Potassium oxide

Potassium oxide
Names
IUPAC name
potassium oxide
Systematic IUPAC name
potassium oxidopotassium
Other names
dipotassium monoxide, potash
Identifiers
12136-45-7 YesY
3D model (Jmol) Interactive image
ChemSpider 23354117 N
ECHA InfoCard 100.032.012
EC Number 235-227-6
MeSH Potassium+oxide
UNII 58D606078H N
Properties
K2O
Molar mass 94.20 g·mol−1
Appearance Pale yellow solid
Odor Odorless
Density 2.32 g/cm3 (20 °C)[1]
2.13 g/cm3 (24 °C)[2]
Melting point 740 °C (1,360 °F; 1,010 K) [2]
decomposes from 300 °C[1]
Reacts[1] forming KOH
Solubility Soluble in EtOH, ether[2]
Structure
Antifluorite cubic, cF12[3]
Fm3m, No. 225[3]
a = 6.436 Å[3]
α = 90°, β = 90°, γ = 90°
Tetrahedral (K+)
Cubic (O2−)
Thermochemistry
83.62 J/mol·K[4]
94.03 J/mol·K[4]
−363.17 kJ/mol[1][4]
−322.1 kJ/mol[1]
Hazards
Main hazards Corrosive, reacts violently with water
Safety data sheet ICSC 0769
Related compounds
Other anions
 Potassium sulfide
Other cations
 Lithium oxide
Sodium oxide
Rubidium oxide
Caesium oxide
Potassium peroxide
Potassium superoxide
Related compounds
 Potassium hydroxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

Potassium oxide (K2O) is an ionic compound of potassium and oxygen. This pale yellow solid, the simplest oxide of potassium, is a rarely encountered, highly reactive compound. Some materials of commerce, such as fertilizers and cements, are assayed assuming the percent composition that would be equivalent to the chemical compoud mixture K2O.

Production

Potassium oxide is produced from the reaction of oxygen and potassium; this reaction affords potassium oxide, K2O. Treatment of the peroxide with potassium produces the oxide:[5]

K2O2 + 2 K → 2 K2O

Alternatively and more conveniently, K2O is synthesized by heating potassium nitrate with metallic potassium:

2 KNO3 + 10 K → 6 K2O + N2

Other possibility is to heat potassium peroxide at 500 °C which decomposes at that temperature giving pure potassium oxide and oxygen.

2 K2O2 2 K2O + O2

Potassium hydroxide cannot be further dehydrated to the oxide but it can react with molten potassium to produce it, releasing hydrogen as a byproduct.

Properties and reactions

K2O crystallises in the antifluorite structure. In this motif the positions of the anions and cations are reversed relative to their positions in CaF2, with potassium ions coordinated to 4 oxide ions and oxide ions coordinated to 8 potassium.[6][7] K2O is a basic oxide and reacts with water violently to produce the caustic potassium hydroxide. It is deliquescent and will absorb water from the atmosphere, initiating this vigorous reaction.

Term use in industry

The chemical formula K2O (or simply 'K') is used in several industrial contexts: the N-P-K numbers for fertilizers, in cement formulas, and in glassmaking formulas. Potassium oxide is often not used directly in these products, but the amount of potassium is reported in terms of the K2O equivalent for whichever type of potash was used, such as potassium carbonate. For example, potassium oxide is about 83% potassium by weight, while potassium chloride is only 52%. Potassium chloride provides less potassium than an equal amount of potassium oxide. Thus, if a fertilizer is 30% potassium chloride by weight, its standard potassium rating, based on potassium oxide, would be only 18.8%.

References

  1. 1 2 3 4 5 Anatolievich, Kiper Ruslan. "potassium oxide". http://chemister.ru. Retrieved 2014-07-04. External link in |website= (help)
  2. 1 2 3 Lide, David R., ed. (2009). CRC Handbook of Chemistry and Physics (90th ed.). Boca Raton, Florida: CRC Press. ISBN 978-1-4200-9084-0.
  3. 1 2 3 Wyckoff, Ralph W.G. (1935). The Structure of Crystals. American Chemical Society (2nd ed.). Reinhold Publishing Corp. p. 25.
  4. 1 2 3 Dipotassium oxide in Linstrom, P.J.; Mallard, W.G. (eds.) NIST Chemistry WebBook, NIST Standard Reference Database Number 69. National Institute of Standards and Technology, Gaithersburg MD. http://webbook.nist.gov (retrieved 2014-07-04)
  5. ↑ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  6. ↑ Zintl, E.; Harder, A.; Dauth B. (1934). "Gitterstruktur der oxyde, sulfide, selenide und telluride des lithiums, natriums und kaliums". Zeitschrift für Elektrochemie und Angewandte Physikalische Chemie. 40: 588–93.
  7. ↑ Wells, A.F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.
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