Barium (Ba), atomic number 56, is a soft, silvery-white metal in Group 2 of the periodic table. This classifies it as an alkaline earth metal, grouping it with elements like magnesium and calcium. Elemental Barium is highly reactive, placing it among the most chemically active metals outside of the alkali metal group. Due to this vigorous nature, Barium is never found in its pure metallic form in nature.
The Basis of Barium’s High Reactivity
Barium’s placement in Group 2 is the reason for its intense chemical activity. All elements in this group possess two valence electrons. Barium achieves its most stable state by shedding these two electrons to adopt the noble gas configuration of Xenon, forming a \(\text{Ba}^{2+}\) ion.
The relatively large atomic size of Barium means its valence electrons are far from the nucleus, reducing the electrostatic attraction. This distance results in low ionization energy, meaning little energy is required to remove those two outermost electrons. This characteristic makes Barium eager to react, easily forming ionic compounds as the divalent cation. Reactivity generally increases moving down the alkaline earth metals group, making Barium more reactive than Strontium, Calcium, and Magnesium.
Demonstrating Reactivity: Reactions with Water and Air
Barium’s high reactivity is evident upon exposure to air. When the fresh, silvery metal is exposed, it rapidly oxidizes, quickly forming a dull, grayish layer of Barium Oxide (BaO) on its surface. Barium metal must be stored under mineral oil or in an inert atmosphere to prevent it from reacting with oxygen and nitrogen in the air.
If Barium metal is heated, it burns vigorously, producing a mixture including Barium Oxide, Barium Nitride (\(\text{Ba}_3\text{N}_2\)), and Barium Peroxide (\(\text{BaO}_2\)). The metal also reacts readily with cold water. When Barium is dropped into water, it produces Barium Hydroxide (\(\text{Ba}(\text{OH})_2\)) and releases hydrogen gas (\(\text{H}_2\)), a reaction quicker than that of Strontium.
The process is exothermic, generating heat that increases the reaction’s vigor, though it is not as explosive as reactions with alkali metals like Sodium or Potassium. The resulting Barium Hydroxide is only slightly soluble, often causing the water to become cloudy as the compound precipitates. These reactions demonstrate Barium’s strong tendency to form stable compounds.
Practical Consequences: Toxicity and Safe Handling
The chemical feature that makes Barium reactive—its propensity to form the \(\text{Ba}^{2+}\) ion—also underlies the toxicity of its soluble compounds. While elemental Barium metal is dangerous due to its reactivity with moisture and air, the primary biological threat comes from water-soluble Barium salts, such as Barium Chloride or Barium Nitrate. These salts dissolve easily, releasing the toxic \(\text{Ba}^{2+}\) ion into the body upon ingestion.
The \(\text{Ba}^{2+}\) ion interferes directly with muscle and nerve function by blocking potassium channels in cell membranes. This interference rapidly leads to a drop in blood potassium levels, known as hypokalemia. Symptoms include cardiac arrhythmias, muscle weakness, and in high doses, paralysis.
Given the reactivity and toxicity, the safe handling of Barium metal requires strict protocols. The metal must be handled in an inert environment, often a glove box filled with Argon gas, or kept submerged under mineral oil to prevent contact with atmospheric oxygen and water. Laboratory personnel must use appropriate protective equipment and ensure adequate ventilation to avoid exposure to dust or fumes.
Barium in Stable Compounds: The Exception of Barium Sulfate
Despite the general toxicity of soluble Barium compounds, Barium Sulfate (\(\text{BaSO}_4\)) is an important exception. This compound is widely used in medicine as a radiocontrast agent for X-ray and CT imaging procedures, such as the Barium swallow. The paradox of using a toxic element internally is explained by its extreme chemical stability.
Barium Sulfate is highly insoluble in water and, critically, in the acidic environment of the stomach. Because the compound does not dissolve, the toxic \(\text{Ba}^{2+}\) ions are never released into the digestive system or absorbed into the bloodstream. The entire compound passes through the gastrointestinal tract and is excreted from the body.
The Barium atom’s high atomic number allows Barium Sulfate to effectively block X-rays, creating a stark, clear contrast on the image. This physical property, combined with its chemical insolubility, makes it an invaluable and safe tool for visualizing the soft tissues of the esophagus, stomach, and intestines. This stable compound effectively neutralizes the toxic nature of the Barium ion.