Selenium (Se)

Selenium has the electronic configuration \([Ar]3d^{10}\,4s^2\,4p^4\). With six valence electrons, it belongs to Group 16 (Chalcogens), along with oxygen, sulfur, tellurium, and polonium. It shows oxidation states of −2, +4, and +6 similar to sulfur.

Selenium exists in several allotropes:

• Amorphous red selenium: obtained by rapid cooling of molten Se.
• Crystalline gray selenium: stable metallic form, a semiconductor with hexagonal rings.
• Black and vitreous forms: intermediate, less stable varieties.

The gray form exhibits photoconductivity—its electrical conductivity increases with light intensity.

Gray selenium conducts electricity better when illuminated. Photons excite electrons from the valence to the conduction band:

\(\mathrm{Se + h\nu \rightarrow Se^* (e^- + h^+)}\)

This property is exploited in photoelectric cells, light meters, and solar rectifiers.

Selenium shows oxidation states −2, +4, and +6:

• −2: hydrogen selenide, \(\mathrm{H_2Se}\), a toxic gas analogous to H₂S.
• +4: selenium dioxide, \(\mathrm{SeO_2}\), used in organic oxidation.
• +6: selenic acid, \(\mathrm{H_2SeO_4}\), a strong oxidizing acid similar to sulfuric acid.

• Both form hydrides (H₂S, H₂Se), oxides, and oxyacids.
• Selenium is more metallic and heavier, with lower electronegativity.
• Se forms stronger Se–Se and Se–metal bonds, and exhibits semiconductivity unlike sulfur.

Thus, Se bridges nonmetallic S and more metallic Te in chemical behavior.

SeO₂ is a white, volatile, acidic oxide formed by burning selenium in air:

\(\mathrm{Se + O_2 \rightarrow SeO_2}\)

It dissolves in water to form selenious acid (H₂SeO₃) and is used in oxidizing organic compounds, especially converting aldehydes to acids.

Yes. Selenium is an essential micronutrient for humans and animals. It is present in the enzyme glutathione peroxidase, which protects cells from oxidative damage:

\(\mathrm{2\,GSH + H_2O_2 \xrightarrow{Se-enzyme} GSSG + 2\,H_2O}\)

However, in large doses, selenium compounds are toxic and can cause selenosis (hair loss, nail brittleness, nervous disorders).

• Electronics: Rectifiers, photocells, solar cells due to photoconductivity.
• Glassmaking: Decolorizing agent and red tints.
• Alloys: Improves machinability and corrosion resistance of metals.
• Pigments: Used in ceramics, plastics, and rubber for bright red coloration.

Formation of SeO₂:

\(\mathrm{Se + O_2 \rightarrow SeO_2}\)

Reduction of SeO₂ back to Se:

\(\mathrm{SeO_2 + 2\,H_2 \rightarrow Se + 2\,H_2O}\)

This redox pair demonstrates Se’s intermediate position between metallic and nonmetallic behavior.

While trace selenium is vital, excess is toxic. Inhalation of selenium dioxide fumes or ingestion of soluble selenites can cause nausea, garlic odor on breath (due to volatile Se compounds), and damage to liver and lungs. Handle Se compounds with gloves, ventilation, and avoid ingestion or prolonged exposure.