Selenium
| Selenium | |
|---|---|
| |
| Atomic number | 34 |
| Symbol | Se |
| Group | 16 (chalcogens) |
| Boiling point | 685 °C |
| Electron configuration | [Ar] 3d10 4s2 4p4 |
| Melting point | 221 °C |
| Period | 4 |
| Cas number | 7782-49-2 |
| Phase STP | Solid |
| Block | p |
| Oxidation states | −2, +4, +6 |
| Wikidata | Q876 |
Selenium (symbol Se, atomic number 34) is a nonmetallic element in the chalcogen (oxygen) family. It sits in period 4 of the periodic table, between sulfur and tellurium, and shares many chemical properties with them. Selenium occurs as a solid at standard conditions and has several allotropes (elemental forms) that range from metallic-gray to red or black. Common oxidation states of selenium are –2, +4 and +6; it can also exhibit +2 in some compounds. Typical compounds include selenides (analogous to sulfides), oxides (such as selenium dioxide), halides (like SeF₆), and acids (selenous and selenic acids). Atoms of selenium have the electron configuration [Ar] 3d¹⁰4s²4p⁴, giving six valence electrons (4s²4p⁴). Because of its intermediate electronegativity (~2.55 on the Pauling scale) and ionic size (~120 picometres covalent radius), selenium has properties between sulfur and tellurium. It is a semiconductor and shows notable photoconductivity, meaning its electrical conductivity increases when illuminated.
Atomic Structure and Properties
A selenium atom has 34 protons and (in the neutral atom) 34 electrons. Its electrons fill the argon-like core Ar, then the 3d and 4s orbitals, leaving four electrons in the 4p subshell. Thus selenium’s valence shell (n=4) has six electrons (4s²4p⁴), typical of Group 16 elements. The filled 3d subshell has little effect on chemical behavior except through shielding. Selenium’s atomic radius is larger than sulfur’s but smaller than tellurium’s. For example, its covalent radius is about 120 pm (about 1.20 Å). Because of this larger size down the group, its first ionization energy (the energy to remove one electron) is lower than sulfur’s; selenium’s first ionization energy is about 941 kJ/mol (9.75 eV), compared to 1000 kJ/mol for sulfur. Electronegativity also decreases down the group: sulfur’s is 2.58, selenium’s about 2.55, and tellurium’s about 2.10. Thus selenium is moderately electronegative, capable of forming covalent bonds but also accepting electrons from very electropositive metals to form selenide (Se²⁻) ions.
Selenium shows typical periodic trends: its atomic and ionic sizes are larger than sulfur’s; it holds valence electrons less tightly than sulfur (so it is more easily oxidized) but still not as readily as tellurium. It is in the p-block of the periodic table (its valence electrons occupy p orbitals). Unlike the noble gas argon, the p⁴ configuration of selenium gives it four unpaired or paired electrons available for bonding. Generally, it behaves as a nonmetal, though it is more metallic in nature than sulfur. For instance, elemental gray selenium conducts electricity (semiconducting behavior), whereas sulfur is a poor conductor.
Isotopes and Nuclear Properties
Naturally occurring selenium is a mix of several isotopes. Five isotopes are stable: ^74Se, ^76Se, ^77Se, ^78Se and ^80Se, with ^80Se being the most abundant (almost 50% of natural Se). A sixth isotope, ^82Se, is not strictly stable but has an extremely long half-life (~9.2×10^19 years) by double-beta decay, so it is often treated as stable for practical purposes. Of the stable isotopes, ^77Se is notable because it has nuclear spin 1/2, making it useful in NMR spectroscopy studies of selenium compounds; the others have spin 0 (even-even nuclei).
Radioactive selenium isotopes are rare in nature but many can be made artificially. For example, selenium-75 (^75Se) is produced by neutron activation of ^74Se; it has a half-life of about 120 days and decays to ^75As by emitting gamma rays. ^75Se’s gamma rays make it useful as an industrial gamma source for radiography (inspecting welds, pipes, etc.) and in medical diagnostics or calibration. Another isotope, ^79Se, is formed in minute amounts in nuclear reactors or nuclear fuel reprocessing; it has a half-life of ~3.3×10^5 years and decays by beta emission to bromine-79. This long-lived radioisotope is mostly of interest in nuclear waste studies. Several other selenium radionuclides (ranging in mass 72–95) have much shorter half-lives (seconds to months) and are used in scientific research but have no common practical applications.
In nuclear physics, ^82Se has been a target of experiments searching for neutrinoless double-beta decay due to its long half-life. Selenium isotopes also have modest neutron-capture cross-sections and have been studied in NMR contexts (especially ^77Se). Overall, selenium’s stable isotopes and long-lived ^82Se mean it does not pose significant natural radioactivity issues, but ^75Se is one deliberately produced for its gamma emission.
Allotropes and Compounds
Elemental selenium exists in several allotropic forms. The most stable allotrope is gray (or trigonal) selenium, which has a helical chain structure of Se atoms and a metallic lustre. This gray form is a semiconductor with a band gap around 1.5–1.7 eV and shows photoconductivity (its electrical conductivity rises under light). Gray selenium crystals are brittle and insoluble in carbon disulfide (CS₂).
Other allotropes include red selenium (which itself has several crystalline forms) and black selenium. Red selenium is typically made by dissolving Se in some solvent and then slowly crystallizing. It involves puckered Se₈ ring molecules (much like sulfur’s S₈ rings). There are α-, β- and γ-red forms, all of which are poor conductors (insulators) and differ slightly in crystal symmetry. Red Se is less dense than gray and can appear as a brick-red powder. Black (vitreous) selenium is an amorphous form, usually produced by rapidly cooling molten selenium. It consists of random polymeric chains (an irregular long-chain network) and also is a poor conductor, but it soften at relatively low temperature (~50°C) and turns into gray selenium at a bit above 100°C. Black selenium is photoconductive (used historically in light-sensing applications before gray form was fully utilized). All these allotropes can interconvert: heating red or black Se gradually converts them to the gray form.
The various allotropes highlight selenium’s bonding: as a chalcogen, Se–Se bonds can form ring or chain molecules. In red forms, Se₈ rings (eight-membered) occur; in gray form, the structure is comprised of long helical chains of Se atoms.
Selenium forms typical compounds analogous to sulfur’s chemistry. In the –2 state, selenium forms selenide anions (Se²⁻). Many metal selenides are known, for example mercury selenide (HgSe), lead selenide (PbSe), zinc selenide (ZnSe) and copper-indium-gallium selenide (CIGS, Cu(In,Ga)Se₂). These selenides are important semiconductors and optical materials. For instance, ZnSe is transparent to infrared light and has been used in IR lasers and optics; PbSe and other lead chalcogenides are used in infrared detectors; CIGS is widely used in thin-film solar cells. When selenides of very electropositive metals are exposed to moisture or weak acids, they hydrolyze to form polyselenides and H₂Se. The compound H₂Se (hydrogen selenide) is Se’s analog of H₂S. It is a colorless, foul-smelling gas that is much more acidic than H₂S (Ka1 ≈ 3×10^–4, pH ~3 in water). H₂Se oxidizes in air to give SeO₂; in water it forms HSe⁻ and Se²⁻. H₂Se is extremely toxic.
Selenium’s oxyacids mirror sulfur’s: Selenous acid (H₂SeO₃) is produced by dissolving SeO₂ in water, analogous to sulfurous acid. The stable oxide SeO₂ is a white to colorless solid that sublimes and forms H₂SeO₃ on hydration. Selenium can be further oxidized to selenic acid (H₂SeO₄), analogous to sulfuric acid. Unlike sulfuric acid, selenic acid is a very strong oxidizer and can dissolve gold. It is often made by oxidizing SeO₂ with strong oxidants. Salts of selenous and selenic acids are called selenites (SeO₃²⁻) and selenates (SeO₄²⁻), respectively. Sodium selenite (Na₂SeO₃) and sodium selenate (Na₂SeO₄) are examples used in laboratories and industry.
Halogen compounds of selenium include selenium hexafluoride (SeF₆), a colorless gas analogous to sulfur hexafluoride, though SeF₆ is even more reactive and toxic. Selenium also forms lower halides: for instance, selenium monochloride “Se₂Cl₂” (sometimes written Se₂Cl₂) and selenium dibromide (Se₂Br₂). These are considered selenium(I) halides, structurally consisting of Se–Se bonded pairs with halogens. Selenium tetrachloride (SeCl₄) and monochloride (SeCl₂) can also exist under specialized conditions. In comparison to sulfur, selenium’s halides and oxyhalides are less well-known, but they demonstrate selenium’s ability to achieve +4 and +6 oxidation states in compounds.
In summary, selenium’s allotropes range from insulating molecular solids to semimetallic chain structures, and its compounds parallel those of sulfur with characteristic selenides, selenites/selenates, and selenic acid chemistry.
Physical Properties
At standard conditions, elemental selenium is a solid. Gray selenium (the stable form) has a metallic-gray luster and is brittle. Its density is relatively high: about 4.8 g/cm³ (similar to metals), whereas the less dense red forms have densities around 4.0 g/cm³. Selenium’s melting point is 221°C (494 K) and its boiling point is 685°C (958 K). These values are lower than those of sulfur (which boils at 444°C) because selenium’s denser structure and heavier atoms result in stronger intermolecular forces in the solid.
Gray selenium crystallizes in a hexagonal (trigonal) lattice consisting of long helical Se–Se chains; this anisotropic structure makes its electrical and thermal properties direction-dependent. At room temperature, gray Se is a semiconductor: its electrical resistivity is moderately high (lower conductivity than metals, but many orders of magnitude better than insulating red selenium). Its band gap is about 1.5–1.7 eV, which lies in the red/infrared region of the spectrum. Because of this band gap, selenium conducts more strongly when exposed to light (photoconductive behavior): photons create electron–hole pairs, increasing its conductivity. This photoconductivity is an unusual property for a pure element and underlies many of selenium’s electronic applications. For comparison, red and black (amorphous) allotropes are essentially insulators without appreciable electrical conductivity in the dark.
Thermal behavior: Selenium has relatively low thermal conductivity, similar to nonmetallic solids. It undergoes several allotropic transitions on heating. For example, gray selenium softens above 160°C and transitions to a rubbery state, and eventually melts at 220°C. Upon heating selenium vapor to higher temperature, the vapor pressure increases significantly, approaching 1 atm near its boiling point (686°C), where it boils to a reddish vapor. The molten and gaseous states conduct electricity only modestly, so for practical purposes selenium is treated as solid with negligible vapor pressure below its boiling point.
Spectroscopically, selenium atoms have characteristic optical emission and absorption lines in the visible and ultraviolet, but these are not widely used as elemental signature lines (they are overshadowed by more common elements). However, selenium does impart distinctive colors to its compounds and solutions. For example, solutions of selenous acid are colorless, but many insoluble selenides (with heavy metals) have deep red, brown or black colors. In colored glass, colloidal selenium produces an intense ruby red hue. In short, selenium’s observable physical traits include a metallic-gray solid form that is brittle and semiconducting, with characteristic red colors in its more polymeric or fine-particle forms.
Chemical Reactivity and Trends
Selenium’s chemistry shows parallels and contrasts with sulfur’s. Like sulfur, elemental selenium is relatively inert at room temperature but burns in air when heated. When heated in oxygen, selenium combusts with a blue flame to produce selenium dioxide:
<code>Se (s) + O₂ (g) → SeO₂ (s). </code>
Solid SeO₂ can sublime or dissolve in water to give selenous acid (H₂SeO₃). Further oxidation (for example by strong oxidizers such as ozone or hot nitric acid) can produce selenic acid (H₂SeO₄), the fully oxidized form (+6 oxidation state). In this respect, selenium resembles sulfur’s oxide chemistry, although SeO₃ is unstable and decomposes back to SeO₂ unless conditions are strongly oxidative.
In acid–base chemistry, selenium exhibits chalcogen behavior. Its hydride, H₂Se, is a strong acid (stronger than H₂S) and liberates H₂Se gas when many metal selenides are treated with strong acids:
<code>Na₂Se (aq) + 2 HCl → H₂Se (g) + 2 NaCl. </code>
Conversely, elemental Se reacts with alkali metals to form selenides (Se²⁻), similar to how sulfur forms sulfides. For example, aluminum reacts with heated selenium to give aluminum selenide (Al₂Se₃), which hydrolyzes to produce H₂Se. Selenium does not dissolve in non-oxidizing acids (e.g. HCl alone) nor in water.
Selenium’s oxidation states range from –2 through +6 in stable compounds. In aqueous solution, the selenite ion (SeO₃²⁻) and selenate ion (SeO₄²⁻) are analogous to sulfite and sulfate. Selenious acid (H₂SeO₃) is a weak acid (pKa ~4), whereas selenic acid (H₂SeO₄) is a strong acid and strong oxidizer (ability to, for example, oxidize gold). Selenium in the –2 state (selenides) is strongly reducing and analogous to hydrogen sulfide or sulfide anions; many selenide salts are readily oxidized by oxygen to sulfuric or selenic species plus elemental selenium.
In redox terms, elemental selenium is of moderate activity. It is more easily oxidized than tellurium but less so than sulfur. It is less electronegative than sulfur, which partly explains its greater variety of positive oxidation states. For example, unlike sulfur which seldom forms +2 compounds, selenium can adopt +2 in bridged compounds like Se₂Cl₂. However, like sulfur, selenium resists oxidation by non-oxidizing acids and does not react with water. With halogens, selenium can be oxidized; fluorine reacts vigorously to form SeF₆, and chlorine or bromine burn iodine-like to give Se₂Cl₂ or Se₂Br₂ (involves Se–Se bonded species). Selenium tetrachloride (SeCl₄) is also known, but it hydrolyzes and is less stable than the higher fluoride.
Complexation chemistry of selenium is less rich than for some transition metals, but it does form complexes with metal centers by acting as a ligand (for example, coordinating through Se in metal-selenium clusters). In general, the trends in chemical reactivity show that selenium is a nonmetal that readily shares electrons (covalent character) but can also form anions with very electropositive partners. It tends to be oxidized in strongly basic or strongly acidic, oxidizing conditions, and behaves as an acid (protolyzing to H₂Se) under basic or strong reducing conditions. Unlike sulfur, selenium resists oxidation to the +2 state under mild conditions; it often skips from 0 directly to +4 in common reactions.
In corrosion terms, gray selenium is quite inert: it does not tarnish in air at room temperature and is not attacked by HCl or H₂SO₄. Very strong oxidizers (like KMnO₄ or HNO₃) will oxidize it to selenite/selenate. In a reactivity series context, metallic selenium is less reactive than nonmetals but more reactive than most metals. It will displace less electropositive anions (like I⁻ from HI) under the right conditions, reflecting its intermediate position.
Occurrence and Production
Selenium is relatively scarce in the Earth’s crust (about 50 parts per billion by weight) and in the cosmos (tens of ppb in solar abundance). It is not mined as a primary ore because there are no major selenium-only deposits. Instead, selenium is found dispersed at trace levels in many sulfide mineral ores. The most important source is as an impurity in copper, lead, and nickel sulfide ores. In these minerals, selenium replaces sulfur in the lattice to a small extent. When such ores are refined (for example, electrolytic refining of copper), the selenium is left behind in the anode slimes or muds. These sludges are then processed to extract selenium as a byproduct.
In a typical production process, the anode mud is oxidized (for example, with nitric acid or by roasting) to convert the selenium into selenium dioxide (SeO₂), a volatile solid. The SeO₂ is condensed and purified. It is then reduced (commonly by sulfur dioxide gas or hydrogen, or by melting with powdered copper/iron) to yield elemental selenium. In older processes, contact sulfoxide reactor slag from sulfuric acid plants also contained selenium compounds; these were leached and processed before contact-converter methods replaced them. Today, copper refining dominates selenium production.
Global selenium production in the 2020s is on the order of a few thousand metric tons per year. Consumption figures show strong usage especially in industries such as glassmaking and metallurgy, so production typically matches demand. The leading producer and consumer by far is China. Other significant contributors include Japan, Germany, Belgium, South Korea and the United States, reflecting their large copper/nickel refining industries. Production often fluctuates with copper output: more copper production generally yields more selenium byproduct, and vice versa.
Selenium sometimes appears in nature as the pure element (so-called native selenium), but this is very rare and usually as small grains in ores. It also occurs in rare minerals such as clausthalite (PbSe) or tiemannite (HgSe), but these are not abundant enough to be mined.
Most selenium in the environment exists in oxidation states +4 or +6. In soils, selenium is often found as selenate (SeO₄²⁻, water-soluble sulfate analog) or as adsorbed or organoselenium. Selenium-rich soils exist in some regions (e.g. parts of the US Great Plains, certain coal deposits in China); these can transfer selenium into plants and the food chain. Sea water contains only about 0.45 ppb of selenium under normal conditions. Overall, selenium cycles through the lithosphere, hydrosphere, and biosphere largely following the sulfur cycle, with unique aspects of bioaccumulation and volatility.
Applications and Technology
Selenium’s unique electronic, optical, and chemical properties make it useful in a variety of applications, though many uses are specialized or have declined with newer materials. The largest single use (over 50% of selenium consumption) is in glass production. Small amounts of selenium compounds (selenites/selenates) are added to glass melts to neutralize or mask color. Iron impurities often give glass a green or yellow tint; colloidal selenium imparts a red color that cancels out the tint, effectively decolorizing the glass. Conversely, higher selenium content yields a bright ruby-red glass used for artistic and decorative effects. Selenium also absorbs ultraviolet light, so it is used in some UV-protective glass.
In electronic and photonic devices, selenium’s semiconducting and photoconductive nature was harnessed in the 20th century. Amorphous selenium (a-Se) thin films have been used in photodetectors, especially in flat-panel X-ray imaging: the Se layer converts incident X-rays directly into charge. Selenium was also key in early photocopy (xerographic) copiers and laser printers: the photoconductive drum was coated with selenium so that images could be written by light exposure. However, many of these applications have shifted to organic photoconductors or silicon-based devices over time.
Selenium rectifiers (solid-state diodes) were historically important: starting in the 1930s, stacks of thin selenium plates were used to convert alternating current to direct current in power supplies and radio transmitters. These were more efficient than earlier copper-oxide rectifiers. Selenium rectifiers remained in use for decades (in some high-voltage or regulator systems) until silicon diodes largely displaced them. One niche where selenium still appears is in DC surge protectors: selenium surge suppressors (which can absorb transient high voltages) outperform many competing components for surge currents, and are still used in some power systems.
A key photovoltaic use is in thin-film solar cells. Copper indium gallium diselenide (CIGS, Cu(In,Ga)Se₂) is an important absorber layer in some commercial solar modules, converting sunlight to electricity efficiently. (Cadmium telluride and silicon dominate the PV market, but CIGS cells have gained market share). Selenium has also been examined as an alternative in novel battery technologies: for example, a lithium–selenium battery has been proposed, analogous to lithium–sulfur batteries but with selenium as the cathode material, potentially offering high energy density with better conductivity.
In alloys and metallurgy, selenium is a minor additive. It improves the machinability of stainless steel and other metals when added at the level of 0.1–0.3%, similarly to sulfur or lead. Selenium-magnesium alloys were used to deoxidize metals. Granulation: Bismuth and selenium are added to copper and brass to replace toxic lead in plumbing alloys (“biobrass” or “lead-free brass”), since selenium imparts machinability without harming corrosion resistance. Selenium is also used in bearings, tellurium-copper alloys, and in special low-melting alloys.
Selenium finds smaller uses in chemistry and materials science. Organo-selenium compounds serve as catalysts in some chemical reactions (e.g. rubber vulcanization or polymer chemistry), though toxicity limits widespread use. Selenium sulfide (with composition around SeS₂) is an active ingredient in dandruff shampoos and antifungal treatments, exploiting selenium’s biocidal properties. Selenium is used in photographic toning: black-and-white prints can be toned with selenium to prevent fading. Historically, zinc selenide was used as a laser medium (for early blue diode lasers) and cadmium selenide was used in colored quantum dots for displays, though newer materials often replace them.
In small amounts, selenium is added to animal feeds and fertilizers (especially in selenium-poor regions) because it is an essential micronutrient for animals. This is not an “industrial” use per se but can influence selenium production/distribution in agriculture.
Overall, selenium’s technology roles span from glassmaking (the biggest use) to photoconductors (one-time large industry), to photovoltaics and alloying (growing niche), with a few specialty roles in imaging and the laboratory.
Biology, Environment, and Safety
Selenium is an essential trace element for many organisms but is toxic at higher concentrations. In biology, selenium’s importance comes from its role in certain enzymes: in animals and humans, selenium is incorporated as the amino acid selenocysteine in several critical proteins. These include glutathione peroxidases and thioredoxin reductases, which protect cells from oxidative damage, and iodothyronine deiodinases, which activate thyroid hormones. Such selenoenzymes are vital for antioxidant defenses, DNA synthesis, and thyroid function.
The recommended dietary intake of selenium for humans is on the order of 50–70 micrograms per day for adults (for example, 55 µg/day in the US RDA). Selenium is obtained from foods such as meat, seafood, eggs, and grains; the actual content depends strongly on soil selenium levels. Plants take up selenium from soil mostly as selenate; some plants (Brazil nuts, certain grains) can accumulate high levels. Deficiency in selenium can lead to health problems: in humans, very low selenium status has been linked to Keshan disease (a cardiomyopathy) and Kashin-Beck disease (a bone/joint disorder) in selenium-poor regions of China. However, selenium deficiency is uncommon elsewhere, as most diets provide enough (0.07 mg/kg body weight stores can be maintained).
Toxicity: Selenium has a narrow safe window. Chronic intake above ~300–400 µg/day in adults can cause selenosis, with symptoms like hair and nail loss or brittleness, garlic-like odor in breath, gastrointestinal upset, and neurological abnormalities. In the workplace, exposure to selenium dust or fumes is strictly limited — for example, many regulations set exposure limits around 0.2 mg/m³ for an 8-hour time-weighted average. Elemental selenium dust and compounds (especially selenides, selenates) can be irritating to the lungs if inhaled. Ingestion of selenium salts can be severely toxic. Handling selenium and its compounds requires precautions (good ventilation, protective equipment).
Environmentally, selenium can be both a nutrient and a pollutant. In soils and water, selenium is mobile when in the form of water-soluble selenate or selenite. Certain plants (e.g. some Astragalus species) are selenium-accumulators, storing high Se levels and volatilizing it into dimethyl selenide, a gaseous compound with a weak odor. Selenium can biomagnify in food chains, especially aquatic ones: algae absorb selenium which is passed to fish and then to birds. Excess selenium in water from irrigation runoff or industrial waste (from mining or coal plants) can cause wildlife toxicity. Famous cases include fish and bird deformities at Kesterson Reservoir (CA, 1980s) and fish kills in irrigation waters: concentrations on the order of micrograms per liter (µg/L) can be harmful (U.S. EPA aquatic criteria are around 5 µg/L for continuous exposure).
Nonetheless, in natural cycling low levels of selenium play roles in ecosystems. Microbes can convert selenium oxyanions into organic forms, and the element is recycled through soils, plants, and animals. Because both deficiency and excess of selenium are important issues, environmental regulations monitor selenium discharges from industry and agriculture. In summary, selenium is biologically essential but potentially toxic, requiring balance: its chemistry mirrors that of sulfur in biology (replacing S in amino acids), yet its handling demands caution due to toxicity at higher dose.
History and Etymology
Selenium was discovered in 1817 by Swedish chemist Jöns Jacob Berzelius (who is credited with naming it) and his assistant Johan Gottlieb Gahn. They found it while investigating residues from the lead-chamber process of sulfuric acid production (using pyrite, FeS₂, from the Falun mine). A red powder formed in their apparatus was initially mistaken for an arsenic compound. Berzelius noticed this powder emitted a distinct odor like horseradish when burned — a trait known for tellurium compounds. Realizing it could not be tellurium (since their ores lacked that element), Berzelius identified it as a new element. In 1817 he announced “selenium” (Greek selḗnē “Moon”), choosing a name to parallel tellurium (Latin tellus “Earth”), since the two elements are chemically similar. Thus, selenium’s name reflects its connection to the moon (echoing the way tellurium is linked to earth).
The original isolation of selenium was from chemical residues, but it was not isolated as the pure metal until later. In 1818 Berzelius published a second letter confirming selenium’s properties. Over the 19th century, chemists characterized its allotropes and compounds. Notably, in 1873 physicist Willoughby Smith discovered selenium’s photoconductivity by observing that its electrical resistance changed with light exposure – the first demonstration of photoconductivity in any material. This finding would later underpin selenium’s use in devices like photocells and photocopiers.
In the late 19th and early 20th centuries, selenium rectifiers were invented for converting AC to DC electricity. By 1933, selenium diodes made from stacks of thin selenium plates were commercially used for power rectification, replacing older metal-oxide rectifiers. This was a major development in electronics until silicon diodes largely supplanted selenium rectifiers in the latter 20th century.
Selenium entered biological and nutritional awareness in the 20th century when its role in enzymes was discovered. By mid-20th century, it was recognized as an essential dietary trace element for animals; awards like the 1998 Nobel Prize in Chemistry highlighted selenoenzyme functions. Meanwhile, selenium’s ability to add and remove color in glass was exploited in art and architecture. In recent decades, selenium’s role in high-efficiency solar cells and modern X-ray imaging has given it new technological relevance.
In popular culture and industry, selenium’s name has appeared in contexts ranging from “selenosis” (selenium poisoning) to brands (selenium sulfide shampoos). The element’s story — from a mistaken byproduct to an enabling material in electronics and health — illustrates the interplay of chemistry and human technology through history.
Data Table
| Property | Value |
|---|---|
| Symbol | Se |
| Atomic number (Z) | 34 |
| Atomic weight (average) | 78.96 |
| Element category | Nonmetal (chalcogen) |
| Group (periodic table) | 16 (oxygen family) |
| Period | 4 |
| Block | p-block |
| Valence shell configuration | 4s² 4p⁴ (total 6 valence electrons) |
| Electron configuration | [Ar] 3d¹⁰ 4s² 4p⁴ |
| Common oxidation states | –2, +2 (rare), +4, +6 |
| Standard state | Solid (at STP) |
| Phase at STP | Solid (gray crystalline form) |
| Density (20°C, gray Se) | ~4.8 g/cm³ |
| Melting point | 221 °C (494 K) |
| Boiling point | 685 °C (958 K) |
| Crystal structure | Hexagonal/trigonal (gray Se chains) |
| Electronegativity (Pauling) | ~2.55 |
| First ionization energy | ~9.75 eV (~941 kJ/mol) |
| Atomic/Covalent radius | ~115 pm (empirical) / 120 pm (covalent) |
| Electron affinity | ~195 kJ/mol |
| Isotopes (naturally occurring) | ⁷⁴Se, ⁷⁶Se, ⁷⁷Se, ⁷⁸Se, ⁸⁰Se (stable); ⁸²Se (very long-lived) |
| Number of stable isotopes | 5 (or 6 if ⁸²Se is counted) |
| Notable radioisotopes | ⁷⁵Se (t₁/₂=120 d, emits γ), ⁷⁹Se (t₁/₂≈3.3×10⁵ y) |
| Color (solid) | Gray (metallic, stable); red or black (amorphous forms) |
| Magnetic property | Diamagnetic |
| Electrical resistivity (gray Se) | ~1×10⁻⁵ Ω·cm (semi\-conducting) |
| Band gap (gray Se) | ~1.5–1.7 eV |
| Thermal conductivity (gray Se) | Low (∼0.46 W/m·K at 20°C) |