Potassium azide

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Potassium azide
K+.svg
Azid-Ion.svg
KN3viewCropped.tif
Names
IUPAC name
Potassium azide
Identifiers
3D model (Jmol)
ECHA InfoCard 100.039.997
Properties
KN
3
Molar mass 81.1184 g/mol
Appearance Colorless crystals[1]
Density 2.038 g/cm3
[1]
Melting point 350 °C (662 °F; 623 K) (in vacuum)[1]
Boiling point decomposes
41.4 g/100 mL (0 °C)
50.8 g/100 mL (20 °C)
105.7 g/100 mL (100 °C)
Solubility soluble in ethanol
insoluble in ether
Thermochemistry
-1.7 kJ/mol
Hazards
Main hazards Very Toxic, explosive if strongly heated
NFPA 704
Flammability code 3: Liquids and solids that can be ignited under almost all ambient temperature conditions. Flash point between 23 and 38 °C (73 and 100 °F). E.g., gasoline) Health code 4: Very short exposure could cause death or major residual injury. E.g., VX gas Reactivity code 3: Capable of detonation or explosive decomposition but requires a strong initiating source, must be heated under confinement before initiation, reacts explosively with water, or will detonate if severely shocked. E.g., fluorine Special hazards (white): no codeNFPA 704 four-colored diamond
Lethal dose or concentration (LD, LC):
27 mg/kg (oral, rat)[2]
Related compounds
Other cations
Sodium azide, copper(II) azide, lead(II) azide, silver azide
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 azide is the inorganic compound having the formula KN
3
. It is a white, water-soluble salt. It is used as a reagent in the laboratory.

It has been found to act as a nitrification inhibitor in soil.[3]

Structure[edit]

KN3, RbN3, CsN3, and TlN3 adopt the same structures. They crystallize in a tetragonal habit.[4] The azide is bound to eight cations in an eclipsed orientation. The cations are bound to eight terminal N centers.[5]

Coordination sphere of azide in K,Rb,Cs,TlN3.

Synthesis and reactions[edit]

KN3 is prepared by treating potassium carbonate with hydrazoic acid, which is generated in situ.[6] In contrast, the analogous sodium azide is prepared (industrially) by the "Wislicenus process," which proceeds via the reaction sodium amide with nitrous oxide.[7]

Upon heating or upon irradiation with ultraviolet light, it decomposes into potassium metal and nitrogen gas.[8] The decomposition temperatures of the alkali metal azides are: NaN3 (275 °C), KN3 (355 °C), RbN3 (395 °C), CsN3 (390 °C).[9]

Health hazards[edit]

Like sodium azide, potassium azide is very toxic. The TLV of the related sodium azide is 0.07 ppm. The toxicity of azides arise from their ability to inhibit cytochrome c oxidase.[7]

References[edit]

  1. ^ a b c Dale L. Perry; Sidney L. Phillips (1995). Handbook of inorganic compounds. CRC Press. p. 301. ISBN 0-8493-8671-3. 
  2. ^ http://chem.sis.nlm.nih.gov/chemidplus/rn/20762-60-1
  3. ^ T. D. Hughes; L. F. Welch (1970). "Potassium Azide as a Nitrification Inhibitor". Agronomy Journal. American Society of Agronomy. 62: 595–599. doi:10.2134/agronj1970.00021962006200050013x. 
  4. ^ Khilji, M. Y.; Sherman, W. F.; Wilkinson, G. R. (1982). "Variable temperature and pressure Raman spectra of potassium azide KN
    3
    ". Journal of Raman Spectroscopy. 12 (3): 300–303. Bibcode:1982JRSp...12..300K. doi:10.1002/jrs.1250120319.
     
  5. ^ Ulrich Müller "Verfeinerung der Kristallstrukturen von KN3, RbN3, CsN3 und TIN3" Zeitschrift für anorganische und allgemeine Chemie 1972, Volume 392, 159–166. doi:10.1002/zaac.19723920207
  6. ^ P. W. Schenk "Alkali Azides from Carbonates" in Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 475.
  7. ^ a b Horst H. Jobelius, Hans-Dieter Scharff "Hydrazoic Acid and Azides" in Ullmann's Encyclopedia of Industrial Chemistry, 2005, Wiley-VCH, Weinheim. doi:10.1002/14356007.a13_193
  8. ^ Tompkins, F. C.; Young, D. A. (1982). "The Photochemical and Thermal Formation of Colour Centres in Potassium Azide Crystals". Proceedings of the Royal Society of London. Series A, Mathematical and Physical Sciences. 236 (1204): 10–23. 
  9. ^ E. Dönges "Alkali Metals" in Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 475.