Nitric acid

Nitric acid (HNO3), 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.

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
Names
IUPAC name
Nitric acid
Other names
Aqua fortis, Spirit of niter, Eau forte, Hydrogen nitrate, Acidum nitricum
Identifiers
3D model (JSmol)
3DMet B00068
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.028.832
EC Number 231-714-2
1576
KEGG
MeSH Nitric+acid
RTECS number QU5775000
UNII
UN number 2031
Properties
HNO3
Molar mass 63.012 g·mol−1
Appearance Colorless, yellow or red fuming liquid[1]
Odor acrid, suffocating[1]
Density 1.51 g cm−3, 1.41 g cm−3 [68% w/w]
Melting point −42 °C (−44 °F; 231 K)
Boiling point 83 °C (181 °F; 356 K) 68% solution boils at 121 °C (250 °F; 394 K)
Completely miscible
log P −0.13[2]
Vapor pressure 48 mmHg (20 °C)[1]
Acidity (pKa) −1.4[3]
Conjugate base Nitrate
−1.99×10−5 cm3/mol
1.397 (16.5 °C)
2.17 ± 0.02 D
Thermochemistry
146 J·mol−1·K−1[4]
−207 kJ·mol−1[4]
Hazards
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 skull-and-crossbones pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS) The health hazard pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS) The environment pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)
GHS signal word DANGER
H272, H300, H310, H330, H373, H411
P210, P220, P260, P305+351+338, P310, P370+378
NFPA 704
Flash point 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).

Physical and chemical properties

Commercially available nitric acid is an azeotrope with water at a concentration of 68% HNO3, which is the ordinary concentrated nitric acid of commerce. This solution has a boiling temperature of 120.5 °C at 1 atm. Two solid hydrates are known; the monohydrate (HNO3·H2O or [H3O]NO3) and the trihydrate (HNO3·3H2O).

Nitric acid 70 percent
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). The density of concentrated nitric acid is 1.35 g/cm3 (68% conc). This is considered concentrated or technical grade, while reagent grades are specified at 70% HNO3. An older density scale is occasionally seen, with concentrated nitric acid specified as 42° Baumé.[5]

Contamination with nitrogen dioxide

Fuming nitric acid 40ml
Fuming nitric acid contaminated with yellow nitrogen dioxide.

Nitric acid is subject to thermal 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) 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 name "red fuming acid" or "fuming nitric acid" – the most concentrated form of nitric acid at Standard Temperature and Pressure (STP). Nitrogen oxides (NOx) are soluble in nitric acid.

Fuming nitric acid

A commercial grade of 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.

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

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. One specification for white fuming nitric acid is that it has a maximum of 2% water and a maximum of 0.5% dissolved NO2. Anhydrous nitric acid has a density of 1.513 g/cm3 and has the approximate concentration of 24 molar. Anhydrous nitric acid is a colorless mobile liquid with a density of 1.512 g/cm3, which solidifies at −42 °C to form white crystals. As it decomposes to NO2 and water, it obtains a yellow tint. It boils at 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

Nitric-acid-resonance-A
Two major resonance representations of HNO3

The molecule is planar. Two of the N–O bonds are equivalent and relatively short (this can be explained by theories of resonance; the canonical forms show double-bond character in these two bonds, causing them to be shorter than typical N–O bonds), and the third N–O bond is elongated because the O atom is also attached to a proton.[6][7]

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 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 H2SO4NO+
2
+ H3O+ + 2 HSO
4
; Equilibrium constant: K ≈ 22

The nitronium ion, NO+
2
, 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 HNO3NO+
2
+ NO
3
+ 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 H2:

Mg + 2 HNO3 → Mg(NO3)2 + H2 (Magnesium nitrate)
Mn + 2 HNO3 → Mn(NO3)2 + H2 (Manganese 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:

3 Cu + 8 HNO3 → 3 Cu2+ + 2 NO + 4 H2O + 6 NO
3

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:

Cu + 4 H+ + 2 NO
3
→ Cu2+ + 2 NO2 + 2 H2O

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

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

or

3 C + 4 HNO3 → 3 CO2 + 4 NO + 2 H2O

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

Xanthoproteic test

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.

Production

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

4 NO2 + 2 H2O → 2 HNO3 + NO + NO2 + H2O

Or, shortened:

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 atm.

4 NH3 (g) + 5 O2 (g) → 4 NO (g) + 6 H2O (g) (ΔH = −905.2 kJ)

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

2 NO (g) + O2 (g) → 2 NO2 (g) (ΔH = −114 kJ/mol)

This is subsequently absorbed in water to form nitric acid and nitric oxide.

3 NO2 (g) + H2O (l) → 2 HNO3 (aq) + NO (g) (ΔH = −117 kJ/mol)

The nitric oxide is cycled back for reoxidation. Alternatively, if the last step is carried out in air:

4 NO2 (g) + O2 (g) + 2 H2O (l) → 4 HNO3 (aq)

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

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.

2 Cu(NO3)2 → 2 CuO (s) + 4 NO2 (g) + O2 (g)

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 minutes 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.

Dilute nitric acid may be concentrated by distillation up to 68% acid, which is a maximum boiling azeotrope containing 32% water. In the laboratory, further concentration involves distillation with either sulfuric acid or magnesium nitrate which act as dehydrating agents. Such distillations must be done with all-glass apparatus at reduced pressure, to prevent decomposition of the acid. Industrially, highly concentrated nitric acid is produced by dissolving additional nitrogen dioxide in 68% nitric acid in an absorption tower.[13] Dissolved nitrogen oxides are either stripped in the case of white fuming nitric acid, or remain in solution to form red fuming nitric acid. More recently, electrochemical means have been developed to produce anhydrous acid from concentrated nitric acid feedstock.[14]

Uses

Nitric acid lab
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 26 million tonnes 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

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 HNO3C6H2(NO2)3CH3 + 3 H2O

Either concentrated sulfuric acid or oleum absorbs the excess water.

H2S2O7 + H2O → 2 H2SO4

Use as an oxidant

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

Rocket propellant

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

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.[18] 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 HNO
3
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.

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).[19]

Etchant and cleaning agent

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.

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. [21] It is also used to clean glass prior to silvering in the production of silver mirrors. [22]

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.

Nitric acid can be used as a spot test for alkaloids like LSD, giving a variety of colours depending on the alkaloid.[23]

Safety

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.[24] Systemic effects are unlikely, however, and the substance is not considered a carcinogen or mutagen.[25]

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 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.

History

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).[26]

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.[26]

References

  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".
  3. ^ Bell, R. P. (1973), The Proton in Chemistry (2nd ed.), Ithaca, NY: Cornell University Press
  4. ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22. ISBN 978-0-618-94690-7.
  5. ^ Dean, John (1992). Lange's Handbook of Chemistry (14 ed.). McGraw-Hill. pp. 2.79–2.80. ISBN 978-0-07-016194-8.
  6. ^ Luzzati, V. (1951). "Structure cristalline de l'acide nitrique anhydre". Acta Crystallographica (in French). 4 (2): 120–131. doi:10.1107/S0365110X51000404.
  7. ^ a b Allan, D. R.; Marshall, W. G.; Francis, D. J.; Oswald, I. D. H.; Pulham, C. R.; Spanswick, C. (2010). "The crystal structures of the low-temperature and high-pressure polymorphs of nitric acid" (PDF). Dalton Trans. (Submitted manuscript). 39 (15): 3736–3743. doi:10.1039/B923975H. PMID 20354626.
  8. ^ IUPAC SC-Database A comprehensive database of published data on equilibrium constants of metal complexes and ligands
  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 978-1-4086-2802-7.
  11. ^ Knowles, Frank (2007). A practical course in agricultural chemistry. Read Books. p. 76. ISBN 978-1-4067-4583-2.
  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.
  14. ^ US 6200456, Harrar, Jackson E.; Roland Quong & Lester P. Rigdon 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
  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 978-3527306732.
  16. ^ Clark, John D (1972). Ignition!. Rutgers University Press. ISBN 978-0-8135-0725-5.
  17. ^ "BOMARC Summary". BILLONY.COM. Retrieved 2009-05-28.
  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.
  20. ^ Muraoka, Hisashi (1995) "Silicon wafer cleaning fluid with HNO3, HF, HCl, surfactant, and water" U.S. Patent 5,635,463
  21. ^ Zeller, Rolf; Rose, Jack; Jacobson, Kenneth A.; Fischer, Andrew H. (2008-05-01). "Preparation of Slides and Coverslips for Microscopy". Cold Spring Harbor Protocols. 2008 (5): pdb.prot4988. doi:10.1101/pdb.prot4988. ISSN 1559-6095. PMID 21356831.
  22. ^ Curtis, Heber D. (February 1911). "METHODS OF SILVERING MIRRORS". Publications of the Astronomical Society of the Pacific. 23 (135): 13. doi:10.1086/122040. ISSN 1538-3873.
  23. ^ 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.
  24. ^ May, Paul (November 2007). "Nitric acid". Retrieved 2009-05-28.
  25. ^ "Nitric acid: Toxicological overview". Health Protection Agency. Retrieved 2011-12-07.
  26. ^ a b Wikisource Chisholm, Hugh, ed. (1911). "Nitric Acid" . Encyclopædia Britannica. 19 (11th ed.). Cambridge University Press. pp. 711–712.

External links

Aluminium nitrate

Aluminium nitrate is a white, water-soluble salt of aluminium and nitric acid, most commonly existing as the crystalline hydrate, aluminium nitrate nonahydrate, Al(NO3)3·9H2O.

Aqua regia

Aqua regia (; from Latin, lit. "regal water" or "king's water") is a mixture of nitric acid and hydrochloric acid, optimally in a molar ratio of 1:3. Aqua regia is a yellow-orange (sometimes red) fuming liquid, so named by alchemists because it can dissolve the noble metals gold and platinum, though not all metals.

Barium nitrate

Barium nitrate is the inorganic compound with the chemical formula Ba(NO3)2. It, like most barium salts, is colorless, toxic, and water-soluble. It burns with a green flame and is an oxidizer; the compound is commonly used in pyrotechnics.

Cobalt(II) nitrate

Cobalt Nitrate is the Inorganic compound with the formula Co(NO3)2.xH2O. It is cobalt(II) salt. The most common form is the hexahydrate Co(NO3)2·6H2O, which is a red-brown deliquescent salt that is soluble in water and other polar solvents.

Dinitrogen tetroxide

Dinitrogen tetroxide, commonly referred to as nitrogen tetroxide, is the chemical compound N2O4. It is a useful reagent in chemical synthesis. It forms an equilibrium mixture with nitrogen dioxide.

Dinitrogen tetroxide is a powerful oxidizer that is hypergolic (spontaneously reacts) upon contact with various forms of hydrazine, which has made the pair a common bipropellant for rockets.

Ethyl nitrate

Ethyl nitrate is the ethyl ester of nitric acid and has the chemical formula C2H5NO3. It is a colourless, volatile, highly flammable liquid. It is used in organic synthesis and as an intermediate in the preparation of some drugs, dyes, and perfumes.Ethyl nitrate is found in the atmosphere, where it can react with other gases to form smog. Originally thought to be a pollutant, formed mainly by the combustion of fossil fuels, recent analysis of ocean water samples reveal that in places where cool water rises from the deep, the water is saturated with alkyl nitrates, likely formed by natural processes.

NOx

In atmospheric chemistry, NO x is a generic term for the nitrogen oxides that are most relevant for air pollution, namely nitric oxide (NO) and nitrogen dioxide (NO 2 ). These gases contribute to the formation of smog and acid rain, as well as affecting tropospheric ozone.

NO x gases are usually produced from the reaction among nitrogen and oxygen during combustion of fuels, such as hydrocarbons, in air; especially at high temperatures, such as occur in car engines. In areas of high motor vehicle traffic, such as in large cities, the nitrogen oxides emitted can be a significant source of air pollution. NO x gases are also produced naturally by lightning.

The term NO x is chemistry shorthand for molecules containing one nitrogen and one or more oxygen atom. It is generally meant to include nitrous oxide (N2O), although nitrous oxide is a fairly inert oxide of nitrogen that has many uses as an oxidizer for rockets and car engines, an anesthetic, and a propellant for aerosol sprays and whipped cream. Nitrous oxide plays hardly any role in air pollution, although it may have a significant impact on the ozone layer, and is a significant greenhouse gas.

NO y is defined as the sum of NO x plus the NO z compounds produced from the oxidation of NO x which include nitric acid, nitrous acid(HONO), dinitrogen pentoxide(N2O5), peroxyacetyl nitrate(PAN), alkyl nitrates (RONO2), peroxyalkyl nitrates (ROONO2), the nitrate radical (NO3), and peroxynitric acid(HNO4).

Nitration

Nitration is a general class of chemical process for the introduction of a nitro group into an organic chemical compound. More loosely the term also is applied incorrectly to the different process of forming nitrate esters between alcohols and nitric acid, as occurs in the synthesis of nitroglycerin. The difference between the resulting structure of nitro compounds and nitrates is that the nitrogen atom in nitro compounds is directly bonded to a non-oxygen atom, typically carbon or another nitrogen atom, whereas in nitrate esters, also called organic nitrates, the nitrogen is bonded to an oxygen atom that in turn usually is bonded to a carbon atom (nitrito group).

There are many major industrial applications of nitration in the strict sense; the most important by volume are for the production of Nitroaromatic compounds such as nitrobenzene. Nitration reactions are notably used for the production of explosives, for example the conversion of guanidine to nitroguanidine and the conversion of toluene to trinitrotoluene. However, they are of wide importance as chemical intermediates and precursors. Millions of tons of nitroaromatics are produced annually.

Nitrocellulose

Nitrocellulose (also known as cellulose nitrate, flash paper, flash cotton, guncotton, and flash string) is a highly flammable compound formed by nitrating cellulose through exposure to nitric acid or another powerful nitrating agent. When used as a propellant or low-order explosive, it was originally known as guncotton.

Partially nitrated cellulose has found uses as a plastic film and in inks and wood coatings. In 1862, the first man-made plastic, nitrocellulose (branded Parkesine), was created by Alexander Parkes from cellulose treated with nitric acid and a solvent. In 1868, American inventor John Wesley Hyatt developed a plastic material he named Celluloid, improving on Parkes' invention by plasticizing the nitrocellulose with camphor so it could be processed into finished form and used as a photographic film. Celluloid was used by Kodak, and other suppliers, from the late 1880s as a film base in photography, X-ray films, and motion-picture films, and was known as nitrate film. After numerous fires caused by unstable nitrate films, "safety film" (cellulose acetate film) started to be used from the 1930s in the case of X-ray stock and from 1948 for motion-picture film.

Nitrogen

Nitrogen is a chemical element with symbol N and atomic number 7. It was first discovered and isolated by Scottish physician Daniel Rutherford in 1772. Although Carl Wilhelm Scheele and Henry Cavendish had independently done so at about the same time, Rutherford is generally accorded the credit because his work was published first. The name nitrogène was suggested by French chemist Jean-Antoine-Claude Chaptal in 1790, when it was found that nitrogen was present in nitric acid and nitrates. Antoine Lavoisier suggested instead the name azote, from the Greek ἀζωτικός "no life", as it is an asphyxiant gas; this name is instead used in many languages, such as French, Russian, Romanian and Turkish, and appears in the English names of some nitrogen compounds such as hydrazine, azides and azo compounds.

Nitrogen is the lightest member of group 15 of the periodic table, often called the pnictogens. The name comes from the Greek πνίγειν "to choke", directly referencing nitrogen's asphyxiating properties. It is a common element in the universe, estimated at about seventh in total abundance in the Milky Way and the Solar System. At standard temperature and pressure, two atoms of the element bind to form dinitrogen, a colourless and odorless diatomic gas with the formula N2. Dinitrogen forms about 78% of Earth's atmosphere, making it the most abundant uncombined element. Nitrogen occurs in all organisms, primarily in amino acids (and thus proteins), in the nucleic acids (DNA and RNA) and in the energy transfer molecule adenosine triphosphate. The human body contains about 3% nitrogen by mass, the fourth most abundant element in the body after oxygen, carbon, and hydrogen. The nitrogen cycle describes movement of the element from the air, into the biosphere and organic compounds, then back into the atmosphere.

Many industrially important compounds, such as ammonia, nitric acid, organic nitrates (propellants and explosives), and cyanides, contain nitrogen. The extremely strong triple bond in elemental nitrogen (N≡N), the second strongest bond in any diatomic molecule after carbon monoxide (CO), dominates nitrogen chemistry. This causes difficulty for both organisms and industry in converting N2 into useful compounds, but at the same time means that burning, exploding, or decomposing nitrogen compounds to form nitrogen gas releases large amounts of often useful energy. Synthetically produced ammonia and nitrates are key industrial fertilisers, and fertiliser nitrates are key pollutants in the eutrophication of water systems.

Apart from its use in fertilisers and energy-stores, nitrogen is a constituent of organic compounds as diverse as Kevlar used in high-strength fabric and cyanoacrylate used in superglue. Nitrogen is a constituent of every major pharmacological drug class, including antibiotics. Many drugs are mimics or prodrugs of natural nitrogen-containing signal molecules: for example, the organic nitrates nitroglycerin and nitroprusside control blood pressure by metabolizing into nitric oxide. Many notable nitrogen-containing drugs, such as the natural caffeine and morphine or the synthetic amphetamines, act on receptors of animal neurotransmitters.

Nitrogen dioxide

Nitrogen dioxide is the chemical compound with the formula NO2. It is one of several nitrogen oxides. NO2 is an intermediate in the industrial synthesis of nitric acid, millions of tons of which are produced each year which is used primarily in the production of fertilizers. At higher temperatures it is a reddish-brown gas that has a characteristic sharp, biting odor and is a prominent air pollutant. Nitrogen dioxide is a paramagnetic, bent molecule with C2v point group symmetry.

Nitroglycerin

Nitroglycerin (NG), also known as nitroglycerine, trinitroglycerin (TNG), nitro, glyceryl trinitrate (GTN), or 1,2,3-trinitroxypropane, is a dense, colorless, oily, explosive liquid most commonly produced by nitrating glycerol with white fuming nitric acid under conditions appropriate to the formation of the nitric acid ester. Chemically, the substance is an organic nitrate compound rather than a nitro compound, yet the traditional name is often retained. Invented in 1847, nitroglycerin has been used as an active ingredient in the manufacture of explosives, mostly dynamite, and as such it is employed in the construction, demolition, and mining industries. Since the 1880s, it has been used by the military as an active ingredient, and a gelatinizer for nitrocellulose, in some solid propellants, such as cordite and ballistite.

Nitroglycerin is a major component in double-based smokeless gunpowders used by reloaders. Combined with nitrocellulose, hundreds of powder combinations are used by rifle, pistol, and shotgun reloaders.

In medicine for over 130 years, nitroglycerin has been used as a potent vasodilator (dilation of the vascular system) to treat heart conditions, such as angina pectoris and chronic heart failure. Though it was previously known that these beneficial effects are due to nitroglycerin being converted to nitric oxide, a potent venodilator, the enzyme for this conversion was not discovered to be mitochondrial aldehyde dehydrogenase (ALDH2) until 2002. Nitroglycerin is available in sublingual tablets, sprays, and patches.

Nitrous acid

Nitrous acid (molecular formula HNO2) is a weak and monobasic acid known only in solution and in the form of nitrite (NO−2) salts. Nitrous acid is used to make diazonium salts from amines. The resulting diazonium salts are reagents in azo coupling reactions to give azo dyes.

Oleum

Oleum (Latin oleum, meaning oil), or fuming sulfuric acid, is a solution of various compositions of sulfur trioxide in sulfuric acid, or sometimes more specifically to disulfuric acid (also known as pyrosulfuric acid). Oleum is identified by the CAS number 8014-95-7 (EC/List number: 616-954-1 ; ECHA InfoCard: 100.116.872).

Oleums can be described by the formula ySO3.H2O where y is the total molar sulfur trioxide content. The value of y can be varied, to include different oleums. They can also be described by the formula H2SO4.xSO3 where x is now defined as the molar free sulfur trioxide content. Oleum is generally assessed according to the free SO3 content by mass. It can also be expressed as a percentage of sulfuric acid strength; for oleum concentrations, that would be over 100%. For example, 10% oleum can also be expressed as H2SO4.0.13611SO3, 1.0225SO3.H2O or 102.25% sulfuric acid. The conversion between % acid and % oleum is: % acid = 100 + 18/80 × % oleum

A value for x of 1 gives the empirical formula H2S2O7 for disulfuric (pyrosulfuric) acid. Pure disulfuric acid is a solid at room temperature, melting at 36 °C and rarely used either in the laboratory or industrial processes.

Polar stratospheric cloud

Polar stratospheric clouds (PSCs), also known as nacreous clouds (, from nacre, or mother of pearl, due to its iridescence), are clouds in the winter polar stratosphere at altitudes of 15,000–25,000 m (49,000–82,000 ft). They are best observed during civil twilight, when the Sun is between 1 and 6 degrees below the horizon, as well as in winter and in more northerly latitudes. They are implicated in the formation of ozone holes. The effects on ozone depletion arise because they support chemical reactions that produce active chlorine which catalyzes ozone destruction, and also because they remove gaseous nitric acid, perturbing nitrogen and chlorine cycles in a way which increases ozone depletion.

Potassium nitrate

Potassium nitrate is a chemical compound with the chemical formula KNO3. It is an ionic salt of potassium ions K+ and nitrate ions NO3−, and is therefore an alkali metal nitrate.

It occurs in nature as a mineral, niter. It is a source of nitrogen, from which it derives its name. Potassium nitrate is one of several nitrogen-containing compounds collectively referred to as saltpeter or saltpetre.

Major uses of potassium nitrate are in fertilizers, tree stump removal, rocket propellants and fireworks. It is one of the major constituents of gunpowder (black powder). In processed meats, potassium nitrate reacts with hemoglobin and generates a pink color.

Red fuming nitric acid

Red fuming nitric acid (RFNA) is a storable oxidizer used as a rocket propellant. It consists of 84% nitric acid (HNO3), 13% dinitrogen tetroxide and 1–2% water. The color of red fuming nitric acid is due to the dinitrogen tetroxide, which breaks down partially to form nitrogen dioxide. The nitrogen dioxide dissolves until the liquid is saturated, and evaporates off into fumes with a suffocating odor. RFNA increases the flammability of combustible materials and is highly exothermic when reacting with water.

It is usually used with an inhibitor (with various, sometimes secret, substances, including hydrogen fluoride; any such combination is called inhibited RFNA, IRFNA) because nitric acid attacks most container materials.

It can also be a component of a monopropellant; with substances like amine nitrates dissolved in it, it can be used as the sole fuel in a rocket. It is not normally used this way however.

During World War II, the German military used RFNA in some rockets. The mixtures used were called S-Stoff (96% nitric acid with 4% ferric chloride as an ignition catalyst) and SV-Stoff (94% nitric acid with 6% dinitrogen tetroxide) and nicknamed Salbei (sage).

Inhibited RFNA was the oxidizer of the world's most-launched light orbital rocket, the Kosmos-3M.

Other uses for RFNA include fertilizers, dye intermediates, explosives, and pharmaceutic aid as acidifier. It can also be used as a laboratory reagent in photoengraving and metal etching.

Soil acidification

Soil acidification is the buildup of hydrogen cations, also called protons, reducing the soil pH. Chemically, this happens when a proton donor gets added to the soil. The donor can be an acid, such as nitric acid and sulfuric acid (these acids are common components of acid rain). It can also be a compound such as aluminium sulfate, which reacts in the soil to release protons. Many nitrogen compounds, which are added as fertilizer, also acidify soil over the long term because they produce nitrous and nitric acid when oxidized in the process of nitrification.

Acidification also occurs when base cations such as calcium, magnesium, potassium and sodium are leached from the soil. This leaching increases with increasing precipitation. Acid rain accelerates the leaching of these bases. Plants use bases as they grow, and when plant material is removed through activities such as logging or crop harvesting, the base elements within are permanently lost from the soil.

White fuming nitric acid

White fuming nitric acid (WFNA) is a storable liquid oxidizer used with kerosene and hydrazine rocket fuel. It consists of nearly pure nitric acid (HNO3). WFNA is commonly specified as containing no more than 2% water and less than 0.5% dissolved nitrogen dioxide or dinitrogen tetroxide.

WFNA was sometimes used with an inhibitor compound to reduce corrosiveness, often hydrogen fluoride. Without inhibitors, WFNA will corrode nearly all structural metals. Inhibited WFNA is often called IWFNA. The hydrogen fluoride addition causes formation of protective layer of fluoride on the metal surfaces.

WFNA as an oxidizer has somewhat less performance than red fuming nitric acid (RFNA) but is considerably safer (though extremely corrosive), as it has little to no dissolved nitrogen tetroxide, which is an extremely toxic and volatile chemical. If not inhibited, it will form nitrogen tetroxide on contact with most metals and some organic materials. RFNA can be converted from WFNA by simply leaving the RFNA out in low temperature for a couple of hours.

WFNA is also used in the manufacture of nitroglycerin, an explosive, by mixing it with concentrated sulfuric acid to produce the nitro group ester, and then by the slow addition of glycerol. This has now mostly been replaced by the less expensive procedure that uses a nearly 1:1 solution of oleum and azeotropic nitric acid (70%).

WFNA and IWFNA are hypergolic with a long list of other propellants, including: UDMH, hydrazine, furfuryl alcohol, and aniline.

Hydrogen compounds
Nitrogen species

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