Radium is a highly radioactive alkaline earth metal. Fresh metal is silvery but quickly blackens in air; its intense radioactivity causes self-luminescence. It decays to radon gas and was historically used in luminous paints.
Radium was discovered in 1898 by Marie Curie and Pierre Curie while investigating the radioactivity of pitchblende (uraninite). Their work isolated highly radioactive fractions that contained the new element, which they named radium due to its intense radiation.
Pure radium metal does not intrinsically glow; the observed self-luminescence historically came from radium salts mixed with phosphors (e.g., zinc sulfide). Alpha, beta, and gamma radiation excite the phosphor, which then emits visible light. Moist air can also cause faint bluish luminescence from excitation of surrounding gases.
The most notable isotope is \(^{226}\mathrm{Ra}\) with a half-life of about 1600 years. It decays by alpha emission to the radioactive noble gas radon:
\(\mathrm{^{226}Ra \;\rightarrow\; ^{222}Rn + \alpha}\)
Other isotopes include \(^{223}\mathrm{Ra}\) (half-life ~11.4 days), used in targeted cancer therapy.
Radium is highly radiotoxic. Ingested or inhaled radium behaves chemically like calcium and can be incorporated into bones, where alpha radiation damages marrow and surrounding tissue. Decay to radon gas adds an inhalation hazard; prolonged exposure increases cancer risk.
As an alkaline earth metal (Group 2), radium is almost exclusively +2 in compounds. It forms ionic salts (RaCl2, RaBr2), a sparingly soluble sulfate (RaSO4), and a hydroxide that is strongly basic (Ra(OH)2).
Radium metal reacts with water to yield radium hydroxide and hydrogen gas, analogous to barium:
\(\mathrm{Ra(s) + 2\,H_2O(l) \;\rightarrow\; Ra(OH)_2(aq) + H_2(g)}\)
In air, fresh radium quickly blackens as surface oxides/nitrides form (e.g., RaO, Ra3N2), and accumulated decay products can also discolor the surface.
Historic uses included luminous paints on instrument dials and watch faces (radium + phosphor), and early radiotherapy sources. Modern practice has phased out radium for safety reasons; safer alternatives (tritium, photoluminescent pigments) are used for lighting, while medical use favors sealed gamma sources or radionuclides like \(^{223}\mathrm{Ra}\) in strictly controlled therapies.
Radium occurs at trace levels in uranium ores (e.g., pitchblende). Industrially it was isolated by fractional crystallization/precipitation of radium salts (such as radium bromide or radium chloride) from barium-rich fractions. Today, production is extremely limited and tightly regulated.
Freshly prepared radium is a silvery, dense, soft metal (denser than barium). It blackens in air, is electrically conductive, and forms ionic salts. Because samples are minuscule and radioactive, many values are inferred by analogy with barium and by indirect measurement.
Work with radium demands licensed facilities, remote handling tools, shielding, and continuous monitoring. Strict contamination control prevents ingestion/inhalation. Storage uses sealed, labeled containers, often with lead shielding, and engineered ventilation to manage any radon.
Yes—typical examples include water reaction and sulfate precipitation, plus alpha decay:
\(\mathrm{Ra + 2\,H_2O \rightarrow Ra(OH)_2 + H_2}\)
\(\mathrm{Ra^{2+} + SO_4^{2-} \rightarrow RaSO_4\downarrow}\)
\(\mathrm{^{226}Ra \rightarrow ^{222}Rn + \alpha}\)