Nitric acid is aninorganic compound with the formulaHNO3. It is a highlycorrosivemineral acid.[6] The compound is colorless, but samples tend to acquire a yellow cast over time due to decomposition intooxides of nitrogen. Most commercially available nitric acid has a concentration of 68% in water. When the solution contains more than 86%HNO3, it is referred to asfuming nitric acid. Depending on the amount ofnitrogen dioxide present, fuming nitric acid is further characterized asred fuming nitric acid at concentrations above 86%, orwhite fuming nitric acid at concentrations above 95%.
Nitric acid is the primary reagent used fornitration – the addition of anitro group, typically to anorganic molecule. While some resultingnitro compounds are shock- and thermally-sensitiveexplosives, a few are stable enough to be used in munitions and demolition, while others are still more stable and used as synthetic dyes and medicines (e.g.metronidazole). Nitric acid is also commonly used as astrong oxidizing agent.
History
Medieval alchemy
The discovery ofmineral acids such as nitric acid is generally believed to go back to 13th-century Europeanalchemy.[7] The conventional view is that nitric acid was first described inpseudo-Geber'sDe inventione veritatis ("On the Discovery of Truth", afterc. 1300).[8]
The recipe in theṢundūq al-ḥikma attributed to Jabir has been translated as follows:[10][11]
Take five parts ofpure flowers of nitre, three parts ofCyprus vitriol and two parts of Yemenalum. Powder them well, separately, until they are like dust and then place them in a flask. Plug the latter with a palm fibre and attach a glass receiver to it. Then invert the apparatus and heat the upper portion (i.e. the flask containing the mixture) with a gentle fire. There will flow down by reason of the heat an oil like cow's butter.
In the 17th century,Johann Rudolf Glauber devised a process to obtain nitric acid by distilling potassium nitrate with sulfuric acid. In 1776Antoine Lavoisier citedJoseph Priestley's work to point out that it can be converted from nitric oxide (which he calls "nitrous air"), "combined with an approximately equal volume of the purest part of common air, and with a considerable quantity of water."[15][a] In 1785Henry Cavendish determined its precise composition and showed that it could be synthesized by passing a stream ofelectric sparks through moistair.[12] In 1806,Humphry Davy reported the results of extensive distilled water electrolysis experiments concluding that nitric acid was produced at the anode from dissolved atmospheric nitrogen gas. He used a high voltage battery and non-reactive electrodes and vessels such as gold electrode cones that doubled as vessels bridged by damp asbestos.[16]
The industrial production of nitric acid from atmospheric air began in 1905 with theBirkeland–Eyde process, also known as the arc process.[17] This process is based upon the oxidation of atmospheric nitrogen by atmospheric oxygen to nitric oxide with a very high temperature electric arc. Yields of up to approximately 4–5% nitric oxide were obtained at 3000 °C, and less at lower temperatures.[17][18] The nitric oxide was cooled and oxidized by the remaining atmospheric oxygen to nitrogen dioxide, and this was subsequently absorbed in water in a series ofpacked column orplate column absorption towers to produce dilute nitric acid. The first towers bubbled the nitrogen dioxide through water and non-reactive quartz fragments. About 20% of the produced oxides of nitrogen remained unreacted so the final towers contained an alkali solution to neutralize the rest.[19] The process was very energy intensive and was rapidly displaced by the Ostwald process once cheap ammonia became available.
Another early production method was invented by French engineer Albert Nodon around 1913. His method produced nitric acid from electrolysis of calcium nitrate converted by bacteria from nitrogenous matter in peat bogs. An earthenware pot surrounded by limestone was sunk into the peat and staked with tarred lumber to make a compartment for the carbon anode around which the nitric acid is formed. Nitric acid was pumped out from an earthenware[20] pipe that was sunk down to the bottom of the pot. Fresh water was pumped into the top through another earthenware pipe to replace the fluid removed. The interior was filled withcoke. Cast iron cathodes were sunk into the peat surrounding it. Resistance was about 3 ohms per cubic meter and the power supplied was around 10 volts. Production from one deposit was 800 tons per year.[20][21]
Once theHaber process for the efficient production of ammonia was introduced in 1913, nitric acid production from ammonia using theOstwald process overtook production from the Birkeland–Eyde process. This method of production is still in use today.
Physical and chemical properties
Commercially available nitric acid is anazeotrope with water at a concentration of 68%HNO3. This solution has a boiling temperature of 120.5 °C (249 °F) at 1 atm. It is known as "concentrated nitric acid". The azeotrope of nitric acid and water is a colourless liquid at room temperature.
Two solid hydrates are known: the monohydrateHNO3·H2O or oxonium nitrate[H3O]+[NO3]− and the trihydrateHNO3·3H2O.
An older density scale is occasionally seen, with concentrated nitric acid specified as 42 Baumé.[22]
Contamination with nitrogen dioxide
Fuming nitric acid contaminated with yellow nitrogen dioxide
Nitric acid is subject tothermal or light decomposition and for this reason it was often stored in brown glass bottles:
4 HNO3 → 2 H2O + 4 NO2 + O2
This reaction may give rise to some non-negligible variations in the vapor pressure above the liquid because the nitrogen oxides produced dissolve partly or completely in the acid.
The nitrogen dioxide (NO2) and/or dinitrogen tetroxide (N2O4) remains dissolved in the nitric acid coloring it yellow or even red at higher temperatures. While the pure acid tends to give off white fumes when exposed to air, acid with dissolved nitrogen dioxide gives off reddish-brown vapors, leading to the common names "red fuming nitric acid" and "white fuming nitric acid". Nitrogen oxides (NOx) are soluble in nitric acid.
Commercial-grade fuming nitric acid contains 98%HNO3 and has a density of 1.50 g/cm3. This grade is often used in the explosives industry. It is not as volatile nor as corrosive as the anhydrous acid and has the approximate concentration of 21.4 M.
Red fuming nitric acid, or RFNA, contains substantial quantities of dissolved nitrogen dioxide (NO2) leaving the solution with a reddish-brown color. Due to the dissolved nitrogen dioxide, the density of red fuming nitric acid is lower at 1.490 g/cm3.
Aninhibited fuming nitric acid, either white inhibited fuming nitric acid (IWFNA), or red inhibited fuming nitric acid (IRFNA), can be made by the addition of 0.6 to 0.7%hydrogen fluoride (HF). This fluoride is added forcorrosion resistance in metal tanks. The fluoride creates a metal fluoride layer that protects the metal.
Anhydrous nitric acid
White fuming nitric acid, pure nitric acid or WFNA, is very close to anhydrous nitric acid. It is available as 99.9% nitric acid by assay, or about 24 molar. One specification for white fuming nitric acid is that it has a maximum of 2% water and a maximum of 0.5% dissolvedNO2.Anhydrous nitric acid is a colorless, low-viscosity (mobile) liquid with a density of 1.512–3 g/cm3 that solidifies at −42 °C (−44 °F) to form white crystals.[citation needed] Its dynamic viscosity under standard conditions is 0.76 cP.[23] As it decomposes toNO2 and water, it obtains a yellow tint. It boils at 83 °C (181 °F). It is usually stored in a glass shatterproof amber bottle with twice the volume of head space to allow for pressure build up, but even with those precautions the bottle must be vented monthly to release pressure.
Structure and bonding
Two major resonance representations ofHNO3
The two terminal N–O bonds are nearly equivalent and relatively short, at 1.20 and 1.21 Å.[24] This can be explained by theories ofresonance; the two majorcanonical forms show somedouble bond character in these two bonds, causing them to be shorter than N–Osingle bonds. The third N–O bond is elongated because its O atom is bonded to H atom,[25][26] with abond length of 1.41 Å in the gas phase.[24] The molecule is slightly aplanar (theNO2 and NOH planes are tilted away from each other by 2°) and there isrestricted rotation about the N–OH single bond.[6][27]
Reactions
Acid-base properties
Nitric acid is normally considered to be astrong acid at ambient temperatures. There is some disagreement over the value of the acid dissociation constant, though thepKa value is usually reported as less than −1. This means that the nitric acid in diluted solution is fully dissociated except in extremely acidic solutions. The pKa value rises to 1 at a temperature of 250 °C.[28]
Nitric acid can act as a base with respect to an acid such assulfuric acid:
Nitric acid reacts with most metals, but the details depend on the concentration of the acid and the nature of the metal. Dilute nitric acid behaves as a typicalacid in its reaction with most metals.Magnesium,manganese, andzinc liberateH2:
Nitric acid can oxidize non-active metals such ascopper andsilver. With these non-active or less electropositive metals the products depend on temperature and the acid concentration. For example, copper reacts with dilute nitric acid at ambient temperatures with a 3:8 stoichiometry:
3 Cu + 8 HNO3 → 3 Cu(NO3)2 + 2 NO + 4 H2O
Thenitric oxide produced may react with atmosphericoxygen to givenitrogen dioxide. With more concentrated nitric acid, nitrogen dioxide is produced directly in a reaction with 1:4 stoichiometry:
Someprecious metals, such as puregold and platinum-group metals do not react with nitric acid, though pure gold does react withaqua regia, a mixture of concentrated nitric acid andhydrochloric acid. However, some less noble metals (Ag,Cu, ...) present in somegold alloys relatively poor in gold such ascolored gold can be easily oxidized and dissolved by nitric acid, leading to colour changes of the gold-alloy surface. Nitric acid is used as a cheap means injewelry shops to quickly spot low-gold alloys (< 14karats) and to rapidly assess the gold purity.
Being a powerful oxidizing agent, nitric acid reacts with many non-metallic compounds, sometimes explosively. Depending on the acid concentration, temperature and thereducing agent involved, the end products can be variable. Reaction takes place with all metals except thenoble metals series and certainalloys. As a general rule, oxidizing reactions occur primarily with the concentrated acid, favoring the formation of nitrogen dioxide (NO2). However, the powerful oxidizing properties of nitric acid arethermodynamic in nature, but sometimes its oxidation reactions are ratherkinetically non-favored. The presence of small amounts ofnitrous acid (HNO2) greatly increases the rate of reaction.[29]
Althoughchromium (Cr),iron (Fe), andaluminium (Al) readily dissolve in dilute nitric acid, the concentrated acid forms a metal-oxide layer that protects the bulk of the metal from further oxidation. The formation of this protective layer is calledpassivation.[29] Typical passivation concentrations range from 20% to 50% by volume.[30][full citation needed] Metals that are passivated by concentrated nitric acid areiron,cobalt,chromium,nickel, andaluminium.[29]
Reactions with non-metals
Being a powerfuloxidizing acid, nitric acid reacts with many organic materials, and the reactions may be explosive. Thehydroxyl group will typically strip a hydrogen from the organic molecule to form water, and the remaining nitro group takes the hydrogen's place. Nitration of organic compounds with nitric acid is the primary method of synthesis of many common explosives, such asnitroglycerin andtrinitrotoluene (TNT). As very many less stable byproducts are possible, these reactions must be carefully thermally controlled, and the byproducts removed to isolate the desired product.
Reaction with non-metallic elements, with the exceptions of nitrogen, oxygen,noble gases,silicon, andhalogens other than iodine, usually oxidizes them to their highestoxidation states as acids with the formation of nitrogen dioxide for concentrated acid andnitric oxide for dilute acid.
C (graphite) + 4 HNO3 → CO2 + 4 NO2 + 2 H2O
3 C (graphite) + 4 HNO3 → 3 CO2 + 4 NO + 2 H2O
Concentrated nitric acid oxidizesI2,P4, andS8 intoHIO3,H3PO4, andH2SO4, respectively.[29] Although it reacts with graphite and amorphous carbon, it does not react with diamond; it can separate diamond from the graphite that it oxidizes.[31]
Xanthoproteic test
Nitric acid reacts withproteins to form yellow nitrated products. This reaction is known as thexanthoproteic reaction. This test is carried out by adding concentrated nitric acid to the substance being tested, and then heating the mixture. If proteins that containamino acids witharomatic rings are present, the mixture turns yellow. Upon adding a base such asammonia, the color turns orange. These color changes are caused by nitrated aromatic rings in the protein.[32][33]Xanthoproteic acid is formed when the acid contactsepithelial cells. Respective local skin color changes are indicative of inadequate safety precautions when handling nitric acid.
The dioxide then disproportionates inwater to nitric acid and the nitric oxide feedstock:
3 NO2 + H2O → 2 HNO3 + NO
The net reaction is maximal oxidation of ammonia:
NH3 + 2 O2 → HNO3 + H2O
Dissolved nitrogen oxides are either stripped (in the case of white fuming nitric acid) or remain in solution to formred fuming nitric acid.
Commercial grade nitric acid solutions are usually between 52% and 68% nitric acid by mass, themaximum distillable concentration. Furtherdehydration to 98% can be achieved with concentratedH2SO4.[34][36] Historically, higher acid concentrations were also produced by dissolving additional nitrogen dioxide in the acid, but the last plant in theUnited States ceased using that process in 2012.[36]
More recently, electrochemical means have been developed to produce anhydrous acid from concentrated nitric acid feedstock.[37]
Laboratory synthesis
Laboratory-scale nitric acid syntheses abound. Most take inspiration from the industrial techniques.
Distillation at nitric acid's 83 °C boiling point then separates the solid metal-salt residue.[26] The resulting acid solution is the 68.5% azeotrope, and can be further concentrated (as in industry) with eithersulfuric acid ormagnesium nitrate.[36]
Alternatively, thermal decomposition ofcopper(II) nitrate gives nitrogen dioxide and oxygen gases; these are then passed through water orhydrogen peroxide[38] as in the Ostwald process:
2 Cu(NO3)2 → 2 CuO + 4 NO2 + O2
2 NO2 + H2O → HNO2 + HNO3or2 NO2 + H2O2 → 2 HNO3
Uses
Nitric acid in a laboratory
The main industrial use of nitric acid is for the production offertilizers. Nitric acid is neutralized with ammonia to giveammonium nitrate. This application consumes 75–80% of the 26 million tonnes produced annually (1987). The other main applications are for the production of explosives, nylon precursors, and specialty organic compounds.[39]
Nitric acid has been used in various forms as theoxidizer inliquid-fueled rockets. These forms include red fuming nitric acid, white fuming nitric acid, mixtures with sulfuric acid, and these forms with HF inhibitor.[40] IRFNA (inhibitedred fuming nitric acid) was one of three liquid fuel components for theBOMARC missile.[41]
Niche uses
Metal processing
Nitric acid can be used to convert metals to oxidized forms, such as converting copper metal tocupric nitrate. It can also be used in combination withhydrochloric acid asaqua regia to dissolve noble metals such asgold (aschloroauric acid). These salts can be used to purify gold and other metals beyond 99.9% purity by processes ofrecrystallization andselective precipitation. Its ability to dissolve certain metals selectively or be a solvent for many metal salts makes it useful ingold parting processes.
Analytical reagent
Inelemental analysis byICP-MS,ICP-AES, GFAA, and Flame AA, dilute nitric acid (0.5–5.0%) is used as a matrix compound for determining metal traces in solutions.[42] Ultrapure trace metal grade acid is required for such determination, because small amounts of metal ions could affect the result of the analysis.
It is also typically used in the digestion process of turbid water samples, sludge samples, solid samples as well as other types of unique samples which require elemental analysis viaICP-MS,ICP-OES,ICP-AES, GFAA and flameatomic absorption spectroscopy. Typically these digestions use a 50% solution of the purchasedHNO3 mixed with Type 1 DI Water.
In a low concentration (approximately 10%), nitric acid is often used to artificially agepine andmaple. The color produced is a grey-gold very much like very old wax- or oil-finished wood (wood finishing).[43]
Etchant and cleaning agent
The corrosive effects of nitric acid are exploited for some specialty applications, such asetching in printmaking,pickling stainless steel or cleaning silicon wafers in electronics.[44]
A solution of nitric acid, water and alcohol,nital, is used for etching metals to reveal the microstructure. ISO 14104 is one of the standards detailing this well known procedure.[45]
Nitric acid is used either in combination with hydrochloric acid or alone to clean glass cover slips and glass slides for high-end microscopy applications.[46] It is also used to clean glass before silvering when making silver mirrors.[47]
Commercially available aqueous blends of 5–30% nitric acid and 15–40%phosphoric acid are commonly used for cleaning food and dairy equipment primarily to remove precipitated calcium and magnesium compounds (either deposited from the process stream or resulting from the use of hard water during production and cleaning). The phosphoric acid content helps to passivateferrous alloys against corrosion by the dilute nitric acid.[citation needed]
Nitric acid can be used as a spot test foralkaloids likeLSD, giving a variety of colours depending on the alkaloid.[48]
Nuclear fuel reprocessing
Nitric acid plays a key role inPUREX and othernuclear fuel reprocessing methods, where it can dissolve many differentactinides. The resulting nitrates are converted to various complexes that can be reacted and extracted selectively in order to separate the metals from each other.
The standard first-aid treatment for acid spills on the skin is, as for other corrosive agents, irrigation with large quantities of water. Washing is continued for at least 10–15 minutes to cool the tissue surrounding the acid burn and to prevent secondary damage. Contaminated clothing is removed immediately and the underlying skin washed thoroughly.
Being a strong oxidizing agent, nitric acid can react violently with many compounds.
Use in acid attacks
Nitric acid is one of the most common types of acid used inacid attacks.[51]
Notes
^He goes on to point out that "nitrous air" is the reverse, or "nitric acid deprived of air and water."[15]
^Bell, R. P. (1973),The Proton in Chemistry (2nd ed.), Ithaca, NY: Cornell University Press
^abZumdahl, Steven S. (2009).Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22.ISBN978-0-618-94690-7.
^"Safety Data Sheet"(PDF).fishersci.com. Fisher Scientific International. 23 March 2015. p. 2.Archived(PDF) from the original on 10 September 2022. Retrieved4 October 2022.
^Karpenko & Norris 2002, p. 1002. As Karpenko & Norris note, the uncertain dating of the pseudo-Geber corpus (which was probably written by more than one author) renders the date of its description of nitric acid equally uncertain. According toAl-Hassan 2001, p. 62, recipes for the preparation of nitric acid also occur in theLiber Luminis luminum, a Latin treatise usually attributed toMichael Scot (died before 1236) but perhaps translated by him from the Arabic. One of the manuscripts of theLiber Luminis luminum mentions that it was translated by Michael Scot; seeMoureau, Sébastien (2020)."Min al-kīmiyāʾ ad alchimiam. The Transmission of Alchemy from the Arab-Muslim World to the Latin West in the Middle Ages".Micrologus.28:87–141.hdl:2078.1/211340. p. 115 (no. 22). Al-Hassan 2001 mentionsAbu Bakr al-Razi as the work's author, but this is likely a conflation with several other Latin treatises calledLiber Luminis luminum that were sometimes attributed to al-Razi; see Moureau 2020, p. 107 (no. 5), p. 114 (no. 20), pp. 114–115 (no. 21).
^IUPAC SC-Database A comprehensive database of published data on equilibrium constants of metal complexes and ligands
^abcdeCatherine E. Housecroft; Alan G. Sharpe (2008). "Chapter 15: The group 15 elements".Inorganic Chemistry (3rd ed.). Pearson.ISBN978-0-13-175553-6.
^US 6200456, Harrar, Jackson E.; Quong, Roland & Rigdon, Lester P. et al., "Large-scale production of anhydrous nitric acid and nitric acid solutions of dinitrogen pentoxide", published April 13, 1987, issued March 13, 2001, assigned to United States Department of Energy
^Eaton, Andrew D.; Greenberg, Arnold E.; Rice, Eugene W.; Clesceri, Lenore S.; Franson, Mary Ann H., eds. (2005).Standard Methods For the Examination of Water and Wastewater (21 ed.). American Public Health Association.ISBN978-0-87553-047-5. Also available on CD-ROM andonline by subscription.[page needed]
^Fischer, A. H.; Jacobson, K. A.; Rose, J.; Zeller, R. (1 May 2008). "Preparation of Slides and Coverslips for Microscopy".Cold Spring Harbor Protocols.2008 (6): pdb.prot4988.doi:10.1101/pdb.prot4988.PMID21356831.
^O’Neal, Carol L; Crouch, Dennis J; Fatah, Alim A (April 2000). "Validation of twelve chemical spot tests for the detection of drugs of abuse".Forensic Science International.109 (3):189–201.doi:10.1016/S0379-0738(99)00235-2.PMID10725655.
^May, Paul (November 2007)."Nitric acid". Retrieved2009-05-28.
