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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 since become known, and less dangerous isotopes are used instead in radioluminescent devices.
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.
When radium decays, ionizing radiation is a product, which can excite fluorescent chemicals and cause radioluminescence.
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).
Pure radium is a volatile silvery-white metal, although its lighter congeners calcium, strontium, and barium have a slight yellow tint.
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 Ra F) is a colorless, luminous compound.
Radium is a chemical element with 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.