Radium (symbol Ra, atomic number 88) is a chemical element classified as an alkaline earth metal. It is positioned in group 2 of the periodic table and is the heaviest known member of this group. Radium is a silvery-white metal that exhibits a faint bluish glow due to intense radioactivity.
Discovery
Radium was discovered in 1898 by Marie and Pierre Curie while they were processing pitchblende (uranium ore). The Curies isolated radium as a distinct element by precipitating it as radium chloride after extensive chemical separation from other radioactive substances.
Physical and Chemical Properties
- Appearance: Silvery‑white, metallic, tarnishes rapidly in air.
- Density: Approximately 5.5 g cm⁻³ at 25 °C.
- Melting point: 700 °C (1,292 °F).
- Boiling point: 1,737 °C (3,159 °F).
- Radioactivity: All isotopes are radioactive; the most stable isotope, ^226Ra, has a half‑life of 1,600 years and decays by alpha emission to radon‑222.
Radium reacts readily with water, producing radium hydroxide and hydrogen gas, and it forms a range of compounds such as radium chloride (RaCl₂) and radium bromide (RaBr₂).
Occurrence
Radium is not found in its elemental form in nature due to its radioactivity. It occurs in trace amounts within uranium and thorium ores, most notably in pitchblende and monazite. Commercial extraction of radium is no longer practiced because of the health hazards and the availability of alternative radioactive sources.
Isotopes
More than 30 isotopes of radium have been identified. The most significant are:
- ^226Ra (half‑life ≈ 1,600 years) – the principal isotope found in nature.
- ^224Ra (half‑life ≈ 3.66 days) – produced in the decay chain of thorium‑232.
All isotopes emit alpha particles; several also emit gamma radiation, making radium a potent source of ionizing radiation.
Applications
Historically, radium was used in:
- Luminous paints for watch dials, aircraft instrument panels, and clocks.
- Medical therapy (brachytherapy) for certain cancers, exploiting its emitted radiation.
These uses have been largely discontinued or replaced by safer isotopes (e.g., cobalt‑60, cesium‑137) and non‑radioactive phosphors because of radium’s health risks.
Health and Safety
Radium’s radioactivity poses severe health hazards:
- Inhalation or ingestion of radium salts can lead to bone accumulation, increasing the risk of bone sarcoma and leukemia.
- External exposure to its gamma emissions can cause radiation burns and increase cancer risk.
- Historical cases, such as the “Radium Girls,” highlighted occupational exposure risks in factories that painted luminous dials.
Strict regulations now govern the handling, transport, and disposal of radium. Protective measures include sealed containment, shielding, and monitoring of radiation levels.
Environmental Impact
Radium released into the environment can be incorporated into groundwater and soils through the decay of uranium and thorium minerals. It can contribute to background radiation levels, though typical concentrations are low. Remediation of radium‑contaminated sites requires specialized radioactive waste management practices.
Scientific Significance
Radium's discovery contributed to the development of nuclear physics and radiochemistry. The element’s intense radioactivity was pivotal in early studies of radioactive decay series, leading to the identification of radon gas and the elucidation of nuclear transmutation processes.
