Barium

• 1. Introduction

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  Barium is a chemical element carrying the symbol Ba and atomic number 56, and its story begins not in a laboratory but in a volcanic field near Bologna, Italy, sometime in the early Middle Ages. Alchemists stumbled across smooth, pebble-like stones there and noticed something uncanny: expose them to light, and they would glow in the dark for years afterward. Those stones were baryte, one of barium's two principal minerals, and they set off centuries of investigation into what, exactly, made them so strange.

  What makes barium difficult to understand at first glance is that it never appears in nature as a pure metal. Its chemical reactivity is so high that it bonds immediately with whatever surrounds it. The name itself comes from the Greek word "barys," meaning heavy, transmitted through the alchemical term "baryta" into the chemical vocabulary we use today.

  By 1772, the Swedish chemist Carl Scheele had concluded that baryte contained a new element, though he could not pull it free. That would take another thirty-six years. When isolation finally came in 1808, it arrived in England, through the work of Sir Humphry Davy, and it required a technology that barely existed at the time: electrolysis. The questions worth carrying forward are these: why does a metal this reactive find so little industrial use, and yet turn up in such unlikely places, from oil wells to fireworks to the imaging of the human gut?

• 2. A Metal Born From Electricity

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  Sir Humphry Davy achieved the first isolation of barium in 1808, working in England with electrolysis of molten barium salts. He chose the name "barium" by direct analogy with calcium, taking the root from "baryta" and appending the "-ium" ending that signals a metallic element. The route from discovery to pure metal would take far longer still.

  Robert Bunsen and Augustus Matthiessen later refined the process, electrolyzing a molten mixture of barium chloride and ammonium chloride to obtain pure barium. Yet electrolysis is no longer the commercial method of choice. Barium dissolves readily in molten halides during the process, leaving an impure product. The modern industrial route runs through aluminium instead: barium oxide is reduced by aluminium at 1100 degrees Celsius, producing barium vapor, which is then condensed and packed into molds under argon gas to prevent oxidation. The result is ultrapure barium; the commonly sold form is about 99 percent pure, with strontium and calcium as the chief impurities, each present at fractions of one percent.

  Before that industrial method dominated, barium peroxide had one significant large-scale role. In the 1880s, the Brin process used barium oxide to manufacture pure oxygen: the oxide reacted with air at 500-600 degrees Celsius to form barium peroxide, which then released that oxygen when heated above 700 degrees Celsius. The Brin process was displaced only in the early 1900s, when electrolysis and fractional distillation of liquefied air became the preferred routes to pure oxygen. The transition marks how barium applications tend to work: a brief period of industrial prominence, followed by replacement once a cleaner technique emerges.

• 3. Reactive by Nature

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  Barium sits fifth in group 2 of the periodic table, making it an alkaline earth metal in the same family as magnesium, calcium, and strontium. It is more reactive than any of them. Metallic barium is often stored under oil or in an inert atmosphere for this reason: it reacts with atmospheric oxygen even at room temperature, the silvery-white surface quickly darkening beneath a layer of gray oxide.

  The metal is soft, rating only 1.25 on the Mohs hardness scale. Its melting point of 1000 Kelvin falls between the lighter strontium at 1050 K and the heavier radium at 973 K, but its boiling point of 2170 K outstrips strontium's 1655 K by a wide margin. Density follows a similar pattern: at 3.62 grams per cubic centimeter, barium is denser than strontium at 2.36 g/cm3 but lighter than radium's approximate 5 g/cm3.

  Barium's compounds appear almost exclusively in the plus-two oxidation state. Reactions with chalcogens, the oxygen-family elements, are highly exothermic, releasing significant energy. The metal also reacts with water and alcohols, again releasing heat and hydrogen gas. Sulfuric acid is the notable exception: it actually stops attacking barium by forming a coating of insoluble barium sulfate on the surface, a process called passivation. Barium also combines with aluminium, zinc, lead, and tin to form intermetallic alloys. The organobarium compounds are a newer frontier: scientists have recently identified dialkylbariums and alkylhalobariums, though the field is still developing.

• 4. Minerals, Mines, and Markets

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  Barium is present in the Earth's crust at 0.0425 percent, and in seawater at 13 micrograms per liter. Its two commercially important minerals are baryte, which is barium sulfate, and witherite, which is barium carbonate. The discovery of witherite as a distinct mineral came in the 18th century, when the English mineralogist William Withering identified a heavy mineral in the lead mines of Cumberland; the mineral was later named in his honor.

  Baryte dominates production. Reserves are estimated somewhere between 0.7 and 2 billion tonnes. The single highest annual output on record was 8.3 million tonnes, reached in 1981, though only 7 to 8 percent of that was used to make barium metal or compounds. Output declined and then climbed again: from 5.6 million tonnes in 1996 to 7.6 million in 2005 and 7.8 million in 2011. China accounts for more than half of global production. India contributed about 14 percent in 2011, followed by Morocco at 8.3 percent and the United States at 8.2 percent.

  Mined baryte goes through washing, crushing, and classification before quartz is separated out. When quartz penetrates deeply into the ore, or when iron, zinc, or lead levels are abnormally high, froth flotation is used. The finished product must be at least 95 percent pure baryte by mass. From barium sulfide, the water-soluble starting point produced by carbon reduction, manufacturers can generate the sulfate, nitrate, carbonate, and other compounds by treating with the appropriate reagents. Barium metal itself is produced separately, through the aluminium reduction route at 1100 degrees Celsius.

• 5. Quiet Industrial Life

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  Barium's most recognized industrial role in the 20th century was as a getter inside vacuum tubes, including television picture tubes. Barium or a barium-aluminium alloy was introduced into the tube to absorb oxygen, nitrogen, carbon dioxide, and water vapor that would otherwise degrade the vacuum. Its low vapor pressure and extreme reactivity toward gases made it effective; it could even absorb traces of noble gases by incorporating them into its crystal lattice. The rise of LCD, LED, and plasma displays has rendered this application largely obsolete.

  Barium sulfate carries its own industrial weight. Known in the pigment trade as blanc fixe, a French phrase meaning "permanent white," the precipitated compound serves as a filler in paints, varnishes, rubber, and plastics, and as a coating pigment for paper. In the petroleum industry, baryte is a key ingredient in the drilling fluids used in oil and gas wells: it adds density to the fluid, which helps control pressure in the borehole. The compound is also combined with zinc sulfide to make lithopone, a white pigment that does not darken when exposed to sulfides.

  Barium imparts color in other applications. Added as barium nitrate to fireworks in the absence of chlorine donors, it produces a yellow or apple-green flame. When chlorine donors are present, barium chloride forms in the reaction and generates emerald greens. Barium peroxide serves as a catalyst in the aluminothermic reaction used to weld rail tracks, and it appears as a green flare in tracer ammunition. Barium compounds are also added to steel and cast iron to reduce the size of carbon grains within the microstructure. Each of these uses is narrow and specific, reflecting an element whose industrial reach is limited but not negligible.

• 6. Medicine, Superconductors, and the Ocean Floor

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  Barium sulfate found its way into medicine beginning in 1908, when it was first applied as a radiocontrast agent in X-ray imaging of the digestive system. Because the compound has low toxicity and a relatively high density of approximately 4.5 grams per cubic centimeter, it is opaque to X-rays while being safe to swallow. Patients undergoing imaging of the gastrointestinal tract are given "barium meals" or "barium enemas" containing this compound, which outlines the internal structures against the X-ray beam. The water-soluble barium compounds are a different matter entirely: they are poisonous, and barium carbonate has been used as a rodenticide, a practice considered obsolete though possibly still in use in some countries.

  Barium plays a structural role in YBCO, the first high-temperature superconductor that could be cooled using liquid nitrogen. YBCO reached a transition temperature of 93 K, which exceeds the boiling point of nitrogen at 77 K, making liquid-nitrogen cooling practical for the first time. Barium titanate is under investigation as a promising electroceramic, and barium fluoride is used in infrared optics because it remains transparent across a wide range of wavelengths from 0.15 to 12 micrometers.

  In the ocean, barium behaves differently from most elements. Its average concentration in seawater is 109 nanomoles per kilogram, and it carries a nutrient-like profile with a residence time of 10,000 years. Scientists have discovered strong correlations between dissolved barium and silicic acid, as well as between particulate barium and particulate organic carbon. Those correlations make barium a useful proxy for studying the ocean's biological pump, the carbon cycle, and past climate. Barite particles deposited in seafloor sediments, as well as barite from hydrothermal vents, preserve chemical signatures that geologists read to reconstruct ocean chemistry across geological time.

• 7. Isotopes and Radioactive Edges

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  Natural barium found in the Earth's crust is a mixture of seven primordial isotopes: barium-130, 132, and 134 through 138. Among the stable forms, barium-138 is dominant, making up 71.7 percent of all natural barium. The others decrease in abundance with decreasing mass number, with a probable reversal for the two heaviest of the p-nuclei, barium-130 and barium-132.

  Barium-130 is not quite stable. It undergoes double beta-plus decay very slowly, transforming into xenon-130 with a half-life estimated at between 0.5 and 2.7 times ten to the power of 21 years. That figure is approximately 100 billion times the age of the universe. Its abundance in natural barium is only about 0.11 percent. Barium-132 could theoretically undergo the same decay to xenon-132, but no experimental evidence has detected it doing so.

  Barium has 41 known isotopes in total, spanning mass numbers 114 to 154. The most stable artificial radioisotope is barium-133, with a half-life of 10.538 years. Among metastable nuclear isomers, barium-133m lasts 38.90 hours and barium-135m1 lasts 28.11 hours. One isomer, barium-137m1, has a half-life of only 2.552 minutes, but it arises from the decay of caesium-137, one of the most common products of nuclear fission. The connection between this short-lived isomer and caesium-137 means barium-137m1 is relevant anywhere nuclear reactors or spent fuel are studied.