Editing Nitric acid

{{short description|Highly corrosive mineral acid}}
{{Distinguish|nitrous acid}}
{{pp-semi-indef|small=yes}}
{{Chembox
| Watchedfields  = changed
| verifiedrevid  = 476996354
| Name           = 
| ImageFile      = White Fuming Nitric Acid.png
| ImageSize      = 150px
| ImageCaption   = Pure nitric acid
| ImageFile1     = Nitric acid resonance.svg
| ImageClass1 = skin-invert-image
| ImageName1     = Resonance description of the bonding in the nitric acid molecule
| ImageFileL2    = Nitric-acid-3D-balls-B.png
| ImageClassL2 = bg-transparent
| ImageNameL2    = Ball-and-stick model of nitric acid
| ImageFileR2    = Nitric-acid-3D-vdW-A.png
| ImageClassR2 = bg-transparent
| ImageNameR2    = Resonance space-filling model of nitric acid
| IUPACName      = Nitric acid
| OtherNames     = {{ubl|''Aqua fortis''|Spirit of niter|[[Etching|''Eau forte'']]|Hydrogen nitrate|''Acidum nitricum''}}
| SystematicName = 
| Section1       = {{Chembox Identifiers
| CASNo = 7697-37-2
| CASNo_Ref = {{cascite|correct|CAS}}
| PubChem = 944
| ChemSpiderID = 919
| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
| UNII = 411VRN1TV4
| UNII_Ref = {{fdacite|correct|FDA}}
| EINECS = 231-714-2
| UNNumber = 2031
| KEGG_Ref = {{keggcite|correct|kegg}}
| KEGG = D02313
| MeSHName = Nitric+acid
| ChEBI_Ref = {{ebicite|correct|EBI}}
| ChEBI = 48107
| RTECS = QU5775000
| Gmelin = 1576
| 3DMet = B00068
| SMILES = [N+](=O)(O)[O-]
| SMILES1 = ON(=O)=O
| ChEMBL_Ref = {{ebicite|correct|EBI}}
| ChEMBL = 1352
| StdInChI = 1S/HNO3/c2-1(3)4/h(H,2,3,4)
| StdInChI_Ref = {{stdinchicite|correct|chemspider}}
| InChI = 1/HNO3/c2-1(3)4/h(H,2,3,4)
| StdInChIKey = GRYLNZFGIOXLOG-UHFFFAOYSA-N
| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
| InChIKey = GRYLNZFGIOXLOG-UHFFFAOYAO}}
| Section2       = {{Chembox Properties
| Formula = {{chem2|HNO3}}
| H=1|N=1|O=3
| Appearance = Colorless liquid<ref name=PGCH/>
| Odor = Acrid, suffocating<ref name=PGCH/>
| Density = 1.51 g/cm<sup>3</sup>, 1.41 g/cm<sup>3</sup> [68% w/w]
| Solubility = Miscible
| MeltingPtC = -42
| BoilingPtC = 83
| BoilingPt_notes = 68% solution boils at {{convert|121|C|F K}}
| ConjugateBase = [[Nitrate]]
| pKa = −1.4<ref>{{Citation |last=Bell |first=R. P. |title=The Proton in Chemistry |edition=2nd |publisher=Cornell University Press |location=Ithaca, NY |year=1973 }}</ref>
| RefractIndex = 1.397 (16.5&nbsp;°C)
| Dipole = 2.17 ± 0.02 D
| VaporPressure = 48 mmHg (20&nbsp;°C)<ref name=PGCH/>
| MagSus = {{val|-1.99e-5|u=cm<sup>3</sup>/mol}}
| LogP = −0.13<ref name="chemsrc">{{Cite web|url=https://www.chemsrc.com/en/cas/7697-37-2_895647.html|title=nitric acid_msds}}</ref>
  }}
| Section3       = 
| Section4       = {{Chembox Thermochemistry
| DeltaHf = −207 kJ/mol<ref name=b1>{{cite book| author = Zumdahl, Steven S.|title =Chemical Principles 6th Ed.| publisher = Houghton Mifflin Company| year = 2009| isbn = 978-0-618-94690-7|page=A22}}</ref>
| Entropy = 146 J/(mol·K)<ref name=b1/>
  }}
| Section5       = 
| Section6       = 
| Section7       = {{Chembox Hazards
| ExternalSDS = [http://www.inchem.org/documents/icsc/icsc/eics0183.htm ICSC 0183]
| GHSPictograms = {{GHS03}}{{GHS05}}{{GHS06}}
| GHSSignalWord = '''DANGER'''
| HPhrases = {{H-phrases|272|290|314|331}}
| PPhrases = {{P-phrases|210|220|280|303+361+353|304+340+310|305+351+338}}
| NFPA-H = 3
| NFPA-R = 2
| NFPA-F = 0
| NFPA-S = OX
| NFPA_ref = <ref>{{cite web |title=Safety Data Sheet |url=https://www.fishersci.com/content/dam/fishersci/en_US/documents/programs/education/regulatory-documents/sds/chemicals/chemicals-n/S25860.pdf |website=fishersci.com |publisher=Fisher Scientific International |access-date=4 October 2022 |archive-url=https://web.archive.org/web/20220910202346/https://www.fishersci.com/content/dam/fishersci/en_US/documents/programs/education/regulatory-documents/sds/chemicals/chemicals-n/S25860.pdf |archive-date=10 September 2022 |page=2 |date=23 March 2015 |url-status=live}}</ref>
| FlashPt = Non-flammable
| PEL = TWA 2 ppm (5 mg/m<sup>3</sup>)<ref name=PGCH>{{PGCH|0447}}</ref>
| IDLH = 25 ppm<ref name=PGCH/>
| LC50 = 138 ppm (rat, 30 min)<ref name=PGCH/>
| REL = TWA 2 ppm (5 mg/m<sup>3</sup>)<br />ST 4 ppm (10 mg/m<sup>3</sup>)<ref name=PGCH/>
  }}
| Section8       = {{Chembox Related
| OtherAnions = [[Nitrous acid]]
| OtherCations = {{ubl|[[Sodium nitrate]]|[[Potassium nitrate]]|[[Ammonium nitrate]]}}
| OtherCompounds = {{ubl|[[Dinitrogen trioxide]]|[[Dinitrogen tetroxide]]|[[Dinitrogen pentoxide]]|[[Nitrogen oxide]]|[[Nitrogen monoxide]]|[[Nitrogen dioxide]]}}
}}
}}

'''Nitric acid''' is an [[inorganic compound]] with the formula {{chem2|HNO3|auto=1}}. It is a highly [[corrosive]] [[mineral acid]].<ref name="G&E">{{Greenwood&Earnshaw2nd|pages=465–471}}</ref> The compound is colorless, but samples tend to acquire a yellow cast over time due to decomposition into [[nitrogen oxide|oxides of nitrogen]]. Most commercially available nitric acid has a concentration of 68% in water. When the solution contains more than 86% {{chem2|HNO3}}, it is referred to as ''fuming nitric acid''. Depending on the amount of [[nitrogen dioxide]] present, fuming nitric acid is further characterized as [[red fuming nitric acid]] at concentrations above 86%, or '''white fuming nitric acid''' at concentrations above 95%.

Nitric acid is the primary reagent used for [[nitration]] – the addition of a [[nitro group]], typically to an [[organic molecule]]. While some resulting [[nitro compound]]s are shock- and thermally-sensitive [[explosive]]s, 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 a [[oxidizing agent|strong oxidizing agent]].

== History ==
=== Medieval alchemy ===
The discovery of [[mineral acids]] such as nitric acid is generally believed to go back to 13th-century European [[alchemy]].<ref>Examples: 
*{{cite book |last=Multhauf |first=Robert P. |title=The Origins of Chemistry |publisher=Oldbourne |year=1966 |location=London |author-link=Robert P. Multhauf |pages=140–141 |quote=But among them we find the rudiments of processes which were finally to lead to the discovery of the mineral acids, sulphuric, hydrochloric and nitric. The mineral acids manifest themselves clearly only about three centuries after [[Abu Bakr al-Razi|al-Razi]], in the works of Europeans [...]}}
*{{Cite book |last1=Needham |first1=Joseph |url=https://books.google.com/books?id=xrNDwP0pS8sC&pg=PA195 |title=Science and Civilisation in China |volume=5: Chemistry and Chemical Technology |at=Part IV: Spagyrical Discovery and Invention: Apparatus, Theories and Gifts, p. 195 |last2=Ping-Yü |first2=Ho |last3=Gwei-Djen |first3=Lu |last4=Sivin |first4=Nathan |date=1980 |publisher=Cambridge University Press |isbn=978-0-521-08573-1 |location=Cambridge |author1-link=Joseph Needham |quote=It is generally accepted that mineral acids were quite unknown both to the ancients in the West and to the Arabic alchemists.}}
*{{Cite book |last=Al-Hassan |first=Ahmad Y. |url=https://books.google.com/books?id=h2g1qte4iegC&pg=PA59 |title=Science and Technology in Islam: Technology and applied sciences |date=2001 |publisher=UNESCO |isbn=978-92-3-103831-0 |author-link=Ahmad Y. al-Hassan |page=59 |quote=The text is given here in full because of the prevailing notion that Islamic chemists did not produce mineral acids.}}
*{{Cite journal |last1=Karpenko |first1=Vladimír |last2=Norris |first2=John A. |year=2002 |title=Vitriol in the History of Chemistry |url=http://www.chemicke-listy.cz/ojs3/index.php/chemicke-listy/article/view/2266 |journal=Chemické listy |volume=96 |issue=12 |pages=997–1005 |quote=[...] dating the discovery of nitric acid is likewise uncertain. It is estimated that this discovery took place after 1300 [...] A passage from the second part of Pseudo-Geber's ''[[Summa perfectionis]]'' [...] was long considered to be the earliest known recipe for sulfuric acid [...]}}
*{{cite book |last=Newman |first=William R. |url=https://books.google.com/books?id=hy5sxK7pHGIC&pg=PA98 |title=Atoms and Alchemy: Chymistry and the Experimental Origins of the Scientific Revolution |publisher=University of Chicago Press |year=2006 |isbn=978-0226576961 |location=Chicago |author-link=William R. Newman |page=98 |quote=[...] between the time when the ''Summa perfectionis'' was written and the seventeenth century, the mineral acids–sulfuric, hydrochloric, nitric, and the mixture of the latter two, called ''aqua regia'', had been discovered.}}</ref> The conventional view is that nitric acid was first described in [[pseudo-Geber]]'s ''De inventione veritatis'' ("On the Discovery of Truth", after {{circa|1300}}).<ref>{{harvnb|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 to {{harvnb|Al-Hassan|2001|p=62}}, recipes for the preparation of nitric acid also occur in the {{lang|la|Liber Luminis luminum}}, a Latin treatise usually attributed to [[Michael Scot]] (died before 1236) but perhaps translated by him from the Arabic. One of the manuscripts of the {{lang|la|Liber Luminis luminum}} mentions that it was translated by Michael Scot; see {{cite journal |last1=Moureau |first1=Sébastien |date=2020 |title=Min al-kīmiyāʾ ad alchimiam. The Transmission of Alchemy from the Arab-Muslim World to the Latin West in the Middle Ages |url=http://hdl.handle.net/2078.1/211340 |journal=Micrologus |volume=28 |issue=22 |pages=87–141 |hdl=2078.1/211340}} Al-Hassan 2001 mentions [[Abu Bakr al-Razi]] as the work's author, but this is likely a conflation with several other Latin treatises called {{lang|la|Liber 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).</ref>

However, according to [[Eric John Holmyard]] and [[Ahmad Y. al-Hassan]], the nitric acid also occurs in various earlier [[Islamicate alchemy|Arabic works]] such as the {{transliteration|ar|Ṣundūq al-ḥikma}} ("Chest of Wisdom") attributed to [[Jabir ibn Hayyan]] (8th century) or the {{transliteration|ar|Taʿwīdh al-Ḥākim}} attributed to the Fatimid caliph [[al-Hakim bi-Amr Allah]] (985–1021).<ref>For the claims regarding the {{transliteration|ar|Ṣundūq al-ḥikma}}, see {{harvnb|Al-Hassan|2001|p=62}}; {{Cite book |last1=Holmyard |first1=John Eric |url=https://archive.org/details/makersofchemistr029725mbp/page/n79/mode/2up |title=Makers of Chemistry |date=1931 |publisher=Clarendon Press |location=Oxford |author1-link=Eric John Holmyard |page=60}} For the claim regarding the {{transliteration|ar|Taʿwīdh al-Ḥākim}}, see {{harvnb|Al-Hassan|2001|p=62}}.</ref>

The recipe in the {{transliteration|ar|Ṣundūq al-ḥikma}} attributed to Jabir has been translated as follows:<ref>{{Cite book |url=https://books.google.com/books?id=440FAAAAMAAJ&q=%22Take+five+parts+of+pure+flowers+of+nitre%22 |title=Discovery: A Monthly Popular Journal of Knowledge |volume=5 |date=1924 |page=215 |editor-first=Hugh |editor-last=Pollard |publisher=John Murray |language=en}}</ref><ref>{{Cite book |last1=Ḥasan |first1=Aḥmad Yūsuf |url=https://books.google.com/books?id=it2E29EkCkUC&q=%22Take+five+parts+of+pure+flowers+of+nitre%22 |title=Islamic Technology: An Illustrated History |last2=Hill |first2=Donald Routledge |date=1986 |publisher=Cambridge University Press |isbn=978-92-3-102294-4 |page=147 |language=en}}</ref>

{{quote|Take five parts of [[sodium nitrate|pure flowers of nitre]], three parts of [[Copper(II) sulfate|Cyprus vitriol]] and two parts of Yemen [[alum]]. 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.}}

Nitric acid is also found in post-1300 works [[Pseudepigrapha|falsely attributed]] to [[Albert the Great]] and [[Ramon Llull]] (both 13th century). These works describe the distillation of a mixture containing niter and [[Iron(II) sulfate|green vitriol]], which they call {{lang|fr|eau forte}} (aqua fortis).<ref name="Britannica19112">{{cite EB1911|wstitle=Nitric Acid|volume=19|pages=711–712}}</ref><ref>{{Cite book |last=Thomson |first=Thomas |url=http://archive.org/details/historyofchemist01unse |title=The history of chemistry |date=1830 |location=London |publisher=H. Colburn, and R. Bentley |volume=1 |page=40 |author-link=Thomas Thomson (chemist)}}</ref><ref>{{cite book |last1=Katz |first1=David A.|title=An Illustrated History of Alchemy and Early Chemistry |date=2008 |page=23 |url=http://www.chymist.com/History%20Alchemy.pdf |access-date=21 October 2023}}</ref>

=== Modern era ===
In the 17th century, [[Johann Rudolf Glauber]] devised a process to obtain nitric acid by distilling potassium nitrate with sulfuric acid. In 1776 [[Antoine Lavoisier]] cited [[Joseph 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."<ref name=":0">{{cite book |last=Gillispie |first=Charles Coulston |url=https://archive.org/details/edgeofobjectivit00char |title=The Edge of Objectivity: An Essay in the History of Scientific Ideas |publisher=Princeton University Press |year=1960 |isbn=0-691-02350-6 |location=Princeton, NJ |pages=223–24 |author-link=Charles Coulston Gillispie}}</ref>{{Efn|He goes on to point out that "nitrous air" is the reverse, or "nitric acid deprived of air and water."<ref name=":0" />}} In 1785 [[Henry Cavendish]] determined its precise composition and showed that it could be synthesized by passing a stream of [[electric arc|electric sparks]] through moist [[air]].<ref name="Britannica19112"/> 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.<ref name="Davy1839">{{cite book |title=The Collected Works of Sir Humphry Davy |year=1839 |editor-last=Davy |editor-first=John |volume=5 |pages=1–12 |chapter=On Some Chemical Agencies of Electricity |chapter-url=https://archive.org/details/collectedworks05davy}}</ref>

The industrial production of nitric acid from atmospheric air began in 1905 with the [[Birkeland–Eyde process]], also known as the arc process.<ref name="Mell1918">{{cite book |author=Mellor, J. W. |url=https://archive.org/details/cu31924055328623 |title=Modern Inorganic Chemistry |publisher=Longmans, Green and Co. |year=1918 |page=[https://archive.org/details/cu31924055328623/page/n532 509]}}</ref> 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&nbsp;°C, and less at lower temperatures.<ref name="Mell1918" /><ref name="Geof1915">{{cite book |author1=Martin, Geoffrey |url=https://archive.org/details/IndustrialNitrogenCompoundsAndExplosives |title=Industrial Nitrogen Compounds and Explosives |author2=Barbour, William |publisher=Crosby Lockwood and Son |year=1915 |page=[https://archive.org/details/IndustrialNitrogenCompoundsAndExplosives/page/n24 21]}}</ref> 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 of [[Packed bed#Packed column|packed column]] or [[plate 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.<ref name="Knox1914">{{cite book |last=Knox |first=Joseph |url=https://archive.org/details/fixationatmosph00knoxgoog |title=The Fixation of Atmospheric Nitrogen |publisher=D. Van Nostrand Company |year=1914 |pages=[https://archive.org/details/fixationatmosph00knoxgoog/page/n56 45]–50}}</ref> 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<ref name="Dary1913">{{cite journal |author=Dary, G. |year=1913 |title=The Production of Nitrates by the Direct Electrolysis of Peat Deposits |url=https://archive.org/details/electricalreview73lond |journal=London Electrical Review |volume=73 |pages=1020–1021}}</ref> 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 with [[Coke (fuel)|coke]]. 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.<ref name="Dary1913" /><ref name="Hale1919">{{cite book |last=Hale |first=Arthur |url=https://archive.org/details/manufacturechem00halegoog |title=The Manufacture of Chemicals by Electrolysis |publisher=D. Van Nostrand Co. |year=1919 |pages=[https://archive.org/details/manufacturechem00halegoog/page/n46 30]–32 |access-date=2019-09-15}}</ref>

Once the [[Haber process]] for the efficient production of ammonia was introduced in 1913, nitric acid production from ammonia using the [[Ostwald 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 an [[azeotrope]] with water at a concentration of 68% {{chem2|HNO3}}. This solution has a boiling temperature of {{cvt|120.5|C|F}} at {{cvt|1|atm|kPa psi}}. 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 monohydrate {{chem2|HNO3*H2O}} or oxonium nitrate {{chem2|[H3O]+[NO3]-}} and the trihydrate {{chem2|HNO3*3H2O}}.

An older density scale is occasionally seen, with concentrated nitric acid specified as 42&nbsp;[[Baumé scale|Baumé]].<ref>{{Cite book| last = Dean| first = John| title = Lange's Handbook of Chemistry| edition = 14| publisher = [[McGraw-Hill]]| year = 1992| pages = [https://archive.org/details/langeshandbookof00lang_0/page/2 2.79–2.80]| isbn = 978-0-07-016194-8| url = https://archive.org/details/langeshandbookof00lang_0/page/2}}</ref>

===Contamination with nitrogen dioxide===
[[File:Fuming nitric acid 40ml.jpg|thumb|left|upright=0.9|Fuming nitric acid contaminated with yellow nitrogen dioxide]]
Nitric acid is subject to [[Heat|thermal]] or light decomposition and for this reason it was often stored in brown glass bottles:

{{block indent|{{chem2|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 ({{chem2|NO2}}) and/or dinitrogen tetroxide ({{chem2|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 ({{chem2|NO_{''x''}|}}) are soluble in nitric acid.

===Fuming nitric acid===
{{main|Red fuming nitric acid}}
Commercial-grade fuming nitric acid contains 98% {{chem2|HNO3}} and has a density of 1.50&nbsp;g/cm<sup>3</sup>. 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&nbsp;M.

[[Red fuming nitric acid]], or RFNA, contains substantial quantities of dissolved nitrogen dioxide ({{chem2|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&nbsp;g/cm<sup>3</sup>.

An ''inhibited'' 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 for [[corrosion 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&nbsp;[[molarity|molar]]. One specification for white fuming nitric acid is that it has a maximum of 2% water and a maximum of 0.5% dissolved {{chem2|NO2}}. [[Anhydrous]] nitric acid is a colorless, low-[[viscosity]] (mobile) liquid with a density of 1.512&ndash;3&nbsp;g/cm<sup>3</sup> that solidifies at {{convert|-42|C}} to form white crystals.<ref>{{cite web |url=https://www.sigmaaldrich.com/US/en/product/mm/100455 |title=Nitric acid fuming 100% |publisher=Sigma-Aldrich |access-date=May 15, 2025}}</ref> Its dynamic viscosity under standard conditions is 0.76&nbsp;mPa·s.<ref>[https://www.wolframalpha.com/input?i=nitric+acid+viscosity "nitric acid viscosity"]. Wolfram Alpha Knowledgebase (2002). Champaign, Illinois.</ref> As it decomposes to {{chem2|NO2}} and water, it obtains a yellow tint. It boils at {{convert|83|C}}. 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===
[[Image:Nitric acid resonance.svg|thumb|upright=1.2|Two major resonance representations of {{chem2|HNO3}}]]
The two terminal N–O bonds are nearly equivalent and relatively short, at 1.20 and 1.21 Å.<ref name="CEAF">{{ cite journal | title = Microwave spectrum of DNO<sub>3</sub>, and average structures of nitric and nitrous acids | first1 = A. P. | last1 = Cox | first2 = M. C. | last2 = Ellis | first3 = C. J. | last3 = Attfield | first4 = A. C. | last4 = Ferris | journal = [[Journal of Molecular Structure]] | year = 1994 | volume = 320 | issue = 1–2 | pages = 91–106 | doi = 10.1016/0022-2860(93)08008-R | bibcode = 1994JMoSt.320...91C }}</ref> This can be explained by theories of [[Resonance (chemistry)|resonance]]; the two major [[Resonance (chemistry)|canonical forms]] show some [[double bond]] character in these two bonds, causing them to be shorter than N–O [[single bond]]s. The third N–O bond is elongated because its O atom is bonded to H atom,<ref>{{cite journal |first=V. |last=Luzzati |title=Structure cristalline de l'acide nitrique anhydre |language=fr |journal=Acta Crystallographica |year=1951 |volume=4 |issue= 2|pages=120–131 |doi=10.1107/S0365110X51000404|bibcode=1951AcCry...4..120L |doi-access=free }}</ref><ref name=Allan>{{cite journal |first1=D. R. |last1=Allan |first2=W. G. |last2=Marshall |first3=D. J. |last3=Francis |first4=I. D. H. |last4=Oswald |first5=C. R. |last5=Pulham |first6=C. |last6=Spanswick |title=The crystal structures of the low-temperature and high-pressure polymorphs of nitric acid |journal=Dalton Transactions |year=2010 |volume=39 |issue=15 |pages=3736–3743 |doi=10.1039/B923975H |pmid=20354626 |url=https://strathprints.strath.ac.uk/26164/1/The_crystal_structures_of_the_low-temperature_and_high-pressure_polymorphs.pdf |type=Submitted manuscript }}</ref> with a [[bond length]] of 1.41 Å in the gas phase.<ref name="CEAF" /> The molecule is slightly aplanar (the [[nitro group|{{chem2|NO2}}]] and NOH planes are tilted away from each other by 2°) and there is [[Conformational isomerism|restricted rotation]] about the N–OH single bond.<ref name="G&E" /><ref>{{cite journal | title = Microwave Spectrum and Structure of Nitric Acid | first1 = A. P. | last1 = Cox | first2 = J. M. | last2 = Riveros | journal = The Journal of Chemical Physics | year = 1965 | volume = 42 | issue = 9 | page = 3106 | doi = 10.1063/1.1696387 | bibcode = 1965JChPh..42.3106C }}</ref>

==Reactions==
===Acid-base properties===
Nitric acid is normally considered to be a [[strong acid]] at ambient temperatures. There is some disagreement over the value of the acid dissociation constant, though the [[acid dissociation constant|p''K''<sub>a</sub>]] 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 p''K''<sub>a</sub> value rises to 1 at a temperature of 250&nbsp;°C.<ref name=scdb>[http://www.acadsoft.co.uk/scdbase/scdbase.htm IUPAC SC-Database] A comprehensive database of published data on equilibrium constants of metal complexes and ligands</ref>

Nitric acid can act as a base with respect to an acid such as [[sulfuric acid]]:

{{block indent|{{chem2|HNO3 + 2 H2SO4 ⇌ [NO2]+ + [H3O]+ + 2 HSO4−}};}}
{{block indent|{{pad|3em}}[[Equilibrium constant]]: ''K'' ≈ 22}}

The [[nitronium ion]], {{chem2|[NO2]+}}, is the active reagent in [[aromatic nitration]] reactions. Since nitric acid has both acidic and basic properties, it can undergo an autoprotolysis reaction, similar to the [[self-ionization of water]]:

{{block indent|{{chem2|2 HNO3 ⇌ [NO2]+ + NO3- + H2O}}}}

===Reactions with metals===
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 typical [[acid]] in its reaction with most metals. [[Magnesium]], [[manganese]], and [[zinc]] liberate [[Hydrogen|{{chem2|H2}}]]:

{{block indent|{{chem2|Mg + 2 HNO3 → [[Magnesium nitrate|Mg(NO3)2]] + H2}}}}
{{block indent|{{chem2|Mn + 2 HNO3 → [[Manganese(II) nitrate|Mn(NO3)2]] + H2}}}}
{{block indent|{{chem2|Zn + 2 HNO3 → [[Zinc nitrate|Zn(NO3)2]] + H2}}}}

Nitric acid can oxidize non-active metals such as [[copper]] and [[silver]]. 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:

{{block indent|{{chem2|3 Cu + 8 HNO3 → 3 Cu(NO3)2 + 2 NO + 4 H2O}}}}

The [[nitric oxide]] produced may react with atmospheric [[oxygen]] to give [[nitrogen dioxide]]. With more concentrated nitric acid, nitrogen dioxide is produced directly in a reaction with 1:4 stoichiometry:

{{block indent|{{chem2|Cu + 4 H+ + 2 NO3- → Cu(2+) + 2 NO2 + 2 H2O}}}}

Upon reaction with nitric acid, most metals give the corresponding [[nitrates]]. Some [[metalloids]] and [[metals]] give the [[oxides]]; for instance, [[Tin|Sn]], [[Arsenic|As]], [[Antimony|Sb]], and [[Titanium|Ti]] are oxidized into [[Tin(IV) oxide|{{chem2|SnO2}}]], [[Arsenic pentoxide|{{chem2|As2O5}}]], [[Antimony pentoxide|{{chem2|Sb2O5}}]], and [[Titanium dioxide|{{chem2|TiO2}}]] respectively.<ref name="InorgChem">{{cite book |title=Inorganic Chemistry |edition=3rd |publisher=Pearson|year=2008|isbn=978-0-13-175553-6|chapter=Chapter 15: The group 15 elements|first1=Catherine E. |last1=Housecroft|first2=Alan G. |last2=Sharpe}}</ref>

Some [[precious metal]]s, such as pure [[gold]] and platinum-group metals do not react with nitric acid, though pure gold does react with ''[[aqua regia]]'', a mixture of concentrated nitric acid and [[hydrochloric acid]]. However, some less noble metals ([[Silver|Ag]], [[Copper|Cu]], ...) present in some [[gold alloy]]s relatively poor in gold such as [[colored 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 in [[jewelry]] shops to quickly spot low-gold alloys (<&nbsp;14 [[karat (purity)|karats]]) 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 the [[reducing agent]] involved, the end products can be variable. Reaction takes place with all metals except the [[noble metal]]s series and certain [[alloy]]s. As a general rule, oxidizing reactions occur primarily with the concentrated acid, favoring the formation of nitrogen dioxide ({{chem2|NO2}}). However, the powerful oxidizing properties of nitric acid are [[thermodynamic]] in nature, but sometimes its oxidation reactions are rather [[chemical kinetics|kinetically]] non-favored. The presence of small amounts of [[nitrous acid]] ({{chem2|HNO2}}) greatly increases the rate of reaction.<ref name="InorgChem" />

Although [[chromium]] (Cr), [[iron]] (Fe), and [[aluminium]] (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 called [[passivation (chemistry)|passivation]].<ref name="InorgChem" /> Typical passivation concentrations range from 20% to 50% by volume.<ref>[[ASTM]] standard A967-05</ref>{{full citation needed|date=October 2022}} Metals that are passivated by concentrated nitric acid are [[iron]], [[cobalt]], [[chromium]], [[nickel]], and [[aluminium]].<ref name="InorgChem" />

===Reactions with non-metals===
Being a powerful [[oxidizing acid]], nitric acid reacts with many organic materials, and the reactions may be explosive. The [[hydroxyl]] 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 as [[nitroglycerin]] and [[trinitrotoluene]] (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 gas]]es, [[silicon]], and [[halogen]]s other than iodine, usually oxidizes them to their highest [[Oxidation number|oxidation state]]s as acids with the formation of nitrogen dioxide for concentrated acid and [[nitric oxide]] for dilute acid.

{{block indent|{{chem2|C (graphite) + 4 HNO3 → CO2 + 4 NO2 + 2 H2O}}}}
{{block indent|{{chem2|3 C (graphite) + 4 HNO3 → 3 CO2 + 4 NO + 2 H2O}}}}

Concentrated nitric acid oxidizes [[Iodine|{{chem2|I2}}]], [[White phosphorus|{{chem2|P4}}]], and [[Octasulfur|{{chem2|S8}}]] into [[Iodic acid|{{chem2|HIO3}}]], [[Phosphoric acid|{{chem2|H3PO4}}]], and [[Sulfuric acid|{{chem2|H2SO4}}]], respectively.<ref name="InorgChem" /> Although it reacts with graphite and amorphous carbon, it does not react with diamond; it can separate diamond from the graphite that it oxidizes.<ref>{{cite journal |last1=Ōsawa |first1=Eiji |title=Recent progress and perspectives in single-digit nanodiamond |journal=Diamond and Related Materials |date=December 2007 |volume=16 |issue=12 |pages=2018–2022 |doi=10.1016/j.diamond.2007.08.008 |bibcode=2007DRM....16.2018O }}</ref>

===Xanthoproteic test===
Nitric acid reacts with [[protein]]s to form yellow nitrated products. This reaction is known as the [[xanthoproteic reaction]]. This test is carried out by adding concentrated nitric acid to the substance being tested, and then heating the mixture. If proteins that contain [[amino acid]]s with [[aromaticity|aromatic]] rings are present, the mixture turns yellow. Upon adding a base such as [[ammonia]], the color turns orange. These color changes are caused by nitrated aromatic rings in the protein.<ref>{{Cite book| title = Methods of organic analysis| first1 = Henry Clapp| last1 = Sherman
| publisher = Read Books| year = 2007| isbn = 978-1-4086-2802-7| page = 315}}</ref><ref>{{Cite book| title = A practical course in agricultural chemistry | first1 = Frank| last1 = Knowles| publisher = Read Books| year = 2007| isbn = 978-1-4067-4583-2| page = 76}}</ref> [[Xanthoproteic acid]] is formed when the acid contacts [[epithelial cell]]s. Respective local skin color changes are indicative of inadequate safety precautions when handling nitric acid.

==Production==
Industrial nitric acid production uses the [[Ostwald process]].  The combined Ostwald and [[Haber process]]es are extremely efficient, requiring only air and natural gas [[feedstock]]s.<ref name=Considine/>

The Ostwald process' technical innovation is the proper conditions under which anhydrous [[ammonia]] burns to [[nitric oxide]] (NO) instead of [[dinitrogen]] ({{chem2|N2}}).<ref name=Considine>{{Cite book |editor-last=Considine |editor-first=Douglas M. |title=Chemical and process technology encyclopedia |year=1974 |publisher=McGraw-Hill |location=New York |isbn=978-0-07-012423-3 |pages=[https://archive.org/details/chemicalprocesst00newy/page/769 769–72] |url=https://archive.org/details/chemicalprocesst00newy/page/769 }}</ref><ref>{{cite web |last1=Foist |first1=Laura |title=The Ostwald Process & Catalytic Oxidation of Ammonia |url=https://study.com/academy/lesson/the-ostwald-process-catalytic-oxidation-of-ammonia.html |website=Study.com |access-date=5 January 2019}}</ref>  The nitric oxide is then oxidized, often with [[atmospheric oxygen]], to [[nitrogen dioxide]] ({{chem2|NO2}}):

{{block indent|{{chem2|2 NO + O2 → 2 NO2}}}}

The dioxide then disproportionates in [[water]] to nitric acid and the nitric oxide feedstock:

{{block indent|{{chem2|3 NO2 + H2O → 2 HNO3 + NO}}}}

The net reaction is maximal oxidation of ammonia:

{{block indent|{{chem2|NH3 + 2 O2 → HNO3 + H2O}}}}

Dissolved nitrogen oxides are either stripped (in the case of white fuming nitric acid) or remain in solution to form [[red fuming nitric acid]].

Commercial grade nitric acid solutions are usually between 52% and 68% nitric acid by mass, the [[azeotrope|maximum distillable concentration]]. Further [[desiccation|dehydration]] to 98% can be achieved with concentrated [[sulfuric acid|{{chem2|H2SO4}}]].<ref name=Considine/><ref name="Wiley Nitric Acid">{{Cite book|publisher=Wiley|title=Kirk-Othmer Encyclopedia of Chemical Technology |chapter=Nitric acid|date=2020 |pages=1–37 |doi=10.1002/0471238961.1409201803120118.a01.pub3 |isbn=9780471484943 |s2cid=260923593 |chapter-url=https://onlinelibrary.wiley.com/doi/10.1002/0471238961.1409201803120118.a01.pub3|access-date=2023-08-09|language=en |last1=Groves |first1=Michael C.E }}</ref> Historically, higher acid concentrations were also produced by dissolving additional nitrogen dioxide in the acid, but the last plant in the [[United States]] ceased using that process in 2012.<ref name="Wiley Nitric Acid"/>

More recently, electrochemical means have been developed to produce anhydrous acid from concentrated nitric acid feedstock.<ref>{{cite patent |inventor-last=Harrar |inventor-first=Jackson E. |inventor2-last=Quong |inventor2-first=Roland |inventor3-last=Rigdon |inventor3-first=Lester P. |inventor4-last=McGuire |inventor4-first=Raymond R. |assign1=United States Department of Energy |title=Large-scale production of anhydrous nitric acid and nitric acid solutions of dinitrogen pentoxide |country-code=US |patent-number=6200456 |publication-date=April 13, 1987 |issue-date=March 13, 2001 }}</ref>

===Laboratory synthesis===
Laboratory-scale nitric acid syntheses abound. Most take inspiration from the industrial techniques.

A wide variety of [[nitrate]] salts metathesize with [[sulfuric acid]] ({{chem2|H2SO4}}) – for example, [[sodium nitrate]]:

{{block indent|{{chem2|NaNO3 + H2SO4 → HNO3 + NaHSO4}}}}

[[Distillation]] at nitric acid's 83&nbsp;°C boiling point then separates the solid metal-salt residue.<ref name=Allan/> The resulting acid solution is the 68.5% azeotrope, and can be further concentrated (as in industry) with either [[sulfuric acid]] or [[magnesium nitrate]].<ref name="Wiley Nitric Acid"/>

Alternatively, thermal decomposition of [[copper(II) nitrate]] gives nitrogen dioxide and oxygen gases; these are then passed through water or [[hydrogen peroxide]]<ref>{{Cite journal |last=Dong |first=Kai |date=April 19, 2024 |title=H2O2-mediated electrosynthesis of nitrate from air |url=https://www.nature.com/articles/s44160-024-00522-8 |journal=[[Nature (journal)|Nature]]}}</ref> as in the Ostwald process:

{{block indent|{{chem2|2 Cu(NO3)2 → 2 CuO + 4 NO2 + O2}}}}
{{block indent|{{chem2|2 NO2 + H2O → HNO2 + HNO3}}{{pad|2em}}or{{pad|2em}}{{chem2|2 NO2 + H2O2 → 2 HNO3}}}}

==Uses==
[[Image:Nitric acid lab.jpg|thumb|upright=1.25|Nitric acid in a laboratory]]
The main industrial use of nitric acid is for the production of [[fertilizer]]s. Nitric acid is neutralized with ammonia to give [[ammonium 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.<ref name=Ullmann>{{Ullmann|last1=Thiemann |first1=Michael |last2=Scheibler |first2=Erich |last3=Wiegand |first3=Karl Wilhelm |date=2000 |title=Nitric Acid, Nitrous Acid, and Nitrogen Oxides |doi=10.1002/14356007.a17_293|isbn=978-3527306732 }}</ref>

===Precursor to organic nitrogen compounds===
{{See also|Nitration}}
In [[organic synthesis]], industrial and otherwise, the nitro group is a versatile [[functional group]]. A mixture of nitric and sulfuric acids introduces a nitro [[substituent]] onto various [[aromatic compound]]s by [[electrophilic aromatic substitution]]. Many explosives, such as [[Trinitrotoluene|TNT]], are prepared this way:

{{block indent|{{chem2|[[Toluene|C6H5CH3]] + 3 HNO3 → [[TNT|C6H2(NO2)3CH3]] + 3 [[Water|H2O]]}}}}

Either concentrated sulfuric acid or oleum absorbs the excess water.

{{block indent|{{chem2|[[Disulfuric acid|H2S2O7]] + [[Water|H2O]] → 2 [[Sulfuric acid|H2SO4]]}}}}

The nitro group can be [[reduction reaction|reduced]] to give an [[amine group]], allowing synthesis of [[aniline]] compounds from various [[nitrobenzene]]s:{{alt text missing|date=April 2025}}

[[File:Aniline from Nitrobenzene.svg|upright=1.8|center]]

===Use as an oxidant===
The precursor to [[nylon]], [[adipic acid]], is produced on a large scale by oxidation of "KA oil"—a mixture of [[cyclohexanone]] and [[cyclohexanol]]—with nitric acid.<ref name=Ullmann/>

===Rocket propellant===
Nitric acid has been used in various forms as the [[oxidizer]] in [[liquid-fueled rocket]]s. These forms include red fuming nitric acid, white fuming nitric acid, mixtures with sulfuric acid, and these forms with HF inhibitor.<ref>{{Cite book|last=Clark|first=John D.|title=Ignition!|isbn=978-0-8135-0725-5|publisher=Rutgers University Press|year=1972}}</ref> IRFNA (inhibited [[red fuming nitric acid]]) was one of three liquid fuel components for the [[BOMARC]] missile.<ref>{{Cite web|title=BOMARC Summary| url=http://www.themilitarystandard.com/missile/bomarc/summary.php| access-date=2025-04-07| publisher=TheMilitaryStandard}}</ref>{{unreliable source|reason=no authorship information|date=April 2025}}

===Niche uses===

====Metal processing====

Nitric acid can be used to convert metals to oxidized forms, such as converting copper metal to [[cupric nitrate]]. It can also be used in combination with [[hydrochloric acid]] as [[aqua regia]] to dissolve noble metals such as [[gold]] (as [[chloroauric acid]]). These salts can be used to purify gold and other metals beyond 99.9% purity by processes of [[Recrystallization (chemistry)|recrystallization]] and [[Precipitation (chemistry)|selective precipitation]]. Its ability to dissolve certain metals selectively or be a solvent for many metal salts makes it useful in [[gold parting]] processes.

====Analytical reagent====
In [[elemental analysis]] by [[ICP-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.<ref>{{cite book |title=Standard Methods For the Examination of Water and Wastewater |edition=21 |editor-last1=Eaton |editor-first1=Andrew D. |editor-last2=Greenberg |editor-first2=Arnold E. |editor-last3=Rice |editor-first3=Eugene W. |editor-last4=Clesceri |editor-first4=Lenore S. |editor-last5=Franson |editor-first5=Mary Ann H. |year=2005 |publisher=American Public Health Association |isbn=978-0-87553-047-5}} Also available on CD-ROM and [http://www.standardmethods.org/ online] by subscription.{{page needed|date=November 2022}}</ref> 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 via [[ICP-MS]], [[ICP-OES]], [[ICP-AES]], GFAA and flame [[atomic absorption spectroscopy]]. Typically these digestions use a 50% solution of the purchased {{Chem2|HNO3}} mixed with Type 1 DI Water.

In [[electrochemistry]], nitric acid is used as a chemical doping agent for [[organic semiconductor]]s, and in purification processes for raw [[carbon nanotube]]s.

====Woodworking====
In a low concentration (approximately 10%), nitric acid is often used to artificially age [[pine]] and [[maple]]. The color produced is a grey-gold very much like very old wax- or oil-finished wood ([[wood finishing]]).<ref>{{Cite book|last=Jewitt|first=Jeff|title=Hand-applied finishes|url=https://archive.org/details/handappliedfinis0000jewi|url-access=registration|access-date=2009-05-28|year=1997|publisher=Taunton Press|isbn=978-1-56158-154-2}}</ref>

====Etchant and cleaning agent====
The corrosive effects of nitric acid are exploited for some specialty applications, such as [[etching]] in printmaking, [[Pickling (metal)|pickling stainless steel]] or cleaning silicon wafers in electronics.<ref>Muraoka, Hisashi (1995). "Silicon wafer cleaning fluid with HNO<sub>3</sub>, HF, HCl, surfactant, and water" {{US Patent|5635463}}</ref>

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.<ref>{{cite ISO standard|csnumber=70396|title=ISO 14104:2017 – Gears – Surface temper etch inspection after grinding, chemical method}}</ref>

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.<ref>{{cite journal |last1=Fischer |first1=A. H. |last2=Jacobson |first2=K. A. |last3=Rose |first3=J. |last4=Zeller |first4=R. |title=Preparation of Slides and Coverslips for Microscopy |journal=Cold Spring Harbor Protocols |date=1 May 2008 |volume=2008 |issue=6 |pages=pdb.prot4988 |doi=10.1101/pdb.prot4988 |pmid=21356831 }}</ref> It is also used to clean glass before silvering when making silver mirrors.<ref>{{cite journal |last1=Curtis |first1=Heber D. |title=Methods of Silvering Mirrors |journal=Publications of the Astronomical Society of the Pacific |date=February 1911 |volume=23 |issue=135 |pages=13 |doi=10.1086/122040 |bibcode=1911PASP...23...13C |s2cid=120665136 |url=https://zenodo.org/record/1431273 |doi-access=free |hdl=2027/mdp.39015018047608 |hdl-access=free }}</ref>

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 passivate [[iron alloys|ferrous alloys]] against corrosion by the dilute nitric acid.{{Citation needed|date=September 2011}}

Nitric acid can be used as a spot test for [[alkaloid]]s like [[LSD]], giving a variety of colours depending on the alkaloid.<ref name="validation paper">{{cite journal |last1=O’Neal |first1=Carol L |last2=Crouch |first2=Dennis J |last3=Fatah |first3=Alim A |title=Validation of twelve chemical spot tests for the detection of drugs of abuse |journal=Forensic Science International |date=April 2000 |volume=109 |issue=3 |pages=189–201 |doi=10.1016/S0379-0738(99)00235-2 |pmid=10725655 }}</ref>

==== Nuclear fuel reprocessing ====
Nitric acid plays a key role in [[PUREX]] and other [[nuclear fuel reprocessing]] methods, where it can dissolve many different [[actinide]]s. The resulting nitrates are converted to various complexes that can be reacted and extracted selectively in order to separate the metals from each other.

==Safety==
Nitric acid is a [[corrosive]] [[acid]] and a powerful [[oxidizing agent]]. The major hazard posed by it is [[chemical burn]]s, as it carries out [[acid hydrolysis]] with [[protein]]s ([[amide]]) and fats ([[ester]]), which consequently decomposes [[tissue (biology)|living tissue]] (e.g. [[skin]] and [[flesh]]). Concentrated nitric acid stains [[human skin]] yellow due to its reaction with the [[keratin]]. These yellow stains turn orange when neutralized.<ref>{{Cite web|title=Nitric acid |url=http://www.chm.bris.ac.uk/motm/nitric/nitrich.htm|first=Paul|last=May|access-date=2009-05-28|date=November 2007}}</ref> Systemic effects are unlikely, and the substance is not considered a carcinogen or mutagen.<ref>{{cite web | title = Nitric acid: Toxicological overview | url = http://www.hpa.org.uk/webc/HPAwebFile/HPAweb_C/1194947355794 | publisher = [[Health Protection Agency]] | access-date = 2011-12-07}}</ref>

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 in [[acid attacks]].<ref>{{cite news |last1=Rees |first1=Anna |title=Freeze mob to highlight the issue of acid attacks |url=https://en.reset.org/blog/freeze-mob-highlight-issue-acid-attacks |access-date=25 June 2021 |work=RESET.to |date=1 October 2013 |language=en}}</ref>

== Notes ==
{{Notelist}}

==References==
{{Reflist|30em}}

==External links==
* [https://www.cdc.gov/niosh/npg/npgd0447.html NIOSH Pocket Guide to Chemical Hazards]
* [https://www.dcceew.gov.au/environment/protection/npi/substances/fact-sheets/nitric-acid National Pollutant Inventory – Nitric Acid Fact Sheet]
*Calculators: [http://www.aim.env.uea.ac.uk/aim/surftens/surftens.php surface tensions] {{Webarchive|url=https://web.archive.org/web/20200222134012/http://www.aim.env.uea.ac.uk/aim/surftens/surftens.php |date=2020-02-22 }}, and [http://www.aim.env.uea.ac.uk/aim/density/density_electrolyte.php densities, molarities and molalities] {{Webarchive|url=https://web.archive.org/web/20200222134007/http://www.aim.env.uea.ac.uk/aim/density/density_electrolyte.php |date=2020-02-22 }} of aqueous nitric acid

{{Hydrogen compounds}}
{{Nitrogen compounds}}
{{Nitrates}}
{{Authority control}}

{{DEFAULTSORT:Nitric acid}}
[[Category:Pnictogen oxoacids]]
[[Category:Nitrogen oxoacids]]
[[Category:Mineral acids]]
[[Category:Photographic chemicals]]
[[Category:Drug testing reagents]]
[[Category:Oxidizing acids]]
[[Category:Nitrogen(V) compounds]]