Properties and Overview of Radium

Overview:

Radium (Ra) is a chemical element with the atomic number 88 and the symbol Ra. It is an alkaline earth metal in Group 2 of the periodic table. Marie and Pierre Curie discovered radium in 1898 while they were investigating the radioactive properties of uranium ore. The element's name comes from the Latin word "radius," meaning "ray," due to its intense radioactive glow. Radium is a rare element that is primarily found in trace amounts in uranium and thorium ores, such as pitchblende and carnotite. As a member of the alkaline earth metals, radium shares some chemical properties with other group members like barium and calcium but is distinguished by its high radioactivity. Physically, radium is a silvery-white metal that tarnishes in air, turning black due to the formation of radium nitride on its surface. Radium is relatively dense, about 5.5g/cm³. It has a melting point of approximately 700°C and a boiling point of about 1,737°C. Radium is notable for its intense radioactivity, which causes it to glow faintly in the dark with a blue or bluish-green hue. This luminescence results from radium's alpha particle emissions exciting the surrounding air molecules. The element is highly reactive, especially with water, forming radium hydroxide (Ra(OH)₂) and releasing hydrogen gas. Under standard conditions, it also reacts with other nonmetals, such as oxygen and nitrogen.
Chemically, radium is the heaviest member of the alkaline earth metals, and its properties are similar to those of barium. In chemical reactions, radium typically exhibits a +2 oxidation state, forming compounds like radium chloride (RaCl₂) and radium sulfate (RaSO₄). Radium's highly radioactive compounds emit alpha particles, helium nuclei composed of two protons and two neutrons. This radioactivity is a defining feature of radium's chemistry. The element is also known to form various coordination compounds with different ligands. Due to its similarity to calcium, radium can replace calcium in specific biological processes, which has implications for its biological effects and toxicity.
Safety considerations for radium are significant due to its extreme radioactivity and the harmful effects of its decay products. Radium decays into radon-222, a radioactive gas that poses severe health risks if inhaled, as it can cause lung cancer. The alpha radiation emitted by radium itself is not deeply penetrating and can be stopped by a sheet of paper or the outer layer of human skin. However, if radium is ingested or inhaled, it can accumulate in bones, where its alpha radiation can destroy bone marrow and lead to serious health issues, including anemia, leukemia, and other cancers. During the early 20th century, radium was used in luminous paints for watches, instrument dials, and other applications. This led to significant health problems for workers who handled the substance without proper safety measures. Today, the handling of radium is strictly regulated, requiring specialized facilities, protective equipment, and rigorous safety protocols to prevent exposure.

Production:

Radium is produced through the refinement of uranium ore, where it is found in small concentrations. The extraction process involves isolating radium from other decay products of uranium using a series of chemical reactions, including precipitation and solvent extraction. The radium is typically separated as radium chloride or radium bromide, which can then be converted into other radium compounds or metallic radium through reduction processes. Due to its scarcity and the availability of safer, more effective radioactive materials, radium is no longer produced in large quantities. Most existing supplies come from historical stockpiles or the small amounts naturally present in uranium and thorium ores.

Applications:

Historically, radium has had several applications, primarily due to its radioactive properties. In the early 20th century, it was used in self-luminous paints for watch dials, aircraft switches, and instrument panels. The glowing paint, made by mixing radium with phosphorescent materials, allowed these items to be visible in the dark. However, this use was discontinued after the health hazards of radium exposure became evident. Radium was also used in medicine, particularly in radiotherapy for cancer treatment, as its intense radioactivity could target and destroy cancerous cells. However, with the development of safer and more effective radioactive isotopes, such as Cobalt-60 and cesium-137, radium's use in medical treatments has significantly declined. Today, radium has limited applications, mainly in research settings where its radioactive properties are utilized to study the behavior of radioactive materials or to generate neutron sources in certain types of scientific experiments.

Summary:

Radium is a highly radioactive alkaline earth metal with a fascinating history and a significant impact on the fields of medicine, industry, and science. Its intense radioactivity and ability to emit alpha particles made it useful in various applications during the early 20th century, although these uses have largely been replaced by safer alternatives. The health risks associated with radium exposure have led to strict safety regulations and a decline in its use. However, radium's importance in understanding radioactive decay processes and the properties of heavy alkaline earth metals ensures its ongoing relevance in the field of chemistry and radioactivity.

See a comprehensive list of atomic, electrical, mechanical, physical and thermal properties for radium below: