Radium (Ra, atomic number 88) is a chemical element discovered in 1898 by Marie and Pierre Curie. Its most defining characteristic is its intense radioactivity. Radium is unequivocally a metal.
Classification as an Alkaline Earth Metal
Radium is situated in Group 2 of the periodic table, designating it as an alkaline earth metal alongside elements like magnesium, calcium, and barium. All elements in this group possess two electrons in their outermost energy shell. This configuration dictates Radium’s highly metallic chemical behavior.
Metals readily lose outer-shell electrons, a trait Radium exhibits by forming a cation with a \(+2\) oxidation state. This propensity to donate electrons means Radium is highly reactive and is never found in its pure state in nature.
Defining Physical and Chemical Characteristics
When freshly isolated, pure Radium is a lustrous, silvery-white metal. This bright appearance is fleeting, as the element is highly reactive when exposed to air. Radium quickly tarnishes, developing a black surface layer from reacting with atmospheric nitrogen to form radium nitride.
Radium exists as a solid at standard room temperature, possessing a melting point of approximately \(700^\circ\text{C}\) and a boiling point around \(1140^\circ\text{C}\). It exhibits strong chemical reactivity with water, decomposing the water molecule to form radium hydroxide and release hydrogen gas. Furthermore, its compounds emit a faint, characteristic carmine-red color when introduced into a flame.
The Unique Factor of Intense Radioactivity
While its chemical classification is metallic, Radium’s most defining feature is its intense radioactivity, stemming from the instability of its atomic nucleus. Radium is a byproduct of the radioactive decay chain of uranium and thorium; all thirty-four known isotopes are unstable. The most common naturally occurring isotope is Radium-226, which has a half-life of approximately 1,600 years.
Radioactivity is the spontaneous process where an unstable nucleus transforms, releasing energy and radiation. Radium-226 primarily decays through alpha emission, ejecting an alpha particle and transforming into the noble gas Radon-222. The decay rate is high; one gram of Radium-226 undergoes \(3.7 \times 10^{10}\) nuclear disintegrations every second. This rate was historically used to define the curie (Ci) unit of radioactivity. The energy released as Radium interacts with surrounding matter generates heat and causes its compounds to spontaneously emit light, a phenomenon called radioluminescence.
Historical and Contemporary Uses
Radium’s intensely radioactive properties have led to both historical and specific contemporary applications. Historically, one widespread use was in the production of luminous paint for instrument panels and dials. When mixed with zinc sulfide powder, the alpha particles emitted by the Radium excite the zinc sulfide, causing it to glow in the dark.
In medicine, Radium was recognized for its therapeutic potential in cancer treatment, known as brachytherapy. Early applications involved placing small, sealed containers of Radium salts directly near tumors to destroy malignant cells. Today, while many former uses have been replaced by safer radioisotopes like tritium or cobalt-60, Radium still sees specialized medical use. The isotope Radium-223 is currently used in targeted cancer therapy, particularly for prostate cancer that has metastasized to the bone.