Arsenic is a chemical element with the symbol As, and atomic number 33. Arsenic occurs in many minerals, usually in combination with sulfur and metals, but also as a pure elemental crystal. Arsenic is a metalloid. It has various allotropes, but only the gray form, which has a metallic appearance, is important to industry.

Arsenic Description

Arsenic is a chemical element with the symbol As, and atomic number 33. Arsenic occurs in many minerals, usually in combination with sulfur and metals, but also as a pure elemental crystal. Arsenic is a metalloid. It has various allotropes, but only the gray form, which has a metallic appearance, is important to industry.

The primary use of arsenic is in alloys of lead (for example, in car batteries and ammunition). Arsenic is a common n-type dopant in semiconductor electronic devices, and the optoelectronic compound gallium arsenide is the second most commonly used semiconductor after doped silicon. Arsenic and its compounds, especially the trioxide, are used in the production of pesticides, treated wood products, herbicides, and insecticides. These applications are declining due to the increasing recognition of the toxicity of arsenic and its compounds.

Arsenic Compounds

Compounds of arsenic resemble in some respects those of phosphorus which occupy the same group (column) of the periodic table. The most common oxidation states for arsenic are −3 in the arsenides, which are alloy-like intermetallic compounds, +3 in the arsenites, and +5 in the arsenates and most organoarsenic compounds. Arsenic also bonds readily to itself as seen in the square As3−4 ions in the mineral skutterudite. In the +3 oxidation state, arsenic is typically pyramidal owing to the influence of the lone pair of electrons.

Inorganic compounds

One of the simplest arsenic compounds is the trihydride, the highly toxic, flammable, pyrophoric arsine (AsH3). This compound is generally regarded as stable since at room temperature it decomposes only slowly. At temperatures of 250–300 °C decomposition to arsenic and hydrogen is rapid.[25] Several factors, such as humidity, presence of light, and certain catalysts (namely aluminum) facilitate the rate of decomposition. It oxidizes readily in the air to form arsenic trioxide and water, and analogous reactions take place with sulfur and selenium instead of oxygen.

Arsenic forms colorless, odorless, crystalline oxides As2O3 ("white arsenic") and As2O5 which are hygroscopic and readily soluble in water to form acidic solutions. Arsenic(V) acid is a weak acid and the salts are called arsenates, the most common arsenic contamination of groundwater, and a problem that affects many people. Synthetic arsenates include Scheele's Green (cupric hydrogen arsenate, acidic copper arsenate), calcium arsenate, and lead hydrogen arsenate. These three have been used as agricultural insecticides and poisons.

The protonation steps between the arsenate and arsenic acid are similar to those between phosphate and phosphoric acid. Unlike phosphorous acid, arsenous acid is genuinely tribasic, with the formula As(OH)3.

A broad variety of sulfur compounds of arsenic are known. Orpiment (As2S3) and realgar (As4S4) are somewhat abundant and were formerly used as painting pigments. In As4S10, arsenic has a formal oxidation state of +2 in As4S4 which features As-As bonds so that the total covalency of As is still 3. Both orpiment and realgar, as well as As4S3, have selenium analogs; the analogous As2Te3 is known as the mineral kalgoorlieite, and the anion As2Te− is known as a ligand in cobalt complexes.

All trihalides of arsenic(III) are well known except the astatide, which is unknown. Arsenic pentafluoride (AsF5) is the only important pentahalide, reflecting the lower stability of the +5 oxidation state; even so, it is a very strong fluorinating and oxidizing agent. (The pentachloride is stable only below −50 °C, at which temperature it decomposes to the trichloride, releasing chlorine gas.


Arsenic is used as the group 5 element in the III-V semiconductors gallium arsenide, indium arsenide, and aluminum arsenide. The valence electron count of GaAs is the same as a pair of Si atoms, but the band structure is completely different which results in distinct bulk properties. Other arsenic alloys include the II-V semiconductor cadmium arsenide.

Arsenic Applications


The toxicity of arsenic to insects, bacteria, and fungi led to its use as a wood preservative. Arsenic was also used in various agricultural insecticides and poisons. Arsenic is used as a feed additive in poultry and swine production, in particular in the U.S. to increase weight gain, improve feed efficiency, and to prevent disease


During the 18th, 19th, and 20th centuries, a number of arsenic compounds were used as medicines, including arsphenamine (by Paul Ehrlich) and arsenic trioxide (by Thomas Fowler). Arsenic trioxide has been used in a variety of ways over the past 500 years, most commonly in the treatment of cancer, but also in medications as diverse as Fowler's solution in psoriasis. Recently, researchers have been locating tumors using arsenic-74 (a positron emitter). This isotope produces clearer PET scan images than the previous radioactive agent, iodine-124, because the body tends to transport iodine to the thyroid gland producing signal noise. Nanoparticles of arsenic have shown the ability to kill cancer cells with lesser cytotoxicity than other arsenic formulations


The main use of arsenic is in alloying with lead. Lead components in car batteries are strengthened by the presence of a very small percentage of arsenic. Dezincification of brass (a copper-zinc alloy) is greatly reduced by the addition of arsenic. "Phosphorus Deoxidized Arsenical Copper" with an arsenic content of 0.3% has increased corrosion stability in certain environments. Gallium arsenide is an important semiconductor material, used in integrated circuits. Circuits made from GaAs are much faster (but also much more expensive) than those made from silicon. Unlike silicon, GaAs has a direct bandgap and can be used in laser diodes and LEDs to convert electrical energy directly into light.


After World War I, the United States built a stockpile of 20,000 tons of weaponized lewisite (ClCH=CHAsCl2), an organoarsenic vesicant (blister agent), and lung irritant. The stockpile was neutralized with bleach and dumped into the Gulf of Mexico in the 1950s. During the Vietnam War, the United States used Agent Blue, a mixture of sodium cacodylate and its acid form, as one of the rainbow herbicides to deprive North Vietnamese soldiers of foliage cover and rice.


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