Wikipedia

Main article:
Isotopes of arsenic
iso	NA	half-life	DM	DE (MeV)	DP
⁷³As	syn	80.3 d	ε	-	⁷³Ge
⁷³As	syn	80.3 d	γ	0.05D, 0.01D, e	-
⁷⁴As	syn	17.78 d	ε	-	⁷⁴Ge
			β⁺	0.941	⁷⁴Ge
			γ	0.595, 0.634	-
			β^-	1.35, 0.717	⁷⁴Se
⁷⁵As	100%	As is stable with 42 neutrons

References

Arsenic (pronounced /ˈɑrs^ənɪk/) is a chemical element that has the symbol As and atomic number 33. Arsenic was discovered by Albertus Magnus (Germany) in 1250. Its Atomic Mass is 74.92. Its Ionic Charge is (3-) Its position in the periodic table is shown at right. This is a notoriously poisonous metalloid that has many allotropic forms: yellow (molecular non-metallic) and several black and gray forms (metalloids) are a few that are seen. Three metalloidal forms of arsenic with different crystal structures are found free in nature (the minerals arsenic sensu stricto and the much rarer arsenolamprite and pararsenolamprite), but it is more commonly found as arsenide and arsenate compounds. Several hundred such mineral species are known. Arsenic and its compounds are used as pesticides, herbicides, insecticides and various alloys.

The most common oxidation states for arsenic are -3 (arsenides: usually alloy-like intermetallic compounds), +3 (arsenates(III) or arsenites, and most organoarsenic compounds), and +5 (arsenates(V): the most stable inorganic arsenic oxycompounds). Arsenic also bonds readily to itself, forming, for instance, As-As pairs in the red sulfide realgar and square As₄^3- ions in the arsenide skutterudite. In the +3 oxidation state, the stereochemistry of arsenic is affected by possession of a lone pair of electrons.

[edit] Notable characteristics

Arsenic is very similar chemically to its predecessor, phosphorus. Similar to phosphorus, it forms colourless, odourless, crystalline oxides As₂O₃ and As₂O₅ which are hygroscopic and readily soluble in water to form acidic solutions. Arsenic (V) acid, like phosphoric acid, is a weak acid. Like phosphorus, arsenic forms an unstable, gaseous hydride: arsine (AsH₃). The similarity is so great that arsenic will partly substitute for phosphorus in biochemical reactions and is thus poisonous. However, in subtoxic doses, soluble arsenic compounds act as stimulants, and were once popular in small doses as medicinals by people in the mid 18th century.

When heated in air it 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. Arsenic (and some arsenic compounds) sublimes upon heating at atmospheric pressure, converting directly to a gaseous form without an intervening liquid state. The liquid state appears at 20 atmospheres and above, which explains why the melting point is higher than the boiling point ^[3]. Elemental arsenic is found in many solid forms: the yellow form is soft, waxy and unstable, and is made of tetrahedral As ₄ molecules similar to the molecules of white phosphorus. The gray, black or 'metallic' forms have somewhat layered crystal structures with bonds extending throughout the crystal. They are brittle semiconductors with a metallic luster. The density of the yellow form is 1.97 g/cm³; rhombohedral 'gray arsenic' is much denser with a density of 5.73 g/cm³; the other metalloidal forms are similarly dense.

[edit] Applications

Lead hydrogen arsenate has been used, well into the 20th century, as an insecticide on fruit trees (sometimes resulting in brain damage to those working the sprayers), and Scheele's Green (a copper arsenate) has even been recorded in the 19th century as a coloring agent in sweets. In the last half century, monosodium methyl arsenate (MSMA), a less toxic organic form of arsenic, has replaced lead arsenate's role in agriculture.

The application of most concern to the general public is probably that of wood which has been treated with chromated copper arsenate ("CCA", or "Tanalith", and the vast majority of older " pressure treated" wood). CCA timber 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, where there was a risk of rot, or insect infestation in untreated timber. Although widespread bans followed the publication of studies which showed low-level leaching from in-situ timbers (such as children's playground equipment) into surrounding soil, the most serious ^{[ citation needed]} risk is presented by the burning of CCA timber. Recent years have seen fatal animal poisonings, and serious human poisonings resulting from the ingestion - directly or indirectly - of wood ash from CCA timber (the lethal human dose is approximately 20 grams of ash). Scrap CCA construction timber continues to be widely burnt through ignorance, in both commercial and domestic fires. Protocols for safe disposal of CCA timber are still in place only patchily; there is concern in some quarters about the widespread landfill disposal of such timber.

During the 18th, 19th, and 20th centuries, a number of arsenic compounds have been used as medicines, including arsphenamine (by Paul Ehrlich) and arsenic trioxide (by Thomas Fowler). Arsphenamine as well as Neosalvarsan was indicated for syphilis and trypanosomiasis, but has been superseded by modern antibiotics. Arsenic trioxide has been used in a variety of ways over the past 200 years, but most commonly in the treatment of cancer. The US Food and Drug Administration in 2000 approved this compound for the treatment of patients with acute promyelocytic leukemia that is resistant to ATRA.^[4] It was also used as Fowler 's solution in psoriasis.^[5]

Copper acetoarsenite was used as a green pigment known under many different names, including ' Paris Green' and 'Emerald Green'. It caused numerous arsenic poisonings.

Other uses;

Various agricultural insecticides, termination and poisons.
Used in animal feed, particularly in the US as a method of disease prevention and growth stimulation.
Gallium arsenide is an important semiconductor material, used in integrated circuits. Circuits made using the compound are much faster (but also much more expensive) than those made in silicon. Unlike silicon it is direct bandgap, and so can be used in laser diodes and LEDs to directly convert electricity into light.
Also used in bronzing and pyrotechny.

Occupational Exposures

Exposure to higher-than-average levels of arsenic can occur in some occupations placing workers at risk. Industries that use inorganic arsenic and its compounds include 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. ^[6]

[edit] History

The word arsenic is borrowed from the Persian word زرنيخ Zarnikh meaning "yellow orpiment". Zarnikh was borrowed by Greek as arsenikon wich means masculine or potent. Arsenic has been known and used in Persia and elsewhere since ancient times. As the symptoms of arsenic poisoning were somewhat ill-defined, it was frequently used for murder until the advent of the Marsh test, a sensitive chemical test for its presence. (Another less sensitive but more general test is the Reinsch test.) Due 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.

During the Bronze Age, arsenic was often included in bronze, which made the alloy harder (so-called " arsenical bronze").

Albertus Magnus (Albert the Great, 1193-1280) is believed to have been the first to isolate the element in 1250^[7]. In 1649 Johann Schröder published two ways of preparing arsenic.

Alchemical symbol for arsenic

In the Victorian era, 'arsenic' (colourless, crystalline, soluble 'white arsenic') was mixed with vinegar and chalk and eaten by women to improve the complexion of their faces, making their skin paler to show they did not work in the fields. Arsenic was also rubbed into the faces and arms of women to 'improve their complexion'. The accidental use of arsenic in the adulteration of foodstuffs led to the Bradford sweet poisoning in 1858, which resulted in approximately 20 deaths and 200 people taken ill with arsenic poisoning.

[edit] Occurrence

A large sample of native arsenic.

The most important compounds of arsenic are arsenic (III) oxide, As₂O₃, ('white arsenic'), the yellow sulfide orpiment (As₂S₃) and red realgar (As₄S₄), Paris Green, calcium arsenate, and lead hydrogen arsenate. The latter three have been used as agricultural insecticides and poisons. Orpiment and realgar were formerly used as painting pigments, though they have fallen out of use due to their toxicity and reactivity. Although arsenic is sometimes found native in nature, its main economic source is the mineral arsenopyrite mentioned above; it is also found in arsenides of metals such as silver, cobalt (cobaltite: CoAsS and skutterudite: CoAs₃) and nickel, as sulfides, and when oxidised as arsenate minerals such as mimetite, Pb₅(AsO₄)₃Cl and erythrite, Co₃(AsO₄)₂. 8H₂O, and more rarely arsenites ('arsenite' = arsenate(III), AsO₃^3- as opposed to arsenate (V), AsO₄^3- ). In addition to the inorganic forms mentioned above, arsenic also occurs in various organic forms in the environment. Inorganic arsenic and its compounds, upon entering the food chain, are progressively metabolised to a less toxic form of arsenic through a process of methylation.

Nickernuts are said to contain arsenic. See also Arsenide minerals, Arsenate minerals.

[edit] Toxicity

Main article: Arsenic poisoning

Arsenic and many of its compounds are especially potent poisons. Arsenic disrupts ATP production through several mechanisms. At the level of the citric acid cycle, arsenic inhibits pyruvate dehydrogenase and by competing with phosphate it uncouples oxidative phosphorylation, thus inhibiting energy-linked reduction of NAD+, mitochondrial respiration, and ATP synthesis. Hydrogen peroxide production is also increased, which might form reactive oxygen species and oxidative stress. These metabolic interferences lead to death from multi-system organ failure (see arsenic poisoning) probably from necrotic cell death, not apoptosis. A post mortem reveals brick red colored mucosa, due to severe hemorrhage. Although arsenic causes toxicity, it can also play a protective role. ^[8].

Elemental arsenic and arsenic compounds are classified as "toxic" and "dangerous for the environment" in the European Union under directive 67/548/EEC.

The IARC recognizes arsenic and 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 due to its manifestation in drinking water, "the most common species being arsenate [HAsO₄^2- ; As(V)] and arsenite [H₃AsO₃ ; As(III)]". The ability of arsenic to undergo redox conversion between As(III) and As(V) makes its availability in the environment possible. 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.^[9]

[edit] Arsenic in drinking water

Main article: Arsenic contamination of groundwater

Arsenic contamination of groundwater has led to a massive epidemic of arsenic poisoning in Bangladesh^[10] and neighbouring countries. It is estimated that approximately 57 million people are drinking groundwater with arsenic concentrations elevated above the World Health Organization's standard of 10 parts per billion. The arsenic in the groundwater is of natural origin, and is released from the sediment into the groundwater due to the anoxic conditions of the subsurface. This groundwater began to be used after western NGOs instigated a massive tube well drinking-water program in the late twentieth century. This program was designed to prevent drinking of bacterially contaminated surface waters, but failed to test for arsenic in the groundwater.(2) Many other countries and districts in South East Asia, such as Vietnam, Cambodia, and Tibet, China, are thought to have geological environments similarly conducive to generation of high-arsenic groundwaters. Arsenicosis was reported in Nakhon Si Thammarat, Thailand in 1987, and the dissolved arsenic in the Chao Phraya River is suspected of containing high levels of naturally occurring arsenic, but has not been a public health problem due to the use of bottled water.^[11]

The northern United States, including parts of Michigan, Wisconsin, Minnesota and the Dakotas are known to have significant concentrations of arsenic in ground water.

Arsenic can be removed from drinking water through coprecipitation of iron minerals by oxidation and filtering. When this treatment fails to produce acceptable results, adsorptive arsenic removal media may be utilized. Several adsorptive media systems have been approved for point of service use in a study funded by the United States Environmental Protection Agency (U.S.EPA) and the National Science Foundation (NSF).

Magnetic separations of arsenic at very low magnetic field gradients have been demonstrated in point-of-use water purification with high–surface area and monodisperse magnetite (Fe₃O₄) nanocrystals. Using the high specific surface area of Fe₃O₄ nanocrystals the mass of waste associated with arsenic removal from water has been dramatically reduced. ^[12]

[edit] Compounds

Arsenic acid (H₃AsO₄)
Arsenous acid (H₃AsO₃)
Arsenic trioxide (As₂O₃)
Arsine (Arsenic Trihydride AsH₃)
Cadmium arsenide (Cd₃As₂)
Gallium arsenide (GaAs)
Lead hydrogen arsenate (PbHAsO₄)

Barium · Science/Tech

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For other uses, see Barium (disambiguation).

caesium ← barium → lanthanum

Sr
↑
Ba
↓
Ra

Periodic Table - Extended Periodic Table

General

Name, Symbol , Number

barium, Ba, 56

Chemical series

alkaline earth metals

Group, Period , Block

2, 6, s

Appearance

silvery white

Standard atomic weight

137.327 (7) g·mol⁻¹

Electron configuration

[Xe] 6s²

Electrons per shell

2, 8, 18, 18, 8, 2

Physical properties

Phase

solid

Density (near r.t.)

3.51 g·cm⁻³

Liquid density at m.p.

3.338 g·cm⁻³

Melting point

1000 K
(727 °C, 1341 °F)

Boiling point

2170 K
(1897 °C, 3447 °F)

Heat of fusion

7.12 kJ·mol⁻¹

Heat of vaporization

140.3 kJ·mol⁻¹

Heat capacity

(25 °C) 28.07 J·mol⁻¹·K⁻¹

Vapor pressure
P(Pa)	1	10	100	1 k	10 k	100 k
at T(K)	911	1038	1185	1388	1686	2170

Atomic properties

Crystal structure

cubic body centered

Oxidation states

2
(strongly basic oxide)

Electronegativity

0.89 (Pauling scale)

Ionization energies

1st: 502.9 kJ/mol

2nd: 965.2 kJ/mol

3rd: 3600 kJ/mol

Atomic radius

215 pm

Atomic radius (calc.)

253 pm

Covalent radius

198 pm

Miscellaneous

Magnetic ordering

paramagnetic

Electrical resistivity

(20 °C) 332 n Ω·m

Thermal conductivity

(300 K) 18.4 W·m⁻¹·K⁻¹

Thermal expansion

(25 °C) 20.6 µm·m⁻¹·K⁻¹

Speed of sound (thin rod)

(20 °C) 1620 m/s

13 GPa

4.9 GPa

9.6 GPa

1.25

7440-39-3

Selected isotopes

Main article:
Isotopes of barium
iso	NA	half-life	DM	DE (MeV)	DP
¹³⁰Ba	0.106%	Ba is stable with 74 neutrons
¹³²Ba	0.101%	Ba is stable with 76 neutrons
¹³³Ba	syn	10.51 y	ε	0.517	¹³³Cs
¹³⁴Ba	2.417%	Ba is stable with 78 neutrons
¹³⁵Ba	6.592%	Ba is stable with 79 neutrons
¹³⁶Ba	7.854%	Ba is stable with 80 neutrons
¹³⁷Ba	11.23%	Ba is stable with 81 neutrons
¹³⁸Ba	71.7%	Ba is stable with 82 neutrons

References

Barium (pronounced /ˈbɛəriəm/) is a chemical element. It has the symbol Ba, and atomic number 56. Barium is a soft silvery metallic alkaline earth metal. It is never found in nature in its pure form due to its reactivity with air. Its oxide is historically known as baryta but it reacts with water and carbon dioxide and is not found as a mineral. The most common naturally occurring minerals are the very insoluble barium sulfate, BaSO₄ (barite), and barium carbonate, BaCO₃ (witherite). Benitoite is a rare gem containing barium.

[edit] Notable characteristics

Barium is a metallic element that is chemically similar to calcium but more reactive. This metal oxidizes very easily when exposed to air and is highly reactive with water or alcohol, producing hydrogen gas. Burning in air or oxygen produces not just barium oxide (BaO) but also the peroxide. Simple compounds of this heavy element are notable for their high specific gravity. This is true of the most common barium-bearing mineral, its sulfate barite BaSO₄, also called 'heavy spar' due to the high density (4.5 g/cm³).

[edit] Applications

Barium has some medical and many industrial uses:

Barium compounds, and especially barite (BaSO₄), are extremely important to the petroleum industry. Barite is used in drilling mud , a weighting agent in drilling new oil wells.
Barium sulfate is used as a radiocontrast agent for X-ray imaging of the digestive system (" barium meals" and "barium enemas").
Barium carbonate is a useful rat poison and can also be used in making bricks. Unlike the sulfate, the carbonate dissolves in stomach acid, allowing it to be poisonous.
An alloy with nickel is used in spark plug wire.
Barium oxide is used in a coating for the electrodes of fluorescent lamps , which facilitates the release of electrons.
The metal is a "getter" in vacuum tubes, to remove the last traces of oxygen.
Barium carbonate is used in glassmaking. Being a heavy element, barium increases the refractive index and luster of the glass.
Barite is used extensively in rubber production.
Barium nitrate and chlorate give green colors in fireworks.
Impure barium sulfide phosphoresces after exposure to the light.
Lithopone, a pigment that contains barium sulfate and zinc sulfide, is a permanent white that has good covering power, and does not darken in when exposed to sulfides.
Barium peroxide can be used as a catalyst to start an aluminothermic reaction when welding rail tracks together. It can also be used in green tracer ammunition.
Barium titanate was proposed in 2007[1] to be used in next generation battery technology for electric cars.
Barium Fluoride is used in infrared applications.
Barium is a key element in YBCO superconductors.

[edit] History

Barium (Greek barys, meaning "heavy") was first identified in 1774 by Carl Scheele and extracted in 1808 by Sir Humphry Davy in England. The oxide was at first called barote, by Guyton de Morveau, which was changed by Antoine Lavoisier to baryta, from which "barium" was derived to describe the metal.

[edit] Occurrence

Because barium quickly becomes oxidized in air, it is difficult to obtain this metal in its pure form. It is primarily found in and extracted from the mineral barite which is crystalized barium sulfate. Barium is commercially produced through the electrolysis of molten barium chloride (BaCl₂) Isolation (* follow):

(cathode) Ba²⁺* + 2e^- → Ba ( anode) Cl^-* → ½Cl₂ (g) + e^-

[edit] Compounds

The most important compounds are barium peroxide, barium chloride, sulfate, carbonate, nitrate, and chlorate.

[edit] Isotopes

Main article: isotopes of barium

Naturally occurring barium is a mix of seven stable isotopes. There are twenty-two isotopes known, but most of these are highly radioactive and have half-lives in the several millisecond to several minute range. The only notable exceptions are ¹³³Ba which has a half-life of 10.51 years, and ^137mBa (2.55 minutes).

[edit] Precautions

All water or acid soluble barium compounds are extremely poisonous. At low doses, barium acts as a muscle stimulant, while higher doses affect the nervous system, causing cardiac irregularities, tremors, weakness, anxiety, dyspnea and paralysis. This may be due to its ability to block potassium ion channels which are critical to the proper function of the nervous system.

Barium sulfate can be taken orally because it is highly insoluble in water, and is eliminated completely from the digestive tract. Unlike other heavy metals, barium does not bioaccumulate.^[1] However, inhaled dust containing barium compounds can accumulate in the lungs, causing a benign condition called baritosis.

Oxidation occurs very easily and, to remain pure, barium should be kept under a petroleum-based fluid (such as kerosene) or other suitable oxygen-free liquids that exclude air.

Barium acetate could lead to death in high doses. Marie Robards poisoned her father with the substance in Texas in 1993. She was tried and convicted in 1996.

[edit] References