Magnesium (Mg)

The atomic number of magnesium is 12, and atomic number is defined by the number of protons in the nucleus. That is why every magnesium atom has 12 protons: if the nucleus had a different proton count, it would be a different element.

A neutral atom has equal numbers of positive protons and negative electrons, so neutral Mg has 12 electrons. Isotopes keep the same 12 protons but contain different numbers of neutrons. The common stable isotopes are ²⁴Mg with 12 neutrons, ²⁵Mg with 13, and ²⁶Mg with 14.

The ground-state electron configuration ends in 3s². Those two outer electrons are the electrons most directly involved when magnesium forms ordinary chemical bonds.

Removing those two electrons produces Mg²⁺, leaving a 10-electron arrangement equivalent in electron count to neon. This helps explain why the +2 oxidation state is so characteristic of magnesium and why magnesium belongs to Group 2, the alkaline-earth metals.

The first shell contains 2 electrons, the second contains 8, and the third contains the remaining 2, giving the familiar shell pattern 2-8-2. In orbital notation, the 1s, 2s and 2p subshells are filled before the two outer electrons occupy 3s.

The shorthand [Ne] represents the filled 1s² 2s² 2p⁶ core. Writing [Ne] 3s² makes the chemically important outer electrons easy to see.

Magnesium starts with two relatively outer 3s electrons. Removing the first electron gives Mg⁺, but one 3s electron is still left. Removing the second completes the loss of the outer 3s shell and leaves the closed-shell electron configuration [Ne].

Removing a third electron would require taking an electron from that stable inner-shell core and therefore demands much more energy. That sharp change is one reason the +2 state dominates ordinary magnesium chemistry.

Magnesium is reactive enough that native metallic magnesium is not the normal form found in nature. Instead, it occurs in minerals such as dolomite, magnesite, brucite, carnallite and olivine, and large quantities of magnesium ions are dissolved in seawater and natural brines.

This is why commercial magnesium production begins with a magnesium-bearing mineral, salt solution or brine and then uses chemical separation followed by electrolysis or thermal-reduction processes to obtain the metal.

Its position is closely connected to its electron structure. Period 3 tells us that occupied electron shells extend through the third principal shell, while Group 2 reflects the two outer 3s electrons.

Magnesium therefore shares important patterns with beryllium above it and calcium below it, while comparisons across Period 3 with sodium and aluminium help students see trends such as atomic radius and ionization energy.

Metallic magnesium has a low density, so magnesium alloys are useful where reducing mass matters, including selected transport and engineering components. Magnesium also plays roles in metallurgical processing and alloy manufacture.

Its intense white light when burning explains its association with flares and pyrotechnic applications. Magnesium compounds have very different uses: magnesium oxide is important in heat-resistant materials, while other salts and compounds are used across chemical, industrial and laboratory applications.

When magnesium burns in oxygen, it forms magnesium oxide in a strongly exothermic reaction. The high temperature excites atoms, ions and particles in the flame, and the hot products emit a broad, intense visible light that appears brilliant white to the eye.

The brightness is precisely why burning magnesium became useful historically in photographic flash powders and remains memorable in chemistry demonstrations. It should be observed only under appropriate laboratory safety controls because burning magnesium is extremely bright and hot.