Potassium oxide

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Potassium oxide
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Names
IUPAC name
potassium oxide
Systematic IUPAC name
potassium oxidopotassium
Other names
potash, dipotassium oxide
Identifiers
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3D model (JSmol)
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MeSH Potassium+oxide
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UN number 2033
  • InChI=1S/2K.O ☒N
    Key: NOTVAPJNGZMVSD-UHFFFAOYSA-N ☒N
  • O([K])[K]
Properties
K2O
Molar mass 94.196 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]
Reacts[1] forming KOH
Solubility Soluble in diethyl 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
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Safety data sheet (SDS) ICSC 0769
Related compounds
Other anions
 Potassium sulfide
Potassium selenide
Potassium telluride
Potassium polonide
Other cations
 Lithium oxide
Sodium oxide
Rubidium oxide
Caesium oxide
Potassium peroxide
Potassium superoxide
Potassium ozonide
Related compounds
 Potassium hydroxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Template:Chembox Footer/trackingTemplate:Short description

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

Production

Potassium oxide is produced from the reaction of oxygen and potassium; this reaction affords potassium peroxide, K2O2. Treatment of the peroxide with potassium produces the oxide:[5]

Page Module:Chem2/styles.css has no content.K2O2 + 2 K → 2 K2O

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

Page Module:Chem2/styles.css has no content.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.

Page Module:Chem2/styles.css has no content.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.

Page Module:Chem2/styles.css has no content.2 KOH + 2 K ⇌ 2 K2O + H2 ↑

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 whatever 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

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  1. ^ a b c d Page Module:Citation/CS1/styles.css has no content.Anatolievich, Kiper Ruslan. "potassium oxide". chemister.ru. Retrieved 2014-07-04.
  2. ^ a b c Page Module:Citation/CS1/styles.css has no content.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. ^ a b c Page Module:Citation/CS1/styles.css has no content.Wyckoff, Ralph W.G. (1935). The Structure of Crystals (2nd ed.). Reinhold Publishing Corp. p. 25. {{cite book}}: |work= ignored (help)
  4. ^ a b c Template:Nist
  5. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. Template:ISBN.
  6. ^ Page Module:Citation/CS1/styles.css has no content.Zintl, E.; Harder, A.; Dauth B. (1934). "Gitterstruktur der oxide, sulfide, selenide und telluride des lithiums, natriums und kaliums". Zeitschrift für Elektrochemie und Angewandte Physikalische Chemie. 40 (8): 588–93. doi:10.1002/bbpc.19340400811.
  7. ^ Wells, A.F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. Template:ISBN.

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