Strontium (Sr) is a silvery-white alkaline earth metal. While often overshadowed by neighbors like Calcium or Magnesium, its chemical and physical properties give it a surprising range of applications. Strontium plays roles in producing vibrant colors in pyrotechnics and, controversially, in both bone health and nuclear fallout.
Strontium’s Periodic Table Identity
Strontium is element number 38, meaning each atom contains 38 protons. It has a standard atomic mass of 87.62 atomic mass units and exists as a solid metal at room temperature. Strontium is positioned in Group 2 of the periodic table, known as the alkaline earth metals, alongside Beryllium, Magnesium, Calcium, and Barium.
All Group 2 elements have two valence electrons in their outermost shell. This configuration drives Strontium’s chemical behavior. Since Strontium readily gives up these two electrons to achieve stability, it exhibits a consistent oxidation state of +2 in its compounds.
Chemical Behavior and Reactivity
Strontium is a highly reactive metal, a common trait among alkaline earth metals. Pure Strontium is never found naturally and must be stored under kerosene or mineral oil to prevent immediate reaction. Freshly cut metal quickly tarnishes, turning yellowish as it reacts with atmospheric oxygen to form strontium oxide.
Strontium also reacts vigorously with water, producing strontium hydroxide and releasing hydrogen gas. It forms strong ionic bonds with non-metals, including oxygen, sulfur, and the halogens. Volatile Strontium salts burn with a brilliant crimson flame, a property central to its use in pyrotechnics.
Natural Sources and Industrial Applications
Strontium comprises about 0.04 percent of the Earth’s crust, making it the 15th most abundant element. It is primarily mined from two major mineral ores: celestine (\(\text{SrSO}_4\)) and strontianite (\(\text{SrCO}_3\)). Celestine is the most commercially significant source globally.
The most well-known use for Strontium compounds is in pyrotechnics, where Strontium salts produce a vivid red color. Strontium carbonate and strontium nitrate are used in fireworks, flares, and tracer ammunition to generate this intense crimson light. There is currently no substitute that can produce the same brilliant color, ensuring Strontium remains in demand for these applications.
Strontium compounds were historically important in glass manufacturing. Strontium carbonate was incorporated into the faceplate glass of cathode ray tube (CRT) televisions and monitors. The resulting Strontium oxide blocked X-ray emissions generated by the CRT.
While the use in CRT technology has declined, Strontium is also used in metallurgy. It serves as an alloying agent to improve the strength and machinability of aluminum castings, such as engine blocks and wheels.
Biological Role and Health Considerations
Stable Strontium interacts with the body due to its chemical similarity to Calcium. Strontium ions are nearly the same size as Calcium ions, allowing stable Strontium to be absorbed and deposited in bones and teeth. This bone-seeking property is leveraged medically, sometimes using stable Strontium compounds in supplements to support bone density.
The most significant health concern involves the radioactive isotope Strontium-90 (\(\text{Sr-90}\)), a byproduct of nuclear fission. \(\text{Sr-90}\) has a long half-life of about 29 years and behaves chemically identically to stable Strontium. When ingested, \(\text{Sr-90}\) accumulates in the bone marrow where its beta radiation can damage cells, potentially leading to cancers like leukemia.
Naturally occurring Strontium is not radioactive and is of low toxicity. The bone-seeking characteristic that makes \(\text{Sr-90}\) hazardous also makes the stable form useful in medical contexts, such as using \(\text{Sr-89}\) for palliative care for bone cancer. Children are more susceptible to high levels of stable Strontium because their growing bones incorporate it more readily.