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 K2SO4, a white water-Solubility solid. It is commonly used in , providing both potassium and sulfur.
History
Potassium sulfate (K
2SO
4) has been known since early in the 14th century. It was studied by Glauber,
Robert Boyle, and
Otto Tachenius. In the 17th century, it was named
arcanuni or
sal duplicatum, as it was a combination of an
acid salt with an
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, HNO3) from nitre (potassium nitrate, KNO3) and oil of vitriol (sulphuric acid, H2SO4) 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 by Charles Frederick, Duke of Holstein-Gottorp. Schroder, the duke's physician, wrote wonders of its great uses in cases, continued and intermitting , 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, MgSO4·H2O, 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-K
2SO
4 is the common form, but it converts to α-K
2SO
4 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 (Specific gravity 1.35), or in absolute ethanol.
Uses
The dominant use of potassium sulfate is as a fertilizer. K2SO4 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 Colored fire.
A 5% solution of potassium sulfate was used in the beginning of the 20th century as a topical mosquito repellent.
Reactions
Acidification
Potassium hydrogen sulfate (also known as potassium bisulfate), KHSO4, is readily produced by reacting K2SO4 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, K2SO4 and H2SO4, 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 , , 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.
External links
