Potassium sulfate
Potassium sulfate (US) or potassium sulphate (UK), also called sulphate of potash (SOP), arcanite, or archaically potash of sulfur, is the inorganic compound with formula KSO, a white water-soluble solid. It is commonly used in fertilizers, providing both potassium and sulfur. History. Potassium sulfate (KSO) has been known since early in the 14th century. It was studied by Glauber, Boyle, and Tachenius. In the 17th century, it was named "arcanuni" or "sal duplicatum", as it was a combination of an acid salt with an alkaline salt. It was also known as "vitriolic tartar" and "Glaser's salt  or sal polychrestum Glaseri" after the pharmaceutical chemist Christopher Glaser who prepared it and used medicinally. Known as "arcanum duplicatum" ("double secret") or "panacea duplicata" in pre-modern medicine, it was prepared from the residue ("caput mortuum") left over from the production of aqua fortis (nitric acid, HNO) from nitre (potassium nitrate, KNO) and oil of vitriol (sulphuric acid, HSO) via Glauber's process: The residue was dissolved in hot water, filtered, and evaporated to a cuticle. It was then left to crystallise. It was used as a diuretic and sudorific. According to Chambers's "Cyclopedia", the recipe was purchased  for five hundred thalers by Charles Frederick, Duke of Holstein-Gottorp. Schroder, the duke's physician, wrote wonders of its great uses in hypochondriacal cases, continued and intermitting fevers, stone, scurvy, and more. Natural resources. The mineral form of potassium sulfate, arcanite, is relatively rare. Natural resources of potassium sulfate are minerals abundant in the Stassfurt salt. These are cocrystallizations of potassium sulfate and sulfates of magnesium, calcium, and sodium. Relevant minerals are: The potassium sulfate can be separated from some of these minerals, like kainite, because the corresponding salt is less soluble in water. Kieserite, MgSO·HO, can be combined  with a solution of potassium chloride to produce potassium sulfate. Production. Approximately 1.5 million tons were produced in 1985, typically by the reaction of potassium chloride with sulfuric acid, analogous to the Mannheim process for producing sodium sulfate. The process involves intermediate formation of potassium bisulfate, an exothermic reaction that occurs at room temperature: The second step of the process is endothermic, requiring energy input: Structure and properties. Two crystalline forms are known. Orthorhombic β-KSO is the common form, but it converts to α-KSO above 583 °C. These structures are complex, although the sulfate adopts the typical  tetrahedral geometry. It does not form a hydrate, unlike sodium sulfate. The salt crystallizes as double six-sided pyramids, classified as rhombic. They are transparent, very hard and have a bitter, salty taste. The salt is soluble in water, but insoluble in solutions of potassium hydroxide (sp. gr. 1.35), or in absolute ethanol. Uses. The dominant use of potassium sulfate is as a fertilizer. KSO does not contain chloride, which can be harmful to some crops. Potassium sulfate is preferred for these crops, which include tobacco and some fruits and vegetables. Crops that are less sensitive may still require potassium  sulfate for optimal growth if the soil accumulates chloride from irrigation water. The crude salt is also used occasionally in the manufacture of glass. Potassium sulfate is also used as a flash reducer in artillery propellant charges. It reduces muzzle flash, flareback and blast overpressure. It is sometimes used as an alternative blast media similar to soda in soda blasting as it is harder and similarly water-soluble. Potassium sulfate can also be used in pyrotechnics in combination with potassium nitrate to generate a purple flame. Reactions. Acidification. Potassium hydrogen sulfate (also known as potassium bisulfate), KHSO, is readily produced  by reacting KSO with sulfuric acid. It forms rhombic pyramids, which melt at . It dissolves in three parts of water at . The solution behaves much as if its two congeners, KSO and HSO, were present side by side of each other uncombined; an excess of ethanol the precipitates normal sulfate (with little bisulfate) with excess acid remaining. The behavior of the fused dry salt is similar when heated to several hundred degrees; it acts on silicates, titanates, etc., the same way as sulfuric acid that is heated beyond its natural boiling point does. Hence it is frequently used in analytical  chemistry as a disintegrating agent. For information about other salts that contain sulfate, see sulfate.