Magnesium (Mg, Z = 12): A Light Metal with Remarkable Properties

Model of the magnesium atom — Simplified atomic model of the magnesium atom. The most abundant isotope is \,^{24}\mathrm{Mg}\, with its 12 protons and 12 neutrons.
Image source: astronoo.com (new window) — AI-generated image, public domain.

History of the Discovery of Magnesium

Magnesium takes its name from the Magnesia region in Greece, where a white mineral called magnesia (magnesium oxide) was found in abundance. In 1755, Joseph Black (1728-1799) recognized magnesia as a substance distinct from lime. In 1808, Humphry Davy (1778-1829) succeeded in isolating metallic magnesium by electrolyzing a mixture of magnesium oxide and mercury oxide. However, it was Antoine Bussy (1794-1882) who, in 1831, produced pure metallic magnesium by reducing magnesium chloride with potassium.

Structure and Fundamental Properties

Magnesium (symbol Mg, atomic number 12) is an alkaline earth metal located in the second column of the periodic table. Its atom has 12 protons, 12 electrons, and usually 12 neutrons in its most abundant isotope (\(\,^{24}\mathrm{Mg}\)). Other stable isotopes exist: magnesium-25 (\(\,^{25}\mathrm{Mg}\)) and magnesium-26 (\(\,^{26}\mathrm{Mg}\)).
At room temperature, magnesium is a solid, silvery-white, lightweight metal (density ≈ 1.738 g/cm³), malleable, and a good conductor of heat and electricity. The melting point of magnesium: 923 K (650 °C). The boiling point: 1,363 K (1,090 °C). Magnesium oxidizes easily in contact with air, forming a thin protective layer of magnesium oxide.

Table of Magnesium Isotopes

Magnesium isotopes (key physical properties)
Isotope / Notation	Protons (Z)	Neutrons (N)	Atomic mass (u)	Natural abundance	Half-life / Stability	Decay / Remarks
Magnesium-24 — \(\,^{24}\mathrm{Mg}\,\)	12	12	23.985042 u	≈ 78.99 %	Stable	Most abundant isotope of natural magnesium.
Magnesium-25 — \(\,^{25}\mathrm{Mg}\)	12	13	24.985837 u	≈ 10.00 %	Stable	Second stable isotope; used in isotopic research.
Magnesium-26 — \(\,^{26}\mathrm{Mg}\)	12	14	25.982593 u	≈ 11.01 %	Stable	Third stable isotope; decay product of aluminum-26.
Magnesium-28 — \(\,^{28}\mathrm{Mg}\)	12	16	27.983877 u	Not natural	20.915 hours	Radioactive β\(^-\) decaying into aluminum-28. Used in nuclear research.
Other isotopes — \(\,^{20}\mathrm{Mg}\) to \(\,^{40}\mathrm{Mg}\)	12	8 — 28	— (variable)	Not natural	Milliseconds to minutes	Unstable isotopes produced artificially; used in nuclear physics.

Electronic Configuration and Electron Shells of Magnesium

N.B.:
Electron shells: How electrons organize around the nucleus (new window).

Magnesium has 12 electrons distributed across three electron shells. Its full electronic configuration is: 1s² 2s² 2p⁶ 3s², or simplified as: [Ne] 3s². This configuration can also be written as: K(2) L(8) M(2).

Detailed Structure of the Shells

K Shell (n=1): contains 2 electrons in the 1s subshell. This inner shell is complete and highly stable.
L Shell (n=2): contains 8 electrons distributed as 2s² 2p⁶. This shell is also complete, forming a noble gas configuration (neon).
M Shell (n=3): contains 2 electrons in the 3s subshell. The 3p and 3d orbitals remain empty. These two valence electrons are relatively easily lost during chemical reactions.

Valence Electrons and Oxidation States

The 2 electrons in the outer shell (3s²) are the valence electrons of magnesium. This configuration explains its chemical properties:
By losing its 2 electrons in the 3s subshell, magnesium forms the Mg²⁺ ion (oxidation state +2), its unique and systematic oxidation state in all its compounds.
The Mg²⁺ ion then adopts an electronic configuration identical to that of neon [Ne], a noble gas, which gives this ion great stability.
Magnesium exhibits no other stable oxidation state; only the +2 state is observed in chemistry.

The electronic configuration of magnesium, with its valence shell containing 2 electrons in the 3s subshell, classifies it among the alkaline earth metals (Group 2 of the periodic table). This structure gives it characteristic properties: significant chemical reactivity (it oxidizes in air and reacts with water, especially when hot), exclusive formation of ionic compounds with an oxidation state of +2, and the ability to form metallic bonds in its crystalline structure. Magnesium spontaneously forms a thin layer of magnesium oxide (MgO) in the air, which slows down further oxidation, although this protection is less effective than that of aluminum. Its tendency to lose its valence electrons makes magnesium a good reducing agent. Its importance is considerable in both biology and industry: magnesium is essential for the functioning of living cells (enzyme cofactor, stabilization of DNA and RNA, chlorophyll in plants). In industry, it is used to produce lightweight, high-strength alloys (particularly with aluminum) for aeronautics and automobiles, as a reducing agent in metallurgy, and in the manufacture of fireworks due to its intense combustion producing a bright white light.

Chemical Reactivity

Magnesium is a moderately reactive metal. It burns with an intense white flame in the presence of oxygen, producing magnesium oxide (MgO). It reacts slowly with cold water but vigorously with hot water or steam, releasing dihydrogen (H₂). Magnesium forms ionic compounds with non-metals and can act as a reducing agent in many chemical reactions. Its main compounds include magnesium chloride (MgCl₂), magnesium sulfate (MgSO₄), magnesium carbonate (MgCO₃), and magnesium hydroxide (Mg(OH)₂).

Industrial and Technological Applications of Magnesium

in lightweight alloys for aeronautics, automotive, and electronics (laptops, smartphones);
as a deoxidizing and desulfurizing agent in steel and cast iron metallurgy;
in fireworks and distress signals due to its bright combustion;
to produce titanium and other metals by chemical reduction;
as a sacrificial anode to protect metal structures from corrosion;
in the production of refractory cement resistant to high temperatures;
in traditional photography as a source of intense light (magnesium flash);
in the chemical industry to synthesize organic compounds via Grignard reagents;
as a dietary supplement and medicine (citrate, oxide, or magnesium sulfate);
in construction materials, especially lightweight and fireproof panels;
to manufacture tools and mechanical parts requiring lightness and strength;
in agriculture, to enrich magnesium-poor soils and improve plant growth;
in rechargeable magnesium batteries, potential alternatives to lithium-ion batteries;
as a hydrogen storage material in emerging energy technologies.

Essential Biological Role

Magnesium is the fourth most abundant cation in the human body and plays a crucial role in more than 300 enzymatic reactions. It is involved in protein synthesis, nerve transmission, muscle contraction, blood sugar regulation, and energy production (ATP). In plants, magnesium is at the heart of the chlorophyll molecule, essential for photosynthesis. A magnesium deficiency can lead to fatigue, muscle cramps, heart disorders, and in plants, yellowing of the leaves (chlorosis).

Abundance and Natural Sources

Magnesium is the eighth most abundant element in the Earth's crust (about 2.3% by mass) and the third most abundant dissolved element in seawater. It is mainly found in minerals such as dolomite (CaMg(CO₃)₂), magnesite (MgCO₃), carnallite (KMgCl₃·6H₂O), and olivine ((Mg,Fe)₂SiO₄). Industrial extraction is mainly done by electrolysis of molten magnesium chloride or by thermal reduction of magnesium oxide.

Role in Astrophysics

Magnesium is synthesized in massive stars during the fusion of oxygen and carbon. During supernova explosions, magnesium is dispersed into the interstellar medium, contributing to the chemical enrichment of subsequent generations of stars and planets. Its relative abundance in the universe and its presence in meteorites make it an important tracer of galactic chemical evolution. Astronomers use magnesium spectral lines to study the composition of distant stars and galaxies.

N.B.:
Metallic magnesium can be difficult to ignite in bulk form, but once lit, it burns with such intensity that it is almost impossible to extinguish with water. Indeed, at high temperatures, magnesium reacts with water by extracting oxygen from H₂O molecules, which further fuels the combustion. This property makes magnesium a formidable material in case of fire, requiring the use of special sands or powders to smother it.