Hydrazoic acid, also known as hydrogen azide, azic acid or azoimide,[2] is a compound with the chemical formula HN3.[3] It is a colorless, volatile, and explosive liquid at room temperature and pressure. It is a compound of nitrogen and hydrogen, and is therefore a pnictogen hydride. It was first isolated in 1890 by Theodor Curtius.[4] The acid has few applications, but its conjugate base, the azide ion, is useful in specialized processes.
Hydrazoic acid, like its fellow mineral acids, is soluble in water. Undiluted hydrazoic acid is dangerously explosive[5] with a standard enthalpy of formation ΔfHo (l, 298K) = +264 kJ/mol.[6] When dilute, the gas and aqueous solutions (<10%) can be safely prepared but should be used immediately; because of its low boiling point, hydrazoic acid is enriched upon evaporation and condensation such that dilute solutions incapable of explosion can form droplets in the headspace of the container or reactor that are capable of explosion.[7][8]
Production
The acid is usually formed by acidification of an azide salt like sodium azide. Normally solutions of sodium azide in water contain trace quantities of hydrazoic acid in equilibrium with the azide salt, but introduction of a stronger acid can convert the primary species in solution to hydrazoic acid. The pure acid may be subsequently obtained by fractional distillation as an extremely explosive colorless liquid with an unpleasant smell.[2]
This reaction is unusual in that it involves compounds with nitrogen in four different oxidation states.[11]
Reactions
In its properties hydrazoic acid shows some analogy to the halogen acids, since it forms poorly soluble (in water) lead, silver and mercury(I) salts. The metallic salts all crystallize in the anhydrous form and decompose on heating, leaving a residue of the pure metal.[2] It is a weak acid (pKa = 4.75.[6]) Its heavy metal salts are explosive and readily interact with the alkyl iodides. Azides of heavier alkali metals (excluding lithium) or alkaline earth metals are not explosive, but decompose in a more controlled way upon heating, releasing spectroscopically-pure N2 gas.[12] Solutions of hydrazoic acid dissolve many metals (e.g. zinc, iron) with liberation of hydrogen and formation of salts, which are called azides (formerly also called azoimides or hydrazoates).
Hydrazoic acid may react with carbonyl derivatives, including aldehydes, ketones, and carboxylic acids, to give an amine or amide, with expulsion of nitrogen. This is called Schmidt reaction or Schmidt rearrangement.
Dissolution in the strongest acids produces explosive salts containing the aminodiazonium ion [H2N=N=N]+ ⇌ [H2N−N≡N]+, for example:[12]
The decomposition of hydrazoic acid, triggered by shock, friction, spark, etc. produces nitrogen and hydrogen:
2 HN3 → H2 + 3 N2
Hydrazoic acid undergoes unimolecular decomposition at sufficient energy:
HN3 → NH + N2
The lowest energy pathway produces NH in the triplet state, making it a spin-forbidden reaction. This is one of the few reactions whose rate has been determined for specific amounts of vibrational energy in the ground electronic state, by laser photodissociation studies.[13] In addition, these unimolecular rates have been analyzed theoretically, and the experimental and calculated rates are in reasonable agreement.[14]
Toxicity
Hydrazoic acid is volatile and highly toxic. It has a pungent smell and its vapor can cause violent headaches. The compound acts as a non-cumulative poison.
Applications
2-Furonitrile, a pharmaceutical intermediate and potential artificial sweetening agent has been prepared in good yield by treating furfural with a mixture of hydrazoic acid (HN3) and perchloric acid (HClO4) in the presence of magnesium perchlorate in the benzene solution at 35 °C.[15][16]
^Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN0-07-049439-8
^ abcChisholm, Hugh, ed. (1911). "Azoimide" . Encyclopædia Britannica. Vol. 3 (11th ed.). Cambridge University Press. pp. 82–83. This also contains a detailed description of the contemporaneous production process.
^Dictionary of Inorganic and Organometallic Compounds. Chapman & Hall.
^Furman, David; Dubnikova, Faina; van Duin, Adri C. T.; Zeiri, Yehuda; Kosloff, Ronnie (2016-03-10). "Reactive Force Field for Liquid Hydrazoic Acid with Applications to Detonation Chemistry". The Journal of Physical Chemistry C. 120 (9): 4744–4752. Bibcode:2016APS..MARH20013F. doi:10.1021/acs.jpcc.5b10812. ISSN1932-7447. S2CID102029987.
^ abCatherine E. Housecroft; Alan G. Sharpe (2008). "Chapter 15: The group 15 elements". Inorganic Chemistry, 3rd Edition. Pearson. p. 449. ISBN978-0-13-175553-6.
^Gonzalez-Bobes, F. et al Org. Process Res. Dev. 2012, 16, 2051-2057.
^Treitler, D. S. et al Org. Process Res. Dev. 2017, 21, 460-467.
^L . F. Audrieth, C. F. Gibbs Hydrogen Azide in Aqueous and Ethereal Solution" Inorganic Syntheses 1939, vol. 1, pp. 71-79.
^ abEgon Wiberg; Nils Wiberg; Arnold Frederick Holleman (2001). "The Nitrogen Group". Inorganic chemistry. Academic Press. p. 625. ISBN978-0-12-352651-9.
^Foy, B.R.; Casassa, M.P.; Stephenson, J.C.; King, D.S. (1990). "Overtone-excited HN 3 (X1A') - Anharmonic resonance, homogeneous linewidths, and dissociation rates". Journal of Chemical Physics. 92: 2782–2789. doi:10.1063/1.457924.
^Besora, M.; Harvey, J.N. (2008). "Understanding the rate of spin-forbidden thermolysis of HN 3 and CH 3N 3". Journal of Chemical Physics. 129 (4): 044303. doi:10.1063/1.2953697. PMID18681642.
^P. A. Pavlov; Kul'nevich, V. G. (1986). "Synthesis of 5-substituted furannitriles and their reaction with hydrazine". Khimiya Geterotsiklicheskikh Soedinenii. 2: 181–186.
^B. Bandgar; Makone, S. (2006). "Organic reactions in water. Transformation of aldehydes to nitriles using NBS under mild conditions". Synthetic Communications. 36 (10): 1347–1352. doi:10.1080/00397910500522009. S2CID98593006.
