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Nitric acid
""Resonance description of the bonding in the nitric acid molecule
""Ball-and-stick model of nitric acid
""Resonance space-filling model of nitric acid
"IUPAC name
Nitric acid
Other names
Aqua fortis, Spirit of niter, "Eau forte, Hydrogen nitrate, Acidum nitricum
3D model ("JSmol)
"3DMet B00068
"ECHA InfoCard 100.028.832
"EC Number 231-714-2
"MeSH Nitric+acid
"PubChem CID
"RTECS number QU5775000
"UN number 2031
"Molar mass 63.01 g·mol−1
Appearance Colorless, yellow or red fuming liquid[1]
"Odor acrid, suffocating[1]
"Density 1.5129 g cm−3
"Melting point −42 °C (−44 °F; 231 K)
Completely miscible
"log P -0.13[2]
"Vapor pressure 48 mmHg (20 °C)[1]
"Acidity (pKa) -1.4[3]
−19.9·10−6 cm3/mol
1.397 (16.5 °C)
2.17 ± 0.02 D
146 J·mol−1·K−1[4]
−207 kJ·mol−1[4]
"Safety data sheet ICSC 0183
PCTL Safety Website
"GHS pictograms ""The flame-over-circle pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS) ""The corrosion pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS) ""The skull-and-crossbones pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)
"GHS signal word DANGER
H272, H290, H300, H314, H332, H335
"NFPA 704
"Flash point Non-flammable
Lethal dose or concentration (LD, LC):
138 ppm (rat, 30 min)[1]
US health exposure limits ("NIOSH):
"PEL (Permissible)
TWA 2 ppm (5 mg/m3)[1]
"REL (Recommended)
TWA 2 ppm (5 mg/m3)
ST 4 ppm (10 mg/m3)[1]
"IDLH (Immediate danger)
25 ppm[1]
Related compounds
Other "anions
"Nitrous acid
Other "cations
"Sodium nitrate
"Potassium nitrate
"Ammonium nitrate
Related compounds
"Dinitrogen pentoxide
Except where otherwise noted, data are given for materials in their "standard state (at 25 °C [77 °F], 100 kPa).
YesY verify ("what is YesYNo ?)
"Infobox references

Nitric acid ("H"NO3), also known as aqua fortis ("Latin for "strong water") and spirit of niter, is a highly "corrosive "mineral acid.

The pure compound is colorless, but older samples tend to acquire a yellow cast due to decomposition into "oxides of nitrogen and water. Most commercially available nitric acid has a concentration of 68% in water. When the solution contains more than 86% HNO3, it is referred to as fuming nitric acid. Depending on the amount of "nitrogen dioxide present, fuming nitric acid is further characterized as "white fuming nitric acid at concentrations above 95%, or "red fuming nitric acid at concentrations above 86%.

Nitric acid is the primary reagent used for "nitration – the addition of a "nitro group, typically to an "organic molecule. While some resulting "nitro compounds are shock- and thermally-sensitive "explosives, a few are stable enough to be used in munitions and demolition, while others are still more stable and used as pigments in inks and dyes. Nitric acid is also commonly used as a "strong oxidizing agent.


Physical and chemical properties[edit]


Nitric acid 70%

Nitric acid of commercial interest usually consists of the maximum boiling azeotrope of nitric acid and water, which is approximately 68% HNO3, (approx. 15 molar). This is considered concentrated or technical grade, while reagent grades are specified at 70% HNO3. The density of concentrated nitric acid is 1.35 g/cm3 (68% conc).["inconsistent] An older density scale is occasionally seen, with concentrated nitric acid specified as 42° "Baumé.[5]

Contamination with nitrogen dioxide[edit]

Fuming nitric acid contaminated with yellow nitrogen dioxide.



Fuming nitric acid[edit]



An inhibited fuming nitric acid (either IWFNA, or "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[edit]


Structure and bonding[edit]

Two major resonance representations of HNO3



Acid-base properties[edit]

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 "pKa 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.[8]

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

HNO3 + 2 H2SO4 ⇌ NO2+ + H3O+ + 2HSO4; "Equillibrium constant: K ~ 22

The nitronium ion, 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:

2 HNO3 ⇌ NO2+ + NO3 + H2O

Reactions with metals[edit]

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 H2:

Mg + 2 HNO3 → Mg(NO3)2 + H2 ("Magnesium nitrate)
Mn + 2 HNO3 → Mn(NO3)2 + H2 ("Manganese nitrate)
Zn + 2 HNO3 → Zn(NO3)2 + H2 ("Zinc nitrate)

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:




Upon reaction with nitric acid, most metals give the corresponding nitrates. Some "metalloids and metals give the oxides; for instance, "Sn, "As, "Sb, and "Ti are oxidized into SnO2, As2O5, Sb2O5, and TiO2 respectively.[9]

Some "precious metals, 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 ("Ag, "Cu, ...) present in some "gold alloys 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 (< 14 "carats) and to rapidly assess the gold purity.

Being a powerful oxidizing agent, nitric acid reacts violently with many non-metallic compounds, and the reactions may be explosive. 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 metals series and certain "alloys. 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 are "thermodynamic in nature, but sometimes its oxidation reactions are rather "kinetically non-favored. The presence of small amounts of "nitrous acid (HNO2) greatly enhance the rate of reaction.[9]

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. Typical passivation concentrations range from 20% to 50% by volume (see ASTM A967-05). Metals that are passivated by concentrated nitric acid are "iron, "cobalt, "chromium, "nickel, and "aluminium.[9]

Reactions with non-metals[edit]

Being a powerful "oxidizing acid, nitric acid reacts violently 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 gases, "silicon, and "halogens other than iodine, usually oxidizes them to their highest "oxidation states as acids with the formation of nitrogen dioxide for concentrated acid and "nitric oxide for dilute acid.

C + 4 HNO3 → CO2 + 4 NO2 + 2 H2O



Concentrated nitric acid oxidizes I2, P4, and S8 into "HIO3, "H3PO4, and "H2SO4, respectively.[9]

Xanthoproteic test[edit]

Nitric acid reacts with "proteins 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 acids with "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.[10][11] "Xanthoproteic acid is formed when the acid contacts "epithelial cells. Respective local skin color changes are indicative of inadequate safety precautions when handling nitric acid.


Nitric acid is made by reaction of nitrogen dioxide (NO2) with "water.

3 NO2 + H2O → 2 HNO3 + NO

Normally, the "nitric oxide produced by the reaction is reoxidized by the oxygen in air to produce additional nitrogen dioxide.

Bubbling nitrogen dioxide through "hydrogen peroxide can help to improve acid yield.

2 NO2 + H2O2 → 2 HNO3

Commercial grade nitric acid solutions are usually between 52% and 68% nitric acid. Production of nitric acid is via the "Ostwald process, named after German chemist "Wilhelm Ostwald. In this process, anhydrous "ammonia is oxidized to "nitric oxide, in the presence of "platinum or "rhodium gauze "catalyst at a high temperature of about 500 K and a pressure of 9 bar.


"Nitric oxide is then reacted with oxygen in air to form nitrogen dioxide.






The aqueous HNO3 obtained can be concentrated by "distillation up to about 68% by mass. Further concentration to 98% can be achieved by "dehydration with concentrated "H2SO4. By using ammonia derived from the "Haber process, the final product can be produced from nitrogen, hydrogen, and oxygen which are derived from air and natural gas as the sole feedstocks.[12]

Prior to the introduction of the "Haber process for the production of ammonia in 1913, nitric acid was produced using the "Birkeland–Eyde process, also known as the arc process. This process is based upon the oxidation of atmospheric nitrogen by atmospheric oxygen to "nitric oxide at very high temperatures. An electric arc was used to provide the high temperatures, and yields of up to 4% nitric oxide were obtained. The nitric oxide was cooled and oxidized by the remaining atmospheric oxygen to nitrogen dioxide, and this was subsequently absorbed in dilute nitric acid. The process was very energy intensive and was rapidly displaced by the Ostwald process once cheap ammonia became available.

Laboratory synthesis[edit]

In laboratory, nitric acid can be made by thermal decomposition of "copper(II) nitrate, producing nitrogen dioxide and oxygen gases, which are then passed through water to give nitric acid.


An alternate route is by reaction of approximately equal masses of any nitrate salt such as "sodium nitrate with 96% "sulfuric acid (H2SO4), and "distilling this mixture at nitric acid's boiling point of 83 °C. A nonvolatile residue of the metal hydrogen sulfate remains in the distillation vessel. The red fuming nitric acid obtained may be converted to the white nitric acid.[7]

"NaNO3 + H2SO4 → HNO3 + "NaHSO4

The dissolved NOx is readily removed using reduced pressure at room temperature (10–30 min at 200 "mmHg or 27 "kPa) to give white fuming nitric acid. This procedure can also be performed under reduced pressure and temperature in one step in order to produce less nitrogen dioxide gas.["citation needed]



Nitric acid in a laboratory.

The main industrial use of nitric acid is for the production of "fertilizers. Nitric acid is neutralized with ammonia to give "ammonium nitrate. This application consumes 75–80% of the 26M tons produced annually (1987). The other main applications are for the production of explosives, nylon precursors, and specialty organic compounds.[15]

Precursor to organic nitrogen compounds[edit]

In "organic synthesis, industrial and otherwise, the nitro group is a versatile "functional group. Most derivatives of "aniline are prepared via nitration of "aromatic compounds followed by reduction. Nitrations entail combining nitric and sulfuric acids to generate the "nitronium ion, which electrophilically reacts with aromatic compounds such as "benzene. Many explosives, such as "TNT, are prepared this way:

"C6H5CH3 + 3 HNO3 → "C6H2(NO2)3CH3 + "3 H2O

Either concentrated sulfuric acid or oleum absorbs the excess water.

"H2S2O7 + "H2O → 2 "H2SO4

Use as an oxidant[edit]

The precursor to "nylon, "adipic acid, is produced on a large scale by oxidation of "cyclohexanone and "cyclohexanol with nitric acid.[15]

Rocket propellant[edit]

Nitric acid has been used in various forms as the "oxidizer in "liquid-fueled rockets. These forms include red fuming nitric acid, white fuming nitric acid, mixtures with sulfuric acid, and these forms with HF inhibitor.[16] IRFNA (inhibited "red fuming nitric acid) was one of 3 liquid fuel components for the "BOMARC missile.[17]

Niche uses[edit]

Analytical reagent[edit]


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 HNO
mixed with Type 1 DI Water.

In "electrochemistry, nitric acid is used as a chemical doping agent for organic "semiconductors, and in purification processes for raw "carbon nanotubes.


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).[19]

Etchant and cleaning agent[edit]

The corrosive effects of nitric acid are exploited for a number of specialty applications, such as "etching in printmaking, "pickling stainless steel or cleaning silicon wafers in electronics.[20]

A solution of nitric acid, water and alcohol, "Nital, is used for etching of metals to reveal the microstructure. ISO 14104 is one of the standards detailing this well known procedure.

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 "ferrous alloys against corrosion by the dilute nitric acid.["citation needed]

Nitric acid can be used as a spot test for "alkaloids like "LSD, giving a variety of colours depending on the alkaloid.[21]


Second-degree burn caused by nitric acid

Nitric acid is a "corrosive "acid and a powerful "oxidizing agent. The major hazard posed by it is "chemical burns, as it carries out "acid hydrolysis with "proteins ("amide) and fats ("ester), which consequently decomposes "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.[22] Systemic effects are unlikely, however, and the substance is not considered a carcinogen or mutagen.[23]


Being a strong oxidizing agent, nitric acid can react with compounds such as "cyanides, "carbides, or metallic powders "explosively and with many "organic compounds, such as "turpentine, violently and "hypergolically (i.e. self-igniting). Hence, it should be stored away from bases and organics.


The first mention of nitric acid is in "Pseudo-Geber's De Inventione Veritatis, wherein it is obtained by "calcining a mixture of "niter, "alum and blue "vitriol. It was again described by "Albert the Great in the 13th century and by "Ramon Lull, who prepared it by heating niter and clay and called it "eau forte" (aqua fortis).[24]

"Glauber devised a process to obtain it by distilling potassium nitrate with sulfuric acid. In 1776 "Lavoisier showed that it contained "oxygen, and in 1785 "Henry Cavendish determined its precise composition and showed that it could be synthesized by passing a stream of "electric sparks through moist "air.[24]


  1. ^ a b c d e f g "NIOSH Pocket Guide to Chemical Hazards #0447". "National Institute for Occupational Safety and Health (NIOSH). 
  2. ^ "nitric acid_msds". 
  4. ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22. "ISBN "0-618-94690-X. 
  5. ^ Dean, John (1992). Lange's Handbook of Chemistry (14 ed.). "McGraw-Hill. pp. 2.79–2.80. "ISBN "0-07-016194-1. 
  6. ^ Luzzati, V. (1951). "Structure cristalline de l'acide nitrique anhydre". Acta Crystallographica (in French). 4 (2): 120–131. "doi:10.1107/S0365110X51000404. 
  9. ^ a b c d Catherine E. Housecroft; Alan G. Sharpe (2008). "Chapter 15: The group 15 elements". Inorganic Chemistry, 3rd Edition. Pearson. "ISBN "978-0-13-175553-6. 
  10. ^ Sherman, Henry Clapp (2007). Methods of organic analysis. Read Books. p. 315. "ISBN "1-4086-2802-3. 
  11. ^ Knowles, Frank (2007). A practical course in agricultural chemistry. Read Books. p. 76. "ISBN "1-4067-4583-9. 
  12. ^ Considine, Douglas M., ed. (1974). Chemical and process technology encyclopedia. New York: McGraw-Hill. pp. 769–72. "ISBN "978-0-07-012423-3. 
  13. ^ Urbanski, Tadeusz (1965). Chemistry and technology of explosives. Oxford: Pergamon Press. pp. 85–86. "ISBN "978-0-08-010239-9. 
  15. ^ a b Thiemann, Michael; Scheibler, Erich; Wiegand, Karl Wilhelm (2000). "Nitric Acid, Nitrous Acid, and Nitrogen Oxides". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. "doi:10.1002/14356007.a17_293. "ISBN "3527306730. 
  16. ^ Clark, John D. Ignition!. Rutgers University Press. "ISBN "0-8135-0725-1. 
  18. ^ Clesceri, Lenore S.; Greenberg, Arnold E.; Eaton, Andrew D., eds. (1998). Standard methods for the examination of water and wastewater (20th ed.). American Public Health Association, American Water Works Association, Water Environment Federation. "ISBN "978-0-87553-235-6. 
  19. ^ Jewitt, Jeff (1997). Hand-applied finishes. Taunton Press. "ISBN "978-1-56158-154-2. Retrieved 2009-05-28. 
  21. ^ O’Neal, C. L.; Crouch, D. J.; Fatah, A. A. (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. "PMID 10725655. 
  23. ^ "Nitric acid: Toxicological overview". "Health Protection Agency. Retrieved 2011-12-07. 
  24. ^ a b Encyclopædia Britannica 1911 edition, Nitric Acid

External links[edit]
