Editing Nitric acid (section)

==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.