|Radium in the periodic table|
|silvery white metallic
|Name, symbol, number||radium, Ra, 88|
|Element category||alkaline earth metal|
|Group, period, block||2 (alkaline earth metals), 7, s|
|Standard atomic weight||(226)|
|Electron configuration||[Rn] 7s2
2, 8, 18, 32, 18, 8, 2
|Density (near r.t.)||5.5 g·cm−3|
|Melting point||973 K, 700 °C, 1292 °F|
|Boiling point||2010 K, 1737 °C, 3158.6 °F|
|Heat of fusion||8.5 kJ·mol−1|
|Heat of vaporization||113 kJ·mol−1|
|Oxidation states||2 (strongly basic oxide)|
|Electronegativity||0.9 (Pauling scale)|
|Ionization energies||1st: 509.3 kJ·mol−1|
|2nd: 979.0 kJ·mol−1|
|Covalent radius||221±2 pm|
|Van der Waals radius||283 pm|
|Crystal structure||body-centered cubic|
|Electrical resistivity||(20 °C) 1 µΩ·m|
|Thermal conductivity||18.6 W·m−1·K−1|
|CAS registry number||7440-14-4|
|Discovery||Pierre Curie and Marie Curie (1898)|
|First isolation||Marie Curie (1902)|
|Most stable isotopes|
|Main article: Isotopes of radium|
Radium is a chemical element with symbol Ra and atomic number 88. Radium is an almost pure-white alkaline earth metal, but it readily oxidizes on exposure to air, becoming black in color. All isotopes of radium are highly radioactive, with the most stable isotope being radium-226, which has a half-life of 1601 years and decays into radon gas. Because of such instability, radium is luminescent, glowing a faint blue.
Radium, in the form of radium chloride, was discovered by Marie Curie and Pierre Curie in 1898. 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 1910. Since its discovery, it has given names like radium A and radium C2 to several isotopes of other elements that are decay products of radium-226.
In nature, radium is found in uranium 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.
- 1 Characteristics
- 2 Production
- 3 History
- 4 Historical applications
- 5 Precautions
- 6 See also
- 7 References
- 8 Bibliography
- 9 Further reading
- 10 External links
Although radium is not as well studied as its stable lighter homologue barium, the two elements have very similar properties. Their first two ionization energies are very similar: 509.3 and 979.0 kJ·mol−1 for radium and 502.9 and 965.2 kJ·mol−1 for barium. Such low figures yield both elements' high reactivity and the formation of the very stable Ra2+ ion and similar Ba2+.
Pure radium is a white, silvery, solid metal, melting at 700 °C (1292 °F) and boiling at 1737 °C (3159 °F), similar to barium. Radium has a density of 5.5 g/cm3; the radium-barium density ratio is comparable to the radium-barium atomic mass ratio, as these elements have very similar body-centered cubic structures.
Chemical characteristics and compounds
Radium is the heaviest known alkaline earth metal; its chemical properties mostly resemble those of barium. When exposed to air, radium reacts violently with it, forming radium nitride, which causes blackening of this white metal. It exhibits only the +2 oxidation state in solution. Radium ions do not form complexes easily, due to the highly basic character of the ions. Most radium compounds coprecipitate with all barium, most strontium, and most lead compounds, and are ionic salts. The radium ion is colorless, making radium salts white when freshly prepared, turning yellow and ultimately dark with age owing to self-decomposition from the alpha radiation. Compounds of radium flame red-purple and give a characteristic spectrum. Like other alkaline earth metals, radium reacts violently with water to form radium hydroxide and is slightly more volatile than barium. Because of its geologically short half-life and intense radioactivity, radium compounds are quite rare, occurring almost exclusively in uranium ores.
Radium chloride, radium bromide, radium hydroxide, and radium nitrate are soluble in water, with solubilities slightly lower than those of barium analogs for bromide and chloride, and higher for nitrate. Radium hydroxide is more soluble than hydroxides of other alkaline earth metals, actinium, and thorium, and more basic than barium hydroxide. It can be separated from these elements by their precipitation with ammonia. Insoluble radium compounds include radium sulfate, radium chromate, radium iodate, radium carbonate, and radium tetrafluoroberyllate; the radium sulfate is the most insoluble known sulfate. Radium oxide remains uncharacterized, despite the fact that oxides are common compounds for other alkaline-earth metals. The 6s and 6p electrons participate in the bonding in radium fluoride and radium astatide, making the bonding there more covalent in character.
Radium has 25 different known isotopes, four of which are found in nature, with 226Ra being the most common. 223Ra, 224Ra, 226Ra and 228Ra are all generated naturally in the decay of either uranium (U) or thorium (Th). 226Ra is a product of 238U decay, and is the longest-lived isotope of radium with a half-life of 1601 years; next longest is 228Ra, a product of 232Th breakdown, with a half-life of 5.75 years.
Radium has no stable isotopes; however, four isotopes of radium are present in decay chains,all of which are present in trace amounts. The most abundant and the longest-living one is radium-226, with a half-life of 1601 years. To date, 34 isotopes of radium have been synthesized, ranging in mass number from 202 to 234.
At least 12 nuclear isomers have been reported; the most stable of them is radium-205m, with a half-life of between 130 and 230 milliseconds. All ground states of isotopes from radium-205 to radium-214, and from radium-221 to radium-234, have longer ones.
Three other natural radioisotopes had received historical names in the early 20th century: radium-223 was known as actinium X, radium-224 as thorium X and radium-228 as mesothorium I. Radium-226 has given historical names to its decay products after the whole element, such as radium A for polonium-218.
Radium-226 is 2.7 million times more radioactive then the same molar amount of natural uranium (mostly U238), due to its proportionally shorter half-life. Both are components of the (4n+2) uranium/radium decay series, so all the radium-226 in the world today is the product of uranium-238 decay, hence its occurrence only in ores of uranium. Radium's decay occurs in the last nine steps of the fourteen step uranium series; the successive decay products were studied and were called radium emanation or "exradio" (now identified as radon-222), radium A (polonium-218), radium B (lead-214), radium C (bismuth-214), and so on. Radon is a heavy gas, and the later products are solids. These products are themselves radioactive elements until stable lead-206 is reached, each with an weight four atomic weight units lower than its predecessor, if the decay is by alpha particle. If by beta particle, the weight doesn't change, but the element identity changes to the next lower numbered element.
Radium-226 loses about 1% of its activity in 25 years, being transformed into elements of lower atomic weight, with lead-206 being the final product of disintegration, just as uranium-238 decays down to radium-226.
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, radium itself emits only alpha particles, but other steps in the decay chain emit alpha or beta particles, and almost all particle emissions are accompanied by gamma rays.
All radium occurring today is produced by the decay of heavier elements, being present in decay chains. Owing to such short half-lives of its isotopes, radium is not primordial but trace. It cannot occur in large quantities due both to the fact that isotopes of radium have short half-lives and that parent nuclides have very long ones. Radium is found in tiny quantities in the uranium ore uraninite and various other uranium minerals, and in even tinier quantities in thorium minerals.
Radium-226 is a decay product of uranium and is therefore found in all uranium-bearing ores. (One ton of pitchblende typically yields about one seventh of a gram of radium). All other isotopes of radium, produced by the other two active decay chains and by the occasional neutron capture, have much shorter half lives than radium-226, so it is the most common, predominant isotope of the element.
Uranium had no large scale application in the late 18th century and therefore no large uranium mines existed. In the beginning the only larger source for uranium ore was the silver mines at Joachimsthal (now Jáchymov) in the Austrian Empire. The uranium ore was only a by-product of the mining activities. 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 took over the ownership of the mines and stopped exporting raw ore. For some time the radium availability was low.
The formation of an Austrian monopoly and the strong urge of other countries to have access to radium led to a world wide 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. Radium can also be extracted from the waste from nuclear reactors. Large radium-containing uranium deposits are located in Russia, Canada (the Northwest Territories), the United States (New Mexico, Utah and Colorado, for example) and Australia. Neither of the deposits is mined for radium but the uranium content makes mining profitable.
The amounts produced were always relatively small; for example, in 1918 13.6 g of radium were produced in the United States. As of 1954, the total worldwide supply of purified radium amounted to about 5 pounds (2.3 kg).
Radium (Latin radius, ray) was discovered by Marie Skłodowska-Curie and her husband Pierre on December 21, 1898, in a uraninite sample. While studying the mineral, the Curies removed uranium from it and found that the remaining material was still radioactive. They then separated out a radioactive mixture consisting mostly of compounds of barium which gave a brilliant green flame color and crimson carmine spectral lines that had never been documented before. The Curies announced their discovery to the French Academy of Sciences on 26 December 1898. The naming of radium dates to about 1899, from the French word radium, formed in Modern Latin from radius (ray), called for its power of emitting energy in the form of rays.
In 1910, radium was isolated as a pure metal by Curie and André-Louis Debierne through the electrolysis of a pure radium chloride solution using a mercury cathode and distilling in an atmosphere of hydrogen gas. The same year, E. Eoler produced radium by heating its azide, Ra(N3)2. The Curies' new element was first industrially produced in the beginning of the 20th century by Biraco, a subsidiary company of Union Minière du Haut Katanga (UMHK) in its Olen plant in Belgium. UMHK offered to Marie Curie her first gram of radium. It gave historical names to the decay products of radium, such as radium A, B, C, etc., now known to be isotopes of other elements.
On 4 February 1936, radium E (bismuth-210) became the first radioactive element to be made synthetically in the United States. Dr. John Jacob Livingood, at the radiation lab at University of California, Berkeley, was bombarding several elements with 5-MeV deuterons. He noted that irradiated bismuth emits fast electrons with a 5-day half-life, which matched the behavior of radium E.
Some of the few practical uses of radium are derived from its radioactive properties. More recently discovered radioisotopes, such as 60Co and 137Cs, 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.
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. In the mid-1920s, a lawsuit was filed against the United States Radium Corporation by five dying "Radium Girl" dial painters who had painted radium-based luminous paint on the dials of watches and clocks. The dial painters routinely licked their brushes to give them a fine point, thereby ingesting radium. Their exposure to radium caused serious health effects which included sores, anemia, and bone cancer. This is because radium is treated as calcium by the body, and deposited in the bones, where radioactivity degrades marrow and can mutate bone cells.
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. Worse, 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.
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.
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. Where indefinite self-luminosity in darkness was required, safer radioactive promethium paint was initially used, later replaced by tritium which continues to be used today. Tritium emits beta radiation which cannot penetrate the skin, rather than the penetrating gamma radiation of radium and is regarded as safer. It has a half-life of 12 years.
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. 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.
Radium was once an additive in products such as toothpaste, hair creams, and even food items due to its supposed curative powers. 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".) 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.
Radium (usually in the form of radium chloride) 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. The isotope 223Ra (under the trade name Xofigo) was approved by the FDA in 2013 for use in medicine as a cancer treatment of bone metastasis.
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. His first patient was his own aunt in 1904, who died shortly after surgery. Kelly was known to use excessive amounts of radium to treat various cancers and tumors. As a result, some of his patients died from high amounts of radium exposure. His method of radium application was inserting a radium capsule near the affected area then sewing the radium "points" directly to the tumor. This was the same method used to treat Henrietta Lacks, the host of the original HeLa cells, for cervical cancer.
In 1909, the famous Rutherford experiment used radium as an alpha source to probe the atomic structure of gold. This experiment led to the Rutherford model of the atom and revolutionized the field of nuclear physics. When mixed with beryllium, it is a neutron source. This type of neutron source were for a long time the main source for neutrons in research.
Radium is highly radioactive and its decay product, radon gas, is also radioactive. Since radium is chemically similar to calcium, it has the potential to cause great harm by replacing calcium in bones. Exposure to radium can cause cancer and other disorders, because radium and its decay product radon emit alpha particles upon their decay, which kill and mutate cells. At the time of the Manhattan Project in 1944, the "tolerance dose" for workers was set at 0.1 microgram of ingested radium.
Some of the biological effects of radium were apparent from the start. The first case of so-called "radium-dermatitis" was reported in 1900, only 2 years after the element's discovery. The French physicist Antoine Becquerel carried a small ampoule of radium in his waistcoat pocket for 6 hours and reported that his skin became ulcerated. Marie Curie experimented with a tiny sample that she kept in contact with her skin for 10 hours, and noted that an ulcer appeared several days later. Handling of radium has been blamed for Curie's death due to aplastic anemia. Stored radium should be ventilated to prevent accumulation of radon. Emitted energy from the decay of radium also ionizes gases, fogs photographic emulsions, and produces many other detrimental effects.
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|Periodic table (Large version)|