|Jmol-3D images||Image 1|
|Molar mass||174.259 g/mol|
|Melting point||1,069 °C (1,956 °F; 1,342 K)|
|Boiling point||1,689 °C (3,072 °F; 1,962 K)|
|Solubility in water||111 g/L (20 °C)
120 g/L (25 °C)
240 g/L (100 °C)
|Solubility||slightly soluble in glycerol
insoluble in acetone, alcohol, CS2
|Refractive index (nD)||1.495|
|EU Index||Not listed|
|LD50||6600 mg/kg (oral, rat)|
|Other anions||Potassium selenate
|Other cations||Lithium sulfate
|Related compounds||Potassium hydrogen sulfate
|Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)|
|(what is: / ?)|
Potassium sulfate (K2SO4) (in British English potassium sulphate, also called sulphate of potash, arcanite, or archaically known as potash of sulfur) is a non-flammable white crystalline salt which is soluble in water. The chemical is commonly used in fertilizers, providing both potassium and sulfur.
Potassium sulfate (K2SO4) has been known since early in the 14th century, and 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.
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.
The minerals are:
- Kainite, MgSO4·KCl·H2O
- Schönite, K2SO4·MgSO4·6H2O
- Leonite, K2SO4·MgSO4·4H2O
- Langbeinite, K2Mg2(SO4)3
- Glaserite, K3Na(SO4)2
- Polyhalite, K2SO4·MgSO4·2CaSO4·2H2O
From some of the minerals like kainite, the potassium sulfate can be separated, because the corresponding salt is less soluble in water.
The process for manufacturing potassium sulfate is similar to that used for the manufacture of sodium sulfate.
- 2 KCl + H2SO4 → 2 HCl + K2SO4
The Hargreaves process uses sulfur dioxide, oxygen and water and potassium chloride as the starting materials to produce potassium sulfate. Hydrochloric acid evaporates off. SO2 is produced through the burning of sulfur.
The anhydrous crystals form a double six-sided pyramid, but are in fact 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. It melts at 1,067 °C (1,953 °F).
The principal 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.
Potassium hydrogen sulfate
Potassium hydrogen sulfate or bisulfate, KHSO4, is readily produced by mixing K2SO4 with an equivalent number of moles of sulfuric acid. It forms rhombic pyramids, which melt at 197 °C (387 °F). It dissolves in three parts of water at 0 °C (32 °F). 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 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.
- Arcanum duplicatum, the term for this substance in alchemy
- Patnaik, Pradyot (2002). Handbook of Inorganic Chemicals. McGraw-Hill. ISBN 0-07-049439-8.
- De Milt, Clara (1942). "Christopher Glaser". Journal of Chemical Education 19 (2): 53. doi:10.1021/ed019p53.
- Klooster, van (1959). "Three centuries of Rochelle salt". Journal of Chemical Education 36 (7): 314. doi:10.1021/ed036p314.
- Windholtz, M (Ed.) & Budavari, S (Ed.), 1983. The Merck Index, Rahway: Merck & Co.
- organization, United Nations industrial development, UNIDO, International Fertilizer Development Center, IFDC (1998). Fertilizer manual (3rd ed. ed.). Dordrecht: Kluwer academic publ. p. 615. ISBN 0-7923-5032-4.
|Salts and the ester of the Sulfate ion|