Selenium

History

(Gr. Selene: moon) Discovered by Berzelius in 1817, who found it associated with tellurium (named for the earth).[1]

Properties

General
Name : selenium
Symbol : Se
Atomic Number : 34
Chemical Series : Other Nonmetal
Block, Period : 16, 4
Appearance : grey, metallic lustre
Atomic Properties
Atomic Weight (amu) : 78.96
Covalent Radius (pm) : 116
Physical Properties
Matter : solid (diamagnetic)
Density (kg/) : 4790
Hardness : 2
Melting Point (K) : 494
Boiling Point (K) : 957.8
Evaporation Heat (kJ/mol) : 26.3
Fusion Heat (kJ/mol) : 6.694
Specific Heat (J/(kg*K) ) : 320
Miscellaneous
Electrical Conductivity (/m ohm) :
Thermal Conductivity (W/(m*K) ) : 2.04

Selenium exists in several allotropic forms, although three are generally recognized. Selenium can be prepared with either an amorphous or a crystalline structure. The color of amorphous selenium is either red (in powder form) or black (in vitreous form). Crystalline monoclinic selenium is a deep red; crystalline hexagonal selenium, the most stable variety, is a metallic gray.

Selenium exhibits both photovoltaic action, where light is converted directly into electricity, and photoconductive action, where the electrical resistance decreases with increased illumination. These properties make selenium useful in the production of photocells and exposure meters for photographic use, as well as solar cells. Selenium is also able to convert a.c. electricity to d.c., and is extensively used in rectifiers. Below its melting point, selenium is a p-type semiconductor and has many uses in electronic and solid-state applications.

Elemental selenium has been said to be practically nontoxic and is considered to be an essential trace element; however, hydrogen selenide and other selenium compounds are extremely toxic, and resemble arsenic in their physiological reactions.[1]

Production

Selenium is found in a few rare minerals such as crooksite and clausthalite. In years past it has been obtained from flue dusts remaining from processing copper sulfide ores, but the anode metal from electrolytic copper refineries now provide the source of most of the world's selenium. Selenium is recovered by roasting the mud with soda or sulfuric acid, or by smelting them with soda and niter.[1]

Isotopes

Naturally selenium contains six stable isotopes. Fifteen other isotopes have been characterized. The element is a member of the sulfur family and resembles sulfur both in its various forms and in its compounds.[1]

Uses

Selenium is used in Xerography for reproducing and copying documents, letters, etc. It is used by the glass industry to decolorize glass and to make ruby-colored glasses and enamels. It is also used as a photographic toner, and as an additive to stainless steel.[1]

Handling

Hydrogen selenide at a concentration of 1.5 ppm is intolerable to man. Selenium occurs in some solid in amounts sufficient to produce serious effects on animals feeding on plants, such as locoweed, grown in such soils. Exposure to selenium compounds (as Se) in air should not exceed 0.2 mg/m3 (8-hour time-weighted average - 40-hour week).[1]

Notes

[1] From Los Alamos National Laboratory's Chemistry Division Website