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Argon | External links | External links
Argon at The Periodic Table of Videos (University of Nottingham)
USGS Periodic Table – Argon
Diving applications: Why Argon?
Category:Chemical elements
Category:E-number additives
Category:Noble gases
Category:Industrial gases |
Argon | Table of Content | About, Characteristics, History, Occurrence, Isotopes, Compounds, Production, Applications, Industrial processes, Scientific research, Preservative, Laboratory equipment, Medical use, Lighting, Miscellaneous uses, Safety, See also, References, Further reading, External links |
Arsenic | <!-- {{cs1 config | Arsenic is a chemical element; it has symbol As and atomic number 33. It is a metalloid and one of the pnictogens, and therefore shares many properties with its group 15 neighbors phosphorus and antimony. Arsenic is notoriously toxic. It occurs naturally in many minerals, usually in combination with sulfur and metals, but also as a pure elemental crystal. It has various allotropes, but only the grey 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 also a common n-type dopant in semiconductor electronic devices, and a component of the III–V compound semiconductor gallium arsenide. Arsenic and its compounds, especially the trioxide, are used in the production of pesticides, treated wood products, herbicides, and insecticides. These applications are declining with the increasing recognition of the persistent toxicity of arsenic and its compounds.
Arsenic has been known since ancient times to be poisonous to humans. However, a few species of bacteria are able to use arsenic compounds as respiratory metabolites. Trace quantities of arsenic have been proposed to be an essential dietary element in rats, hamsters, goats, and chickens. Research has not been conducted to determine whether small amounts of arsenic may play a role in human metabolism. However, arsenic poisoning occurs in multicellular life if quantities are larger than needed. Arsenic contamination of groundwater is a problem that affects millions of people across the world.
The United States' Environmental Protection Agency states that all forms of arsenic are a serious risk to human health. The United States Agency for Toxic Substances and Disease Registry ranked arsenic number 1 in its 2001 prioritized list of hazardous substances at Superfund sites. Arsenic is classified as a group-A carcinogen. |
Arsenic | Characteristics | Characteristics |
Arsenic | Physical characteristics | Physical characteristics
thumb|left|Crystal structure common to Sb, AsSb and grey As
thumb|left|Gray arsenic nodule
The three most common arsenic allotropes are grey, yellow, and black arsenic, with grey being the most common. Grey arsenic (α-As, space group Rm No. 166) adopts a double-layered structure consisting of many interlocked, ruffled, six-membered rings. Because of weak bonding between the layers, grey arsenic is brittle and has a relatively low Mohs hardness of 3.5. Nearest and next-nearest neighbors form a distorted octahedral complex, with the three atoms in the same double-layer being slightly closer than the three atoms in the next. This relatively close packing leads to a high density of 5.73 g/cm3. Grey arsenic is a semimetal, but becomes a semiconductor with a bandgap of 1.2–1.4 eV if amorphized. Grey arsenic is also the most stable form.
Yellow arsenic is soft and waxy, and somewhat similar to tetraphosphorus (). Both have four atoms arranged in a tetrahedral structure in which each atom is bound to each of the other three atoms by a single bond. This unstable allotrope, being molecular, is the most volatile, least dense, and most toxic. Solid yellow arsenic is produced by rapid cooling of arsenic vapor, . It is rapidly transformed into grey arsenic by light. The yellow form has a density of 1.97 g/cm3. Black arsenic is similar in structure to black phosphorus.
Black arsenic can also be formed by cooling vapor at around 100–220 °C and by crystallization of amorphous arsenic in the presence of mercury vapors. It is glassy and brittle. Black arsenic is also a poor electrical conductor.Arsenic Element Facts. chemicool.com
Arsenic sublimes upon heating at atmospheric pressure, converting directly to a gaseous form without an intervening liquid state at . The triple point is at 3.63 MPa and . |
Arsenic | Isotopes | Isotopes
Arsenic occurs in nature as one stable isotope, 75As, and is therefore called a monoisotopic element. As of 2024, at least 32 radioisotopes have also been synthesized, ranging in atomic mass from 64–95. The most stable of these is 73As with a half-life of 80.30 days. The majority of the other isotopes have half-lives of under one day, with the exceptions being
71As ( 65.30 hours),
72As ( 26.0 hours),
74As ( 17.77 days),
76As ( 26.26 hours),
77As ( 38.83 hours).
Isotopes that are lighter than the stable 75As tend to decay by β+ decay, and those that are heavier tend to decay by β− decay, with some exceptions.
At least 10 nuclear isomers have been described, ranging in atomic mass from 66 to 84. The most stable of arsenic's isomers is 68mAs with a half-life of 111 seconds. |
Arsenic | Chemistry | Chemistry
Arsenic has a similar electronegativity and ionization energies to its lighter pnictogen congener phosphorus and therefore readily forms covalent molecules with most of the nonmetals. Though stable in dry air, arsenic forms a golden-bronze tarnish upon exposure to humidity which eventually becomes a black surface layer.Greenwood and Earnshaw, pp. 552–4 When heated in air, arsenic oxidizes to arsenic trioxide; the fumes from this reaction have an odor resembling garlic. This odor can be detected on striking arsenide minerals such as arsenopyrite with a hammer. It burns in oxygen to form arsenic trioxide and arsenic pentoxide, which have the same structure as the more well-known phosphorus compounds, and in fluorine to give arsenic pentafluoride. Arsenic makes arsenic acid with concentrated nitric acid, arsenous acid with dilute nitric acid, and arsenic trioxide with concentrated sulfuric acid; however, it does not react with water, alkalis, or non-oxidising acids. Arsenic reacts with metals to form arsenides, though these are not ionic compounds containing the As3− ion as the formation of such an anion would be highly endothermic and even the group 1 arsenides have properties of intermetallic compounds. Like germanium, selenium, and bromine, which like arsenic succeed the 3d transition series, arsenic is much less stable in the +5 oxidation state than its vertical neighbors phosphorus and antimony, and hence arsenic pentoxide and arsenic acid are potent oxidizers. |
Arsenic | Compounds | Compounds
Compounds of arsenic resemble, in some respects, those of phosphorus, which occupies 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 ions in the mineral skutterudite. In the +3 oxidation state, arsenic is typically pyramidal owing to the influence of the lone pair of electrons. |
Arsenic | Inorganic compounds | 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.Greenwood and Earnshaw, pp. 557–558 Several factors, such as humidity, presence of light and certain catalysts (namely aluminium) facilitate the rate of decomposition. It oxidises readily in 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 its salts, known as arsenates, are a major source of arsenic contamination of groundwater in regions with high levels of naturally-occurring arsenic minerals. 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.Greenwood and Earnshaw, pp. 572–578
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.Greenwood and Earnshaw, pp. 578–583
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 | Alloys | Alloys
Arsenic is used as the group 5 element in the III-V semiconductors gallium arsenide, indium arsenide, and aluminium 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 | Organoarsenic compounds | Organoarsenic compounds
left|upright=0.4|thumb|Trimethylarsine
A large variety of organoarsenic compounds are known. Several were developed as chemical warfare agents during World War I, including vesicants such as lewisite and vomiting agents such as adamsite. Cacodylic acid, which is of historic and practical interest, arises from the methylation of arsenic trioxide, a reaction that has no analogy in phosphorus chemistry. Cacodyl was the first organometallic compound known (even though arsenic is not a true metal) and was named from the Greek κακωδία "stink" for its offensive, garlic-like odor; it is very toxic.Greenwood, p. 584 |
Arsenic | Occurrence and production | Occurrence and production
upright=0.9|thumb|A large sample of native arsenic from Sainte-Marie-aux-Mines, France
Arsenic is the 53rd most abundant element in the Earth's crust, comprising about 1.5 parts per million (0.00015%). Typical background concentrations of arsenic do not exceed 3 ng/m3 in the atmosphere; 100 mg/kg in soil; 400 μg/kg in vegetation; 10 μg/L in freshwater and 1.5 μg/L in seawater. Arsenic is the 22nd most abundant element in seawater and ranks 41st in abundance in the universe.
Minerals with the formula MAsS and MAs2 (M = Fe, Ni, Co) are the dominant commercial sources of arsenic, together with realgar (an arsenic sulfide mineral) and native (elemental) arsenic. An illustrative mineral is arsenopyrite (FeAsS), which is structurally related to iron pyrite. Many minor As-containing minerals are known. Arsenic also occurs in various organic forms in the environment.
thumb|upright=1.15|Arsenic output in 2006
In 2014, China was the top producer of white arsenic with almost 70% world share, followed by Morocco, Russia, and Belgium, according to the British Geological Survey and the United States Geological Survey. Most arsenic refinement operations in the US and Europe have closed over environmental concerns. Arsenic is found in the smelter dust from copper, gold, and lead smelters, and is recovered primarily from copper refinement dust. Arsenic is the main of impurity found in copper concentrates to enter copper smelting facilities. There has been an increase in arsenic in copper concentrates over the years since copper mining has moved into deep high-impurity ores as shallow, low-arsenic copper deposits have been progressively depleted.
On roasting arsenopyrite in air, arsenic sublimes as arsenic(III) oxide leaving iron oxides, while roasting without air results in the production of gray arsenic. Further purification from sulfur and other chalcogens is achieved by sublimation in vacuum, in a hydrogen atmosphere, or by distillation from molten lead-arsenic mixture.
Rank Country 2014 As2O3 Production 1 25,000 T 2 8,800 T 3 1,500 T 4 1,000 T 5 52 T 6 45 T — World Total (rounded) 36,400 T |
Arsenic | History | History
thumb|Realgar
upright=0.35|thumb|Alchemical symbol for arsenic
The word arsenic has its origin in the Syriac word zarnika, from Arabic al-zarnīḵ 'the orpiment', based on Persian zar ("gold") from the word zarnikh, meaning "yellow" (literally "gold-colored") and hence "(yellow) orpiment". It was adopted into Greek (using folk etymology) as arsenikon () – a neuter form of the Greek adjective arsenikos (), meaning "male", "virile".
Latin-speakers adopted the Greek term as , which in French ultimately became , whence the English word "arsenic".
Arsenic sulfides (orpiment, realgar) and oxides have been known and used since ancient times. Zosimos () describes roasting sandarach (realgar) to obtain cloud of arsenic (arsenic trioxide), which he then reduces to gray arsenic. As the symptoms of arsenic poisoning are not very specific, the substance was frequently used for murder until the advent in the 1830s of the Marsh test, a sensitive chemical test for its presence. (Another less sensitive but more general test is the Reinsch test.) Owing to its use by the ruling class to murder one another and its potency and discreetness, arsenic has been called the "poison of kings" and the "king of poisons". Arsenic became known as "the inheritance powder" due to its use in killing family members in the Renaissance era.
thumb|left|The arsenic labyrinth, part of Botallack Mine, Cornwall
During the Bronze Age, arsenic was melted with copper to make arsenical bronze.
Jabir ibn Hayyan described the isolation of arsenic before 815 AD.
Albertus Magnus (Albert the Great, 1193–1280) later isolated the element from a compound in 1250, by heating soap together with arsenic trisulfide. In 1649, Johann Schröder published two ways of preparing arsenic. Crystals of elemental (native) arsenic are found in nature, although rarely.
Cadet's fuming liquid (impure cacodyl), often claimed as the first synthetic organometallic compound, was synthesized in 1760 by Louis Claude Cadet de Gassicourt through the reaction of potassium acetate with arsenic trioxide.
thumb|Satirical cartoon by Honoré Daumier of a chemist giving a public demonstration of arsenic, 1841
In the Victorian era, women would eat "arsenic" ("white arsenic" or arsenic trioxide) mixed with vinegar and chalk to improve the complexion of their faces, making their skin paler (to show they did not work in the fields). The accidental use of arsenic in the adulteration of foodstuffs led to the Bradford sweet poisoning in 1858, which resulted in 21 deaths. From the late 18th century wallpaper production began to use dyes made from arsenic,
which was thought to increase the pigment's brightness. One account of the illness and 1821 death of Napoleon implicates arsenic poisoning involving wallpaper.
Two arsenic pigments have been widely used since their discovery – Paris Green in 1814 and Scheele's Green in 1775. After the toxicity of arsenic became widely known, these chemicals were used less often as pigments and more often as insecticides. In the 1860s, an arsenic byproduct of dye production, London Purple, was widely used. This was a solid mixture of arsenic trioxide, aniline, lime, and ferrous oxide, insoluble in water and very toxic by inhalation or ingestion But it was later replaced with Paris Green, another arsenic-based dye. With better understanding of the toxicology mechanism, two other compounds were used starting in the 1890s. Arsenite of lime and arsenate of lead were used widely as insecticides until the discovery of DDT in 1942.
In small doses, soluble arsenic compounds act as stimulants, and were once popular as medicine by people in the mid-18th to 19th centuries; this use was especially prevalent for sport animals such as race horses or work dogs and continued into the 20th century.
A 2006 study of the remains of the Australian racehorse Phar Lap determined that its 1932 death was caused by a massive overdose of arsenic. Sydney veterinarian Percy Sykes stated,
"In those days, arsenic was quite a common tonic, usually given in the form of a solution (Fowler's Solution) ... It was so common that I'd reckon 90 per cent of the horses had arsenic in their system." |
Arsenic | Applications | Applications |
Arsenic | Agricultural | Agricultural
thumb|Roxarsone is a controversial arsenic compound used as a feed ingredient for chickens.
The toxicity of arsenic to insects, bacteria, and fungi led to its use as a wood preservative. In the 1930s, a process of treating wood with chromated copper arsenate (also known as CCA or Tanalith) was invented, and for decades, this treatment was the most extensive industrial use of arsenic. An increased appreciation of the toxicity of arsenic led to a ban of CCA in consumer products in 2004, initiated by the European Union and United States. However, CCA remains in heavy use in other countries (such as on Malaysian rubber plantations).
Arsenic was also used in various agricultural insecticides and poisons. For example, lead hydrogen arsenate was a common insecticide on fruit trees, but contact with the compound sometimes resulted in brain damage among those working the sprayers. In the second half of the 20th century, monosodium methyl arsenate (MSMA) and disodium methyl arsenate (DSMA) – less toxic organic forms of arsenic – replaced lead arsenate in agriculture. These organic arsenicals were in turn phased out in the United States by 2013 in all agricultural activities except cotton farming.
The biogeochemistry of arsenic is complex and includes various adsorption and desorption processes. The toxicity of arsenic is connected to its solubility and is affected by pH. Arsenite () is more soluble than arsenate () and is more toxic; however, at a lower pH, arsenate becomes more mobile and toxic. It was found that addition of sulfur, phosphorus, and iron oxides to high-arsenite soils greatly reduces arsenic phytotoxicity.
Arsenic is used as a feed additive in poultry and swine production, in particular it was used in the U.S. until 2015 to increase weight gain, improve feed efficiency, and prevent disease. An example is roxarsone, which had been used as a broiler starter by about 70% of U.S. broiler growers. In 2011, Alpharma, a subsidiary of Pfizer Inc., which produces roxarsone, voluntarily suspended sales of the drug in response to studies showing elevated levels of inorganic arsenic, a carcinogen, in treated chickens. A successor to Alpharma, Zoetis, continued to sell nitarsone until 2015, primarily for use in turkeys. |
Arsenic | Medical use | Medical use
During the 17th, 18th, and 19th centuries, a number of arsenic compounds were used as medicines, including arsphenamine (by Paul Ehrlich) and arsenic trioxide (by Thomas Fowler), for treating diseases such as cancer or psoriasis. Arsphenamine, as well as neosalvarsan, was indicated for syphilis, but has been superseded by modern antibiotics. However, arsenicals such as melarsoprol are still used for the treatment of trypanosomiasis in spite of their severe toxicity, since the disease is almost uniformly fatal if untreated. In 2000 the US Food and Drug Administration approved arsenic trioxide for the treatment of patients with acute promyelocytic leukemia that is resistant to all-trans retinoic acid.
A 2008 paper reports success in 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 ability to kill cancer cells with lesser cytotoxicity than other arsenic formulations. |
Arsenic | Alloys | Alloys
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 an 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. |
Arsenic | Military | Military
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. Lewisite, the chemical warfare agent, is known for its acute toxicity to aquatic organisms. However, studies assessing the environmental impact of this disposal in the Gulf are lacking. 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. |
Arsenic | Other uses | Other uses
Copper acetoarsenite was used as a green pigment known under many names, including Paris Green and Emerald Green. It caused numerous arsenic poisonings. Scheele's Green, a copper arsenate, was used in the 19th century as a coloring agent in sweets.
Arsenic is used in bronzing.
As much as 2% of produced arsenic is used in lead alloys for lead shot and bullets.
Arsenic is added in small quantities to alpha-brass to make it dezincification-resistant. This grade of brass is used in plumbing fittings and other wet environments.
Arsenic is also used for taxonomic sample preservation. It was also used in embalming fluids historically.
Arsenic was used in the taxidermy process up until the 1980s.
Arsenic was used as an opacifier in ceramics, creating white glazes.
Until recently, arsenic was used in optical glass. Modern glass manufacturers have ceased using both arsenic and lead. |
Arsenic | Biological role | Biological role |
Arsenic | Bacteria | Bacteria
Some species of bacteria obtain their energy in the absence of oxygen by oxidizing various fuels while reducing arsenate to arsenite. Under oxidative environmental conditions some bacteria use arsenite as fuel, which they oxidize to arsenate. The enzymes involved are known as arsenate reductases (Arr).
In 2008, bacteria were discovered that employ a version of photosynthesis in the absence of oxygen with arsenites as electron donors, producing arsenates (just as ordinary photosynthesis uses water as electron donor, producing molecular oxygen). Researchers conjecture that, over the course of history, these photosynthesizing organisms produced the arsenates that allowed the arsenate-reducing bacteria to thrive. One strain, PHS-1, has been isolated and is related to the gammaproteobacterium Ectothiorhodospira shaposhnikovii. The mechanism is unknown, but an encoded Arr enzyme may function in reverse to its known homologues.
In 2011, it was postulated that the Halomonadaceae strain GFAJ-1 could be grown in the absence of phosphorus if that element were substituted with arsenic, exploiting the fact that the arsenate and phosphate anions are similar structurally. The study was widely criticised and subsequently refuted by independent research groups. |
Arsenic | Potential role in higher animals | Potential role in higher animals
Arsenic may be an essential trace mineral in birds, involved in the synthesis of methionine metabolites. However, the role of arsenic in bird nutrition is disputed, as other authors state that arsenic is toxic in small amounts.
Some evidence indicates that arsenic is an essential trace mineral in mammals.Anke M. (1986) "Arsenic", pp. 347–372 in Mertz W. (ed.), Trace elements in human and Animal Nutrition, 5th ed. Orlando, FL: Academic Press |
Arsenic | Heredity | Heredity
Arsenic has been linked to epigenetic changes, heritable changes in gene expression that occur without changes in DNA sequence. These include DNA methylation, histone modification, and RNA interference. Toxic levels of arsenic cause significant DNA hypermethylation of tumor suppressor genes p16 and p53, thus increasing risk of carcinogenesis. These epigenetic events have been studied in vitro using human kidney cells and in vivo using rat liver cells and peripheral blood leukocytes in humans. Inductively coupled plasma mass spectrometry (ICP-MS) is used to detect precise levels of intracellular arsenic and other arsenic bases involved in epigenetic modification of DNA. Studies investigating arsenic as an epigenetic factor can be used to develop precise biomarkers of exposure and susceptibility.
The Chinese brake fern (Pteris vittata) hyperaccumulates arsenic from the soil into its leaves and has a proposed use in phytoremediation. |
Arsenic | Biomethylation | Biomethylation
thumb|Arsenobetaine
Inorganic arsenic and its compounds, upon entering the food chain, are progressively metabolized through a process of methylation. For example, the mold Scopulariopsis brevicaulis produces trimethylarsine if inorganic arsenic is present. The organic compound arsenobetaine is found in some marine foods such as fish and algae, and also in mushrooms in larger concentrations. The average person's intake is about 10–50 μg/day. Values about 1000 μg are not unusual following consumption of fish or mushrooms, but there is little danger in eating fish because this arsenic compound is nearly non-toxic. |
Arsenic | Environmental issues | Environmental issues |
Arsenic | Exposure | Exposure
Naturally occurring sources of human exposure include volcanic ash, weathering of minerals and ores, and mineralized groundwater. Arsenic is also found in food, water, soil, and air. Arsenic is absorbed by all plants, but is more concentrated in leafy vegetables, rice, apple and grape juice, and seafood. An additional route of exposure is inhalation of atmospheric gases and dusts.
During the Victorian era, arsenic was widely used in home decor, especially wallpapers. In Europe, an analysis based on 20,000 soil samples across all 28 countries show that 98% of sampled soils have concentrations less than 20 mg/kg. In addition, the arsenic hotspots are related to both frequent fertilization and close distance to mining activities. Chronic exposure to arsenic, particularly through contaminated drinking water and food, has also been linked to long-term impacts on cognitive function, including reduced verbal IQ and memory. |
Arsenic | Occurrence in drinking water | Occurrence in drinking water
Extensive arsenic contamination of groundwater has led to widespread arsenic poisoning in Bangladesh and neighboring countries. It is estimated that approximately 57 million people in the Bengal basin are drinking groundwater with arsenic concentrations elevated above the World Health Organization's standard of 10 parts per billion (ppb). However, a study of cancer rates in Taiwan suggested that significant increases in cancer mortality appear only at levels above 150 ppb. The arsenic in the groundwater is of natural origin, and is released from the sediment into the groundwater, caused by the anoxic conditions of the subsurface. This groundwater was used after local and western NGOs and the Bangladeshi government undertook a massive shallow tube well drinking-water program in the late twentieth century. This program was designed to prevent drinking of bacteria-contaminated surface waters, but failed to test for arsenic in the groundwater. Many other countries and districts in Southeast Asia, such as Vietnam and Cambodia, have geological environments that produce groundwater with a high arsenic content. Arsenicosis was reported in Nakhon Si Thammarat, Thailand, in 1987, and the Chao Phraya River probably contains high levels of naturally occurring dissolved arsenic without being a public health problem because much of the public uses bottled water. In Pakistan, more than 60 million people are exposed to arsenic polluted drinking water indicated by a 2017 report in Science. Podgorski's team investigated more than 1200 samples and more than 66% exceeded the WHO contamination limits of 10 micrograms per liter.
Since the 1980s, residents of the Ba Men region of Inner Mongolia, China have been chronically exposed to arsenic through drinking water from contaminated wells. A 2009 research study observed an elevated presence of skin lesions among residents with well water arsenic concentrations between 5 and 10 μg/L, suggesting that arsenic-induced toxicity may occur at relatively low concentrations with chronic exposure. Overall, 20 of China's 34 provinces have high arsenic concentrations in the groundwater supply, potentially exposing 19 million people to hazardous drinking water.
A study by IIT Kharagpur found high levels of Arsenic in groundwater of 20% of India's land, exposing more than 250 million people. States such as Punjab, Bihar, West Bengal, Assam, Haryana, Uttar Pradesh, and Gujarat have highest land area exposed to arsenic.
In the United States, arsenic is most commonly found in the ground waters of the southwest. Parts of New England, Michigan, Wisconsin, Minnesota and the Dakotas are also known to have significant concentrations of arsenic in ground water. Increased levels of skin cancer have been associated with arsenic exposure in Wisconsin, even at levels below the 10 ppb drinking water standard. According to a recent film funded by the US Superfund, millions of private wells have unknown arsenic levels, and in some areas of the US, more than 20% of the wells may contain levels that exceed established limits.
Low-level exposure to arsenic at concentrations of 100 ppb (i.e., above the 10 ppb drinking water standard) compromises the initial immune response to H1N1 or swine flu infection according to NIEHS-supported scientists. The study, conducted in laboratory mice, suggests that people exposed to arsenic in their drinking water may be at increased risk for more serious illness or death from the virus.
Some Canadians are drinking water that contains inorganic arsenic. Private-dug–well waters are most at risk for containing inorganic arsenic. Preliminary well water analysis typically does not test for arsenic. Researchers at the Geological Survey of Canada have modeled relative variation in natural arsenic hazard potential for the province of New Brunswick. This study has important implications for potable water and health concerns relating to inorganic arsenic.
Epidemiological evidence from Chile shows a dose-dependent connection between chronic arsenic exposure and various forms of cancer, in particular when other risk factors, such as cigarette smoking, are present. These effects have been demonstrated at contaminations less than 50 ppb. Arsenic is itself a constituent of tobacco smoke.
Analyzing multiple epidemiological studies on inorganic arsenic exposure suggests a small but measurable increase in risk for bladder cancer at 10 ppb. According to Peter Ravenscroft of the Department of Geography at the University of Cambridge, roughly 80 million people worldwide consume between 10 and 50 ppb arsenic in their drinking water. If they all consumed exactly 10 ppb arsenic in their drinking water, the previously cited multiple epidemiological study analysis would predict an additional 2,000 cases of bladder cancer alone. This represents a clear underestimate of the overall impact, since it does not include lung or skin cancer, and explicitly underestimates the exposure. Those exposed to levels of arsenic above the current WHO standard should weigh the costs and benefits of arsenic remediation.
Early (1973) evaluations of the processes for removing dissolved arsenic from drinking water demonstrated the efficacy of co-precipitation with either iron or aluminium oxides. In particular, iron as a coagulant was found to remove arsenic with an efficacy exceeding 90%. Several adsorptive media systems have been approved for use at point-of-service in a study funded by the United States Environmental Protection Agency (US EPA) and the National Science Foundation (NSF). A team of European and Indian scientists and engineers have set up six arsenic treatment plants in West Bengal based on in-situ remediation method (SAR Technology). This technology does not use any chemicals and arsenic is left in an insoluble form (+5 state) in the subterranean zone by recharging aerated water into the aquifer and developing an oxidation zone that supports arsenic oxidizing micro-organisms. This process does not produce any waste stream or sludge and is relatively cheap.
Another effective and inexpensive method to avoid arsenic contamination is to sink wells 500 feet or deeper to reach purer waters. A recent 2011 study funded by the US National Institute of Environmental Health Sciences' Superfund Research Program shows that deep sediments can remove arsenic and take it out of circulation. In this process, called adsorption, arsenic sticks to the surfaces of deep sediment particles and is naturally removed from the ground water.
Magnetic separations of arsenic at very low magnetic field gradients with high-surface-area and monodisperse magnetite (Fe3O4) nanocrystals have been demonstrated in point-of-use water purification. Using the high specific surface area of Fe3O4 nanocrystals, the mass of waste associated with arsenic removal from water has been dramatically reduced.
Epidemiological studies have suggested a correlation between chronic consumption of drinking water contaminated with arsenic and the incidence of all leading causes of mortality. The literature indicates that arsenic exposure is causative in the pathogenesis of diabetes.
Chaff-based filters have recently been shown to reduce the arsenic content of water to 3 μg/L. This may find applications in areas where the potable water is extracted from underground aquifers. |
Arsenic | San Pedro de Atacama | San Pedro de Atacama
For several centuries, the people of San Pedro de Atacama in Chile have been drinking water that is contaminated with arsenic, and some evidence suggests they have developed some immunity. |
Arsenic | Hazard maps for contaminated groundwater | Hazard maps for contaminated groundwater
Around one-third of the world's population drinks water from groundwater resources. Of this, about 10 percent, approximately 300 million people, obtains water from groundwater resources that are contaminated with unhealthy levels of arsenic or fluoride.Eawag (2015) Geogenic Contamination Handbook – Addressing Arsenic and Fluoride in Drinking Water. C.A. Johnson, A. Bretzler (Eds.), Swiss Federal Institute of Aquatic Science and Technology (Eawag), Duebendorf, Switzerland. (download: www.eawag.ch/en/research/humanwelfare/drinkingwater/wrq/geogenic-contamination-handbook/) These trace elements derive mainly from minerals and ions in the ground. |
Arsenic | Redox transformation of arsenic in natural waters | Redox transformation of arsenic in natural waters
Arsenic is unique among the trace metalloids and oxyanion-forming trace metals (e.g. As, Se, Sb, Mo, V, Cr, U, Re). It is sensitive to mobilization at pH values typical of natural waters (pH 6.5–8.5) under both oxidizing and reducing conditions. Arsenic can occur in the environment in several oxidation states (−3, 0, +3 and +5), but in natural waters it is mostly found in inorganic forms as oxyanions of trivalent arsenite [As(III)] or pentavalent arsenate [As(V)]. Organic forms of arsenic are produced by biological activity, mostly in surface waters, but are rarely quantitatively important. Organic arsenic compounds may, however, occur where waters are significantly impacted by industrial pollution.
Arsenic may be solubilized by various processes. When pH is high, arsenic may be released from surface binding sites that lose their positive charge. When water level drops and sulfide minerals are exposed to air, arsenic trapped in sulfide minerals can be released into water. When organic carbon is present in water, bacteria are fed by directly reducing As(V) to As(III) or by reducing the element at the binding site, releasing inorganic arsenic.How Does Arsenic Get into the Groundwater. Civil and Environmental Engineering. University of Maine
The aquatic transformations of arsenic are affected by pH, reduction-oxidation potential, organic matter concentration and the concentrations and forms of other elements, especially iron and manganese. The main factors are pH and the redox potential. Generally, the main forms of arsenic under oxic conditions are , , , and at pH 2, 2–7, 7–11 and 11, respectively. Under reducing conditions, is predominant at pH 2–9.
Oxidation and reduction affects the migration of arsenic in subsurface environments. Arsenite is the most stable soluble form of arsenic in reducing environments and arsenate, which is less mobile than arsenite, is dominant in oxidizing environments at neutral pH. Therefore, arsenic may be more mobile under reducing conditions. The reducing environment is also rich in organic matter which may enhance the solubility of arsenic compounds. As a result, the adsorption of arsenic is reduced and dissolved arsenic accumulates in groundwater. That is why the arsenic content is higher in reducing environments than in oxidizing environments.Zeng Zhaohua, Zhang Zhiliang (2002). "The formation of As element in groundwater and the controlling factor". Shanghai Geology 87 (3): 11–15.
The presence of sulfur is another factor that affects the transformation of arsenic in natural water. Arsenic can precipitate when metal sulfides form. In this way, arsenic is removed from the water and its mobility decreases. When oxygen is present, bacteria oxidize reduced sulfur to generate energy, potentially releasing bound arsenic.
Redox reactions involving Fe also appear to be essential factors in the fate of arsenic in aquatic systems. The reduction of iron oxyhydroxides plays a key role in the release of arsenic to water. So arsenic can be enriched in water with elevated Fe concentrations. Under oxidizing conditions, arsenic can be mobilized from pyrite or iron oxides especially at elevated pH. Under reducing conditions, arsenic can be mobilized by reductive desorption or dissolution when associated with iron oxides. The reductive desorption occurs under two circumstances. One is when arsenate is reduced to arsenite which adsorbs to iron oxides less strongly. The other results from a change in the charge on the mineral surface which leads to the desorption of bound arsenic.Thomas, Mary Ann (2007). "The Association of Arsenic With Redox Conditions, Depth, and Ground-Water Age in the Glacial Aquifer System of the Northern United States". U.S. Geological Survey, Virginia. pp. 1–18.
Some species of bacteria catalyze redox transformations of arsenic. Dissimilatory arsenate-respiring prokaryotes (DARP) speed up the reduction of As(V) to As(III). DARP use As(V) as the electron acceptor of anaerobic respiration and obtain energy to survive. Other organic and inorganic substances can be oxidized in this process. Chemoautotrophic arsenite oxidizers (CAO) and heterotrophic arsenite oxidizers (HAO) convert As(III) into As(V). CAO combine the oxidation of As(III) with the reduction of oxygen or nitrate. They use obtained energy to fix produce organic carbon from CO2. HAO cannot obtain energy from As(III) oxidation. This process may be an arsenic detoxification mechanism for the bacteria.
Equilibrium thermodynamic calculations predict that As(V) concentrations should be greater than As(III) concentrations in all but strongly reducing conditions, i.e. where sulfate reduction is occurring. However, abiotic redox reactions of arsenic are slow. Oxidation of As(III) by dissolved O2 is a particularly slow reaction. For example, Johnson and Pilson (1975) gave half-lives for the oxygenation of As(III) in seawater ranging from several months to a year. In other studies, As(V)/As(III) ratios were stable over periods of days or weeks during water sampling when no particular care was taken to prevent oxidation, again suggesting relatively slow oxidation rates. Cherry found from experimental studies that the As(V)/As(III) ratios were stable in anoxic solutions for up to 3 weeks but that gradual changes occurred over longer timescales. Sterile water samples have been observed to be less susceptible to speciation changes than non-sterile samples. Oremland found that the reduction of As(V) to As(III) in Mono Lake was rapidly catalyzed by bacteria with rate constants ranging from 0.02 to 0.3-day−1. |
Arsenic | Wood preservation in the US | Wood preservation in the US
As of 2002, US-based industries consumed 19,600 metric tons of arsenic. Ninety percent of this was used for treatment of wood with chromated copper arsenate (CCA). In 2007, 50% of the 5,280 metric tons of consumption was still used for this purpose. In the United States, the voluntary phasing-out of arsenic in production of consumer products and residential and general consumer construction products began on 31 December 2003, and alternative chemicals are now used, such as Alkaline Copper Quaternary, borates, copper azole, cyproconazole, and propiconazole.
Although discontinued, this application is also one of the most concerning to the general public. The vast majority of older pressure-treated wood was treated with CCA. CCA lumber is still in widespread use in many countries, and was heavily used during the latter half of the 20th century as a structural and outdoor building material. Although the use of CCA lumber was banned in many areas after studies showed that arsenic could leach out of the wood into the surrounding soil (from playground equipment, for instance), a risk is also presented by the burning of older CCA timber. The direct or indirect ingestion of wood ash from burnt CCA lumber has caused fatalities in animals and serious poisonings in humans; the lethal human dose is approximately 20 grams of ash. Scrap CCA lumber from construction and demolition sites may be inadvertently used in commercial and domestic fires. Protocols for safe disposal of CCA lumber are not consistent throughout the world. Widespread landfill disposal of such timber raises some concern, but other studies have shown no arsenic contamination in the groundwater. |
Arsenic | Mapping of industrial releases in the US | Mapping of industrial releases in the US
One tool that maps the location (and other information) of arsenic releases in the United States is TOXMAP. TOXMAP is a Geographic Information System (GIS) from the Division of Specialized Information Services of the United States National Library of Medicine (NLM) funded by the US Federal Government. With marked-up maps of the United States, TOXMAP enables users to visually explore data from the United States Environmental Protection Agency's (EPA) Toxics Release Inventory and Superfund Basic Research Programs. TOXMAP's chemical and environmental health information is taken from NLM's Toxicology Data Network (TOXNET),TOXNET – Databases on toxicology, hazardous chemicals, environmental health, and toxic releases. Toxnet.nlm.nih.gov. Retrieved 2011-10-24. PubMed, and from other authoritative sources. |
Arsenic | Bioremediation | Bioremediation
Physical, chemical, and biological methods have been used to remediate arsenic contaminated water. Bioremediation is said to be cost-effective and environmentally friendly. Bioremediation of ground water contaminated with arsenic aims to convert arsenite, the toxic form of arsenic to humans, to arsenate. Arsenate (+5 oxidation state) is the dominant form of arsenic in surface water, while arsenite (+3 oxidation state) is the dominant form in hypoxic to anoxic environments. Arsenite is more soluble and mobile than arsenate. Many species of bacteria can transform arsenite to arsenate in anoxic conditions by using arsenite as an electron donor. This is a useful method in ground water remediation. Another bioremediation strategy is to use plants that accumulate arsenic in their tissues via phytoremediation but the disposal of contaminated plant material needs to be considered.
Bioremediation requires careful evaluation and design in accordance with existing conditions. Some sites may require the addition of an electron acceptor while others require microbe supplementation (bioaugmentation). Regardless of the method used, only constant monitoring can prevent future contamination. |
Arsenic | Arsenic removal | Arsenic removal
Coagulation and flocculation are closely related processes common in arsenate removal from water. Due to the net negative charge carried by arsenate ions, they settle slowly or not at all due to charge repulsion. In coagulation, a positively charged coagulent such as iron and aluminum (commonly used salts: FeCl3, Fe2(SO4)3, Al2(SO4)3) neutralize the negatively charged arsenate, enable it to settle. Flocculation follows where a flocculant bridges smaller particles and allows the aggregate to precipitate out from water. However, such methods may not be efficient on arsenite as As(III) exists in uncharged arsenious acid, H3AsO3, at near-neutral pH.
The major drawbacks of coagulation and flocculation are the costly disposal of arsenate-concentrated sludge, and possible secondary contamination of environment. Moreover, coagulents such as iron may produce ion contamination that exceeds safety levels. |
Arsenic | Toxicity and precautions | Toxicity and precautions
Arsenic and many of its compounds are especially potent poisons (e.g. arsine). Small amount of arsenic can be detected by pharmacopoial methods which includes reduction of arsenic to arsenious with help of zinc and can be confirmed with mercuric chloride paper. |
Arsenic | Classification | Classification
Elemental arsenic and arsenic sulfate and trioxide compounds are classified as "toxic" and "dangerous for the environment" in the European Union under directive 67/548/EEC.
The International Agency for Research on Cancer (IARC) recognizes arsenic and inorganic arsenic compounds as group 1 carcinogens, and the EU lists arsenic trioxide, arsenic pentoxide, and arsenate salts as category 1 carcinogens.
Arsenic is known to cause arsenicosis when present in drinking water, "the most common species being arsenate [; As(V)] and arsenite [; As(III)]". |
Arsenic | Legal limits, food, and drink | Legal limits, food, and drink
In the United States since 2006, the maximum concentration in drinking water allowed by the Environmental Protection Agency (EPA) is 10 ppbArsenic Rule. U.S. Environmental Protection Agency. Adopted 22 January 2001; effective 23 January 2006. and the FDA set the same standard in 2005 for bottled water. The Department of Environmental Protection for New Jersey set a drinking water limit of 5 ppb in 2006. The IDLH (immediately dangerous to life and health) value for arsenic metal and inorganic arsenic compounds is 5 mg/m3 (5 ppb). The Occupational Safety and Health Administration has set the permissible exposure limit (PEL) to a time-weighted average (TWA) of 0.01 mg/m3 (0.01 ppb), and the National Institute for Occupational Safety and Health (NIOSH) has set the recommended exposure limit (REL) to a 15-minute constant exposure of 0.002 mg/m3 (0.002 ppb). The PEL for organic arsenic compounds is a TWA of 0.5 mg/m3. (0.5 ppb).
In 2008, based on its ongoing testing of a wide variety of American foods for toxic chemicals,Total Diet Study and Toxic Elements Program the U.S. Food and Drug Administration set the "level of concern" for inorganic arsenic in apple and pear juices at 23 ppb, based on non-carcinogenic effects, and began blocking importation of products in excess of this level; it also required recalls for non-conforming domestic products. In 2011, the national Dr. Oz television show broadcast a program highlighting tests performed by an independent lab hired by the producers. Though the methodology was disputed (it did not distinguish between organic and inorganic arsenic) the tests showed levels of arsenic up to 36 ppb. In response, the FDA tested the worst brand from the Dr. Oz show and found much lower levels. Ongoing testing found 95% of the apple juice samples were below the level of concern. Later testing by Consumer Reports showed inorganic arsenic at levels slightly above 10 ppb, and the organization urged parents to reduce consumption. In July 2013, on consideration of consumption by children, chronic exposure, and carcinogenic effect, the FDA established an "action level" of 10 ppb for apple juice, the same as the drinking water standard.
Concern about arsenic in rice in Bangladesh was raised in 2002, but at the time only Australia had a legal limit for food (one milligram per kilogram, or 1000 ppb). Concern was raised about people who were eating U.S. rice exceeding WHO standards for personal arsenic intake in 2005. In 2011, the People's Republic of China set a food standard of 150 ppb for arsenic.
In the United States in 2012, testing by separate groups of researchers at the Children's Environmental Health and Disease Prevention Research Center at Dartmouth College (early in the year, focusing on urinary levels in children) and Consumer Reports (in November) found levels of arsenic in rice that resulted in calls for the FDA to set limits.Lawmakers Urge FDA to Act on Arsenic Standards. Foodsafetynews.com (24 February 2012). Retrieved 2012-05-23. The FDA released some testing results in September 2012, and as of July 2013, is still collecting data in support of a new potential regulation. It has not recommended any changes in consumer behavior.
Consumer Reports recommended:
That the EPA and FDA eliminate arsenic-containing fertilizer, drugs, and pesticides in food production;
That the FDA establish a legal limit for food;
That industry change production practices to lower arsenic levels, especially in food for children; and
That consumers test home water supplies, eat a varied diet, and cook rice with excess water, then draining it off (reducing inorganic arsenic by about one third along with a slight reduction in vitamin content).
Evidence-based public health advocates also recommend that, given the lack of regulation or labeling for arsenic in the U.S., children should eat no more than 1.5 servings per week of rice and should not drink rice milk as part of their daily diet before age 5. They also offer recommendations for adults and infants on how to limit arsenic exposure from rice, drinking water, and fruit juice.
A 2014 World Health Organization advisory conference was scheduled to consider limits of 200–300 ppb for rice. |
Arsenic | Reducing arsenic content in rice | Reducing arsenic content in rice
thumb|300px|An improved rice cooking approach to maximise arsenic removal while preserving nutrient elements
In 2020, scientists assessed multiple preparation procedures of rice for their capacity to reduce arsenic content and preserve nutrients, recommending a procedure involving parboiling and water-absorption. |
Arsenic | Occupational exposure limits | Occupational exposure limits
CountryLimitArgentinaConfirmed human carcinogenAustraliaTWA 0.05 mg/m3 – CarcinogenBelgiumTWA 0.1 mg/m3 – CarcinogenBulgariaConfirmed human carcinogenCanadaTWA 0.01 mg/m3ColombiaConfirmed human carcinogenDenmarkTWA 0.01 mg/m3FinlandCarcinogenEgyptTWA 0.2 mg/m3HungaryCeiling concentration 0.01 mg/m3 – Skin, carcinogenIndiaTWA 0.2 mg/m3JapanGroup 1 carcinogenJordanConfirmed human carcinogenMexicoTWA 0.2 mg/m3New ZealandTWA 0.05 mg/m3 – CarcinogenNorwayTWA 0.02 mg/m3PhilippinesTWA 0.5 mg/m3PolandTWA 0.01 mg/m3SingaporeConfirmed human carcinogenSouth KoreaTWA 0.01 mg/m3Korea Occupational Safety & Health Agency . kosha.or.krSwedenTWA 0.01 mg/m3ThailandTWA 0.5 mg/m3TurkeyTWA 0.5 mg/m3United KingdomTWA 0.1 mg/m3United StatesTWA 0.01 mg/m3VietnamConfirmed human carcinogen |
Arsenic | Ecotoxicity | Ecotoxicity
Arsenic is bioaccumulative in many organisms, marine species in particular, but it does not appear to biomagnify significantly in food webs. In polluted areas, plant growth may be affected by root uptake of arsenate, which is a phosphate analog and therefore readily transported in plant tissues and cells. In polluted areas, uptake of the more toxic arsenite ion (found more particularly in reducing conditions) is likely in poorly-drained soils. |
Arsenic | Toxicity in animals | Toxicity in animals
CompoundAnimalLD50RouteArsenicRat763 mg/kgoralArsenicMouse145 mg/kgoralCalcium arsenateRat20 mg/kgoralCalcium arsenateMouse794 mg/kgoralCalcium arsenateRabbit50 mg/kgoralCalcium arsenateDog38 mg/kgoralLead arsenateRabbit75 mg/kgoral
CompoundAnimalLD50RouteArsenic trioxide (As(III))Mouse26 mg/kgoralArsenite (As(III))Mouse8 mg/kgimArsenate (As(V))Mouse21 mg/kgimMMA (As(III))Hamster2 mg/kgipMMA (As(V))Mouse916 mg/kgoralDMA (As(V))Mouse648 mg/kgoralim = injected intramuscularly
ip = administered intraperitoneally |
Arsenic | Biological mechanism | Biological mechanism
Arsenic's toxicity comes from the affinity of arsenic(III) oxides for thiols. Thiols, in the form of cysteine residues and cofactors such as lipoic acid and coenzyme A, are situated at the active sites of many important enzymes.
Arsenic disrupts ATP production through several mechanisms. At the level of the citric acid cycle, arsenic inhibits lipoic acid, which is a cofactor for pyruvate dehydrogenase. By competing with phosphate, arsenate uncouples oxidative phosphorylation, thus inhibiting energy-linked reduction of NAD+, mitochondrial respiration and ATP synthesis. Hydrogen peroxide production is also increased, which, it is speculated, has potential to form reactive oxygen species and oxidative stress. These metabolic interferences lead to death from multi-system organ failure. The organ failure is presumed to be from necrotic cell death, not apoptosis, since energy reserves have been too depleted for apoptosis to occur. |
Arsenic | Exposure risks and remediation | Exposure risks and remediation
Occupational exposure and arsenic poisoning may occur in people working in industries involving the use of inorganic arsenic and its compounds, such as wood preservation, glass production, nonferrous metal alloys, and electronic semiconductor manufacturing. Inorganic arsenic is also found in coke oven emissions associated with the smelter industry.
The conversion between As(III) and As(V) is a large factor in arsenic environmental contamination. According to Croal, Gralnick, Malasarn and Newman, "[the] understanding [of] what stimulates As(III) oxidation and/or limits As(V) reduction is relevant for bioremediation of contaminated sites (Croal). The study of chemolithoautotrophic As(III) oxidizers and the heterotrophic As(V) reducers can help the understanding of the oxidation and/or reduction of arsenic. |
Arsenic | Treatment | Treatment
Treatment of chronic arsenic poisoning is possible. British anti-lewisite (dimercaprol) is prescribed in doses of 5 mg/kg up to 300 mg every 4 hours for the first day, then every 6 hours for the second day, and finally every 8 hours for 8 additional days. However the USA's Agency for Toxic Substances and Disease Registry (ATSDR) states that the long-term effects of arsenic exposure cannot be predicted. Blood, urine, hair, and nails may be tested for arsenic; however, these tests cannot foresee possible health outcomes from the exposure. Long-term exposure and consequent excretion through urine has been linked to bladder and kidney cancer in addition to cancer of the liver, prostate, skin, lungs, and nasal cavity. |
Arsenic | Footnotes | Footnotes |
Arsenic | See also | See also
Aqua Tofana
Arsenic and Old Lace
Grainger challenge
Hypothetical types of biochemistry |
Arsenic | References | References |
Arsenic | Bibliography | Bibliography
|
Arsenic | Further reading | Further reading
|
Arsenic | External links | External links
WHO fact sheet on arsenic
Arsenic Cancer Causing Substances, U.S. National Cancer Institute.
CTD's Arsenic page and CTD's Arsenicals page from the Comparative Toxicogenomics Database
Contaminant Focus: Arsenic by the EPA.
Environmental Health Criteria for Arsenic and Arsenic Compounds, 2001 by the WHO.
National Institute for Occupational Safety and Health – Arsenic Page
Category:Chemical elements
Category:Metalloids
Category:Semimetals
Category:Hepatotoxins
Category:Pnictogens
Category:Endocrine disruptors
Category:IARC Group 1 carcinogens
Category:Trigonal minerals
Category:Minerals in space group 166
Category:Teratogens
Category:Fetotoxicants
Category:Suspected testicular toxicants
Category:Native element minerals
Category:Chemical elements with rhombohedral structure |
Arsenic | Table of Content | <!-- {{cs1 config, Characteristics, Physical characteristics, Isotopes, Chemistry, Compounds, Inorganic compounds, Alloys, Organoarsenic compounds, Occurrence and production, History, Applications, Agricultural, Medical use, Alloys, Military, Other uses, Biological role, Bacteria, Potential role in higher animals, Heredity, Biomethylation, Environmental issues, Exposure, Occurrence in drinking water, San Pedro de Atacama, Hazard maps for contaminated groundwater, Redox transformation of arsenic in natural waters, Wood preservation in the US, Mapping of industrial releases in the US, Bioremediation, Arsenic removal, Toxicity and precautions, Classification, Legal limits, food, and drink, Reducing arsenic content in rice, Occupational exposure limits, Ecotoxicity, Toxicity in animals, Biological mechanism, Exposure risks and remediation, Treatment, Footnotes, See also, References, Bibliography, Further reading, External links |
Antimony | distinguish | Antimony is a chemical element; it has symbol Sb () and atomic number 51. A lustrous grey metal or metalloid, it is found in nature mainly as the sulfide mineral stibnite (). Antimony compounds have been known since ancient times and were powdered for use as medicine and cosmetics, often known by the Arabic name kohl.David Kimhi's Commentary on Isaiah 4:30 and I Chronicles 29:2; Hebrew: פוך/כְּחֻל, Aramaic: כּוּחְלִי/צדידא; Arabic: كحل, and which can also refer to antimony trisulfide. See also Z. Dori, Antimony and Henna (Heb. הפוך והכופר), Jerusalem 1983 (Hebrew). The earliest known description of this metalloid in the West was written in 1540 by Vannoccio Biringuccio.
China is the largest producer of antimony and its compounds, with most production coming from the Xikuangshan Mine in Hunan. The industrial methods for refining antimony from stibnite are roasting followed by reduction with carbon, or direct reduction of stibnite with iron.
The most common applications for metallic antimony are in alloys with lead and tin, which have improved properties for solders, bullets, and plain bearings. It improves the rigidity of lead-alloy plates in lead–acid batteries. Antimony trioxide is a prominent additive for halogen-containing flame retardants. Antimony is used as a dopant in semiconductor devices. |
Antimony | Characteristics | Characteristics |
Antimony | Properties | Properties
thumb|left|alt=A clear vial containing small chunks of a slightly lustrous black solid, labeled "Sb".|A vial containing the metallic allotrope of antimony
left|thumb|alt=An irregular piece of silvery stone with spots of variation in luster and shade.|Native antimony with oxidation products
thumb|left|Crystal structure common to Sb, AsSb and gray As
Antimony is a member of group 15 of the periodic table, one of the elements called pnictogens, and has an electronegativity of 2.05. In accordance with periodic trends, it is more electronegative than tin or bismuth, and less electronegative than tellurium or arsenic. Antimony is stable in air at room temperature but, if heated, it reacts with oxygen to produce antimony trioxide,.Wiberg and Holleman, p. 758
Antimony is a silvery, lustrous gray metalloid with a Mohs scale hardness of 3, which is too soft to mark hard objects. Coins of antimony were issued in China's Guizhou in 1931; durability was poor, and minting was soon discontinued because of its softness and toxicity. Antimony is resistant to attack by acids.
The only stable allotrope of antimony under standard conditions is metallic, brittle, silver-white, and shiny. It crystallises in a trigonal cell, isomorphic with bismuth and the gray allotrope of arsenic, and is formed when molten antimony is cooled slowly. Amorphous black antimony is formed upon rapid cooling of antimony vapor, and is only stable as a thin film (thickness in nanometres); thicker samples spontaneously transform into the metallic form. It oxidizes in air and may ignite spontaneously. At 100 °C, it gradually transforms into the stable form. The supposed yellow allotrope of antimony, generated only by oxidation of stibine () at −90 °C, is also impure and not a true allotrope; above this temperature and in ambient light, it transforms into the more stable black allotrope., [ pp. 50–51] A rare explosive form of antimony can be formed from the electrolysis of antimony trichloride, but it always contains appreciable chlorine and is not really an antimony allotrope. When scratched with a sharp implement, an exothermic reaction occurs and white fumes are given off as metallic antimony forms; when rubbed with a pestle in a mortar, a strong detonation occurs.
Elemental antimony adopts a layered structure (space group Rm No. 166) whose layers consist of fused, ruffled, six-membered rings. The nearest and next-nearest neighbors form an irregular octahedral complex, with the three atoms in each double layer slightly closer than the three atoms in the next. This relatively close packing leads to a high density of 6.697 g/cm3, but the weak bonding between the layers leads to the low hardness and brittleness of antimony. |
Antimony | Isotopes | Isotopes
Antimony has two stable isotopes: with a natural abundance of 57.36% and with a natural abundance of 42.64%. It also has 35 radioisotopes, of which the longest-lived is with a half-life of 2.75 years. In addition, 29 metastable states have been characterized. The most stable of these is with a half-life of 5.76 days. Isotopes that are lighter than the stable tend to decay by β+ decay, and those that are heavier tend to decay by β− decay, with some exceptions. Antimony is the lightest element to have an isotope with an alpha decay branch, excluding and other light nuclides with beta-delayed alpha emission. |
Antimony | Occurrence | Occurrence
thumb|Stibnite, China CM29287 Carnegie Museum of Natural History specimen on display in Hillman Hall of Minerals and Gems|alt=
The abundance of antimony in the Earth's crust is estimated at 0.2 parts per million,Greenwood and Earnshaw, p. 548 comparable to thallium at 0.5 ppm and silver at 0.07 ppm. It is the 63rd most abundant element in the crust. Even though this element is not abundant, it is found in more than 100 mineral species.Antimony minerals. mindat.org Antimony is sometimes found natively (e.g. on Antimony Peak), but more frequently it is found in the sulfide stibnite () which is the predominant ore mineral. |
Antimony | Compounds | Compounds
Antimony compounds are often classified according to their oxidation state: Sb(III) and Sb(V). The +5 oxidation state is more common.Greenwood and Earnshaw, p. 553 |
Antimony | Oxides and hydroxides | Oxides and hydroxides
Antimony trioxide is formed when antimony is burnt in air. In the gas phase, the molecule of the compound is , but it polymerizes upon condensing. Antimony pentoxide () can be formed only by oxidation with concentrated nitric acid. Antimony also forms a mixed-valence oxide, antimony tetroxide (), which features both Sb(III) and Sb(V). Unlike oxides of phosphorus and arsenic, these oxides are amphoteric, do not form well-defined oxoacids, and react with acids to form antimony salts.
Antimonous acid is unknown, but the conjugate base sodium antimonite () forms upon fusing sodium oxide and .Wiberg and Holleman, p. 763 Transition metal antimonites are also known. Antimonic acid exists only as the hydrate , forming salts as the antimonate anion . When a solution containing this anion is dehydrated, the precipitate contains mixed oxides.
The most important antimony ore is stibnite (). Other sulfide minerals include pyrargyrite (), zinkenite, jamesonite, and boulangerite.Wiberg and Holleman, p. 757 Antimony pentasulfide is non-stoichiometric, which features antimony in the +3 oxidation state and S–S bonds. Several thioantimonides are known, such as and . |
Antimony | Halides | Halides
Antimony forms two series of halides: and . The trihalides , , , and are all molecular compounds having trigonal pyramidal molecular geometry. The trifluoride is prepared by the reaction of antimony trioxide with hydrofluoric acid:Wiberg and Holleman, pp. 761–762
It is Lewis acidic and readily accepts fluoride ions to form the complex anions and . Molten antimony trifluoride is a weak electrical conductor. The trichloride is prepared by dissolving stibnite in hydrochloric acid:
Arsenic sulfides are not readily attacked by the hydrochloric acid, so this method offers a route to As-free Sb.
thumb|upright|left|Structure of gaseous
The pentahalides and have trigonal bipyramidal molecular geometry in the gas phase, but in the liquid phase, is polymeric, whereas is monomeric.Wiberg and Holleman, p. 761 Antimony pentafluoride is a powerful Lewis acid used to make the superacid fluoroantimonic acid ().
Oxyhalides are more common for antimony than for arsenic and phosphorus. Antimony trioxide dissolves in concentrated acid to form oxoantimonyl compounds such as SbOCl and .Wiberg and Holleman, p. 764 |
Antimony | Antimonides, hydrides, and organoantimony compounds | Antimonides, hydrides, and organoantimony compounds
Compounds in this class generally are described as derivatives of . Antimony forms antimonides with metals, such as indium antimonide (InSb) and silver antimonide ().Wiberg and Holleman, p. 760 The alkali metal and zinc antimonides, such as and , are more reactive. Treating these antimonides with acid produces the highly unstable gas stibine, :
Stibine can also be produced by treating salts with hydride reagents such as sodium borohydride. Stibine decomposes spontaneously at room temperature. Because stibine has a positive heat of formation, it is thermodynamically unstable and thus antimony does not react with hydrogen directly.Greenwood and Earnshaw, p. 558
Organoantimony compounds are typically prepared by alkylation of antimony halides with Grignard reagents.Elschenbroich, C. (2006) "Organometallics". Wiley-VCH: Weinheim. A large variety of compounds are known with both Sb(III) and Sb(V) centers, including mixed chloro-organic derivatives, anions, and cations. Examples include triphenylstibine () and pentaphenylantimony ().Greenwood and Earnshaw, p. 598 |
Antimony | History | History
upright=0.3|thumb|alt=An unshaded circle surmounted by a cross.|One of the alchemical symbols for antimony
Antimony(III) sulfide, , was recognized in predynastic Egypt as an eye cosmetic (kohl) as early as about 3100 BC, when the cosmetic palette was invented.
An artifact, said to be part of a vase, made of antimony dating to about 3000 BC was found at Telloh, Chaldea (part of present-day Iraq), and a copper object plated with antimony dating between 2500 BC and 2200 BC has been found in Egypt. Austen, at a lecture by Herbert Gladstone in 1892, commented that "we only know of antimony at the present day as a highly brittle and crystalline metal, which could hardly be fashioned into a useful vase, and therefore this remarkable 'find' (artifact mentioned above) must represent the lost art of rendering antimony malleable."
The British archaeologist Roger Moorey was unconvinced the artifact was indeed a vase, mentioning that Selimkhanov, after his analysis of the Tello object (published in 1975), "attempted to relate the metal to Transcaucasian natural antimony" (i.e. native metal) and that "the antimony objects from Transcaucasia are all small personal ornaments." This weakens the evidence for a lost art "of rendering antimony malleable".
The Roman scholar Pliny the Elder described several ways of preparing antimony sulfide for medical purposes in his treatise Natural History, around 77 AD. Pliny the Elder also made a distinction between "male" and "female" forms of antimony; the male form is probably the sulfide, while the female form, which is superior, heavier, and less friable, has been suspected to be native metallic antimony.Pliny, Natural history, 33.33; W.H.S. Jones, the Loeb Classical Library translator, supplies a note suggesting the identifications.
The Greek naturalist Pedanius Dioscorides mentioned that antimony sulfide could be roasted by heating by a current of air. It is thought that this produced metallic antimony.
thumb|right|upright=0.9|The Italian metallurgist Vannoccio Biringuccio described a procedure to isolate antimony.
Antimony was frequently described in alchemical manuscripts, including the Summa Perfectionis of Pseudo-Geber, written around the 14th century. A description of a procedure for isolating antimony is later given in the 1540 book De la pirotechnia by Vannoccio Biringuccio,Vannoccio Biringuccio, De la Pirotechnia (Venice (Italy): Curtio Navo e fratelli, 1540), Book 2, chapter 3: Del antimonio & sua miniera, Capitolo terzo (On antimony and its ore, third chapter), pp. 27–28. [Note: Only every second page of this book is numbered, so the relevant passage is to be found on the 74th and 75th pages of the text.] (in Italian) predating the more famous 1556 book by Agricola, De re metallica. In this context Agricola has been often incorrectly credited with the discovery of metallic antimony. The book Currus Triumphalis Antimonii (The Triumphal Chariot of Antimony), describing the preparation of metallic antimony, was published in Germany in 1604. It was purported to be written by a Benedictine monk, writing under the name Basilius Valentinus in the 15th century; if it were authentic, which it is not, it would predate Biringuccio.
The metal antimony was known to German chemist Andreas Libavius in 1615 who obtained it by adding iron to a molten mixture of antimony sulfide, salt and potassium tartrate. This procedure produced antimony with a crystalline or starred surface.
With the advent of challenges to phlogiston theory, it was recognized that antimony is an element forming sulfides, oxides, and other compounds, as do other metals.
The first discovery of naturally occurring pure antimony in the Earth's crust was described by the Swedish scientist and local mine district engineer Anton von Swab in 1783; the type-sample was collected from the Sala Silver Mine in the Bergslagen mining district of Sala, Västmanland, Sweden. |
Antimony | Etymology | Etymology
The medieval Latin form, from which the modern languages and late Byzantine Greek take their names for antimony, is . The origin of that is uncertain, and all suggestions have some difficulty either of form or interpretation. The popular etymology, from ἀντίμοναχός anti-monachos or French , would mean "monk-killer", which is explained by the fact that many early alchemists were monks, and some antimony compounds were poisonous. Fernando connects the proposed etymology to the story of "Basil Valentine", although antimonium is found two centuries before Valentine's time.
Another popular etymology is the hypothetical Greek word ἀντίμόνος antimonos, "against aloneness", explained as "not found as metal", or "not found unalloyed"."Antimony" in Kirk-Othmer Encyclopedia of Chemical Technology, 5th ed. 2004. However, ancient Greek would more naturally express the pure negative as α- ("not")., which considers the derivation a "popular etymology". Edmund Oscar von Lippmann conjectured a hypothetical Greek word ανθήμόνιον anthemonion, which would mean "floret", and cites several examples of related Greek words (but not that one) which describe chemical or biological efflorescence.von Lippmann, Edmund Oscar (1919) Entstehung und Ausbreitung der Alchemie, teil 1. Berlin: Julius Springer (in German). pp. 642–5
The early uses of antimonium include the translations, in 1050–1100, by Constantine the African of Arabic medical treatises. Several authorities believe antimonium is a scribal corruption of some Arabic form; Meyerhof derives it from ithmid;Meyerhof as quoted in , asserts that ithmid or athmoud became corrupted in the medieval "traductions barbaro-latines". The OED asserts some Arabic form is the origin, and if ithmid is the root, posits athimodium, atimodium, atimonium as intermediates. other possibilities include athimar, the Arabic name of the metalloid, and a hypothetical as-stimmi, derived from or parallel to the Greek.
The standard chemical symbol for antimony (Sb) is credited to Jöns Jakob Berzelius, who derived the abbreviation from stibium.Jöns Jacob Berzelius, "Essay on the cause of chemical proportions, and on some circumstances relating to them: together with a short and easy method of expressing them," Annals of Philosophy, vol. 2, pages 443–454 (1813) and vol. 3, pages 51–62, 93–106, 244–255, 353–364 (1814). On [ p. 52], Berzelius lists the symbol for antimony as "St"; however, starting from [ p. 248], Berzelius consistently uses the symbol "Sb" instead.
The ancient words for antimony mostly have, as their chief meaning, kohl, the sulfide of antimony.
The Egyptians called antimony mśdmt or stm.
The Arabic word for the substance, as opposed to the cosmetic, can appear as ithmid, athmoud, othmod, or uthmod. Littré suggests the first form, which is the earliest, derives from stimmida, an accusative for stimmi. The Greek word στίμμι (stimmi) is used by Attic tragic poets of the 5th century BC, and is possibly a loan word from Arabic or from Egyptian stm. |
Antimony | Production | Production |
Antimony | Process | Process
The extraction of antimony from ores depends on the quality and composition of the ore. Most antimony is mined as the sulfide; lower-grade ores are concentrated by froth flotation, while higher-grade ores are heated to 500–600 °C, the temperature at which stibnite melts and separates from the gangue minerals. Antimony can be isolated from the crude antimony sulfide by reduction with scrap iron:
The sulfide is converted to an oxide by roasting. The product is further purified by vaporizing the volatile antimony(III) oxide, which is recovered. This sublimate is often used directly for the main applications, impurities being arsenic and sulfide., [ p. 45] Antimony is isolated from the oxide by a carbothermal reduction:
The lower-grade ores are reduced in blast furnaces while the higher-grade ores are reduced in reverberatory furnaces.
thumb|upright=1.6|World antimony output in 2010
thumb|upright=1.3|World production trend of antimony |
Antimony | Top producers and production volumes | Top producers and production volumes
In 2022, according to the US Geological Survey, China accounted for 54.5% of total antimony production, followed in second place by Russia with 18.2% and Tajikistan with 15.5%.
+Antimony mining in 2022 Country Tonnes % of total60,00054.520,00018.217,00015.54,0003.64,0003.6Top 5105,00095.5Total world110,000100.0
Chinese production of antimony is expected to decline in the future as mines and smelters are closed down by the government as part of pollution control. Especially due to an environmental protection law having gone into effect in January 2015 and revised "Emission Standards of Pollutants for Stanum, Antimony, and Mercury" having gone into effect, hurdles for economic production are higher.
Reported production of antimony in China has fallen and is unlikely to increase in the coming years, according to the Roskill report. No significant antimony deposits in China have been developed for about ten years, and the remaining economic reserves are being rapidly depleted. |
Antimony | Reserves | Reserves
+World antimony reserves in 2022 Country Reserves (tonnes) 350,000350,000310,000260,000140,000120,000100,00078,00060,00060,00050,000Total world>1,800,000 |
Antimony | Supply risk | Supply risk
For antimony-importing regions, such as Europe and the U.S., antimony is considered to be a critical mineral for industrial manufacturing that is at risk of supply chain disruption. With global production coming mainly from China (74%), Tajikistan (8%), and Russia (4%), these sources are critical to supply.
European Union: Antimony is considered a critical raw material for defense, automotive, construction and textiles. The E.U. sources are 100% imported, coming mainly from Turkey (62%), Bolivia (20%) and Guatemala (7%).
United Kingdom: The British Geological Survey's 2015 risk list ranks antimony second highest (after rare earth elements) on the relative supply risk index.
United States: Antimony is a mineral commodity considered critical to the economic and national security. In 2022, no antimony was mined in the U.S. |
Antimony | Applications | Applications
Approximately 48% of antimony is consumed in flame retardants, 33% in lead–acid batteries, and 8% in plastics. |
Antimony | Flame retardants | Flame retardants
Antimony is mainly used as the trioxide for flame-proofing compounds, always in combination with halogenated flame retardants except in halogen-containing polymers. The flame retarding effect of antimony trioxide is produced by the formation of halogenated antimony compounds, which react with hydrogen atoms, and probably also with oxygen atoms and OH radicals, thus inhibiting fire. Markets for these flame-retardants include children's clothing, toys, aircraft, and automobile seat covers. They are also added to polyester resins in fiberglass composites for such items as light aircraft engine covers. The resin will burn in the presence of an externally generated flame, but will extinguish when the external flame is removed.Grund, Sabina C.; Hanusch, Kunibert; Breunig, Hans J.; Wolf, Hans Uwe (2006) "Antimony and Antimony Compounds" in Ullmann's Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim. |
Antimony | Alloys | Alloys
Antimony forms a highly useful alloy with lead, increasing its hardness and mechanical strength. When casting it increases fluidity of the melt and reduces shrinkage during cooling. For most applications involving lead, varying amounts of antimony are used as alloying metal. In lead–acid batteries, this addition improves plate strength and charging characteristics. For sailboats, lead keels are used to provide righting moment, ranging from 600 lbs to over 200 tons for the largest sailing superyachts; to improve hardness and tensile strength of the lead keel, antimony is mixed with lead between 2% and 5% by volume. Antimony is used in antifriction alloys (such as Babbitt metal), in bullets and lead shot, electrical cable sheathing, type metal (for example, for linotype printing machines), solder (some "lead-free" solders contain 5% Sb), in pewter, and in hardening alloys with low tin content in the manufacturing of organ pipes. |
Antimony | Other applications | Other applications
thumb|upright|InSb infrared detector manufactured by Mullard in the 1960s
Three other applications consume nearly all the rest of the world's supply. One application is as a stabilizer and catalyst for the production of polyethylene terephthalate. Another is as a fining agent to remove microscopic bubbles in glass, mostly for TV screens antimony ions interact with oxygen, suppressing the tendency of the latter to form bubbles. The third application is pigments.
In the 1990s antimony was increasingly being used in semiconductors as a dopant in n-type silicon wafers for diodes, infrared detectors, and Hall-effect devices. In the 1950s, the emitters and collectors of n-p-n alloy junction transistors were doped with tiny beads of a lead-antimony alloy. Indium antimonide (InSb) is used as a material for mid-infrared detectors.
The material is used as for phase-change memory, a type of computer memory.
Biology and medicine have few uses for antimony. Treatments containing antimony, known as antimonials, are used as emetics. Antimony compounds are used as antiprotozoan drugs. Potassium antimonyl tartrate, or tartar emetic, was once used as an anti-schistosomal drug from 1919 on. It was subsequently replaced by praziquantel. Antimony and its compounds are used in several veterinary preparations, such as anthiomaline and lithium antimony thiomalate, as a skin conditioner in ruminants. Antimony has a nourishing or conditioning effect on keratinized tissues in animals.
Antimony-based drugs, such as meglumine antimoniate, are also considered the drugs of choice for treatment of leishmaniasis. Early treatments used antimony(III) species (trivalent antimonials), but in 1922 Upendranath Brahmachari invented a much safer antimony(V) drug, and since then so-called pentavalent antimonials have been the standard first-line treatment. However, Leishmania strains in Bihar and neighboring regions have developed resistance to antimony. Elemental antimony as an antimony pill was once used as a medicine. It could be reused by others after ingestion and elimination.
Antimony(III) sulfide is used in the heads of some safety matches. Antimony sulfides help to stabilize the friction coefficient in automotive brake pad materials. Antimony is used in bullets, bullet tracers, paint, glass art, and as an opacifier in enamel. Antimony-124 is used together with beryllium in neutron sources; the gamma rays emitted by antimony-124 initiate the photodisintegration of beryllium. The emitted neutrons have an average energy of 24 keV. Natural antimony is used in startup neutron sources.
The powder derived from crushed antimony sulfide (kohl) has been used for millennia as an eye cosmetic. Historically it was applied to the eyes with a metal rod and with one's spittle, and was thought by the ancients to aid in curing eye infections. The practice is still seen in Yemen and in other Muslim countries. |
Antimony | Precautions | Precautions
Antimony and many of its compounds are toxic, and the effects of antimony poisoning are similar to arsenic poisoning. The toxicity of antimony is far lower than that of arsenic; this might be caused by the significant differences of uptake, metabolism and excretion between arsenic and antimony. The uptake of antimony(III) or antimony(V) in the gastrointestinal tract is at most 20%. Antimony(V) is not quantitatively reduced to antimony(III) in the cell (in fact antimony(III) is oxidised to antimony(V) instead).
Since methylation of antimony does not occur, the excretion of antimony(V) in urine is the main way of elimination. Like arsenic, the most serious effect of acute antimony poisoning is cardiotoxicity and the resulting myocarditis; however, it can also manifest as Adams–Stokes syndrome, which arsenic does not. Reported cases of intoxication by antimony equivalent to 90 mg antimony potassium tartrate dissolved from enamel has been reported to show only short term effects. An intoxication with 6 g of antimony potassium tartrate was reported to result in death after three days.
Inhalation of antimony dust is harmful and in certain cases may be fatal; in small doses, antimony causes headaches, dizziness, and depression. Larger doses such as prolonged skin contact may cause dermatitis, or damage the kidneys and the liver, causing violent and frequent vomiting, leading to death in a few days.
Antimony is incompatible with strong oxidizing agents, strong acids, halogen acids, chlorine, or fluorine. It should be kept away from heat.Antimony MSDS. Baker
Antimony leaches from polyethylene terephthalate (PET) bottles into liquids. While levels observed for bottled water are below drinking water guidelines, fruit juice concentrates (for which no guidelines are established) produced in the UK were found to contain up to 44.7 μg/L of antimony, well above the EU limits for tap water of 5 μg/L. The guidelines are:
World Health Organization: 20 μg/L
Japan: 15 μg/LWakayama, Hiroshi (2003) "Revision of Drinking Water Standards in Japan", Ministry of Health, Labor and Welfare (Japan); Table 2, p. 84
United States Environmental Protection Agency, Health Canada and the Ontario Ministry of Environment: 6 μg/LScreening assessment antimony-containing substances. Health Canada. July 2020.
EU and German Federal Ministry of Environment: 5 μg/L
The tolerable daily intake (TDI) proposed by WHO is 6 μg antimony per kilogram of body weight. The immediately dangerous to life or health (IDLH) value for antimony is 50 mg/m3. |
Antimony | Toxicity | Toxicity
Certain compounds of antimony appear to be toxic, particularly antimony trioxide and antimony potassium tartrate. Effects may be similar to arsenic poisoning. Occupational exposure may cause respiratory irritation, pneumoconiosis, antimony spots on the skin, gastrointestinal symptoms, and cardiac arrhythmias. In addition, antimony trioxide is potentially carcinogenic to humans.
Adverse health effects have been observed in humans and animals following inhalation, oral, or dermal exposure to antimony and antimony compounds. Antimony toxicity typically occurs either due to occupational exposure, during therapy or from accidental ingestion. It is unclear if antimony can enter the body through the skin. The presence of low levels of antimony in saliva may also be associated with dental decay. |
Antimony | Notes | Notes |
Antimony | References | References |
Antimony | Cited sources | Cited sources |
Antimony | External links | External links
Public Health Statement for Antimony
International Antimony Association vzw (i2a)
Chemistry in its element podcast (MP3) from the Royal Society of Chemistry's Chemistry World: Antimony
Antimony at The Periodic Table of Videos (University of Nottingham)
CDC – NIOSH Pocket Guide to Chemical Hazards – Antimony
Antimony Mineral data and specimen images
Category:Chemical elements
Category:Metalloids
Category:Native element minerals
Category:Nuclear materials
Category:Pnictogens
Category:Trigonal minerals
Category:Minerals in space group 166
Category:Chemical elements with rhombohedral structure |
Antimony | Table of Content | distinguish, Characteristics, Properties, Isotopes, Occurrence, Compounds, Oxides and hydroxides, Halides, Antimonides, hydrides, and organoantimony compounds, History, Etymology, Production, Process, Top producers and production volumes, Reserves, Supply risk, Applications, Flame retardants, Alloys, Other applications, Precautions, Toxicity, Notes, References, Cited sources, External links |
Actinium | Distinguish | Actinium is a chemical element; it has symbol Ac and atomic number 89. It was discovered by Friedrich Oskar Giesel in 1902, who gave it the name emanium; the element got its name by being wrongly identified with a substance André-Louis Debierne found in 1899 and called actinium. The actinide series, a set of 15 elements between actinium and lawrencium in the periodic table, are named for actinium. Together with polonium, radium, and radon, actinium was one of the first non-primordial radioactive elements to be discovered.
A soft, silvery-white radioactive metal, actinium reacts rapidly with oxygen and moisture in air forming a white coating of actinium oxide that prevents further oxidation. As with most lanthanides and many actinides, actinium assumes oxidation state +3 in nearly all its chemical compounds. Actinium is found only in traces in uranium and thorium ores as the isotope 227Ac, which decays with a half-life of 21.772 years, predominantly emitting beta and sometimes alpha particles, and 228Ac, which is beta active with a half-life of 6.15 hours. One tonne of natural uranium in ore contains about 0.2 milligrams of actinium-227, and one tonne of thorium contains about 5 nanograms of actinium-228. The close similarity of physical and chemical properties of actinium and lanthanum makes separation of actinium from the ore impractical. Instead, the element is prepared, in milligram amounts, by the neutron irradiation of in a nuclear reactor. Owing to its scarcity, high price and radioactivity, actinium has no significant industrial use. Its current applications include a neutron source and an agent for radiation therapy. |
Actinium | History | History
André-Louis Debierne, a French chemist, announced the discovery of a new element in 1899. He separated it from pitchblende residues left by Marie and Pierre Curie after they had extracted radium. In 1899, Debierne described the substance as similar to titanium and (in 1900) as similar to thorium. Friedrich Oskar Giesel found in 1902 a substance similar to lanthanum and called it "emanium" in 1904. After a comparison of the substances' half-lives determined by Debierne, Harriet Brooks in 1904, and Otto Hahn and Otto Sackur in 1905, Debierne's chosen name for the new element was retained because it had seniority, despite the contradicting chemical properties he claimed for the element at different times.
Articles published in the 1970s and later suggest that Debierne's results published in 1904 conflict with those reported in 1899 and 1900. Furthermore, the now-known chemistry of actinium precludes its presence as anything other than a minor constituent of Debierne's 1899 and 1900 results; in fact, the chemical properties he reported make it likely that he had, instead, accidentally identified protactinium, which would not be discovered for another fourteen years, only to have it disappear due to its hydrolysis and adsorption onto his laboratory equipment. This has led some authors to advocate that Giesel alone should be credited with the discovery. A less confrontational vision of scientific discovery is proposed by Adloff. He suggests that hindsight criticism of the early publications should be mitigated by the then nascent state of radiochemistry: highlighting the prudence of Debierne's claims in the original papers, he notes that nobody can contend that Debierne's substance did not contain actinium. Debierne, who is now considered by the vast majority of historians as the discoverer, lost interest in the element and left the topic. Giesel, on the other hand, can rightfully be credited with the first preparation of radiochemically pure actinium and with the identification of its atomic number 89.
The name actinium originates from the Ancient Greek aktis, aktinos (ακτίς, ακτίνος), meaning beam or ray. Its symbol Ac is also used in abbreviations of other compounds that have nothing to do with actinium, such as acetyl, acetate and sometimes acetaldehyde. |
Actinium | Properties | Properties
Actinium is a soft, silvery-white,Actinium, in Encyclopædia Britannica, 15th edition, 1995, p. 70 radioactive, metallic element. Its estimated shear modulus is similar to that of lead.Seitz, Frederick and Turnbull, David (1964) Solid state physics: advances in research and applications. Academic Press. pp. 289–291 Owing to its strong radioactivity, actinium glows in the dark with a pale blue light, which originates from the surrounding air ionized by the emitted energetic particles. Actinium has similar chemical properties to lanthanum and other lanthanides, and therefore these elements are difficult to separate when extracting from uranium ores. Solvent extraction and ion chromatography are commonly used for the separation.
The first element of the actinides, actinium gave the set its name, much as lanthanum had done for the lanthanides. The actinides are much more diverse than the lanthanides and therefore it was not until 1945 that the most significant change to Dmitri Mendeleev's periodic table since the recognition of the lanthanides, the introduction of the actinides, was generally accepted after Glenn T. Seaborg's research on the transuranium elements (although it had been proposed as early as 1892 by British chemist Henry Bassett).
Actinium reacts rapidly with oxygen and moisture in air forming a white coating of actinium oxide that impedes further oxidation. As with most lanthanides and actinides, actinium exists in the oxidation state +3, and the Ac3+ ions are colorless in solutions. The oxidation state +3 originates from the [Rn] 6d17s2 electronic configuration of actinium, with three valence electrons that are easily donated to give the stable closed-shell structure of the noble gas radon. Although the 5f orbitals are unoccupied in an actinium atom, it can be used as a valence orbital in actinium complexes and hence it is generally considered the first 5f element by authors working on it. Ac3+ is the largest of all known tripositive ions and its first coordination sphere contains approximately 10.9 ± 0.5 water molecules. |
Actinium | Chemical compounds | Chemical compounds
Due to actinium's intense radioactivity, only a limited number of actinium compounds are known. These include: AcF3, AcCl3, AcBr3, AcOF, AcOCl, AcOBr, Ac2S3, Ac2O3, AcPO4 and Ac(NO3)3. They all contain actinium in the oxidation state +3. In particular, the lattice constants of the analogous lanthanum and actinium compounds differ by only a few percent.
Formula color symmetry space group No Pearson symbol a (pm) b (pm) c (pm) Z density, g/cm3 Ac silvery fcc Fmm 225 cF4 531.1 531.1 531.1 4 10.07 AcH2unknown cubic Fmm 225 cF12 567 567 567 4 8.35 Ac2O3 white trigonal Pm1 164 hP5 408 408 630 1 9.18 Ac2S3 black cubic I3d 220 cI28 778.56 778.56 778.56 4 6.71 AcF3 whiteMeyer, p. 71 hexagonal Pc1 165 hP24 741 741 755 6 7.88 AcCl3white hexagonal P63/m 165 hP8 764 764 456 2 4.8 AcBr3 white hexagonal P63/m 165 hP8 764 764 456 2 5.85 AcOF white cubic Fmm 593.1 8.28 AcOClwhite tetragonal 424 424 707 7.23 AcOBrwhite tetragonal 427 427 740 7.89 AcPO4·0.5H2Ounknown hexagonal 721 721 664 5.48
Here a, b and c are lattice constants, No is space group number and Z is the number of formula units per unit cell. Density was not measured directly but calculated from the lattice parameters. |
Actinium | Oxides | Oxides
Actinium oxide (Ac2O3) can be obtained by heating the hydroxide at or the oxalate at , in vacuum. Its crystal lattice is isotypic with the oxides of most trivalent rare-earth metals. |
Actinium | Halides | Halides
Actinium trifluoride can be produced either in solution or in solid reaction. The former reaction is carried out at room temperature, by adding hydrofluoric acid to a solution containing actinium ions. In the latter method, actinium metal is treated with hydrogen fluoride vapors at in an all-platinum setup. Treating actinium trifluoride with ammonium hydroxide at yields oxyfluoride AcOF. Whereas lanthanum oxyfluoride can be easily obtained by burning lanthanum trifluoride in air at for an hour, similar treatment of actinium trifluoride yields no AcOF and only results in melting of the initial product.Meyer, pp. 87–88
AcF3 + 2 NH3 + H2O → AcOF + 2 NH4F
Actinium trichloride is obtained by reacting actinium hydroxide or oxalate with carbon tetrachloride vapors at temperatures above . Similarly to the oxyfluoride, actinium oxychloride can be prepared by hydrolyzing actinium trichloride with ammonium hydroxide at . However, in contrast to the oxyfluoride, the oxychloride could well be synthesized by igniting a solution of actinium trichloride in hydrochloric acid with ammonia.
Reaction of aluminium bromide and actinium oxide yields actinium tribromide:
Ac2O3 + 2 AlBr3 → 2 AcBr3 + Al2O3
and treating it with ammonium hydroxide at results in the oxybromide AcOBr. |
Actinium | Other compounds | Other compounds
Actinium hydride was obtained by reduction of actinium trichloride with potassium at , and its structure was deduced by analogy with the corresponding LaH2 hydride. The source of hydrogen in the reaction was uncertain.Meyer, p. 43
Mixing monosodium phosphate (NaH2PO4) with a solution of actinium in hydrochloric acid yields white-colored actinium phosphate hemihydrate (AcPO4·0.5H2O), and heating actinium oxalate with hydrogen sulfide vapors at for a few minutes results in a black actinium sulfide Ac2S3. It may possibly be produced by acting with a mixture of hydrogen sulfide and carbon disulfide on actinium oxide at . |
Actinium | Isotopes | Isotopes
Naturally occurring actinium is principally composed of two radioactive isotopes; (from the radioactive family of ) and (a granddaughter of ). decays mainly as a beta emitter with a very small energy, but in 1.38% of cases it emits an alpha particle, so it can readily be identified through alpha spectrometry. Thirty-three radioisotopes have been identified, the most stable being with a half-life of 21.772 years, with a half-life of 10.0 days and with a half-life of 29.37 hours. All remaining radioactive isotopes have half-lives that are less than 10 hours and the majority of them have half-lives shorter than one minute. The shortest-lived known isotope of actinium is (half-life of 69 nanoseconds) which decays through alpha decay. Actinium also has two known meta states. The most significant isotopes for chemistry are 225Ac, 227Ac, and 228Ac.
Purified comes into equilibrium with its decay products after about a half of year. It decays according to its 21.772-year half-life emitting mostly beta (98.62%) and some alpha particles (1.38%); the successive decay products are part of the actinium series. Owing to the low available amounts, low energy of its beta particles (maximum 44.8 keV) and low intensity of alpha radiation, is difficult to detect directly by its emission and it is therefore traced via its decay products.Actinium, Great Soviet Encyclopedia (in Russian) The isotopes of actinium range in atomic weight from 203 u () to 236 u ().
Isotope
Production
Decay
Half-life
221Ac
232Th(d,9n)→225Pa(α)→221Ac
α
52 ms
222Ac
232Th(d,8n)→226Pa(α)→222Ac
α
5.0 s
223Ac
232Th(d,7n)→227Pa(α)→223Ac
α
2.1 min
224Ac
232Th(d,6n)→228Pa(α)→224Ac
α
2.78 hours
225Ac
232Th(n,γ)→233Th(β−)→233Pa(β−)→233U(α)→229Th(α)→225Ra(β−)→225Ac
α
10 days
226Ac
226Ra(d,2n)→226Ac
α, β− electron capture
29.37 hours
227Ac
235U(α)→231Th(β−)→231Pa(α)→227Ac
α, β−
21.77 years
228Ac
232Th(α)→228Ra(β−)→228Ac
β−
6.15 hours
229Ac
228Ra(n,γ)→229Ra(β−)→229Ac
β−
62.7 min
230Ac
232Th(d,α)→230Ac
β−
122 s
231Ac
232Th(γ,p)→231Ac
β−
7.5 min
232Ac
232Th(n,p)→232Ac
β−
119 s
|
Actinium | Occurrence and synthesis | Occurrence and synthesis
upright=0.70|thumb|Uraninite ores have elevated concentrations of actinium.
Actinium is found only in traces in uranium ores – one tonne of uranium in ore contains about 0.2 milligrams of 227Ac – and in thorium ores, which contain about 5 nanograms of 228Ac per one tonne of thorium. The actinium isotope 227Ac is a transient member of the uranium-actinium series decay chain, which begins with the parent isotope 235U (or 239Pu) and ends with the stable lead isotope 207Pb. The isotope 228Ac is a transient member of the thorium series decay chain, which begins with the parent isotope 232Th and ends with the stable lead isotope 208Pb. Another actinium isotope (225Ac) is transiently present in the neptunium series decay chain, beginning with 237Np (or 233U) and ending with thallium (205Tl) and near-stable bismuth (209Bi); even though all primordial 237Np has decayed away, it is continuously produced by neutron knock-out reactions on natural 238U.
The low natural concentration, and the close similarity of physical and chemical properties to those of lanthanum and other lanthanides, which are always abundant in actinium-bearing ores, render separation of actinium from the ore impractical. The most concentrated actinium sample prepared from raw material consisted of 7 micrograms of 227Ac in less than 0.1 milligrams of La2O3, and complete separation was never achieved. Instead, actinium is prepared, in milligram amounts, by the neutron irradiation of in a nuclear reactor.
^{226}_{88}Ra + ^{1}_{0}n -> ^{227}_{88}Ra ->[\beta^-][42.2 \ \ce{min}] ^{227}_{89}Ac
The reaction yield is about 2% of the radium weight. 227Ac can further capture neutrons resulting in small amounts of 228Ac. After the synthesis, actinium is separated from radium and from the products of decay and nuclear fusion, such as thorium, polonium, lead and bismuth. The extraction can be performed with thenoyltrifluoroacetone-benzene solution from an aqueous solution of the radiation products, and the selectivity to a certain element is achieved by adjusting the pH (to about 6.0 for actinium). An alternative procedure is anion exchange with an appropriate resin in nitric acid, which can result in a separation factor of 1,000,000 for radium and actinium vs. thorium in a two-stage process. Actinium can then be separated from radium, with a ratio of about 100, using a low cross-linking cation exchange resin and nitric acid as eluant.
225Ac was first produced artificially at the Institute for Transuranium Elements (ITU) in Germany using a cyclotron and at St George Hospital in Sydney using a linac in 2000. This rare isotope has potential applications in radiation therapy and is most efficiently produced by bombarding a radium-226 target with 20–30 MeV deuterium ions. This reaction also yields 226Ac which however decays with a half-life of 29 hours and thus does not contaminate 225Ac.
Actinium metal has been prepared by the reduction of actinium fluoride with lithium vapor in vacuum at a temperature between . Higher temperatures resulted in evaporation of the product and lower ones lead to an incomplete transformation. Lithium was chosen among other alkali metals because its fluoride is most volatile.Hammond, C. R. The Elements in |
Actinium | Applications | Applications
Owing to its scarcity, high price and radioactivity, 227Ac currently has no significant industrial use, but 225Ac is currently being studied for use in cancer treatments such as targeted alpha therapies.
227Ac is highly radioactive and was therefore studied for use as an active element of radioisotope thermoelectric generators, for example in spacecraft. The oxide of 227Ac pressed with beryllium is also an efficient neutron source with the activity exceeding that of the standard americium-beryllium and radium-beryllium pairs.Russell, Alan M. and Lee, Kok Loong (2005) Structure-property relations in nonferrous metals. Wiley. , pp. 470–471 In all those applications, 227Ac (a beta source) is merely a progenitor which generates alpha-emitting isotopes upon its decay. Beryllium captures alpha particles and emits neutrons owing to its large cross-section for the (α,n) nuclear reaction:
^{9}_{4}Be + ^{4}_{2}He -> ^{12}_{6}C + ^{1}_{0}n + \gamma
The 227AcBe neutron sources can be applied in a neutron probe – a standard device for measuring the quantity of water present in soil, as well as moisture/density for quality control in highway construction.Majumdar, D. K. (2004) Irrigation Water Management: Principles and Practice. p. 108Chandrasekharan, H. and Gupta, Navindu (2006) Fundamentals of Nuclear Science – Application in Agriculture. pp. 202 ff Such probes are also used in well logging applications, in neutron radiography, tomography and other radiochemical investigations.
thumb|upright=0.70|Chemical structure of the DOTA carrier for 225Ac in radiation therapy
225Ac is applied in medicine to produce in a reusable generator or can be used alone as an agent for radiation therapy, in particular targeted alpha therapy (TAT). This isotope has a half-life of 10 days, making it much more suitable for radiation therapy than 213Bi (half-life 46 minutes). Additionally, 225Ac decays to nontoxic 209Bi rather than toxic lead, which is the final product in the decay chains of several other candidate isotopes, namely 227Th, 228Th, and 230U. Not only 225Ac itself, but also its daughters, emit alpha particles which kill cancer cells in the body. The major difficulty with application of 225Ac was that intravenous injection of simple actinium complexes resulted in their accumulation in the bones and liver for a period of tens of years. As a result, after the cancer cells were quickly killed by alpha particles from 225Ac, the radiation from the actinium and its daughters might induce new mutations. To solve this problem, 225Ac was bound to a chelating agent, such as citrate, ethylenediaminetetraacetic acid (EDTA) or diethylene triamine pentaacetic acid (DTPA). This reduced actinium accumulation in the bones, but the excretion from the body remained slow. Much better results were obtained with such chelating agents as HEHA () or DOTA () coupled to trastuzumab, a monoclonal antibody that interferes with the HER2/neu receptor. The latter delivery combination was tested on mice and proved to be effective against leukemia, lymphoma, breast, ovarian, neuroblastoma and prostate cancers.
The medium half-life of 227Ac (21.77 years) makes it a very convenient radioactive isotope in modeling the slow vertical mixing of oceanic waters. The associated processes cannot be studied with the required accuracy by direct measurements of current velocities (of the order 50 meters per year). However, evaluation of the concentration depth-profiles for different isotopes allows estimating the mixing rates. The physics behind this method is as follows: oceanic waters contain homogeneously dispersed 235U. Its decay product, 231Pa, gradually precipitates to the bottom, so that its concentration first increases with depth and then stays nearly constant. 231Pa decays to 227Ac; however, the concentration of the latter isotope does not follow the 231Pa depth profile, but instead increases toward the sea bottom. This occurs because of the mixing processes which raise some additional 227Ac from the sea bottom. Thus analysis of both 231Pa and 227Ac depth profiles allows researchers to model the mixing behavior.
There are theoretical predictions that AcHx hydrides (in this case with very high pressure) are a candidate for a near room-temperature superconductor as they have Tc significantly higher than H3S, possibly near 250 K. |
Actinium | Precautions | Precautions
227Ac is highly radioactive and experiments with it are carried out in a specially designed laboratory equipped with a tight glove box. When actinium trichloride is administered intravenously to rats, about 33% of actinium is deposited into the bones and 50% into the liver. Its toxicity is comparable to, but slightly lower, than that of americium and plutonium. For trace quantities, fume hoods with good aeration suffice; for gram amounts, hot cells with shielding from the intense gamma radiation emitted by 227Ac are necessary. |
Actinium | See also | See also
Actinium series |
Actinium | Notes | Notes |
Actinium | References | References |
Actinium | Bibliography | Bibliography
|
Actinium | External links | External links
Actinium at The Periodic Table of Videos (University of Nottingham)
NLM Hazardous Substances Databank – Actinium, Radioactive
Actinium in
Category:Chemical elements
Category:Chemical elements with face-centered cubic structure
Category:Actinides |
Actinium | Table of Content | Distinguish, History, Properties, Chemical compounds, Oxides, Halides, Other compounds, Isotopes, Occurrence and synthesis, Applications, Precautions, See also, Notes, References, Bibliography, External links |
Americium | good article | Americium is a synthetic chemical element; it has symbol Am and atomic number 95. It is radioactive and a transuranic member of the actinide series in the periodic table, located under the lanthanide element europium and was thus named after the Americas by analogy.
Americium was first produced in 1944 by the group of Glenn T. Seaborg from Berkeley, California, at the Metallurgical Laboratory of the University of Chicago, as part of the Manhattan Project. Although it is the third element in the transuranic series, it was discovered fourth, after the heavier curium. The discovery was kept secret and only released to the public in November 1945. Most americium is produced by uranium or plutonium being bombarded with neutrons in nuclear reactors – one tonne of spent nuclear fuel contains about 100 grams of americium. It is widely used in commercial ionization chamber smoke detectors, as well as in neutron sources and industrial gauges. Several unusual applications, such as nuclear batteries or fuel for space ships with nuclear propulsion, have been proposed for the isotope 242mAm, but they are as yet hindered by the scarcity and high price of this nuclear isomer.
Americium is a relatively soft radioactive metal with a silvery appearance. Its most common isotopes are 241Am and 243Am. In chemical compounds, americium usually assumes the oxidation state +3, especially in solutions. Several other oxidation states are known, ranging from +2 to +7, and can be identified by their characteristic optical absorption spectra. The crystal lattices of solid americium and its compounds contain small intrinsic radiogenic defects, due to metamictization induced by self-irradiation with alpha particles, which accumulates with time; this can cause a drift of some material properties over time, more noticeable in older samples. |
Americium | History | History
thumb|left|The 60-inch cyclotron at the Lawrence Radiation Laboratory, University of California, Berkeley, in August 1939
Although americium was likely produced in previous nuclear experiments, it was first intentionally synthesized, isolated and identified in late autumn 1944, at the University of California, Berkeley, by Glenn T. Seaborg, Leon O. Morgan, Ralph A. James, and Albert Ghiorso. They used a 60-inch cyclotron at the University of California, Berkeley.Obituary of Dr. Leon Owen (Tom) Morgan (1919–2002), Retrieved 28 November 2010 The element was chemically identified at the Metallurgical Laboratory (now Argonne National Laboratory) of the University of Chicago. Following the lighter neptunium, plutonium, and heavier curium, americium was the fourth transuranium element to be discovered. At the time, the periodic table had been restructured by Seaborg to its present layout, containing the actinide row below the lanthanide one. This led to americium being located right below its twin lanthanide element europium; it was thus by analogy named after the Americas: "The name americium (after the Americas) and the symbol Am are suggested for the element on the basis of its position as the sixth member of the actinide rare-earth series, analogous to europium, Eu, of the lanthanide series."Seaborg, G. T.; James, R.A. and Morgan, L. O.: "The New Element Americium (Atomic Number 95)", THIN PPR (National Nuclear Energy Series, Plutonium Project Record), Vol 14 B The Transuranium Elements: Research Papers, Paper No. 22.1, McGraw-Hill Book Co., Inc., New York, 1949. Abstract; Full text (January 1948), Retrieved 28 November 2010Greenwood, p. 1252
The new element was isolated from its oxides in a complex, multi-step process. First plutonium-239 nitrate (239PuNO3) solution was coated on a platinum foil of about 0.5 cm2 area, the solution was evaporated and the residue was converted into plutonium dioxide (PuO2) by calcining. After cyclotron irradiation, the coating was dissolved with nitric acid, and then precipitated as the hydroxide using concentrated aqueous ammonia solution. The residue was dissolved in perchloric acid. Further separation was carried out by ion exchange, yielding a certain isotope of curium. The separation of curium and americium was so painstaking that those elements were initially called by the Berkeley group as pandemonium (from Greek for all demons or hell) and delirium (from Latin for madness).
Initial experiments yielded four americium isotopes: 241Am, 242Am, 239Am and 238Am. Americium-241 was directly obtained from plutonium upon absorption of two neutrons. It decays by emission of a α-particle to 237Np; the half-life of this decay was first determined as years but then corrected to 432.2 years.
The times are half-lives
The second isotope 242Am was produced upon neutron bombardment of the already-created 241Am. Upon rapid β-decay, 242Am converts into the isotope of curium 242Cm (which had been discovered previously). The half-life of this decay was initially determined at 17 hours, which was close to the presently accepted value of 16.02 h.
The discovery of americium and curium in 1944 was closely related to the Manhattan Project; the results were confidential and declassified only in 1945. Seaborg leaked the synthesis of the elements 95 and 96 on the U.S. radio show for children Quiz Kids five days before the official presentation at an American Chemical Society meeting on 11 November 1945, when one of the listeners asked whether any new transuranium element besides plutonium and neptunium had been discovered during the war. After the discovery of americium isotopes 241Am and 242Am, their production and compounds were patented listing only Seaborg as the inventor.Seaborg, Glenn T. "Element", Filing date: 23 August 1946, Issue date: 10 November 1964 The initial americium samples weighed a few micrograms; they were barely visible and were identified by their radioactivity. The first substantial amounts of metallic americium weighing 40–200 micrograms were not prepared until 1951 by reduction of americium(III) fluoride with barium metal in high vacuum at 1100 °C. |
Americium | Occurrence | Occurrence
thumb|Americium was detected in the fallout from the Ivy Mike nuclear test.
The longest-lived and most common isotopes of americium, 241Am and 243Am, have half-lives of 432.2 and 7,370 years, respectively. Therefore, any primordial americium (americium that was present on Earth during its formation) should have decayed by now. Trace amounts of americium probably occur naturally in uranium minerals as a result of neutron capture and beta decay (238U → 239Pu → 240Pu → 241Am), though the quantities would be tiny and this has not been confirmed. Extraterrestrial long-lived 247Cm is probably also deposited on Earth and has 243Am as one of its intermediate decay products, but again this has not been confirmed.
Existing americium is concentrated in the areas used for the atmospheric nuclear weapons tests conducted between 1945 and 1980, as well as at the sites of nuclear incidents, such as the Chernobyl disaster. For example, the analysis of the debris at the testing site of the first U.S. hydrogen bomb, Ivy Mike, (1 November 1952, Enewetak Atoll), revealed high concentrations of various actinides including americium; but due to military secrecy, this result was not published until later, in 1956. Trinitite, the glassy residue left on the desert floor near Alamogordo, New Mexico, after the plutonium-based Trinity nuclear bomb test on 16 July 1945, contains traces of americium-241. Elevated levels of americium were also detected at the crash site of a US Boeing B-52 bomber aircraft, which carried four hydrogen bombs, in 1968 in Greenland.
In other regions, the average radioactivity of surface soil due to residual americium is only about 0.01 picocuries per gram (0.37 mBq/g). Atmospheric americium compounds are poorly soluble in common solvents and mostly adhere to soil particles. Soil analysis revealed about 1,900 times higher concentration of americium inside sandy soil particles than in the water present in the soil pores; an even higher ratio was measured in loam soils.Human Health Fact Sheet on Americium , Los Alamos National Laboratory, Retrieved 28 November 2010
Americium is produced mostly artificially in small quantities, for research purposes. A tonne of spent nuclear fuel contains about 100 grams of various americium isotopes, mostly 241Am and 243Am.Hoffmann, Klaus Kann man Gold machen? Gauner, Gaukler und Gelehrte. Aus der Geschichte der chemischen Elemente (Can you make gold? Crooks, clowns, and scholars. From the history of the chemical elements), Urania-Verlag, Leipzig, Jena, Berlin 1979, no ISBN, p. 233 Their prolonged radioactivity is undesirable for the disposal, and therefore americium, together with other long-lived actinides, must be neutralized. The associated procedure may involve several steps, where americium is first separated and then converted by neutron bombardment in special reactors to short-lived nuclides. This procedure is well known as nuclear transmutation, but it is still being developed for americium.Baetslé, L. Application of Partitioning/Transmutation of Radioactive Materials in Radioactive Waste Management , Nuclear Research Centre of Belgium Sck/Cen, Mol, Belgium, September 2001, Retrieved 28 November 2010Fioni, Gabriele; Cribier, Michel and Marie, Frédéric Can the minor actinide, americium-241, be transmuted by thermal neutrons? , Department of Astrophysics, CEA/Saclay, Retrieved 28 November 2010 The transuranic elements from americium to fermium occurred naturally in the natural nuclear fission reactor at Oklo, but no longer do so.
Americium is also one of the elements that have theoretically been detected in Przybylski's Star. |
Americium | Synthesis and extraction | Synthesis and extraction |
Americium | Isotope nucleosynthesis | Isotope nucleosynthesis
thumb|Chromatographic elution curves revealing the similarity between the lanthanides Tb, Gd, and Eu and the corresponding actinides Bk, Cm, and Am
Americium has been produced in small quantities in nuclear reactors for decades, and kilograms of its 241Am and 243Am isotopes have been accumulated by now.Greenwood, p. 1262 Nevertheless, since it was first offered for sale in 1962, its price, about of 241Am, remains almost unchanged owing to the very complex separation procedure.Smoke detectors and americium , World Nuclear Association, January 2009, Retrieved 28 November 2010 The heavier isotope 243Am is produced in much smaller amounts; it is thus more difficult to separate, resulting in a higher cost of the order .Hammond C. R. "The elements" in
Americium is not synthesized directly from uranium – the most common reactor material – but from the plutonium isotope 239Pu. The latter needs to be produced first, according to the following nuclear process:
^{238}_{92}U ->[\ce{(n,\gamma)}] ^{239}_{92}U ->[\beta^-][23.5 \ \ce{min}] ^{239}_{93}Np ->[\beta^-][2.3565 \ \ce{d}] ^{239}_{94}Pu
The capture of two neutrons by 239Pu (a so-called (n,γ) reaction), followed by a β-decay, results in 241Am:
^{239}_{94}Pu ->[\ce{2(n,\gamma)}] ^{241}_{94}Pu ->[\beta^-][14.35 \ \ce{yr}] ^{241}_{95}Am
The plutonium present in spent nuclear fuel contains about 12% of 241Pu. Because it beta-decays to 241Am, 241Pu can be extracted and may be used to generate further 241Am. However, this process is rather slow: half of the original amount of 241Pu decays to 241Am after about 15 years, and the 241Am amount reaches a maximum after 70 years.BREDL Southern Anti-Plutonium Campaign, Blue Ridge Environmental Defense League, Retrieved 28 November 2010
The obtained 241Am can be used for generating heavier americium isotopes by further neutron capture inside a nuclear reactor. In a light water reactor (LWR), 79% of 241Am converts to 242Am and 10% to its nuclear isomer 242mAm:The "metastable" state is marked by the letter m. article/200410/000020041004A0333355.php Abstract
Americium-242 has a half-life of only 16 hours, which makes its further conversion to 243Am extremely inefficient. The latter isotope is produced instead in a process where 239Pu captures four neutrons under high neutron flux:
^{239}_{94}Pu ->[\ce{4(n,\gamma)}] \ ^{243}_{94}Pu ->[\beta^-][4.956 \ \ce{h}] ^{243}_{95}Am |
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