Nitric acid

Nitric acid is an inorganic compound with the formula. It is a highly corrosive mineral acid. The compound is colorless, but samples tend to acquire a yellow cast over time due to decomposition into oxides of nitrogen. Most commercially available nitric acid has a concentration of 68% in water. When the solution contains more than 86%, 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 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 synthetic dyes and medicines. Nitric acid is also commonly used as a strong oxidizing agent.

History

Medieval alchemy

The discovery of mineral acids such as nitric acid is generally presumed to go back to 13th-century European alchemy. The conventional view is that nitric acid was first described in pseudo-Geber's De inventione veritatis.
However, according to Eric John Holmyard and Ahmad Y. al-Hassan, nitric acid was also referenced in various earlier Arabic works such as the attributed to Jabir ibn Hayyan or the attributed to the Fatimid caliph al-Hakim bi-Amr Allah.
The recipe in the attributed to Jabir has been translated as follows:
Nitric acid is also found in post-1300 works falsely attributed to Albert the Great and Ramon Llull. These works describe the distillation of a mixture containing niter and green vitriol, which they call eau forte.

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, "combined with an approximately equal volume of the purest part of common air, and with a considerable quantity of water." 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. 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.
The industrial production of nitric acid from atmospheric air began in 1905 with 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 with a very high temperature electric arc. Yields of up to approximately 4–5% nitric oxide were obtained at 3,000 °C, and less at lower temperatures. 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 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. 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. A pit was dug into the peat, then lined with tarred timber stakes around the sides. Into this pit was placed a porous earthenware vessel, surrounded by crushed limestone. The interior was filled with coke around a graphite anode. Nitric acid was pumped out via a glass tube that was sunk down nearly to the bottom of the pot, while fresh water was pumped into the top through another glass pipe to replace the fluid removed. 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 a one hectare deposit, 6.5 feet deep, was estimated to be in excess of 600 tons per year.
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%. This solution has a boiling temperature of at. 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 or oxonium nitrate and the trihydrate.
An older density scale is occasionally seen, with concentrated nitric acid specified as 42 Baumé.

Contamination with nitrogen dioxide

Nitric acid is subject to thermal or light decomposition and for this reason it was often stored in brown glass bottles:
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 and/or dinitrogen tetroxide 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 are soluble in nitric acid.

Fuming nitric acid

Commercial-grade fuming nitric acid contains 98% 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 leaving the solution with a reddish-brown color. The presence of dissolved nitrogen dioxide increases the density of the aqueous and the anhydrous acid, anhydrous nitric acid containing 25% NO₂ has a density of around 1.60 g/cm³, the density is lower in water-containing mixtures. The maximum density of anhydrous nitric acid is attained at 40% NO₂.
An inhibited fuming nitric acid, either white inhibited fuming nitric acid, or red inhibited fuming nitric acid, can be made by the addition of 0.6 to 0.7% hydrogen fluoride. 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 molar. One specification for white fuming nitric acid is that it has a maximum of 2% water and a maximum of 0.5% dissolved. Anhydrous nitric acid is a colorless, low-viscosity liquid with a density of 1.512-3 g/cm³ that solidifies at to form white crystals. Its dynamic viscosity under standard conditions is 0.76 mPa·s. As it decomposes to and water, it obtains a yellow tint. It boils at. 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

The two terminal N–O bonds are nearly equivalent and relatively short, at 1.20 and 1.21 Å. This can be explained by theories of resonance; the two major canonical forms show some double bond character in these two bonds, causing them to be shorter than N–O single bonds. The third N–O bond is elongated because its O atom is bonded to H atom, with a bond length of 1.41 Å in the gas phase. The molecule is slightly aplanar and there is restricted rotation about the N–OH single bond.

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 pK_a 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 pK_a value rises to 1 at a temperature of 250 °C.
Nitric acid can act as a base with respect to an acid such as sulfuric acid:
The nitronium ion,, 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:

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|:
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:
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:
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 Tin oxide|, Arsenic pentoxide|, Antimony pentoxide|, and Titanium dioxide| respectively.
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 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 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 metals series and certain alloys. As a general rule, oxidizing reactions occur primarily with the concentrated acid, favoring the formation of nitrogen dioxide. 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 greatly increases the rate of reaction.
Although chromium, iron, and aluminium 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. Metals that are passivated by concentrated nitric acid are iron, cobalt, chromium, nickel, and aluminium.