Barium: Wikis

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caesiumbariumlanthanum
Sr

Ba

Ra
Appearance
silvery gray
General properties
Name, symbol, number barium, Ba, 56
Element category alkaline earth metals
Group, period, block 26, s
Standard atomic weight 137.33g·mol−1
Electron configuration [Xe] 6s2
Electrons per shell 2, 8, 18, 18, 8, 2 (Image)
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
Specific 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
Oxidation states 2
(strongly basic oxide)
Electronegativity 0.89 (Pauling scale)
Ionization energies 1st: 502.9 kJ·mol−1
2nd: 965.2 kJ·mol−1
3rd: 3600 kJ·mol−1
Atomic radius 222 pm
Covalent radius 215±11 pm
Van der Waals radius 268 pm
Miscellanea
Crystal structure body-centered cubic
Magnetic ordering paramagnetic
Electrical resistivity (20 °C) 332 nΩ·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
Young's modulus 13 GPa
Shear modulus 4.9 GPa
Bulk modulus 9.6 GPa
Mohs hardness 1.25
CAS registry number 7440-39-3
Most stable isotopes
Main article: Isotopes of barium
iso NA half-life DM DE (MeV) DP
130Ba 0.106% 130Ba is stable with 74 neutrons
132Ba 0.101% 132Ba is stable with 76 neutrons
133Ba syn 10.51 y ε 0.517 133Cs
134Ba 2.417% 134Ba is stable with 78 neutrons
135Ba 6.592% 135Ba is stable with 79 neutrons
136Ba 7.854% 136Ba is stable with 80 neutrons
137Ba 11.23% 137Ba is stable with 81 neutrons
138Ba 71.7% 138Ba is stable with 82 neutrons

Barium (pronounced /ˈbɛəriəm/, BAIR-ee-əm) is a chemical element. It has the symbol Ba, atomic number 56, and is the fifth element in Group 2. 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). Barium's name originates from Greek barys (βαρύς), meaning "heavy", describing the high density of some common barium-containing ores.

Metallic barium has few industrial uses, but has been historically used to scavenge air in vacuum tubes. Barium compounds impart a green color to flames and have been used in fireworks. Barium sulfate is used for its heaviness, insolubility, and X-ray opacity. It is used as an insoluble heavy mud-like paste when drilling oil wells, and in purer form, as an X-ray radiocontrast agent for imaging the human gastrointestinal tract. Soluble barium compounds are poisonous due to release of the soluble barium ion, and have been used as rodenticides. New uses for barium continue to be found: it is an essential ingredient in "high temperature" YBCO superconductors.

Contents

Characteristics

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Physical

Barium is a soft and ductile metal. Its simple compounds are notable for their relatively high (for an alkaline earth element) 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).

Chemical

Barium reacts exothermically with oxygen at room temperature to form barium oxide and peroxide. The reaction is violent if barium is powdered. It also reacts violently with dilute acids, alcohol and water

Ba + 2 H2O → Ba(OH)2 + H2 (g)

At elevated temperatures, barium combines with chlorine, nitrogen and hydrogen to produce BaCl2, Ba3N2 and BaH2, respectively. Barium reduces oxides, chlorides and sulfides of less reactive metals. For example:

Ba + CdO → BaO + Cd
Ba + ZnCl2 → BaCl2 + Zn
3 Ba + Al2S3 → 3 BaS + 2 Al

When heated with nitrogen and carbon, it forms the cyanide:

Ba + N2 + 2 C → Ba(CN)2

Barium combines with several metals, including aluminium, zinc, led and tin, forming intermetallic compounds and alloys.[1]

Isotopes

Naturally occurring barium is a mix of seven stable isotopes, the most abundant being 138Ba (71.7 %). There are twenty-two isotopes known, but most of these are highly radioactive and have half-lives in the several millisecond to several day range. The only notable exceptions are 133Ba which has a half-life of 10.51 years, and 137mBa (2.55 minutes).[2]

History

Barium's name originates from Greek barys, meaning "heavy", describing the density of some common barium-containing ores. Alchemists in the early Middle Ages knew about some barium minerals. Smooth pebble-like stones of mineral barite found in Bologna, Italy were known as "Bologna stones". After exposed to light they would glow for years that attracted them to witches and alchemists.[3]

Carl Scheele identified barite as containing a new element in 1774, but could not isolate barium. Oxidized barium was at first called barote, by Guyton de Morveau, a name which was changed by Antoine Lavoisier to baryta. Barium was first isolated by electrolysis of molten barium salts in 1808, by Sir Humphry Davy in England. Davy, by analogy with calcium named "barium" after baryta, with the "-ium" ending signifying a metallic element.[3]

Occurrence and production

Barite
Trend in world production of barite

The abundance of barium is 0.0425 % in the Earth's crust and 13 µg/L in sea water. It occurs in the minerals barite (as the sulfate) and witherite (as the carbonate).[1] A rare gem containing barium is known, called benitoite. Large deposits of barite are found in China, Germany, India, Morocco, and in the US.[4]

Because barium quickly oxidizes in air, it is difficult to obtain the free metal and it is never found free in nature. The metal is primarily found in, and extracted from, barite. Because barite is so insoluble, it cannot be used directly for the preparation of other barium compounds, or barium metal. Instead, the ore is heated with carbon to reduce it to barium sulfide:[5]

BaSO4 + 2 C → BaS + 2 CO2

The barium sulfide is then hydrolyzed or treated with acids to form other barium compounds, such as the chloride, nitrate, and carbonate.

Barium is commercially produced through the electrolysis of molten barium chloride (BaCl2):

(cathode) Ba2+ + 2 e → Ba
(anode) 2 Cl → Cl2 (g) + 2 e

Barium metal is also obtained by the reduction of barium oxide with finely divided aluminium at temperatures between 1100 and 1200 °C:

4 BaO + 2 Al → BaO·Al2O3 + 3 Ba (g)

The barium vapor is cooled by means of a water jacket and condensed into the solid metal. The solid block may be cast into rods or extruded into wires. Being a flammable solid, it is packaged under argon in steel containers or plastic bags.[1]

Applications

Amoebiasis as seen in radiograph of barium-filled colon
Green barium fireworks

The most important use of elemental barium is as a scavenger removing last traces of oxygen and other gases in television and other electronic tubes. Besides, an isotope of barium, 133Ba, is routinely used as a standard source in the calibration of gamma-ray detectors in nuclear physics studies.[1]

Barium is an important component of YBCO superconductors. An alloy of barium 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.

Barium compounds, and especially barite (BaSO4), are extremely important to the petroleum industry.

Precautions

Barium powder is pyrophoric: it can explode in contact with air or oxidizing gases. It is likely to explode when combined with halogenated hydrocarbon solvents. It reacts violently with water. Oxidation occurs very easily and metallic barium should be kept under a petroleum-based fluid (such as kerosene) or other suitable oxygen-free liquids that exclude air.

All water or acid soluble barium compounds are 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.[1] However, individual responses to barium salts vary widely, with some being able to handle barium nitrate casually without problems, and others becoming ill from working with it in small quantities. Barium acetate was used by Marie Robards to poison her father in Texas in 1993. She was tried and convicted in 1996.[11]

Barium sulfate can be taken orally because it is highly insoluble in water, and is eliminated completely from the digestive tract.[1] Unlike other heavy metals, barium does not bioaccumulate.[12][13] However, inhaled dust containing barium compounds can accumulate in the lungs, causing a benign condition called baritosis.[14]

See also


References

  1. ^ a b c d e f Patnaik, Pradyot (2003). Handbook of Inorganic Chemical Compounds. McGraw-Hill. pp. 77–78. ISBN 0070494398. http://books.google.com/books?id=Xqj-TTzkvTEC&pg=PA243. Retrieved 2009-06-06. 
  2. ^ David R. Lide, Norman E. Holden (2005). "Section 11, Table of the Isotopes". CRC Handbook of Chemistry and Physics, 85th Edition. Boca Raton, Florida: CRC Press. 
  3. ^ a b Robert E. Krebs (2006). The history and use of our earth's chemical elements: a reference guide. Greenwood Publishing Group. p. 80. ISBN 0313334382. http://books.google.com/books?id=yb9xTj72vNAC&. 
  4. ^ a b c d C. R. Hammond (2000). The Elements, in Handbook of Chemistry and Physics 81st edition. CRC press. ISBN 0849304814. 
  5. ^ "Toxicological Profile for Barium and Barium Compounds. Agency for Toxic Substances and Disease Registry". CDC. 2007.. http://www.atsdr.cdc.gov/toxprofiles/tp24.pdf. 
  6. ^ Chris J. Jones, John Thornback (2007). Medicinal applications of coordination chemistry. Royal Society of Chemistry. p. 102. ISBN 0854045961. http://books.google.co.jp/books?id=uEJHsZWyO-EC&. 
  7. ^ Michael S. Russell, Kurt Svrcula (2008). Chemistry of Fireworks. Royal Society of Chemistry. p. 110. ISBN 0854041273. http://books.google.co.jp/books?id=yxRyOf8jFeQC&. 
  8. ^ Brent, G. F. (1995). "Surfactant coatings for the stabilization of barium peroxide and lead dioxide in pyrotechnic compositions". Propellants Explosives Pyrotechnics 20: 300. doi:10.1002/prep.19950200604. 
  9. ^ "Battery Breakthrough?". http://www.technologyreview.com/Biztech/18086/. Retrieved 2009-06-06. 
  10. ^ "Crystran Ltd. Optical Component Materials". http://www.crystran.co.uk/barium-fluoride-baf2.htm. Retrieved 2010-12-29. 
  11. ^ "Boyfriend fight preceded Roanoke mom's slaying". http://www.buffalo.edu/news/pdf/October08/DallanMorningNewsEwingSlaying.pdf. Retrieved 2009-06-06. 
  12. ^ "Toxicity Profiles, Ecological Risk Assessment". http://www.epa.gov/region5/superfund/ecology/html/toxprofiles.htm#ba. Retrieved 2009-06-06. 
  13. ^ Moore, J. W. (1991). Inorganic Contaminants of Surface Waters, Research and Monitoring Priorities. New York: Springer-Verlag. 
  14. ^ Doig AT (February 1976). "Baritosis: a benign pneumoconiosis". Thorax 31 (1): 30–9. doi:10.1136/thx.31.1.30. PMID 1257935. 

External links


1911 encyclopedia

Up to date as of January 14, 2010

From LoveToKnow 1911

Medical warning!
This article is from the 1911 Encyclopaedia Britannica. Medical science has made many leaps forward since it has been written. This is not a site for medical advice, when you need information on a medical condition, consult a professional instead.

BARIUM (symbol Ba, atomic weight 137.37 [0=,6]), one of the metallic chemical elements included in the group of the alkaline earths. It takes its name from the Greek f3apvs (heavy) on account of its presence in barytes or heavy spar which was first investigated in 1602 by V. Casciorolus, a shoemaker of Bologna, who found that after ignition with combustible substances it became phosphorescent, and on this account it was frequently called Bolognian phosphorus. In 1774 K. W. Scheele, in examining a specimen of pyrolusite, found a new substance to be present in the mineral, for on treatment with sulphuric acid it gave an insoluble salt which was afterwards shown to be identical with that contained in heavy spar. Barium occurs chiefly in the form of barytes or heavy spar, BaS 04, and witherite, BaCO 3, and to a less extent in baryto-calcite, baryto-celestine, and various complex silicates. The metal is difficult to isolate, and until recently it may be doubted whether the pure metal had been obtained. Sir H. Davy tried to electrolyse baryta, but was unsuccessful; later attempts were made by him using barium chloride in the presence of mercury. In this way he obtained an amalgam, from which on distilling off the mercury the barium was obtained as a silver white residue. R. Bunsen in 1854 electrolysed a thick paste of barium chloride and dilute hydrochloric acid in the presence of mercury, at 10o C., obtaining a barium amalgam, from which the mercury was separated by a process of distillation. A. N. Guntz (Comptes rendus, 1901, 133, p. 872) electrolyses a saturated solution of barium chloride using a mercury cathode and obtains a 3% barium amalgam; this amalgam is transferred to an iron boat in a wide porcelain tube and the tube slowly heated electrically, a good yield of pure barium being obtained at about looo C. The metal when freshly cut possesses a silver white lustre, is a little harder than lead, and is extremely easily oxidized on exposure; it is soluble in liquid ammonia, and readily attacks both water and alcohol.

Three oxides of barium are known, namely, the monoxide, BaO, the dioxide, Ba02, and a suboxide, obtained by heating Ba0 with magnesium in a vacuum to 110o (Guntz, loc. cit., 1906, p. 359). The monoxide is formed when the metal burns in air, but is usually prepared by the ignition of the nitrate, oxygen and oxides of nitrogen being liberated. It can also be obtained by the ignition of an intimate mixture of the carbonate and carbon, and in small quantities by the ignition of the iodate. It is a greyish coloured solid, which combines very energetically with water to form the hydroxide, much heat being evolved during the combination; on heating to redness in a current of oxygen it combines with the oxygen to form the dioxide, which at higher temperatures breaks up again into the monoxide and oxygen.

Barium hydroxide, Ba(OH) 2, is a white powder that can be obtained by slaking the monoxide with the requisite quantity of water, but it is usually made on the large scale by heating heavy spar with small coal whereby a crude barium sulphide is obtained. This sulphide is then heated in a current of moist carbon dioxide, barium carbonate being formed, BaS+H 2 O+CO 2 =BaCO 3 +H 2 S, and finally the carbonate is decomposed by a current of superheated steam, BaC03+H20 = Ba(OH) 2 + C02, leavingaresidue of the hydroxide. It is a white powder moderately soluble in cold water, readily soluble in hot water, the solution possessing an alkaline reaction and absorbing carbon dioxide readily. The solution, known as baryta-water, finds an extensive application in practical chemistry, being used in gas-analysis for the determination of the amount of carbon dioxide in the atmosphere; and also being used in organic chemistry as a hydrolysing agent for the decomposition of complex ureides and substituted aceto-acetic esters, while E. Fischer has used it as a condensing agent in the preparation of aand 0-acrose from acrolein dibromide. A saturated solution of the hydroxide deposits on cooling a hydrated form Ba(OH) 2.8H 2 0, as colourless quadratic prisms, which on exposure to air lose seven molecules of water of crystallization.

Barium dioxide, Ba02, can be prepared as shown above, or in the hydrated condition by the addition of excess of barytawater to hydrogen peroxide solution, when it is precipitated in the crystalline condition as Ba0 2.8H 2 O. These crystals on heating to 130° C. lose the water of crystallization and leave a residue of the anhydrous peroxide. In the Brin process for the manufacture of oxygen, barium dioxide is obtained as an intermediate product by heating barium monoxide with air under pressure. It is a grey coloured powder which is readily decomposed by dilute acids with the production of hydrogen peroxide.

Barium chloride, BaCl 2.2H 2 O, can be obtained by dissolving witherite in dilute hydrochloric acid, and also from heavy spar by ignition in a reverberatory furnace with a mixture of coal, limestone and calcium chloride, the barium chloride being extracted from the fused mass by water, leaving a residue of insoluble calcium sulphide. The chloride crystallizes in colourless rhombic tables of specific gravity 3.9 and is readily soluble in water, but is almost insoluble in concentrated hydrochloric acid and in absolute alcohol. It can be obtained in the anhydrous condition by heating it gently to about 120° C. It has a bitter taste and is a strong poison. Barium bromide is prepared by saturating baryta-water or by decomposing barium carbonate with hydrobromic acid. It crystallizes as BaBr 2.2H 2 O isomorphous with barium chloride. Barium bromate, Ba(Br03)2, can be prepared by the action of excess of bromine on barytawater, or by decomposing a boiling aqueous solution of loo parts of potassium bromate with a similar solution of 74 parts of crystallized barium chloride. It crystallizes in the monoclinic system, and separates from its aqueous solution as Ba(Br03)2.H20. On heating, it begins to decompose at 260-265° C. Barium chlorate, Ba(C103)2, is obtained by adding barium chloride to sodium chlorate solution; on concentration of the solution sodium chloride separates first, and then on further evaporation barium chlorate crystallizes out and can be purified by recrystallization. It can also be obtained by suspending barium carbonate in boiling water and passing in chlorine. It crystallizes in monoclinic prisms of composition Ba(C10 3) 2 H 2 O, and begins to decompose on being heated to 250° C. Barium iodate, Ba(103)2, is obtained by the action of excess of iodic acid on hot caustic baryta solution or by adding sodium iodate to barium chloride solution. It crystallizes in monoclinic prisms of composition Ba(103) 2 H 2 O, and is only very sparingly soluble in cold water.

Barium carbide, BaC2, is prepared by a method similar to that in use for the preparation of calcium carbide (see Acetylene). L. Maquenne has also obtained it by distilling a mixture of barium amalgam and carbon in a stream of hydrogen. Barium sulphide, BaS, is obtained by passing sulphuretted hydrogen over heated barium monoxide, or better by fusion of the sulphate with a small coal. It is a white powder which is readily decomposed by water with the formation of the hydroxide and hydrosulphide. The phosphorescence of the sulphide obtained by heating the thiosulphate is much increased by adding uranium, bismuth, or thorium before ignition pr. Chem., 1905, ii. p. 196).

 ?

Barium sulphate, BaSO 4, is the most abundant of the naturally occurring barium compounds (see Barytes) and can be obtained artificially by the addition of sulphuric acid or any soluble sulphate to a solution of a soluble barium salt, when it is precipitated as an amorphous white powder of specific gravity 4.5. It is practically insoluble in water, and is only very slightly soluble in dilute acids; it is soluble to some extent, when freshly prepared, in hot concentrated sulphuric acid, and on cooling the solution, crystals of composition BaSO 4 H 2 SO 4 are deposited. It is used as a pigment under the name of "permanent white" or blanc fixe. Barium nitride, Ba 3 N 2, is obtained as a brownish mass by passing nitrogen over heated barium amalgam. It is decomposed by water with evolution of hydrogen, and on heating in a current of carbonic oxide forms barium cyanide (L. Maquenne). Barium amide, Ba(NH 2) 2, is obtained from potassammonium and barium bromide.

Barium nitrate, Ba(N03)2, is prepared by dissolving either the carbonate or sulphide in dilute nitric acid, or by mixing hot saturated solutions of barium chloride and sodium nitrate. It crystallizes in octahedra, having a specific gravity of 3.2, and melts at 597° C. (T. Carnelley). It is decomposed by heat, and is largely used in pyrotechny for the preparation of green fire. Barium carbonate, BaCO 31 occurs rather widely distributed as witherite, and may be prepared by the addition of barium chloride to a hot solution of ammonium carbonate, when it is precipitated as a dense white powder of specific gravity 4.3; almost insoluble in water.

Barium and its salts can be readily detected by the yellowishgreen colour they give when moistened with hydrochloric acid and heated in the Bunsenflame, or by observation of their spectra, when two characteristic green lines are seen. In solution, barium salts may be detected by the immediate precipitate they give on the addition of calcium sulphate (this serves to distinguish barium salts from calcium salts), and by the yellow precipitate of barium chromate formed on the addition of potassium chromate. Barium is estimated quantitatively by conversion into the sulphate. The atomic weight of the element has been determined by C. Marignac by the conversion of barium chloride into barium sulphate, and also by a determination of the amount of silver required to precipitate exactly a known weight of the chloride; the mean value obtained being 136.84; T. W. Richards (Zeit. anorg. Chem., 1893, 6, p. 89), by determining the equivalent of barium chloride and bromide to silver, obtained the value 137.44 For the relation of barium to radium, see Radioactivity.


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Wiktionary

Up to date as of January 15, 2010

Definition from Wiktionary, a free dictionary

See also barium

German

Chemical Element: Ba (atomical number 56)

Pronunciation

Noun

Barium n

  1. barium

Simple English

File:Barium
Corroded barium metal

Barium is chemical element 56 on the periodic table. Its symbol is Ba. It contains 56 protons and 56 electrons. Its mass number is about 137.3. It is a metal.

Contents

Properties

Physical properties

Barium is part of a group of elements known as the alkaline earth metals. It is a silvery metal that easily turns black. It is soft and ductile. It can form alloys with some metals that are partially alloys and partially chemical compounds.

Chemical properties

Barium is reactive, and if you put pure barium metal in the air, it will react with oxygen. At first it will turn black, then white as barium oxide is formed. Barium reacts with water to make barium hydroxide and hydrogen gas. Barium also reacts very fast with acids to make a barium salt and hydrogen. Barium can form barium peroxide if it is burned in air.

Barium reacts with many other metal oxides and sulfides to make barium oxide or sulfide and the metal. It also reacts with carbon and nitrogen at a high temperature to make barium cyanide.

Chemical compounds

Main article: Barium compounds

Barium is too reactive as a metal, so it is not found in the earth as a metal. It is found in chemical compounds. Barium only occurs in one oxidation state: +2. Most barium compounds are colorless. The ones that dissolve in water or stomach acid are very toxic. Barium sulfate is well known because it does not dissolve in water or acids. Barium compounds are quite heavy. Barium compounds put out a greenish flame when heated red-hot.

Occurrence

[[File:|thumb|Barium sulfate as barite]] Barium is found as barium sulfate (barite) and barium carbonate (witherite) in the ground. Both of these minerals do not dissolve in water. Barium sulfate hardly dissolves in anything. Barium is found mainly in China, Germany, India, Morocco, and the US.

Preparation

It is very hard to get barium from barium sulfate. So barium sulfate is reduced by carbon to make barium sulfide and carbon dioxide. The barium sulfide is dissolved in hydrochloric acid. This makes hydrogen sulfide and barium chloride. The barium chloride is melted and electrolyzed to get liquid barium metal. The barium metal is solidified and stored in oil.

Barium carbonate, the other ore of barium, is dissolved in hydrochloric acid to make barium chloride and carbon dioxide. The barium chloride is melted and electrolyzed, making barium metal.

Uses

As a metal

Barium is used to remove oxygen from cathode ray tubes and vacuum tubes. It is placed inside and reacts with all of the oxygen, using it up. Barium is also used in spark plug wire.

As chemical compounds

Certain compounds of barium, such as barium sulfate, are not toxic and can be put in the body. We can see where the barium travels in the body by X-rays and this can tell us whether there are problems, such as blockages. The barium sulfate stops the X-rays from going and makes a picture. Barium sulfate can be used as a pigment, too.

Other barium compounds have several other uses.

Safety

Barium is a very toxic element, though, and is dangerous. There is a really small amount of barium in our food, and this does not cause problems. If we get barium from other places, though, it can cause many problems. Even 1 gram of barium can kill you. It is dangerous because it acts like other really important elements, such as calcium and magnesium. If barium replaces these elements, it messes up the body.


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