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Chemistry of Selenium (Z=34)

Element number 34, selenium, was discovered by Swedish chemist Jons Jacob Berzelius in 1817. Selenium is a non-metal and can be compared chemically to its other non-metal counterparts found in Group 16: The Oxygen Family, such as sulfur and tellurium.


Chemical Symbol: Se
Atomic Number: 34
Atomic Weight: 78.96
Electron Configuration: [Ar] 4s23d104p4
Melting Point: 493.65 K
Boiling Point: 958 K
Electronegativity: 2.55 (Pauling)
Oxidation States: Se-2, Se+6, Se+4
Ionization Energies: First: 941 kJ/mol
  Second: 2045 kJ/mol
  Third: 2973.7 kJ/mol


Selenium was discovered by Berzelius in 1818. It is named for the Greek for "moon", selene. The discovery of selenium was an important finding, but at the same time seemingly accidental. Fellow scientist Martin Klaproth, discovered a contamination of sulfuric acid creating a red colored product which he believed to be due to the element tellurium. However, Berzelius went on to further analyze the impurity and came to the conclusion that it was an unknown element that shared similar properties to those of tellurium. Based off of the Greek word “selene,” meaning moon, Jons Berzelius decided to call the newly found element selenium.

Allotropes and Physical Properties

Selenium can exist in multiple allotropes that are essentially different molecular forms of an element with varying physical properties. For example, one allotrope of selenium can be seen as an amphorous (“without crystalline shape”) red powder.  Selenium also takes a crystalline hexagonal structure, forming a metallic gray allotrope which is known to be stable.  The most thermodynamically stable allotrope of selenium is trigonal selenium and also appears as a gray. Most selenium is recovered from the electrolytic copper refining process. This is usually in the form of the red allotrope.

Selenium is mostly noted for its important chemical properties, especially those dealing with electricity. Unlike sulfur, selenium is a semiconductor, meaning that it conducts some electricity, but not as well as conductors. Selenium is a photoconductor, which means it has the ability to change light energy into electrical energy. Not only is selenium able to convert light energy into electrical energy, but it also displays the property of photoconductivity.  Photoconductivity is the idea that the electrical conductivity of selenium increases due to the presence of light or in other words, it becomes a better photoconductor as light intensity increases. 


Isotopes of an element are atoms that have the same atomic numbers but a different number of neutrons (different mass numbers) in their nuclei.  Selenium is known to have over 20 different isotopes; however, only 5 of them are stable. The five stable isotopes of selenium are 74Se, 76Se, 77Se, 78Se, 80Se.


Due to selenium’s property of photoconductivity, it is known to be used in photocells, exposure meters in photography, and also in solar cells. Selenium can also be seen in its production in plain-paper photocopiers, laser printers and photographic toners. Besides its uses in the electronic industry, selenium is also popular in the glass-making industry.  When selenium is added to glass, it is able to negate the color of other elements found in the glass and essentially decolorizes it. Selenium is also able to create a ruby-red colored glass when added. The element can also be used in the production of alloys and is an additive to stainless steel.

Health Hazards

Selenium, a trace element, is important in the diet and health of both plants and animals, but can be only taken in very small amounts.  Exposure to an excess amount of selenium is known to be toxic and cause health problems.  With a tolerable upper intake level of 400 micrograms per day, too much selenium can lead to selenosis and may result in health problems and even death.  Compounds of selenium are also known to be carcinogenic.

Chemical Reactivity

Reaction with hydrogen

Selenium forms hydrogen selenide, H2Se, a colorless flammable gas when reacted with hydrogen.

Reaction with oxygen

Selenium burns in air displaying a blue flame and forms solid selenium dioxide.

\[Se_{8(s)} + 8O_{2(g)} \rightarrow 8SeO_{2(s)}\]

Selenium is also known to form selenium trioxide, SeO3.

Reaction with halides

Selenium reacts with fluorine, F2, and burns to form the selenium hexafluoride.

\[Se_{8(s)} + 24F_{2(g)} \rightarrow 8SeF_{6(l)}\]

Selenium also reacts with chlorine and bromine to form diselenium dichloride, \(Se_2Cl_2\) and diselenium dibromide, \(Se_2Br_2\).

\[Se_8 + 4Cl_2 \rightarrow 4Se_2Cl_{2(l)}\]

\[Se_8 + 4Br_2 \rightarrow 4Se_2Br_{2(l)}\]

Selenium also forms \(SeF_4\), \(SeCl_2\) and \(SeCl_4\),


Selenium reacts with metals to form selenides. Example: Aluminum selenide

\[3 Se_8 + 16 Al \rightarrow 8 Al_2Se_3\]


Selenium reacts to form salts called selenites, e.g., silver selenite (Ag2SeO3) and sodium selenite (Na2SeO3)


  1. Describe selenium’s property of photoconductivity.
  2. Does selenium react with hydrogen? If so what compound is produced?
  3. Describe selenium’s purpose as a trace element.
  4. What are some common uses for selenium?
  5. Does selenium react with oxygen?


  1. Selenium’s ability to change light energy into electrical energy increases as light intensity increases
  2. Yes, selenium reacts with hydrogen and forms hydrogen selenide H2Se
  3. Selenium is important in the health of plants and animals, but is only safe in small amounts.  Too much selenium can be toxic and cause serious health problems.
  4. Selenium is used in the glass-making industry and also in electronics.  It is used in photo cells, solar cells, photocopiers, laser printers and also photographic toners.
  5. Selenium burns in air and forms selenium dioxide. It is also able to form selenium trioxide.


  1. Minaev, V. S., S. P. Timoshenkov, and V. V. Kalugin. "Structural and Phase Transformations in Condensed Selenium."  Journal of Optoelectronics and Advanced Materials, volume 7, number 4, 2005, pp. 1717–1741.
  2. Mary Elvira Weeks and Henry M. Leicester. Discovery of the Elements, 7th edition. Easton, PA: Journal of Chemical Education, 1968.
  3. Petrucci, Ralph H. General Chemistry. 9th ed. Upper Saddle River: Prentice Hall, 2007


  • David Jin (UCD)