Arsenic (symbol As, atomic number 33) is a dense, brittle metalloid in group 15 of the periodic table (the pnictogens). At ordinary conditions it is a dark gray solid (often called “gray arsenic”) with a metallic luster, though it has multiple allotropes including soft yellow and black forms. Arsenic’s atomic weight is about 74.92, and its electron configuration is [Ar] 3d¹⁰4s²4p³ (five valence electrons). It commonly exhibits oxidation states of −3, +3 and +5 (with +3 and +5 found in many compounds, and −3 in metal arsenides). Arsenic is famous for its toxicity – virtually all of its compounds are poisonous – and has a long history as a poison. In modern times it finds use in small amounts in alloys (for example, lead-arsenic battery plates or ammunition), and as the group-15 component of compound semiconductors such as gallium arsenide (GaAs).

As a metalloid, arsenic has properties intermediate between metals and nonmetals. It is a semimetal or poor semiconductor: the stable gray form has a rhombohedral (layered) crystal structure and is electrically conductive but brittle. Arsenic typically sublimes at about 614 °C at atmospheric pressure (turning directly into vapor) and melts at about 817 °C under high pressure. Its density is around 5.72 g/cm³, notably heavier than aluminum or silicon but lighter than many metals. The chemical behavior of arsenic resembles that of its neighbors phosphorus and antimony, but it is notably less reactive with oxygen (forming a protective oxide surface) and more prone to exist in multiple oxidation states. In natural waters and biological systems, arsenic usually occurs as inorganic arsenite or arsenate oxyanions (see the environment section below).

Atomic Structure and Electron Configuration

Arsenic
Atomic number	33
Symbol	As
Group	15 (pnictogens)
Electronegativity	2.18 (Pauling)
Electron configuration	[Ar] 3d10 4s2 4p3
Density	5.73 g/cm^3
Period	4
Cas number	7440-38-2
Phase STP	Solid
Block	p
Oxidation states	−3, +3, +5
Wikidata	Q871

Arsenic has atomic number 33, meaning 33 protons in its nucleus and, in the neutral atom, 33 electrons. Its ground-state electronic configuration is [Ar] 4s² 3d¹⁰ 4p³, which can also be written as 1s²2s²2p⁶3s²3p⁶3d¹⁰4s²4p³. The five electrons in the outermost shell (4s²4p³) make arsenic a group-15 element (like nitrogen and phosphorus), with a lone electron pair and three unpaired electrons in its valence shell.

Arsenic’s atomic and ionic radii reflect its position down the group: larger than phosphorus but smaller than antimony. The covalent radius is about 120 pm. Its first ionization energy (the energy to remove one electron) is about 9.79 eV (945 kJ/mol), which is lower than phosphorus (10.49 eV) and higher than antimony (8.64 eV). In general, arsenic’s ionization energy and electronegativity (Pauling scale 2.18) lie between those of phosphorus and antimony, reflecting the periodic trend of decreasing bond energy and attraction for electrons down the group. As a result, arsenic forms stable covalent bonds but can also exhibit ionic character in metal arsenides.

Because all naturally occurring arsenic is the isotope ⁷⁵As, the atomic and nuclear properties of that isotope are important. The nucleus of ⁷⁵As has a spin of 3/2, which means arsenic is NMR-active (often studied by ⁷⁵As NMR in solid-state chemistry). The magnetic moment of ⁷⁵As is about +1.44 nuclear magnetons. In spectroscopy, the arsenic atom (As I) shows prominent lines around 193 nm and other wavelengths in the ultraviolet. Electron affinity of arsenic is relatively small (about 0.8 eV), reflecting that the As– anion is only weakly stable, consistent with arsenic being not very electronegative compared to many nonmetals.

Isotopes and Nuclear Properties

Elemental arsenic found in nature consists of a single stable isotope, arsenic-75 (⁷⁵As), with a natural abundance of 100%. No other long-lived stable isotope exists. There are several synthetic radioisotopes of arsenic, ranging in mass number roughly from 60 to 90, all of which are short-lived (days or less) and are produced in nuclear reactors or accelerators. The longest-lived radioisotopes and their half-lives include ⁷³As (half-life ~80.3 days) and ⁷⁴As (17.8 days). Arsenic-74 decays by a mix of positron emission, electron capture, and beta decay, making it occasionally useful in research. Other radioisotopes like ⁷²As (26.0 h), ⁷⁶As (26.3 h), and ⁷⁷As (38.8 h) have been used in tracer studies or medical applications. For example, ⁷⁴As has been used in nuclear medicine to image brain tumors (exploiting its positron emission), and ⁷²As has been explored as a positron-emitter for PET scans.

Because all arsenic on Earth is ⁷⁵As with nuclear spin 3/2, naturally occurring arsenic has a nonzero nuclear magnetic moment. This makes arsenic atoms detectable by nuclear-magnetic resonance (NMR) and Mössbauer spectroscopy (especially the gamma resonance of ⁷⁵As at 426 keV). In analytical chemistry and environmental studies, radioactive arsenic isotopes have been used as tracers to track arsenic movement and speciation. However, routine applications are limited due to the radiation hazard and the short half-lives. The CFL isotope ⁷⁴As has also occasionally been used as a calibration source for Ge detectors due to its multiple gamma-ray lines.

Allotropes and Typical Compounds

Arsenic exhibits several allotropes (different structural forms of the element). The most important allotrope industrially is gray arsenic, which has a layered rhombohedral structure similar to black phosphorus but with different bond angles. This gray form has a metallic sheen but is brittle. It is the thermodynamically most stable form at room temperature. Yellow arsenic is another allotrope (As₄ tetrahedral molecules, analogous to white phosphorus P₄) – it is soft and waxy with a lemon-yellow color. Yellow arsenic is metastable and extremely toxic; it slowly transforms to gray arsenic in light. There is also a black arsenic allotrope with layered chain structures (somewhat like black phosphorus); black arsenic is a poor electrical conductor and brittle. Both yellow and black forms are produced by rapid cooling of arsenic vapor under controlled conditions, but they rapidly convert to gray arsenic at ambient conditions, so they are mainly of laboratory interest.

Elemental arsenic does not react with most acids or alkalis, but it is oxidized by strong oxidizers (air or acids) to yield arsenic oxides. In moist air it tarnishes, forming a protective layer of arsenic oxide (As₂O₃) which slows further corrosion. When heated in air it burns with a blue flame giving a characteristic garlic-like odor (from As₂O₃ fumes). Arsenic does not dissolve in non-oxidizing acids (e.g. hydrochloric acid), but it reacts with nitric acid to form arsenic acid or with concentrated sulfuric acid to form arsenic trioxide. Metal arsenides (M–As compounds) are often stable intermetallics; for example, iron arsenide (FeAs) and nickel arsenides were known since the 19th century and many more are used today in electronics or as catalysts. True ionic As³⁻ (arsenide) only occurs in certain solids like sodium arsenide (Na₃As), and even these have more covalent character.

The typical oxidation state +3 compounds include arsenic trioxide (As₂O₃, often called “white arsenic”), arsenates of lower oxidation states, and arsenites. Arsenic trioxide (As₂O₃) is one of the best-known arsenic compounds: it is a white crystalline solid that is highly toxic. It can be made by roasting arsenopyrite. Arsenous acid (H₃AsO₃) is the acid dissolving As₂O₃ in water, beyond which arsenite salts (MAsO₂ or M₃AsO₃) can form with bases. Another example is arsine (AsH₃), the hydride of arsenic: a highly toxic, flammable gas analogous to phosphine. Arsenic trichloride (AsCl₃) and triflouride (AsF₃) are volatile As(III) halides; arsenic(III) also forms bromide and iodide (AsBr₃, AsI₃) under certain conditions.

In the +5 oxidation state, arsenic forms arsenic pentoxide (As₂O₅) and arsenic acid (H₃AsO₄). Arsenic pentoxide is a white powder that dissolves in water to give arsenic acid (an oxyacid), H₃AsO₄ (analogous to phosphoric acid). Arsenic acid is a strong oxidizer and corrosive acid, although not used commercially as much as phosphoric acid. Pentavalent arsenic also forms pentachloride (AsCl₅) and pentafluoride (AsF₅). Notably, AsF₅ is a volatile, reactive gas, while AsCl₅ is unstable and disproportionates to AsCl₃ and Cl₂ unless cooled.

Arsenic also forms a variety of organometallic and organic compounds. Some organic arsenic compounds (such as arsenobetaine and arsenocholine found in seafood) are relatively non-toxic, whereas other organoarsenicals (like trimethylarsine or arsinic acids) can be very toxic. Historically, organic arsenic compounds like cacodyl (tetramethyldiarsine) were among the first synthesized organometallic compounds. Arsenic also famously forms pigments: orpiment (As₂S₃, orange-yellow) and realgar (As₄S₄, red) are natural mineral pigments. Mink plants (like Pteris ferns) and some bacteria can biochemically convert inorganic arsenic to organoarsenic species.

Typical bonding patterns in arsenic chemistry often involve covalent networks or discrete covalent molecules, due to the presence of the lone pair on As(III) and As(V). For example, in As(III) compounds the arsenic atom is typically three-coordinate (pyramidal) because the lone pair pretends to a lone pair. In arsenate (AsO₄³⁻) the geometry is tetrahedral, analogous to phosphate. Arsenic can also catenate (bond to itself), as seen in polyarsenides and in the As₄ tetrahedron (yellow As). However, extended As–As single-bond networks are less common than for phosphorus or antimony because arsenic prefers three bonds plus lone pair in the +3 state. Zintl phases containing As clusters (Arsenic chain or cage anions) are found in salts with alkali or alkaline earth metals.

Physical Properties

The most stable form of arsenic (gray arsenic) is a brittle, silvery-gray solid. Its density is about 5.72 g/cm³ at 20 °C, making it heavier than many metals (aluminum 2.7 g/cm³, copper 8.96 g/cm³) but lighter than others (iron 7.87 g/cm³). Gray arsenic has a metallic luster but a Mohs hardness only around 3–4 (so it can be scratched by a knife). The structure consists of layers of puckered six-membered rings; the layers are held together by weak van der Waals forces, which explains the low hardness and easy cleavage (similar to graphite’s layered structure).

Arsenic has a high melting point under pressure but sublimes readily at atmospheric pressure. At 1 atmosphere, it does not have a normal liquid phase: it sublimates at about 614 °C (887 K). Under higher pressure, the melting point is about 817 °C (1090 K) at 3.63 MPa (this defines the triple point). The boiling point (transition to gas below 1 atm) is also around 615 °C if measured as the sublimation temperature. The enthalpy of sublimation is relatively large (around 90 kJ/mol for the solid to gas transition at the triple point pressure), reflecting the strong bonding in the layers.

Arsenic is a semimetal (semiconductor) in its gray form. It has relatively high electrical resistivity compared to true metals: at 20 °C its electrical conductivity is about 3.45×10^4 S/m (resistivity ~ 2.9×10^−5 Ω·m), which is between typical metals and insulators. When glassy or amorphous, arsenic behaves more like a semiconductor with a small band gap (around 1.2–1.4 eV). Thermal conductivity for gray arsenic is moderate – on the order of 50 W/(m·K) (less than good metals like copper (~400 W/m·K) but higher than typical insulators). Arsenic does not have useful magnetic ordering; it is essentially diamagnetic (no unpaired electrons in the core aside from As³⁺ complexes).

Spectroscopically, elemental arsenic vapor at high temperature emits lines in the ultraviolet and visible. In the near-infrared (used in fiber optics), gallium arsenide (GaAs) crystals are famous for their photoluminescent properties. In solid-state physics, the energy band structure of GaAs (with arsenic) has a direct band gap of about 1.43 eV at room temperature, which is why GaAs is widely used in LEDs and laser diodes (emitting infrared and red light). (Pure arsenic itself in solid or vapor form is not typically used directly for light emission, but it is the group-15 participant in these compound semiconductors.)

Arsenic metal also exhibits the unusual phenomenon of amphoteric oxide behavior. Arsenic in the +3 state (arsenous oxide, As₂O₃) reacts with strong acids to give arsenate and with strong bases to give arsenite salts, whereas As₂O₃ does not dissolve in water. This kind of chemical behavior resembles metalloids or amphoteric behavior (like Al₂O₃).

Chemical Reactivity and Trends

Chemically, arsenic is quite reactive when heated or in the presence of strong reagents, but it is relatively unreactive in mild conditions. In air at room temperature, gray arsenic slowly oxidizes on the surface, forming As₂O₃. In moist air it develops a bronze tarnish (arsenic oxide film). At higher temperatures (e.g. flames or hot metal surfaces), arsenic burns to give a mixture of As₂O₃ (arsenic(III) oxide) and As₂O₅ (arsenic(V) oxide). The former predominates in most conditions; As₂O₅ can form if oxygen is plentiful and temperature is suitable. Both oxides dissolve in water to give arsenous acid (H₃AsO₃) and arsenic acid (H₃AsO₄) respectively, making arsenic an “acidic” element by oxide behavior.

Arsenic does not react with non-oxidizing acids or bases under normal conditions. For example, it will not dissolve in hydrochloric acid or in alkali solutions alone. However, oxidizing acids do react: fuming nitric acid or aqua regia will oxidize arsenic to dissolved arsenic acid. Concentrated sulfuric acid dehydrates arsenic compounds to As₂O₃. With base, solid arsenic is inert, but As₂O₃ is amphoteric: it dissolves in sodium hydroxide to form sodium arsenite (NaAsO₂) or in potassium hydroxide to give potassium hydrogen arsenite (KH₂AsO₃). Thus arsenic and its oxides behave somewhat like aluminum or silicon dioxide, being amphoteric.

Arsenic reacts readily with halogens when heated. Chlorine gas attacks arsenic to produce AsCl₃ and AsCl₅ (arsenic trichloride and pentachloride). Bromine and iodine also make AsBr₃ and AsI₃. Fluorine gives arsenic pentafluoride (AsF₅) and arsenic trifluoride (AsF₃). These halides hydrolyze in water (e.g. AsCl₃ gives arsenous acid and HCl). Arsenic does not react with krypton or fluorine at room temperature, but with hot fluorine it yields AsF₅.

With hydrogen, arsenic forms arsine (AsH₃), a highly toxic gas similar to ammonia or phosphine. Arsine is difficult to make by normal chemistry, but in contact with strong acid and certain metals, trace amounts of AsH₃ can form (e.g. reduction of As₂O₃ by nascent hydrogen). Arsine decomposes at modest temperatures to arsenic and hydrogen gas.

Arsenic forms arsenides with many metals, especially with more electropositive metals (alkali, alkaline earth, some transition metals). For example, sodium arsenide (Na₃As) and magnesium arsenide (Mg₃As₂) can form with very active metals. With transition metals, one often gets intermetallic phases like FeAs, Ni₂As, Cu₃As, etc., which are typically non-stoichiometric solids (more like metal alloys) rather than discrete ionic compounds. These arsenides tend to conduct electricity and have metallic luster; they do not contain free As³⁻ ions. In the context of reactivity, arsenic’s –3 state (arsenide) is less stable than the analogous phosphide or antimonide, making some arsenides sensitive to hydrolysis or oxidation.

Arsenic’s redox behavior is important in environmental chemistry. In water, arsenic exists primarily as oxyanions of either arsenite [AsO₃]³⁻ (As(III)) or arsenate [AsO₄]³⁻ (As(V)). Arsenite species (actually uncharged H₃AsO₃ at neutral pH) are relatively more mobile and toxic than arsenates. Arsenate (As(V)) in solution exists as H₂AsO₄⁻ or HAsO₄²⁻ depending on pH, resembling phosphate chemistry. Generally, in oxygen-rich (oxidizing) environments arsenic is present as As(V), while in reducing or anaerobic conditions it is reduced to As(III). The conversion between As(III) and As(V) can be slow chemically, but bacteria can catalyze these transformations. This speciation controls arsenic’s solubility: As(V) tends to adsorb strongly to metal oxides and clays, whereas As(III) is more soluble and harder to remove.

In terms of acidity/basicity, inorganic arsenic compounds show interesting behavior. Arsenious acid H₃AsO₃ (from As(III)) is a weak acid (pKₐ₁ ≈ 9.22), whereas arsenic acid H₃AsO₄ (from As(V)) is a moderately strong acid (pKₐ₁ ≈ 2.2, similar to phosphoric acid). Arsenites are basic relative to arsenic acid (Arsenic(III) oxide dissolves in strong base to form AsO₃³⁻ anion). Arsenates behave analogously to phosphates: ArSO₄²⁻ salts (where R can be hydrogen or metal) are stable, and arsenic acid can be neutralized by bases to give arsenate salts.

Another trend: stability of oxidation states down the group. Like antimony, arsenic favors the +3 state over +5 due to the inert pair effect (the s-electrons resist participation). Thus As(V) compounds (like As₂O₅) are stronger oxidizers and less stable; they tend to decompose or be reduced to As(III). For example, arsenic acid can oxidize iodide or thiosulfate. By contrast, arsenic in the –3 state (arsenide) is a strong reducing agent (arsenides tend to be oxidized by acids or oxygen to arsenic metal or arsenic oxides).

In corrosion terms, gray arsenic in air quickly forms an oxide layer that passivates it – much like aluminum. It is not attacked by plain water or weak acids. In metal alloys, a small amount of arsenic (e.g. in bronzes or solders) can improve hardness and casting properties, but arsenic can also embrittle some metals.

Reactivity series context: If we rank elements by tendency to oxidize, elemental arsenic sits between phosphorus (more reactive) and antimony (less reactive). Unlike alkali/alkaline earth metals, arsenic does not react with water. It does react with molten metals to form intermetallics, and it reacts vigorously with halogens or strong oxidizers. In practice, uncontrolled reactions with arsenic are generally limited by its toxicity – it is handled as a dangerous material in industry, not used for energetic chemistry.

Occurrence and Production

On Earth, arsenic is relatively common for a heavy element. It is the 53rd most abundant element in Earth’s crust, with an average concentration of about 1.5 parts per million (ppm) by weight (0.00015%). This is roughly similar to mercury or nickel. In the oceans, arsenic is also quite widespread: dissolved arsenic in seawater averages about 1.5 micrograms per liter (μg/L), making it the 22nd most abundant element in seawater. In the atmosphere, arsenic concentrations are very low (a few nanograms per m³). Arsenic tends to accumulate in soil (~5–100 mg/kg, typically around 10 mg/kg) and in plants (up to hundreds of μg/kg). Background levels in rivers are usually below 10 μg/L, although in some mineral-rich areas they can be much higher.

Arsenic is mostly found in nature bound up in minerals, often with sulfur or metals. The most common arsenic mineral is arsenopyrite (FeAsS), an iron-arsenic sulfide. This and related M–As–S minerals (M = Fe, Ni, Co) are the principal commercial sources of arsenic. Other important minerals include realgar (As₄S₄) and orpiment (As₂S₃), which are arsenic sulfides; and arsenolite (As₂O₃), a weathering product of arsenic. Native (elemental) arsenic occurs only rarely in nature. These minerals are often found in ore deposits of other metals; for example, arsenic commonly occurs with gold, copper, antimony and lead ores. Geologically, arsenic is often mobilized into groundwater by chemical reactions, which is why arsenic contamination of well water is a global issue.

World production of elemental arsenic (usually recovered as arsenic trioxide, As₂O₃, the “white arsenic” lumps) comes primarily as a byproduct of metal smelting and refining, rather than by mining arsenic itself. Copper smelting and gold mining are major sources of arsenic-bearing waste and flue dust. Historically in the 19th and early 20th centuries, arsenic was produced by roasting arsenopyrite and condensing the fumes. In recent decades, most arsenic comes from processing high-arsenic copper concentrates: the arsenic is roaster off-gassed and collected as As₂O₃, which can then be purified.

The leading arsenic producers today are China (which accounts for a large majority of the world supply, on the order of 50–70%), Morocco, Russia, and others like Peru and Belgium (where older smelters once recovered arsenic). For example, around 2014 Peru produced on the order of 25,000 metric tons of arsenic trioxide, China about 8,800 tons, with Morocco and Russia producing smaller amounts. (Exact figures shift year to year; USGS data and industry surveys track these quantities.) Many arsenic refineries in Europe and North America closed by the 2000s for environmental reasons, so most new arsenic now comes from developing countries. Recovered arsenic is sold as As₂O₃ or sometimes as metal (gray arsenic) and then converted to compounds.

In the environment, natural sources of arsenic include volcanic gases (volcanoes emit arsenic-containing gases), geothermal springs, and erosion of arsenic-bearing rocks. Anthropogenic (man-made) sources include smelting, coal burning (coal can contain a few ppm of arsenic, which is released to the fly ash and atmosphere), use of arsenic-containing pesticides and herbicides (still in use in parts of the world), and even semiconductor manufacturing (GaAs chip fabrication emits arsenic vapors if not controlled). The global arsenic cycle involves atmospheric transport (mainly attached to dust particles), deposition into soil and water, and incorporation into living organisms.

Applications and Technology

Arsenic’s most important high-tech use today is in the field of electronics and photonics as the group-V component of III–V semiconductors. The compound gallium arsenide (GaAs) is a key material for high-speed and optoelectronic devices. GaAs and related alloys (such as AlGaAs, InGaAs, GaAsP, etc.) have a direct band gap, making them ideal for semiconductor lasers, LEDs (especially infrared and red light), high-frequency transistors (for cellphone towers, satellite communications, radar), and high-efficiency solar cells (GaAs cells are widely used in satellites and as multi-junction solar cells). Arsenic is also used in indium arsenide (InAs) and cadmium arsenide (Cd₃As₂) and other III–V materials with specialized uses in detectors and research on topological insulators. In silicon semiconductor manufacturing, arsenic is commonly used as an n-type dopant: introducing a small amount of arsenic atoms into silicon changes its electrical conductivity, since arsenic has one more valence electron than silicon. This is done by diffusion or by implanting arsenic ions. (The precursor gas arsine, AsH₃, is used in epitaxial growth equipment to supply arsenic.)

Arsenic is also added to certain metal alloys. For example, adding a small percentage of arsenic to lead strengthens the lead for battery plates and bullets. Lead–arsenic alloys have better hardness and casting properties than pure lead. Bismuth-arsenide alloys were once used in “wood thermometers” (because bismuth expands uniformly when solidifying). The metal arsenic (especially in alpha grey form) is a component of specialty alloys like Kovar (an iron-nickel alloy used in glass-to-metal seals).

In the mid-20th century, arsenic compounds were widely used in agriculture and wood preservation. Insecticides and herbicides such as lead arsenate and calcium arsenate were used extensively to control pests (for example in orchards) until safer alternatives were developed. Another famous arsenic-based pesticide was Paris Green (copper(II) acetoarsenite), used as a pigment and insecticide in the 19th century (e.g. against mildew, or in street lamp disinfection). Arsenic also was an ingredient in pressure-treated lumber – the CCA process (copper-chrome-arsenic) impregnated wood with arsenic compounds to prevent decay. Though CCA is still used in some industrial applications (utility poles, marine pilings), residential uses have largely been phased out in many countries.

In the chemical industry, arsenic compounds have niche roles. Arsenic trioxide (As₂O₃) is used in glass manufacturing to filter out bubble defects and improve optical quality. Certain organic arsenicals were formerly used as antibiotics in veterinary medicine (e.g. roxarsone in poultry, though many have been banned or limited). Arsenic-based lasers (e.g. GaAs diode lasers) have revolutionized communications and electronics. In metallurgy, additions of arsenic can help remove oxygen impurities from lead through “lead refining”.

Perhaps the most intriguing modern use of arsenic is medical: arsenic trioxide (As₂O₃) is an approved chemotherapy drug (trade name Trisenox) for treating acute promyelocytic leukemia. In this case, the high toxicity of As is harnessed to trigger programmed cell death in leukemia cells. Another medical isotope application is the use of radioactive ⁷²As or ⁷⁴As as tracers for imaging tumors (though that is not routine).

Electronics and solar applications mean that new facilities produce or handle arsenic in the course of manufacturing chips and panels. For example, during the growth of GaAs wafers, AsH₃ gas is carefully managed to incorporate arsenic into the crystal. Researchers are also investigating arsenic-based compounds in emerging fields like spintronics and quantum materials (e.g. arsenides of rare earths or actinides have unusual electronic properties). However, the hazardous nature of arsenic limits some of its potential applications.

Biology, Environment, and Safety

Arsenic is notorious for its toxicity to living organisms. It has no known essential biological role in humans or animals (a few simple organisms can use it, but in general it is regarded as purely toxic). The acute toxicity of arsenic compounds is high: arsenate and arsenite salts can be lethal in milligram-per-kilogram doses, and gases like arsine are lethal at very low parts-per-million concentrations. Chronic exposure to arsenic (especially inorganic As(III) and As(V) compounds) can cause cancer (skin, lung, bladder), skin lesions, cardiovascular disease, and other ailments. Because it interferes with metabolic enzymes (particularly those requiring thiol (-SH) groups), arsenic disrupts cellular respiration and can cause neuropathy, liver damage, and other systemic effects.

Arsenic enters the environment through both natural and human activities. Naturally, volcanoes and geothermal springs release arsenic, and weathering of arsenic-containing minerals releases it into soil and groundwater. Human sources include mining, coal combustion, and industrial emissions. A very serious global health problem is arsenic in drinking water. In places like Bangladesh, India, parts of China, Argentina and the United States (e.g. along some aquifers in the Southwest), groundwater interacts with arsenic-bearing rocks under reducing conditions, releasing As(III) into the well water. In many regions, tens of millions of people have been exposed for decades to drinking water with arsenic well above the World Health Organization guideline of 10 µg/L. Chronic ingestion of such water leads to widespread arsenicosis (skin pigmentation changes, lesions, internal cancers). Arsenic can also accumulate in crops irrigated with contaminated water (rice is especially prone to uptake arsenic), introducing it into the food chain.

In the environmental cycling of arsenic, oxygen levels and microbial action are key. In oxygenated (aerobic) surface waters, arsenate (As(V)) dominates; it behaves chemically like phosphate, binding to iron oxides and staying relatively immobile. In oxygen-poor (anaerobic) conditions, arsenate can be microbially reduced to arsenite (As(III)) and released into solution. Some bacteria oxidize arsenite back to arsenate, obtaining energy. Other microbes detoxify arsenic by methylating it (forming organic arsenic compounds like monomethylarsenate or dimethylarsenate, which are then excreted). These microbial transformations help determine whether arsenic stays in soil or water, is taken up by organisms, or volatilizes (as methylated arsines) into the atmosphere.

Most organic arsenic compounds found in fish and other seafood (e.g. arsenobetaine) are relatively non-toxic and are processed and excreted by mammals. However, inorganic arsenic is dangerous; As(III) species are typically more acutely toxic than As(V). For example, arsenite (AsO₃³⁻) binds strongly to cysteine residues in proteins and disrupts enzymes, whereas arsenate (AsO₄³⁻) can disrupt energy metabolism by mimicking phosphate. Regulatory agencies consider all inorganic arsenic compounds to be carcinogenic (US EPA classifies inorganic arsenic as a known human carcinogen, Group A). Arsenic exposure limits are low: occupational exposure limits for arsenic dust are on the order of 10 µg/m³, and drinking water standards around 10 µg/L.

Because of the health hazards, handling arsenic requires caution. It is toxic by ingestion, inhalation, and even skin contact (some arsenic compounds can penetrate the skin or cause local cell death). Arsine gas is extremely dangerous (it causes hemolysis of red blood cells). Safe work procedures include closed systems for gases, respirators and protective clothing when milling arsenic ore or dust, and robust ventilation in smelters and semiconductor fabs. Arsenic waste (e.g. contaminated soils or sludges) must be disposed of as hazardous waste.

In ecosystems, arsenic can bioaccumulate to a limited extent, but in humans it does not bioaccumulate strongly (it has a short biological half-life of a few days, though chronic exposure can maintain high body burdens). Arsenic in soil is poorly digested by plants; most arsenic in crops is taken up from water or soluble soil fractions. (Interestingly, some ferns and wetlands plants are known “hyperaccumulators” of arsenic, and are studied for phytoremediation.) In animals, arsenic can cross the placenta and is found in breast milk; it also accumulates in hair and nails.

Modern detection and remediation techniques aim to control arsenic exposure. For water, treatments like coagulation/filtration (iron or alum), reverse osmosis, or ion exchange can remove arsenic. For contaminated land, options include soil washing, stabilization, and phytoremediation. Monitoring of arsenic levels in groundwater, soil, and food has become routine in many countries.

History and Etymology

Arsenic has been known to humans since ancient times, prized because of its potent properties. The name “arsenic” ultimately comes from the Greek word arsenikon (ἀρσενικόν), meaning “potent” or “male,” which was used for a yellow pigment (orpiment) thought to contain the substance. Another origin is the Persian word zarnikh (“yellow orpiment”), which passed into Arabic al-zarnīkh and then into European languages. Alchemy and medieval medicine were well aware of arsenic compounds. The element itself was isolated or recognized in the 17th century. German chemist Johann Schröder in 1649 was the first to prepare metal arsenic by heating arsenic sulfide with soap. Earlier alchemists (like Albertus Magnus in the 13th century) had described impure arsenic upon heating sulfides, but Schröder proved it was an element by bringing it to the metallic state.

In alchemical tradition, arsenic’s toxic “magic” was famous. It was sometimes called the “king of poisons.” Paracelsus and other Renaissance figures experimented with arsenic compounds. By the 18th century, physicians had discovered some medical uses of arsenic (treating syphilis or asthma), and Fowler’s solution (a potassium arsenite solution named after chemist Thomas Fowler) was used as a tonic for ailments. But poisoning — whether accidental, suicidal, or homicidal — was far more noted. Arsenic became notorious in history and literature as a slow, undetectable poison (famously featured in stories like Agatha Christie’s “Arsenic and Old Lace”).

Industrial use of arsenic grew in the 19th and early 20th centuries. In 1800s Europe, Scheele’s Green (copper acetoarsenite) was a valuable pigment (and indoor wallpaper trend) before its toxicity was recognized. Arsenic-laced pesticides, like calcium arsenate against cotton pests and lead arsenate against fruit insects, were widely used in farming up until mid-20th century. One of the first semiconductor devices (the “cat’s whisker” detector of the early 1900s) used a junction of galena (PbS) and a thin wire, and even then traces of arsenic in lead and copper ores were being handled (though electronics didn’t use arsenic until late 20th century).

The development of semiconductor technology in the 1950s–1970s gave arsenic a new role. Martin Pogobler and coworkers discovered the properties of GaAs, and the material was used in the first gallium arsenide laser in the 1960s. Over time, arsenic became central to communications technology. Meanwhile, concerns about arsenic poisoning led to regulations (bys the late 20th century many arsenic pesticides were banned in the US and Europe, and arsenic in drinking water became an important standard).

In etymological terms and official nomenclature, arsenic’s symbol As derives from Arsenicum, the Latinized form. Some historical synonyms survived into modern usage (e.g. “arsenolite” for As₂O₃). The etymology reflects arsenic’s dual image: “arsenic” once meant both the element and the highly potent or male virility, underscoring how its name and legend developed together.

Data Table

The table below summarizes key properties of arsenic:

Property	Value
Symbol	As
Atomic Number (Z)	33
Atomic Weight	74.9216 (u)
Electron Configuration	[Ar] 3d¹⁰ 4s² 4p³
Valence Electrons	5
Atomic Radius (covalent)	~1.20 Å
Density (gray, 20 °C)	5.72 g/cm³
Melting Point	817 °C (1090 K at 3.63 MPa)
Sublimation Point (1 atm)	~614 °C (887 K)
Boiling Point	(sublimes at ~615 °C at 1 atm)
Crystal Structure	Rhombohedral (layered) (β-As)
Electrical Conductivity	~3.45×10^4 S/m (gray arsenic, 20 °C)
Thermal Conductivity	~50 W/(m·K)
Common Oxidation States	–3 (in metal arsenides), +3, +5
Electronegativity (Paul.)	2.18
First Ionization Energy	9.79 eV (946 kJ/mol)
Stable Isotope	⁷⁵As (100%; nuclear spin 3/2)
Abundance (Earth’s crust)	~1.5 ppm (by weight)
Element Category	Metalloid (pnictogen)