Barium

Not to be confused with Borium.

caesium ← barium → lanthanum Sr ↑ Ba ↓ Ra 56Ba Periodic table
Appearance
silvery gray
General properties
Name, symbol, number	barium, Ba, 56
Pronunciation	/ˈbɛəriəm/ bair-ee-əm
Element category	alkaline earth metals
Group, period, block	2, 6, s
Standard atomic weight	137.33
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
Molar heat capacity	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	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−1
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 ( /ˈbɛəriəm/ bair-ee-əm) is a chemical element with the symbol Ba and atomic number 56. It is the fifth element in Group 2, a soft silvery metallic alkaline earth metal. Barium 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.

Barium has few industrial applications, but the metal 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 density, insolubility, and X-ray opacity. It is used as an insoluble heavy additive to oil well drilling mud, 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 sought. It is a component of some "high temperature" YBCO superconductors, and electroceramics.

Characteristics

Barite

Physical properties

Barium is a soft, silvery white alkali earth metal, which quickly oxidizes in air. It crystallizes in body centered cubic lattices. It burns with a green to pale green flame, resulting from emission at 524.2 and 513.7 nm. Its simple compounds are notable for their relatively high (for an alkaline earth element) specific gravity. This high density is true of the most common barium-bearing mineral, barite (BaSO₄), also called 'heavy spar' due to the high density (4.5 g/cm³).

Chemical properties

Barium, as for other alkali earth (group II) metals, is highly reducing. It reacts exothermically with oxygen at room temperature to form barium oxide and peroxide. Because of its sensitivity to air, samples are generally stored under protective oils. The reaction is violent if barium is powdered. The metal is readily attacked in most acids, with the notable exception of sulfuric acid, as passivation stops the reaction by forming the insoluble barium sulfate. It also reacts violently with water according to the reaction:

Ba + 2 H₂O → Ba(OH)₂ + H₂↑

Barium combines with several metals, including aluminium, zinc, lead and tin, forming intermetallic phases and alloys.

Isotopes

Main article: Isotopes of barium

Naturally occurring barium is a mix of seven stable isotopes, the most abundant being ¹³⁸Ba (71.7 %). 22 isotopes are 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). ¹³³Ba is a standard calibrant for gamma-ray detectors in nuclear physics studies.

Occurrence

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). Although witherite deposits were mined from the 17th century till 1969 in northern England, for example in the Settlingstones Mine near Newbrough, today nearly all barium is mined as barite.

Large deposits of barite are found in China, Germany, India, Morocco, and in the United States. A rare gem containing barium is known, called benitoite.

Production

Trend in world production of barite

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:

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 + 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

The barium vapor is cooled and condensed to give the solid metal, which can be cast into rods or extruded into wires. Being a flammable solid, it is packaged under argon in steel containers or plastic bags.

Compounds

Ba²⁺ is the dominant oxidation state throughout the chemistry of barium. Its properties generally resemble those of other alkaline earth ions such as strontium and calcium. All halides, pseudohalides and chalcogenides are known, usually as colourless solids. The sulfate is famously insoluble. BaO forms a peroxide when heated in air. The oxide is basic and reacts with acids to give salts. Barium reduces oxides, chlorides and sulfides of less active metals. For example:

Ba + CdO → BaO + Cd

Ba + ZnCl₂ → BaCl₂ + Zn

3 Ba + Al₂S₃ → 3 BaS + 2 Al

At elevated temperatures, barium combines with nitrogen and hydrogen to produce the nitride Ba₃N₂ and hydride BaH₂, respectively. When heated with nitrogen and carbon, it forms the cyanide:

Ba + N₂ + 2 C → Ba(CN)₂

History

Barium's name originates from Greek βαρύς 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". Witches and alchemists were attracted to them because after exposure to light they would glow for years.

Carl Scheele identified barite as containing a new element in 1774, but could not isolate barium, only barium oxide. Johan Gottlieb Gahn also isolated barium oxide two years later in similar studies. Oxidized barium was at first called barote, by Guyton de Morveau, a name which was changed by Antoine Lavoisier to baryta. Also in the 18th century, English mineralogist William Withering noted a heavy mineral in the lead mines of cumberland, now known to be Witherite. 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. Robert Bunsen and Augustus Matthiessen yielded pure barium by electrolysis of a molten mixture of barium chloride and ammonium chloride.

The production of pure oxygen in the Brin process was a large scale application of barium peroxide before electrolysis and fractionally distill liquefied air became the dominant ways to produce oxygen. In this process the barium oxide reacts at 500–600°C with air to form barium peroxide which decomposes at above 700°C by releasing oxygen.

2 BaO + O₂ ⇌ 2 BaO₂

Applications

Amoebiasis as seen in radiograph of barium-filled colon

Green barium fireworks

The dominating application of elemental barium is as a scavenger or "getter" removing the last traces of oxygen and other gases in electronic vacuum tubes such as television cathode ray tubes.

An alloy of barium with nickel is commonly used in automobile ignitions.

Applications of barium sulfate

Barium sulfate (the mineral barite, BaSO₄) is important to the petroleum industry, for example, as a drilling mud weighting agent in drilling new oil and gas wells. It is also a filler in a variety of products such as rubber. Taking advantage of its opacity to X-rays, the sulfate is used as a radiocontrast agent for X-ray imaging of the digestive system ("barium meals" and "barium enemas"). Lithopone, a pigment that contains barium sulfate and zinc sulfide, is a permanent white that has good covering power, and does not darken when exposed to sulfides.

Applications of other barium compounds

Aside from the sulfate, other compounds of barium find only niche applications. Applications are limited by the toxicity of Ba²⁺ ions (Barium carbonate is a rat poison), which is not a problem for the insoluble BaSO₄.

• Barium oxide is used in a coating for the electrodes of fluorescent lamps, which facilitates the release of electrons.
• Barium carbonate is also used in glassmaking. Being a heavy element, barium increases the refractive index and luster of the glass.
• Barium, commonly as barium nitrate, is used to give green colors in fireworks. The species responsible for the brilliant green is barium monochloride; in the absence of a source of chlorine a yellow or "apple" green is produced instead.
• 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 and as a bleaching agent.
• Barium titanate is a promising electroceramic.
• Barium fluoride is used for optics in infrared applications, since it is transparent from about 0.15 to 12 micrometres.

Precautions

Soluble barium compounds are poisonous. At low doses, barium acts as a muscle stimulant, whereas 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. 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. For example, barium acetate was used by Marie Robards to poison her father in 1993.

Non-toxicity of barium sulfate

Because it is highly insoluble in water as well as stomach acids, barium sulfate can be taken orally. It is eliminated completely from the digestive tract. Unlike other heavy metals, barium does not bioaccumulate. However, inhaled dust containing barium compounds can accumulate in the lungs, causing a benign condition called baritosis.

External links

• WebElements.com – Barium
• Elementymology & Elements Multidict
• 3-D Holographic Display Using Strontium Barium Niobate