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Potassium sulfate k2 so4

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Potassium sulfate (K2SO4) is a non-flammable white crystalline salt commonly used as a fertilizer. It has been known since the 14th century and can be produced by reacting potassium chloride with sulfuric acid. Potassium sulfate exists in two crystalline forms and does not form hydrates like sodium sulfate. Its main use is as a fertilizer because it provides both potassium and sulfur without containing harmful chlorides.

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Potassium Sulfate - K2SO4 Prepare : Daryan Rahman – Media Ahmed – Walid Sndi – Salih Mahdi – Rebaz Arafat
Introduction : 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 compound is commonly used in fertilizers, providing both potassium and sulfur.
History • 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] • 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, by dissolving the residue in hot water, filtering, and evaporating it to a cuticle. It was then left to crystallise. It was used as a diuretic and sudorific.
History • 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, etc.
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: Kainite, MgSO4·KCl·H2O Schönite (now known as picromerite), K2SO4·MgSO4·6H2O Leonite, K2SO4·MgSO4·4H2O Langbeinite, K2Mg2(SO4)3 Glaserite (now known as aphthitalite), K3Na(SO4)2 Polyhalite, K2SO4·MgSO4·2CaSO4·2H2O 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 Leblanc process. Potassium sulfate is produced according to the following reaction, which is conducted in so-called Mannheim furnaces: • 2KCl + 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. SO2 is produced through the burning of sulfur.
Structure and properties Two crystalline forms are known. Orthorhombic β-K2SO4 is the common form, but it converts to α-K2SO4 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 crystallize 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. 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 artillerypropellant 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.[10] • Sometimes, when put over a fire, potassium sulfate can make it purple.
Reaction • 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 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.

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Potassium sulfate k2 so4

  • 1. Potassium Sulfate - K2SO4 Prepare : Daryan Rahman – Media Ahmed – Walid Sndi – Salih Mahdi – Rebaz Arafat
  • 2. Introduction : 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 compound is commonly used in fertilizers, providing both potassium and sulfur.
  • 3. History • 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] • 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, by dissolving the residue in hot water, filtering, and evaporating it to a cuticle. It was then left to crystallise. It was used as a diuretic and sudorific.
  • 4. History • 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, etc.
  • 5. 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: Kainite, MgSO4·KCl·H2O Schönite (now known as picromerite), K2SO4·MgSO4·6H2O Leonite, K2SO4·MgSO4·4H2O Langbeinite, K2Mg2(SO4)3 Glaserite (now known as aphthitalite), K3Na(SO4)2 Polyhalite, K2SO4·MgSO4·2CaSO4·2H2O 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.
  • 6. Production • Approximately 1.5 million tons were produced in 1985, typically by the reaction of potassium chloride with sulfuric acid, analogous to the Leblanc process. Potassium sulfate is produced according to the following reaction, which is conducted in so-called Mannheim furnaces: • 2KCl + 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. SO2 is produced through the burning of sulfur.
  • 7. Structure and properties Two crystalline forms are known. Orthorhombic β-K2SO4 is the common form, but it converts to α-K2SO4 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 crystallize 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.
  • 8. 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 artillerypropellant 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.[10] • Sometimes, when put over a fire, potassium sulfate can make it purple.
  • 9. Reaction • 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 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.

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