Radium

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Radium, 88Ra
Radium226.jpg
Radium
Pronunciation /ˈrdiəm/ (RAY-dee-əm)
Appearancesilvery white metallic
Mass number [226]
Radium in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
Ba

Ra

(Ubn)
franciumradiumactinium
Atomic number (Z)88
Group group 2 (alkaline earth metals)
Period period 7
Block   s-block
Electron configuration [ Rn ] 7s2
Electrons per shell2, 8, 18, 32, 18, 8, 2
Physical properties
Phase at  STP solid
Melting point 973  K (700 °C,1292 °F)(disputed)
Boiling point 2010 K(1737 °C,3159 °F)
Density (near r.t.)5.5 g/cm3
Heat of fusion 8.5  kJ/mol
Heat of vaporization 113 kJ/mol
Vapor pressure
P (Pa)1101001 k10 k100 k
at T (K)8199061037120914461799
Atomic properties
Oxidation states +2 (expected to have a strongly basic oxide)
Electronegativity Pauling scale: 0.9
Ionization energies
  • 1st: 509.3 kJ/mol
  • 2nd: 979.0 kJ/mol
Covalent radius 221±2  pm
Van der Waals radius 283 pm
Color lines in a spectral range Radium spectrum visible.png
Color lines in a spectral range
Spectral lines of radium
Other properties
Natural occurrence from decay
Crystal structure body-centered cubic (bcc)
Cubic-body-centered.svg
Thermal conductivity 18.6 W/(m⋅K)
Electrical resistivity 1 µΩ⋅m(at 20 °C)
Magnetic ordering nonmagnetic
CAS Number 7440-14-4
History
Discovery Pierre and Marie Curie (1898)
First isolationMarie Curie(1910)
Main isotopes of radium
Iso­tope Abun­dance Half-life (t1/2) Decay mode Pro­duct
223Ra trace 11.43 d α 219Rn
224Ratrace3.6319 dα 220Rn
225Ratrace14.9 d β 225Ac
226Ratrace1600 yα 222Rn
228Ratrace5.75 yβ 228Ac
Symbol category class.svg   Category: Radium
| references

Radium is a chemical element with the symbol  Ra and atomic number  88. It is the sixth element in group 2 of the periodic table, also known as the alkaline earth metals. Pure radium is silvery-white, but it readily reacts with nitrogen (rather than oxygen) on exposure to air, forming a black surface layer of radium nitride (Ra3N2). All isotopes of radium are highly radioactive, with the most stable isotope being radium-226, which has a half-life of 1600 years and decays into radon gas (specifically the isotope radon-222). When radium decays, ionizing radiation is a by-product, which can excite fluorescent chemicals and cause radioluminescence.

Contents

Radium, in the form of radium chloride, was discovered by Marie and Pierre Curie in 1898 from ore mined at Jáchymov. They extracted the radium compound from uraninite and published the discovery at the French Academy of Sciences five days later. Radium was isolated in its metallic state by Marie Curie and André-Louis Debierne through the electrolysis of radium chloride in 1911. [1]

In nature, radium is found in uranium and (to a lesser extent) thorium ores in trace amounts as small as a seventh of a gram per ton of uraninite. Radium is not necessary for living organisms, and adverse health effects are likely when it is incorporated into biochemical processes because of its radioactivity and chemical reactivity. Currently, other than its use in nuclear medicine, radium has no commercial applications; formerly, it was used as a radioactive source for radioluminescent devices and also in radioactive quackery for its supposed curative powers. Today, these former applications are no longer in vogue because radium's toxicity has become known, and less dangerous isotopes are used instead in radioluminescent devices.

Bulk properties

Radium is the heaviest known alkaline earth metal and is the only radioactive member of its group. Its physical and chemical properties most closely resemble its lighter congener barium. [2]

Pure radium is a volatile silvery-white metal, although its lighter congeners calcium, strontium, and barium have a slight yellow tint. [2] This tint rapidly vanishes on exposure to air, yielding a black layer of radium nitride (Ra3N2). [3] Its melting point is either 700 °C (1,292 °F) or 960 °C (1,760 °F) [lower-alpha 1] and its boiling point is 1,737 °C (3,159 °F). Both of these values are slightly lower than those of barium, confirming periodic trends down the group 2 elements. [4] Like barium and the alkali metals, radium crystallizes in the body-centered cubic structure at standard temperature and pressure: the radium–radium bond distance is 514.8  picometers. [5] Radium has a density of 5.5 g/cm3, higher than that of barium, again confirming periodic trends; the radium-barium density ratio is comparable to the radium-barium atomic mass ratio, [6] due to the two elements' similar crystal structures. [6] [7]

Isotopes

Decay chain of U, the primordial progenitor of Ra Decay chain(4n+2, Uranium series).svg
Decay chain of U, the primordial progenitor of Ra

Radium has 33 known isotopes, with mass numbers from 202 to 234: all of them are radioactive. [8] Four of these – 223Ra (half-life 11.4 days), 224Ra (3.64 days), 226Ra (1600 years), and 228Ra (5.75 years) – occur naturally in the decay chains of primordial thorium-232, uranium-235, and uranium-238 (223Ra from uranium-235, 226Ra from uranium-238, and the other two from thorium-232). These isotopes nevertheless still have half-lives too short to be primordial radionuclides and only exist in nature from these decay chains. [9] Together with the mostly artificial 225Ra (15 d), which occurs in nature only as a decay product of minute traces of 237Np, [10] these are the five most stable isotopes of radium. [11] All other known radium isotopes have half-lives under two hours, and the majority have half-lives under a minute. [8] At least 12 nuclear isomers have been reported; the most stable of them is radium-205m, with a half-life between 130 and 230 milliseconds; this is still shorter than twenty-four ground-state radium isotopes. [8]

In the early history of the study of radioactivity, the different natural isotopes of radium were given different names. In this scheme, 223Ra was named actinium X (AcX), 224Ra thorium X (ThX), 226Ra radium (Ra), and 228Ra mesothorium 1 (MsTh1). [9] When it was realized that all of these are isotopes of the same element, many of these names fell out of use, and "radium" came to refer to all isotopes, not just 226Ra. Some of radium-226's decay products received historical names including "radium", ranging from radium A to radium G, with the letter indicating approximately how far they were down the chain from their parent 226Ra. Radium emanation = 222Rn, RaA = 218Po, RaB = 214Pb, RaC = 214Bi, RaC1 = 214Po, RaC2 = 210Tl, RaD = 210Pb, RaE = 210Bi, RaF = 210Po and RaG = 206Pb. [11] [12]

226Ra is the most stable isotope of radium and is the last isotope in the (4n + 2) decay chain of uranium-238 with a half-life of over a millennium: it makes up almost all of natural radium. Its immediate decay product is the dense radioactive noble gas radon (specifically the isotope 222Rn), which is responsible for much of the danger of environmental radium. [13] It is 2.7 million times more radioactive than the same molar amount of natural uranium (mostly uranium-238), due to its proportionally shorter half-life. [14] [15]

A sample of radium metal maintains itself at a higher temperature than its surroundings because of the radiation it emits alpha particles, beta particles, and gamma rays. More specifically, natural radium (which is mostly 226Ra) emits mostly alpha particles, but other steps in its decay chain (the uranium or radium series) emit alpha or beta particles, and almost all particle emissions are accompanied by gamma rays. [16]

In 2013, it was discovered that the nucleus of radium-224 is pear-shaped. This was the first discovery of an asymmetric nucleus. [17]

Chemistry

Radium, like barium, is a highly reactive metal and always exhibits its group oxidation state of +2. [3] It forms the colorless Ra2+ cation in aqueous solution, which is highly basic and does not form complexes readily. [3] Most radium compounds are therefore simple ionic compounds, [3] though participation from the 6s and 6p electrons (in addition to the valence 7s electrons) is expected due to relativistic effects and would enhance the covalent character of radium compounds such as RaF 2 and RaAt 2. [18] For this reason, the standard electrode potential for the half-reaction Ra2+ (aq) + 2e → Ra (s) is −2.916  V, even slightly lower than the value −2.92 V for barium, whereas the values had previously smoothly increased down the group (Ca: −2.84 V; Sr: −2.89 V; Ba: −2.92 V). [19] The values for barium and radium are almost exactly the same as those of the heavier alkali metals potassium, rubidium, and caesium. [19]

Compounds

Solid radium compounds are white as radium ions provide no specific coloring, but they gradually turn yellow and then dark over time due to self-radiolysis from radium's alpha decay. [3] Insoluble radium compounds coprecipitate with all barium, most strontium, and most lead compounds. [20]

Radium oxide (RaO) has not been characterized well past its existence, despite oxides being common compounds for the other alkaline earth metals. Radium hydroxide (Ra(OH)2) is the most readily soluble among the alkaline earth hydroxides and is a stronger base than its barium congener, barium hydroxide. [21] It is also more soluble than actinium hydroxide and thorium hydroxide: these three adjacent hydroxides may be separated by precipitating them with ammonia. [21]

Radium chloride (RaCl2) is a colorless, luminous compound. It becomes yellow after some time due to self-damage by the alpha radiation given off by radium when it decays. Small amounts of barium impurities give the compound a rose color. [21] It is soluble in water, though less so than barium chloride, and its solubility decreases with increasing concentration of hydrochloric acid. Crystallization from aqueous solution gives the dihydrate RaCl2·2H2O, isomorphous with its barium analog. [21]

Radium bromide (RaBr2) is also a colorless, luminous compound. [21] In water, it is more soluble than radium chloride. Like radium chloride, crystallization from aqueous solution gives the dihydrate RaBr2·2H2O, isomorphous with its barium analog. The ionizing radiation emitted by radium bromide excites nitrogen molecules in the air, making it glow. The alpha particles emitted by radium quickly gain two electrons to become neutral helium, which builds up inside and weakens radium bromide crystals. This effect sometimes causes the crystals to break or even explode. [21]

Radium nitrate (Ra(NO3)2) is a white compound that can be made by dissolving radium carbonate in nitric acid. As the concentration of nitric acid increases, the solubility of radium nitrate decreases, an important property for the chemical purification of radium. [21]

Radium forms much the same insoluble salts as its lighter congener barium: it forms the insoluble sulfate (RaSO4, the most insoluble known sulfate), chromate (RaCrO4), carbonate (RaCO3), iodate (Ra(IO3)2), tetrafluoroberyllate (RaBeF4), and nitrate (Ra(NO3)2). With the exception of the carbonate, all of these are less soluble in water than the corresponding barium salts, but they are all isostructural to their barium counterparts. Additionally, radium phosphate, oxalate, and sulfite are probably also insoluble, as they coprecipitate with the corresponding insoluble barium salts. [22] The great insolubility of radium sulfate (at 20 °C, only 2.1  mg will dissolve in 1  kg of water) means that it is one of the less biologically dangerous radium compounds. [23] The large ionic radius of Ra2+ (148 pm) results in weak complexation and poor extraction of radium from aqueous solutions when not at high pH. [24]

Occurrence

All isotopes of radium have half-lives much shorter than the age of the Earth, so that any primordial radium would have decayed long ago. Radium nevertheless still occurs in the environment, as the isotopes 223Ra, 224Ra, 226Ra, and 228Ra are part of the decay chains of natural thorium and uranium isotopes; since thorium and uranium have very long half-lives, these daughters are continually being regenerated by their decay. [9] Of these four isotopes, the longest-lived is 226Ra (half-life 1600 years), a decay product of natural uranium. Because of its relative longevity, 226Ra is the most common isotope of the element, making up about one part per trillion of the Earth's crust; essentially all natural radium is 226Ra. [25] Thus, radium is found in tiny quantities in the uranium ore uraninite and various other uranium minerals, and in even tinier quantities in thorium minerals. One ton of pitchblende typically yields about one seventh of a gram of radium. [26] One kilogram of the Earth's crust contains about 900  picograms of radium, and one liter of sea water contains about 89  femtograms of radium. [27]

History

Marie and Pierre Curie experimenting with radium, a drawing by Andre Castaigne Curie and radium by Castaigne.jpg
Marie and Pierre Curie experimenting with radium, a drawing by André Castaigne
Glass tube of radium chloride kept by the US Bureau of Standards that served as the primary standard of radioactivity for the United States in 1927. US radium standard 1927.jpg
Glass tube of radium chloride kept by the US Bureau of Standards that served as the primary standard of radioactivity for the United States in 1927.

Radium was discovered by Marie Skłodowska-Curie and her husband Pierre Curie on 21 December 1898, in a uraninite (pitchblende) sample from Jáchymov. [28] While studying the mineral earlier, the Curies removed uranium from it and found that the remaining material was still radioactive. In July 1898 while studying pitchblende they isolated an element similar to bismuth which turned out to be polonium. They then isolated a radioactive mixture consisting mostly of two components: compounds of barium, which gave a brilliant green flame color, and unknown radioactive compounds which gave carmine spectral lines that had never been documented before. The Curies found the radioactive compounds to be very similar to the barium compounds, except that they were less soluble. This made it possible for the Curies to isolate the radioactive compounds and discover a new element in them. The Curies announced their discovery to the French Academy of Sciences on 26 December 1898. [29] [30] The naming of radium dates to about 1899, from the French word radium, formed in Modern Latin from radius (ray): this was in recognition of radium's power of emitting energy in the form of rays. [31] [32] [33]

In September 1910, Marie Curie and André-Louis Debierne announced that they had isolated radium as a pure metal through the electrolysis of a pure radium chloride (RaCl2) solution using a mercury cathode, producing a radium–mercury amalgam. [34] This amalgam was then heated in an atmosphere of hydrogen gas to remove the mercury, leaving pure radium metal. [35] Later that same year, E. Eoler isolated radium by thermal decomposition of its azide, Ra(N3)2. [9] Radium metal was first industrially produced at the beginning of the 20th century by Biraco, a subsidiary company of Union Minière du Haut Katanga (UMHK) in its Olen plant in Belgium. [36]

The common historical unit for radioactivity, the curie, is based on the radioactivity of 226Ra. [37]

Historical applications

Luminescent paint

Self-luminous white paint which contains radium on the face and hand of an old clock. Radium-paint.jpg
Self-luminous white paint which contains radium on the face and hand of an old clock.
Radium watch hands under ultraviolet light Radium 2.jpg
Radium watch hands under ultraviolet light

Radium was formerly used in self-luminous paints for watches, nuclear panels, aircraft switches, clocks, and instrument dials. A typical self-luminous watch that uses radium paint contains around 1 microgram of radium. [38] In the mid-1920s, a lawsuit was filed against the United States Radium Corporation by five dying "Radium Girls" – dial painters who had painted radium-based luminous paint on the dials of watches and clocks. The dial painters were instructed to lick their brushes to give them a fine point, thereby ingesting radium. [39] Their exposure to radium caused serious health effects which included sores, anemia, and bone cancer. This is because the body treats radium as calcium and deposits it in the bones, where radioactivity degrades marrow and can mutate bone cells. [13]

During the litigation, it was determined that the company's scientists and management had taken considerable precautions to protect themselves from the effects of radiation, yet had not seen fit to protect their employees. Additionally, for several years the companies had attempted to cover up the effects and avoid liability by insisting that the Radium Girls were instead suffering from syphilis. This complete disregard for employee welfare had a significant impact on the formulation of occupational disease labor law. [40]

As a result of the lawsuit, the adverse effects of radioactivity became widely known, and radium-dial painters were instructed in proper safety precautions and provided with protective gear. In particular, dial painters no longer licked paint brushes to shape them (which caused some ingestion of radium salts). Radium was still used in dials as late as the 1960s, but there were no further injuries to dial painters. This highlighted that the harm to the Radium Girls could easily have been avoided. [41]

From the 1960s the use of radium paint was discontinued. In many cases luminous dials were implemented with non-radioactive fluorescent materials excited by light; such devices glow in the dark after exposure to light, but the glow fades. [13] Where long-lasting self-luminosity in darkness was required, safer radioactive promethium-147 (half-life 2.6 years) or tritium (half-life 12 years) paint was used; both continue to be used today. [42] These had the added advantage of not degrading the phosphor over time, unlike radium. [43] Tritium emits very low-energy beta radiation (even lower-energy than the beta radiation emitted by promethium) [8] which cannot penetrate the skin, [44] rather than the penetrating gamma radiation of radium and is regarded as safer. [45]

Clocks, watches, and instruments dating from the first half of the 20th century, often in military applications, may have been painted with radioactive luminous paint. They are usually no longer luminous; however, this is not due to radioactive decay of the radium (which has a half-life of 1600 years) but to the fluorescence of the zinc sulfide fluorescent medium being worn out by the radiation from the radium. [46] The appearance of an often thick layer of green or yellowish brown paint in devices from this period suggests a radioactive hazard. The radiation dose from an intact device is relatively low and usually not an acute risk; but the paint is dangerous if released and inhaled or ingested. [47] [48]

Commercial use

Hotel postcard advertising radium baths, c.1940s Radium Water Bath Department, top floor, Hotel Will Rogers, Claremore, Okla., U.S.A (63053).jpg
Hotel postcard advertising radium baths, c.1940s

Radium was once an additive in products such as toothpaste, hair creams, and even food items due to its supposed curative powers. [49] Such products soon fell out of vogue and were prohibited by authorities in many countries after it was discovered they could have serious adverse health effects. (See, for instance, Radithor or Revigator types of "radium water" or "Standard Radium Solution for Drinking".) [46] Spas featuring radium-rich water are still occasionally touted as beneficial, such as those in Misasa, Tottori, Japan. In the U.S., nasal radium irradiation was also administered to children to prevent middle-ear problems or enlarged tonsils from the late 1940s through the early 1970s. [50]

Medical use

Ad for Radior cosmetics which the manufacturer claimed contained radium, that was supposed to have health benefits for one's skin. Powders, skin creams and soap were part of this line. Radior cosmetics containing radium 1918.jpg
Ad for Radior cosmetics which the manufacturer claimed contained radium, that was supposed to have health benefits for one's skin. Powders, skin creams and soap were part of this line.

Radium (usually in the form of radium chloride or radium bromide) was used in medicine to produce radon gas, which in turn was used as a cancer treatment; for example, several of these radon sources were used in Canada in the 1920s and 1930s. [47] [51] However, many treatments that were used in the early 1900s are not used anymore because of the harmful effects radium bromide exposure caused. Some examples of these effects are anaemia, cancer, and genetic mutations. [52] Safer gamma emitters such as 60Co, which is less costly and available in larger quantities, are usually used today to replace the historical use of radium in this application. [24]

Early in the 1900s, biologists used radium to induce mutations and study genetics. As early as 1904, Daniel MacDougal used radium in an attempt to determine whether it could provoke sudden large mutations and cause major evolutionary shifts. Thomas Hunt Morgan used radium to induce changes resulting in white-eyed fruit flies. Nobel-winning biologist Hermann Muller briefly studied the effects of radium on fruit fly mutations before turning to more affordable x-ray experiments. [53]

Howard Atwood Kelly, one of the founding physicians of Johns Hopkins Hospital, was a major pioneer in the medical use of radium to treat cancer. [54] His first patient was his own aunt in 1904, who died shortly after surgery. [55] Kelly was known to use excessive amounts of radium to treat various cancers and tumors. As a result, some of his patients died from radium exposure. [56] His method of radium application was inserting a radium capsule near the affected area, then sewing the radium "points" directly to the tumor. [56] This was the same method used to treat Henrietta Lacks, the host of the original HeLa cells, for cervical cancer. [57] Currently, safer and more available radioisotopes are used instead. [13]

Production

Monument to the Discovery of Radium in Jachymov Pamatnik objevu radia v Jachymove.jpg
Monument to the Discovery of Radium in Jáchymov

Uranium had no large scale application in the late 19th century and therefore no large uranium mines existed. In the beginning the only large source for uranium ore was the silver mines in Jáchymov, Austria-Hungary (now Czech Republic). [28] The uranium ore was only a byproduct of the mining activities. [58]

In the first extraction of radium, Curie used the residues after extraction of uranium from pitchblende. The uranium had been extracted by dissolution in sulfuric acid leaving radium sulfate, which is similar to barium sulfate but even less soluble in the residues. The residues also contained rather substantial amounts of barium sulfate which thus acted as a carrier for the radium sulfate. The first steps of the radium extraction process involved boiling with sodium hydroxide, followed by hydrochloric acid treatment to minimize impurities of other compounds. The remaining residue was then treated with sodium carbonate to convert the barium sulfate into barium carbonate (carrying the radium), thus making it soluble in hydrochloric acid. After dissolution, the barium and radium were reprecipitated as sulfates; this was then repeated to further purify the mixed sulfate. Some impurities that form insoluble sulfides were removed by treating the chloride solution with hydrogen sulfide, followed by filtering. When the mixed sulfates were pure enough, they were once more converted to mixed chlorides; barium and radium thereafter were separated by fractional crystallisation while monitoring the progress using a spectroscope (radium gives characteristic red lines in contrast to the green barium lines), and the electroscope. [59]

After the isolation of radium by Marie and Pierre Curie from uranium ore from Joachimsthal, several scientists started to isolate radium in small quantities. Later, small companies purchased mine tailings from Joachimsthal mines and started isolating radium. In 1904, the Austrian government nationalised the mines and stopped exporting raw ore. For some time radium availability was low. [58]

The formation of an Austrian monopoly and the strong urge of other countries to have access to radium led to a worldwide search for uranium ores. The United States took over as leading producer in the early 1910s. The Carnotite sands in Colorado provide some of the element, but richer ores are found in the Congo and the area of the Great Bear Lake and the Great Slave Lake of northwestern Canada. [28] [60] Neither of the deposits is mined for radium but the uranium content makes mining profitable.

The Curies' process was still used for industrial radium extraction in 1940, but mixed bromides were then used for the fractionation. [61] If the barium content of the uranium ore is not high enough it is easy to add some to carry the radium. These processes were applied to high grade uranium ores but may not work well with low grade ores.

Small amounts of radium were still extracted from uranium ore by this method of mixed precipitation and ion exchange as late as the 1990s, [25] but today they are extracted only from spent nuclear fuel. [62] In 1954, the total worldwide supply of purified radium amounted to about 5 pounds (2.3 kg) [38] and it is still in this range today, while the annual production of pure radium compounds is only about 100 g in total today. [25] The chief radium-producing countries are Belgium, Canada, the Czech Republic, Slovakia, the United Kingdom, and Russia. [25] The amounts of radium produced were and are always relatively small; for example, in 1918, 13.6 g of radium were produced in the United States. [63] The metal is isolated by reducing radium oxide with aluminium metal in a vacuum at 1200 °C. [24]

Modern applications

Atomic, molecular, and optical physics research

Radium is seeing increasing use in the field of atomic, molecular, and optical physics. Symmetry breaking forces scale proportional to , [64] [65] which makes radium, the heaviest alkaline earth element, well suited for constraining new physics beyond the standard model. Some radium isotopes, such as radium-225, have octupole deformed parity doublets that enhance sensitivity to charge parity violating new physics by two-to-three orders of magnitude compared to Hg. [66] [67] [68]

Radium is also promising for a trapped ion optical clock. The radium ion has two subhertz-linewidth transitions from the ground state that could serve as the clock transition in an optical clock. [69] Additionally, radium could be particularly well suited for a transportable optical clock as all transitions necessary for clock operation can be addressed with direct diode lasers. [70]

Though radium has no stable isotopes, there are eleven radium isotopes with half-lives longer than one minute that could be compared with high precision on a King plot. Isotope shifts could be measured with high precision on either of the radium ion subhertz-linewidth transitions from the ground state, or on the to intercombination line in neutral radium. [71] The degree of any potential nonlinearities in such a King plot could set bounds on new physics beyond the standard model. [72]

Some of the few practical uses of radium are derived from its radioactive properties. More recently discovered radioisotopes, such as cobalt-60 and caesium-137, are replacing radium in even these limited uses because several of these isotopes are more powerful emitters, safer to handle, and available in more concentrated form. [73] [74]

The isotope 223Ra (under the trade name Xofigo) was approved by the United States Food and Drug Administration in 2013 for use in medicine as a cancer treatment of bone metastasis. [75] [76] The main indication of treatment with Xofigo is the therapy of bony metastases from castration-resistant prostate cancer due to the favourable characteristics of this alpha-emitter radiopharmaceutical. [77] 225Ra has also been used in experiments concerning therapeutic irradiation, as it is the only reasonably long-lived radium isotope which does not have radon as one of its daughters. [78]

Radium is still used today as a radiation source in some industrial radiography devices to check for flawed metallic parts, similarly to X-ray imaging. [13] When mixed with beryllium, radium acts as a neutron source. [46] [79] Radium-beryllium neutron sources are still sometimes used even today, [13] [80] but other materials such as polonium are now more common: about 1500 polonium-beryllium neutron sources, with an individual activity of 1,850 Ci (68 TBq), have been used annually in Russia. These RaBeF4-based (α, n) neutron sources have been deprecated despite the high number of neutrons they emit (1.84×106 neutrons per second) in favour of 241Am–Be sources. [24] Today, the isotope 226Ra is mainly used to form 227 Ac by neutron irradiation in a nuclear reactor. [24]

Hazards

Radium is highly radioactive, and its immediate daughter, radon gas, is also radioactive. When ingested, 80% of the ingested radium leaves the body through the feces, while the other 20% goes into the bloodstream, mostly accumulating in the bones. [13] Exposure to radium, internal or external, can cause cancer and other disorders, because radium and radon emit alpha and gamma rays upon their decay, which kill and mutate cells. [13] At the time of the Manhattan Project in 1944, the "tolerance dose" for workers was set at 0.1 micrograms of ingested radium. [81] [82]

Some of the biological effects of radium include the first case of "radium-dermatitis", reported in 1900, two years after the element's discovery. The French physicist Antoine Becquerel carried a small ampoule of radium in his waistcoat pocket for six hours and reported that his skin became ulcerated. Pierre and Marie Curie were so intrigued by radiation that they sacrificed their own health to learn more about it. Pierre Curie attached a tube filled with radium to his arm for ten hours, which resulted in the appearance of a skin lesion, suggesting the use of radium to attack cancerous tissue as it had attacked healthy tissue. [83] Handling of radium has been blamed for Marie Curie's death due to aplastic anemia. A significant amount of radium's danger comes from its daughter radon: being a gas, it can enter the body far more readily than can its parent radium. [13]

Today, 226Ra is considered to be the most toxic of the quantity radioelements, and it must be handled in tight glove boxes with significant airstream circulation that is then treated to avoid escape of its daughter 222Rn to the environment. Old ampoules containing radium solutions must be opened with care because radiolytic decomposition of water can produce an overpressure of hydrogen and oxygen gas. [24] The world's largest concentration of 226Ra is stored within the Interim Waste Containment Structure, approximately 9.6 mi (15.4 km) north of Niagara Falls, New York. [84]

See also

Notes

  1. Both values are encountered in sources and there is no agreement among scientists as to the true value of the melting point of radium.

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Actinium is a chemical element with the symbol Ac and atomic number 89. It was first isolated 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 and called actinium. Actinium gave the name to the actinide series, a group of 15 similar elements between actinium and lawrencium in the periodic table. Together with polonium, radium, and radon, actinium was one of the first non-primordial radioactive elements to be isolated.

Alpha decay Emission of alpha particles by a decaying radioactive atom

Alpha decay or α-decay is a type of radioactive decay in which an atomic nucleus emits an alpha particle and thereby transforms or 'decays' into a different atomic nucleus, with a mass number that is reduced by four and an atomic number that is reduced by two. An alpha particle is identical to the nucleus of a helium-4 atom, which consists of two protons and two neutrons. It has a charge of +2 e and a mass of 4 u. For example, uranium-238 decays to form thorium-234. Alpha particles have a charge +2 e, but as a nuclear equation describes a nuclear reaction without considering the electrons – a convention that does not imply that the nuclei necessarily occur in neutral atoms – the charge is not usually shown. Alpha decay typically occurs in the heaviest nuclides. Theoretically, it can occur only in nuclei somewhat heavier than nickel, where the overall binding energy per nucleon is no longer a maximum and the nuclides are therefore unstable toward spontaneous fission-type processes. In practice, this mode of decay has only been observed in nuclides considerably heavier than nickel, with the lightest known alpha emitters being the lightest isotopes of tellurium. Exceptionally, however, beryllium-8 decays to two alpha particles. Alpha decay is by far the most common form of cluster decay, where the parent atom ejects a defined daughter collection of nucleons, leaving another defined product behind. It is the most common form because of the combined extremely high nuclear binding energy and a relatively small mass of the alpha particle. Like other cluster decays, alpha decay is fundamentally a quantum tunneling process. Unlike beta decay, it is governed by the interplay between both the strong nuclear force and the electromagnetic force. Alpha particles have a typical kinetic energy of 5 MeV and have a speed of about 15,000,000 m/s, or 5% of the speed of light. There is surprisingly small variation around this energy, due to the heavy dependence of the half-life of this process on the energy produced. Because of their relatively large mass, the electric charge of +2 e and relatively low velocity, alpha particles are very likely to interact with other atoms and lose their energy, and their forward motion can be stopped by a few centimeters of air. Approximately 99% of the helium produced on Earth is the result of the alpha decay of underground deposits of minerals containing uranium or thorium. The helium is brought to the surface as a by-product of natural gas production.

The actinoid series encompasses the 15 metallic chemical elements with atomic numbers from 89 to 103, actinium through lawrencium. The actinoid series derives its name from the first element in the series, actinium. The informal chemical symbol An is used in general discussions of actinoid chemistry to refer to any actinoid.

Polonium Chemical element, symbol Po and atomic number 84

Polonium is a chemical element with the symbol Po and atomic number 84. Polonium is a chalcogen. A rare and highly radioactive metal with no stable isotopes, polonium is chemically similar to selenium and tellurium, though its metallic character resembles that of its horizontal neighbors in the periodic table: thallium, lead, and bismuth. Due to the short half-life of all its isotopes, its natural occurrence is limited to tiny traces of the fleeting polonium-210 in uranium ores, as it is the penultimate daughter of natural uranium-238. Though slightly longer-lived isotopes exist, they are much more difficult to produce. Today, polonium is usually produced in milligram quantities by the neutron irradiation of bismuth. Due to its intense radioactivity, which results in the radiolysis of chemical bonds and radioactive self-heating, its chemistry has mostly been investigated on the trace scale only.

Radon Chemical element, symbol Rn and atomic number 86

Radon is a chemical element with the symbol Rn and atomic number 86. It is a radioactive, colorless, odorless, tasteless noble gas. It occurs naturally in minute quantities as an intermediate step in the normal radioactive decay chains through which thorium and uranium slowly decay into lead and various other short-lived radioactive elements. Radon itself is the immediate decay product of radium. Its most stable isotope, 222Rn, has a half-life of only 3.8 days, making it one of the rarest elements. Since thorium and uranium are two of the most common radioactive elements on Earth, while also having three isotopes with half-lives on the order of several billion years, radon will be present on Earth long into the future despite its short half-life. The decay of radon produces many other short-lived nuclides, known as radon daughters, ending at stable isotopes of lead.

Alkaline earth metal Group of chemical elements

The alkaline earth metals are six chemical elements in group 2 of the periodic table. They are beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). The elements have very similar properties: they are all shiny, silvery-white, somewhat reactive metals at standard temperature and pressure.

A period 6 element is one of the chemical elements in the sixth row (or period) of the periodic table of the elements, including the lanthanides. The periodic table is laid out in rows to illustrate recurring (periodic) trends in the chemical behaviour of the elements as their atomic number increases: a new row is begun when chemical behaviour begins to repeat, meaning that elements with similar behaviour fall into the same vertical columns. The sixth period contains 32 elements, tied for the most with period 7, beginning with caesium and ending with radon. Lead is currently the last stable element; all subsequent elements are radioactive. For bismuth, however, its only primordial isotope, 209Bi, has a half-life of more than 1019 years, over a billion times longer than the current age of the universe. As a rule, period 6 elements fill their 6s shells first, then their 4f, 5d, and 6p shells, in that order; however, there are exceptions, such as gold.

Radioactive decay Method of decay in atomic nuclei

Radioactive decay is the process by which an unstable atomic nucleus loses energy by radiation. A material containing unstable nuclei is considered radioactive. Three of the most common types of decay are alpha decay, beta decay, and gamma decay, all of which involve emitting one or more particles. The weak force is the mechanism that is responsible for beta decay, while the other two are governed by the usual electromagnetic and strong forces.

Decay chain Series of radioactive decays

In nuclear science, the decay chain refers to a series of radioactive decays of different radioactive decay products as a sequential series of transformations. It is also known as a "radioactive cascade". Most radioisotopes do not decay directly to a stable state, but rather undergo a series of decays until eventually a stable isotope is reached.

A period 7 element is one of the chemical elements in the seventh row of the periodic table of the chemical elements. The periodic table is laid out in rows to illustrate recurring (periodic) trends in the chemical behavior of the elements as their atomic number increases: a new row is begun when chemical behavior begins to repeat, meaning that elements with similar behavior fall into the same vertical columns. The seventh period contains 32 elements, tied for the most with period 6, beginning with francium and ending with oganesson, the heaviest element currently discovered. As a rule, period 7 elements fill their 7s shells first, then their 5f, 6d, and 7p shells in that order, but there are exceptions, such as uranium.

Nuclear chemistry branch of chemistry concerned with radioactivity, transmutation and other nuclear processes

Nuclear chemistry is the sub-field of chemistry dealing with radioactivity, nuclear processes, and transformations in the nuclei of atoms, such as nuclear transmutation and nuclear properties.

Radium chloride Chemical compound

Radium chloride (RaCl2) is a salt of radium and chlorine, and the first radium compound isolated in a pure state. Marie Curie and André-Louis Debierne used it in their original separation of radium from barium. The first preparation of radium metal was by the electrolysis of a solution of this salt using a mercury cathode.

Polonium-210 is an isotope of polonium. It undergoes alpha decay to stable 206Pb with a half-life of 138.376 days, the longest half-life of all naturally occurring polonium isotopes. First identified in 1898, and also marking the discovery of the element polonium, 210Po is generated in the decay chain of uranium-238 and radium-226. 210Po is a prominent contaminant in the environment, mostly affecting seafood and tobacco. Its extreme toxicity is attributed to intense radioactivity, capable of severely harming humans.

Radionuclides which emit gamma radiation are valuable in a range of different industrial, scientific and medical technologies. This article lists some common gamma-emitting radionuclides of technological importance, and their properties.

Radium and radon in the environment

Radium and radon are important contributors to environmental radioactivity. Radon occurs naturally in the environment as a result of decay of radioactive elements in the soil and it can accumulate in houses built on areas where such decay occurs. Radon is among the major causes of cancer; it is estimated to contribute to about 2% of all cancer related deaths in Europe.

Phosphogypsum

Phosphogypsum is the calcium sulfate hydrate formed as a by-product of the production of fertilizer from phosphate rock. It is mainly composed of gypsum (CaSO4·2H2O). Although gypsum is a widely used material in the construction industry, phosphogypsum is usually not used, but is stored indefinitely because of its weak radioactivity caused by the presence of naturally occurring uranium and thorium, and their daughter isotopes radium, radon and polonium. The long-range storage of phosphogypsum is controversial. About five tons of phosphogypsum are generated per ton of phosphoric acid production. Annually, the estimated generation of phosphogypsum worldwide is 100 to 280 Mt.

Naturally occurring radioactive materials (NORM) and technologically enhanced naturally occurring radioactive materials (TENORM) consist of materials, usually industrial wastes or by-products enriched with radioactive elements found in the environment, such as uranium, thorium and plutonium and any of their decay products, such as radium and radon. Produced water discharges and spills are a good example of entering NORMs into the surrounding environment.

Uranium tile

Uranium tiles have been used in the ceramics industry for many centuries, as uranium oxide makes an excellent ceramic glaze, and is reasonably abundant. In addition to its medical usage, radium was used in the 1920s and 1930s for making watch, clock and aircraft dials. Because it takes approximately three metric tons of uranium to extract 1 gram of radium, prodigious quantities of uranium were mined to sustain this new industry. The uranium ore itself was considered a waste product and taking advantage of this newly abundant resource, the tile and pottery industry had a relatively inexpensive and abundant source of glazing material. Vibrant colors of orange, yellow, red, green, blue, black, mauve, etc. were produced, and some 25% of all houses and apartments constructed during that period used bathroom or kitchen tiles that had been glazed with uranium. These can now be detected by a Geiger counter that detects the beta radiation emitted by uranium's decay chain. In most situations, the radiation exposure is not excessive, but there may be exceptions for pure uranium oxide on bathroom floors, which can pose a hazard for infants crawling around.

Radon-222 is the most stable isotope of radon, with a half-life of approximately 3.8 days. It is transient in the decay chain of primordial uranium-238 and is the immediate decay product of radium-226. Radon-222 was first observed in 1899, and was identified as an isotope of a new element several years later. In 1957, the name radon, formerly the name of only radon-222, became the name of the element. Owing to its gaseous nature and high radioactivity, radon-222 is one of the leading causes of lung cancer.

Discovery of nuclear fission 1938 achievement in physics

Nuclear fission was discovered in December 1938 by chemists Otto Hahn and Fritz Strassmann and physicists Lise Meitner and Otto Robert Frisch. Fission is a nuclear reaction or radioactive decay process in which the nucleus of an atom splits into two or more smaller, lighter nuclei and often other particles. The fission process often produces gamma rays and releases a very large amount of energy, even by the energetic standards of radioactive decay. Scientists already knew about alpha decay and beta decay, but fission assumed great importance because the discovery that a nuclear chain reaction was possible led to the development of nuclear power and nuclear weapons.

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Further reading