but coherent oxide layer in the air. The metal below is protected from further oxidation despite its high affinity for oxygen.

Magnesium is ductile and can be rolled into sheets or drawn into wires. The electrical conductivity is about 2/3 that of aluminum and 1/3 that of copper.

The element crystallizes in a hexagonal close packing of spheres - a type of lattice that many other metals also prefer.


The first production of magnesium was achieved in 1809 by H. Davy electrolytically using the amalgam method (analogous to the amalgam method for sodium production). Large-scale production by fused carnallite, KCl · MgCl2, did not begin in Germany until 1886. The name magnesium is derived from the town of Magnesia in Asia Minor (now Turkey).


Above 500 ° C Magnesium ignites in the air and burns with a dazzling white flame to form magnesium oxide (and magnesium nitride, Mg3N2):



This reaction was used in the early days of photography when using flash powder (a mixture of magnesium powder with oxidizing agents such as potassium chlorate, potassium permanganate, manganese dioxide, etc.).

Magnesium is a powerful reducing agent that reacts with very stable oxides at higher temperatures:


In addition to oxygen and nitrogen, magnesium will readily react with most other non-metals. This is how it ignites in halogens (formation of halidesMgX2) and supplies with hydrogen magnesium hydride, MgH2.

With boiling water, the metal forms magnesium hydroxide with evolution of hydrogen, Mg(OH)2, around:


This reaction soon comes to a standstill in cold water, as the hydroxide formed on the surface protects the metal underneath (passivation).

The formation of a protective layer is prevented by adding a mercury salt, as magnesium amalgam forms on the metal surface, which disrupts the lattice of the magnesium atoms and increases the readiness to react:


Magnesium dissolves easily in dilute acids with evolution of hydrogen:


In terms of its reaction behavior, magnesium occupies a middle position between beryllium on the one hand and the - very similar - elements calcium to radium on the other. The different properties of beryllium and magnesium, as in the case of the element pair lithium / sodium, are due to the increase in the radius and the decrease in the electronegativity of the metals in their compounds at the transition from Be toMg traced back.

In contrast to the corresponding beryllium compounds, magnesium salts are heteropolar, i.e. they can essentially be described by bonds with a strongly ionic character. In aqueous solution they are subject to electrolytic dissociation. Im hydrated[Mg(H2O)6]2+-Ion, the water of the first coordination sphere is much less firmly bound in solution than in [Be(H2O)4]2+-Ion. This is one reason why magnesium salts react practically neutral in water and do not hydrolyze (beryllium salts show an acidic reaction as a result of hydrolysis). However, the hydration energy is still high enough that magnesium salts can crystallize out of water as stable hydrates.

Structural Patterns and Dimensionality in Magnesium Borophosphates: The Crystal Structures of Mg2(H2O) [BP3O9(OH)4] and Mg (H2O)2[B.2P.2O8(OH)2]·H2O

Hydrothermal investigations in the system MgO / B2O3/ P2O5(/H2O) yielded two new magnesium borophosphates, Mg2(H2O) [BP3O9(OH)4] and Mg (H2O)2[B.2P.2O8(OH)2]·H2O. The crystal structures were solved by means of single crystal X-ray diffraction. While the acentric crystal structure of Mg2(H2O) [BP3O9(OH)4] (orthorhombic, P.212121 (No. 19), a = 709.44 (5) pm, b = 859.70 (4) pm, c = 1635.1 (1) pm, V = 997.3 (3) × 10 6 pm 3, Z = 4) contains 1D infinite chains of magnesium coordination octahedra interconnected by a borophosphate tetramer, Mg (H.2O)2[B.2P.2O8(OH)2]·H2O (monoclinic, P.21/ c (No. 14), a = 776.04 (5) pm, b = 1464.26 (9) pm, c = 824.10 (4) pm, β = 90.25 (1) °, V = 936.44 (9) × 10 6 pm 3,Z = 4) represents the first layered borophosphate with 6 3 net topology. The structures are discussed and classified in terms of structural systematics.

The underestimated problem of using serum magnesium measurements to exclude magnesium deficiency in adults a health warning is needed for "normal" results

Background: A major use of serum magnesium measurements in clinical practice is to identify patients with deficiency. However, numerous studies have shown that magnesium deficiency is common and may be present in over 10% of hospitalized patients, as well as in the general population. An important cause for under diagnosis of deficiency is that serum magnesium, the most commonly used test, can be normal despite negative body stores. This article focuses on the limitations of "normal" magnesium results and highlights the importance of lifestyle or "modus vivendi" as a pragmatic means of identifying those individuals potentially at risk for negative body magnesium stores.

Methods: Researched peer reviewed articles on magnesium published between 1990 and 2008 in MEDLINE and EMBASE, using database keywords "magnesium, deficiency, diagnosis, treatment and hypomagnesaemia". Bibliographies of retrieved articles have been searched and followed. We have also performed a manual search of each individual issue in which most of these reports have appeared.

Results: In 183 peer reviewed studies published from 1990 to 2008, magnesium deficiency was associated with increased prevalence and risk in 11 major conditions. Similarly, in 68 studies performed over the same period, magnesium deficiency was found to predict adverse events and a decreased risk of pathology was noted when supplementation or treatment was instituted.

Conclusions: The perception that "normal" serum magnesium excludes deficiency is common among clinicians. This perception is probably enforced by the common laboratory practice of highlighting only abnormal results. A health warning is therefore warranted regarding potential misuse of "normal" serum magnesium because restoration of magnesium stores in deficient patients is simple, tolerable, inexpensive and can be clinically beneficial.


The best way to detect magnesium is by means of Magneson II, titanium yellow or quinalizarin.

For detection with Magneson® (4- (4-Nitrophenylazo) -1-naphthol) the original substance is dissolved in water and made alkaline. Then a few drops of a solution of the azo dye Magneson® are added. If magnesium ions are present, a dark blue colored varnish is created. Other alkaline earth metals should first be removed as carbonates by precipitation.

For detection with titanium yellow (thiazole yellow & # 160G), the original substance is dissolved in water and acidified. Then it is mixed with a drop of the titanium yellow solution and made alkaline with dilute sodium hydroxide solution. If magnesium is present, a light red precipitate is formed. Nickel, zinc, manganese and cobalt ions interfere with this detection and should be precipitated as sulphides beforehand.

For detection with quinalizarin, two drops of the dye solution are added to the acidic sample solution. Then dilute sodium hydroxide solution is added until the reaction is basic. A blue color or precipitation indicates magnesium.

The formation of precipitates with phosphate salt solutions can also be used as a detection reaction for magnesium salts. The heavy metal-free sample solution, buffered to pH 8 to 9 with ammonia and ammonium chloride, is mixed with disodium hydrogen phosphate solution. A white, acid-soluble cloudiness from magnesium ammonium phosphate MgNH4PO4 indicates magnesium ions:

$ mathrm + NH_4 ^ + + PO_4 ^ <3-> rightarrow MgNH_4PO_4 downarrow> $

From ammoniacal solution, Mg 2+ can also be detected with oxine as a poorly soluble yellow-greenish compound. This proof is suitable for the cation separation process.