The symbol Sr represents Strontium, a chemical element with the atomic number 38, classified on the periodic table as an alkaline earth metal in Group 2 alongside elements like calcium and magnesium. Strontium exists naturally as a stable element important for various industrial and biological processes. However, its radioactive isotope is a major product of nuclear fission, revealing a substance with properties ranging from creating dazzling colors to posing environmental concerns.
Defining Strontium: Basic Identity and Chemical Nature
Strontium occupies the fifth period and Group 2 of the periodic table. A neutral Strontium atom possesses two valence electrons in the 5s orbital, giving it the electron configuration of [Kr]5s\(^2\). This configuration dictates Strontium’s chemical behavior, as it readily sheds these two outer electrons to achieve a stable, noble gas-like configuration.
The loss of two electrons results in the formation of a cation with a +2 charge (Sr\(^{2+}\)), which is its exclusive and stable oxidation state. Because it easily gives up electrons, Strontium is a powerful reducing agent and is highly reactive, never being found in its pure elemental form in nature. Its chemical properties are notably similar to those of calcium and barium, particularly in how its ions substitute for calcium in biological and mineral structures.
Physical Properties and High Reactivity
Elemental Strontium is a soft, silvery-white metal that often exhibits a pale yellow tint, and is both malleable and ductile. In its pure form, it is a good conductor of electricity, and its density is intermediate between that of calcium and barium.
The metal’s high chemical reactivity means its clean surface quickly tarnishes upon exposure to air, forming a dark oxide layer. This rapid oxidation necessitates that elemental Strontium be stored under a liquid hydrocarbon, such as mineral oil, to prevent reaction with the atmosphere. Strontium also reacts readily with water, producing strontium hydroxide and releasing hydrogen gas.
Practical Uses in Industry and Everyday Life
One visually striking application of Strontium compounds is in pyrotechnics, where its salts produce a brilliant crimson-red color. This intense coloration results from Strontium’s electrons absorbing energy from the heat of the flame and then releasing it as specific wavelengths of red light. Strontium nitrate and strontium carbonate are the compounds most commonly used for this effect in fireworks and signal flares.
Historically, Strontium oxide was used extensively in the glass of cathode ray tubes (CRTs) to absorb X-rays and provide radiation shielding. Today, Strontium compounds find modern utility in the production of ferrite magnets, in “glow-in-the-dark” paints using strontium aluminate, and in medicine, such as strontium ranelate to treat osteoporosis.
The Significance of Strontium Isotopes
Naturally occurring Strontium is a mixture of four stable isotopes, with Strontium-88 (Sr-88) being the most abundant. However, the element is widely known due to the synthetic, radioactive isotope Strontium-90 (Sr-90), a product of nuclear fission largely dispersed as fallout from atmospheric nuclear weapons testing.
Sr-90 is a health concern because its chemical similarity to calcium causes the body to incorporate it into bones and teeth, acting as a “bone seeker.” The isotope undergoes beta decay with a half-life of approximately 29 years. The continuous emission of high-energy beta particles poses risks of radiation injury to the bone marrow, potentially increasing the risk of leukemia and bone cancer.