Magnesium (Mg)

Magnesium has the configuration \([Ne]3s^2\). It readily loses two 3s electrons to form the stable cation \(\mathrm{Mg^{2+}}\), characteristic of Group 2 metals. This ease of oxidation explains its vigorous reactions under suitable conditions and its strong reducing ability.

When heated strongly, Mg reacts rapidly with oxygen, forming magnesium oxide and releasing substantial heat and visible radiation, perceived as a brilliant white flame:

\(\mathrm{2\,Mg(s) + O_2(g) \rightarrow 2\,MgO(s)}\)

Some Mg also reacts with atmospheric nitrogen at high temperature to form magnesium nitride:

\(\mathrm{3\,Mg + N_2 \rightarrow Mg_3N_2}\)

Metallic Mg is protected by a thin, adherent oxide film that slows reaction with cold water. In hot water or steam, reaction is faster, producing magnesium hydroxide and hydrogen:

\(\mathrm{Mg(s) + 2\,H_2O(l) \rightarrow Mg(OH)_2(s) + H_2(g)}\)

With steam, MgO is often formed:

\(\mathrm{Mg + H_2O(g) \rightarrow MgO + H_2}\)

Magnesium dissolves readily in dilute acids, liberating hydrogen gas. For example, with hydrochloric acid:

\(\mathrm{Mg(s) + 2\,HCl(aq) \rightarrow MgCl_2(aq) + H_2(g)}\)

This is a standard lab method to generate hydrogen.

Mg forms lightweight, high specific-strength alloys such as AZ31, AZ91 (Mg–Al–Zn systems) and rare-earth–containing alloys for improved creep resistance. Applications include aerospace components, automotive parts (wheels, housings), electronics casings, and sports equipment, where weight reduction improves efficiency.

Mg²⁺ is the central ion in chlorophyll, enabling photosynthesis. In animals, Mg²⁺ stabilizes ATP and acts as a cofactor for many enzymes, contributing to nerve and muscle function and nucleic acid stability.

Two major routes are used:

• Electrolytic (Dow-type): Electrolysis of molten anhydrous MgCl₂ (often from seawater/brine via precipitation of Mg(OH)₂ then conversion to MgCl₂).
• Thermochemical (Pidgeon process): Silicothermic reduction of calcined dolomite at high temperature under vacuum: \(\mathrm{2\,MgO\cdot CaO + Si \rightarrow 2\,Mg(g) + Ca_2SiO_4}\), followed by condensation of Mg.

Grignard reagents are organomagnesium halides formed by reacting alkyl/aryl halides with Mg in dry ether:

\(\mathrm{R\!\!\;X + Mg \xrightarrow[\text{ether}]{} R\!\!\;MgX}\)

They act as powerful nucleophiles/bases for C–C bond formation (e.g., addition to carbonyls), making Mg indispensable in organic synthesis.

Mg develops a protective MgO layer that inhibits ignition. Light abrasion removes the oxide, exposing fresh metal that ignites more easily and burns completely to MgO:

\(\mathrm{2\,Mg + O_2 \rightarrow 2\,MgO}\)

Do not use water or CO₂ extinguishers—both can intensify the fire (forming H₂ with water or freeing O₂ from CO₂). Use a Class D (dry powder) extinguisher or cover with dry sand. Avoid direct viewing of intense light without protection.