Potassium sulfate

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Potassium sulfate

General
Systematic name	Potassium sulfate
Other names	Potassium sulphate
Molecular formula	K₂SO₄
Molar mass	174.27 g/mol (anhydrous)
Appearance	White crystalline solid
CAS number	[7778-80-5]
Properties
Density	2.66 g/cm³, anhydrous
Solubility in water	11.1 g/100 ml (20 °C)
In ethanol	insoluble
Melting point	1069°C (1342 K)
Boiling point	1689°C (1962 K)
Structure
Coordination geometry	?
Crystal structure	orthorhombic
Hazards
MSDS	External MSDS
Main hazards	Irritant
R/S statement	None
RTECS number
NFPA 704
Supplementary data page
Structure and properties	n, ε_r, etc.
Thermodynamic data	Phase behaviour Solid, liquid, gas
Spectral data	UV, IR, NMR, MS
Related compounds
Other anions	Potassium hydrogen sulfate Potassium sulfite Potassium bisulfite Potassium persulfate
Other cations	Lithium sulfate Sodium sulfate Magnesium sulfate
Except where noted otherwise, data are given for materials in their standard state (at 25°C, 100 kPa) Infobox disclaimer and references

Potassium sulfate (K₂SO₄) (also 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.

1 History
2 Natural resources
3 Manufacture
4 Properties
5 Uses
6 Potassium hydrogen sulfate
7 See also
8 External links

[edit] 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.

[edit] Natural resources

Natural resources of potassium sulfate are minerals abundant in the Stassfurt salt. These are cocrystalisations of potassium sulfate and sulfates of magnesium calcium and sodium. The minerals are

Kainite MgSO₄• KCl• H₂0
Schönite K₂SO₄ • MgSO₄ • 6 H₂0
Leonite K₂SO₄ • MgSO₄ • 4 H₂0
Langbeinite K₂SO₄ • 2 MgSO₄ •
Glaserite K₃Na(SO₄) ₂
Polyhalite K₂SO₄ • MgSO₄ • 2 CaSO₄ • 2 H₂0

From some of the minerals like kainite, the potassium sulfate can be separated, because the corresponding salt is less soluble in water.

With potassium chloride kieserit MgSO₄ • 2 H₂0 can be transformed and then the potassium sulfate can be dissolved in water.

[edit] Manufacture

Corresponding to the first synthesis form Johann Glauber potassium sulfate can be synthesised by the decomposition of potassium chloride with sulfuric acid, hydrogen chloride evaporates and can be used to produce hydrochloric acid.

The Hargreaves method is basically the same process with different starting materials. Sulfur dioxide, oxygen and water (the starting materials for sulfuric acid) are reacted with potassium chloride. Hydrochloric acid evaporates off.

It is obtained as a by-product in many chemical reactions including the production of nitric acid.

To purify the crude product, it can be dissolved in hot water and then filtered and cooled, causing the bulk of the dissolved salt to crystallize with characteristic promptitude.

[edit] Properties

The very beautiful (anhydrous) crystals form a double six-sided pyramid, but are in fact classified as rhombic. They are transparent, very hard and absolutely resistant to corrosion 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 1078 °C.

[edit] Uses

The principal use of potassium sulfate is as a fertilizer. The crude salt is also used occasionally in the manufacture of glass.

[edit] Potassium hydrogen sulfate

Potassium hydrogen sulfate or bisulfate, KHSO₄, is readily produced by mixing K₂SO₄ with an equivalent no. of moles of sulfuric acid. It forms rhombic pyramids, which melt at 197 °C. It dissolves in three parts of water of 0°C. 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.

Similar is the behavior of the fused dry salt 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.