|Molar mass||63.012 g/mol|
|Appearance||Clear, colorless liquid|
-42 °C, 231 K, -44 °F
83 °C, 356 K, 181 °F (bp of pure acid. 68% solution boils at 120.5 °C)
|Solubility in water||completely miscible|
|Refractive index (nD)||1.397 (16.5 °C)|
|Dipole moment||2.17 ± 0.02 D|
|EU classification||Oxidant (O)
|Other anions||Nitrous acid|
|Other cations||Sodium nitrate
|Related compounds||Dinitrogen pentoxide|
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Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Colorless when pure, older samples tend to acquire a yellow cast due to the accumulation of oxides of nitrogen. If the solution contains more than 86% nitric acid, it is referred to as fuming nitric acid. Fuming nitric acid is characterized as white fuming nitric acid and red fuming nitric acid, depending on the amount of nitrogen dioxide present. At concentrations above 95% at room temperature, it tends to rapidly develop a yellow color due to slow decomposition.
Pure anhydrous nitric acid (100%) is a colorless mobile liquid with a density of 1.522 g/cm3 which solidifies at −42 °C to form white crystals and boils at 83 °C. When boiling in light, even at room temperature, there is a partial decomposition with the formation of nitrogen dioxide following the reaction:
which means that anhydrous nitric acid should be stored below 0 °C to avoid decomposition. The nitrogen dioxide (NO2) remains dissolved in the nitric acid coloring it yellow, or 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". Fuming nitric acid is also referred to as 16 molar nitric acid. It is the most concentrated form of nitric acid at Standard Temperature and Pressure (STP).
Nitric acid is miscible with water and distillation gives a maximum-boiling azeotrope with a concentration of 68% HNO3 and a boiling temperature of 120.5 °C at 1 atm, which is the ordinary concentrated nitric acid of commerce. Two solid hydrates are known; the monohydrate (HNO3·H2O) and the trihydrate (HNO3·3H2O).
Nitrogen oxides (NOx) are soluble in nitric acid and this property influences more or less all the physical characteristics depending on the concentration of the oxides. These mainly include the vapor pressure above the liquid and the boiling temperature, as well as the color mentioned above.
Nitric acid is subject to thermal or light decomposition with increasing concentration and this 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.
Being a typical strong acid, nitric acid reacts with alkalis, basic oxides, and carbonates to form salts, such as ammonium nitrate. Due to its oxidizing nature, nitric acid generally does not donate its proton (that is, it does not liberate hydrogen) on reaction with metals and the resulting salts are usually in the higher oxidized states. For this reason, heavy corrosion can be expected and should be guarded against by the appropriate use of corrosion resistant metals or alloys.
Nitric acid has an acid dissociation constant (pKa) of −1.4. In aqueous solution, it almost completely (93% at 0.1 mol/L) ionizes into the nitrate ion NO−3 and a hydrated proton, known as a hydronium ion, H3O+.
Being a powerful oxidizing agent, nitric acid reacts violently with many organic materials 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 a few of the precious metal series and certain alloys. This characteristic has made it a common agent to be used in acid tests. As a general rule, oxidizing reactions occur primarily with the concentrated acid, favoring the formation of nitrogen dioxide (NO2).
The acidic properties tend to dominate with dilute acid, coupled with the preferential formation of nitric oxide (NO). However, when the reaction is carried out in the presence of atmospheric oxygen, the nitric oxide rapidly reacts to form brown nitrogen dioxide (NO2):
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 metal from further oxidation, which is called passivation. Typical passivation concentrations range from 18% to 22% by weight.
Reaction with non-metallic elements, with the exceptions of nitrogen, oxygen, noble gases, silicon and halogens, 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.
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 are present that contains amino acids with aromatic rings, the mixture turns yellow. Upon adding a strong base such as liquid ammonia, the color turns orange. These color changes are caused by nitrated aromatic rings in the protein. Xanthoproteins are formed when the acid contacts epithelial cells and are indicative of inadequate safety precautions when handling nitric acid.
The concentrated nitric acid of commerce consists of the maximum boiling azeotrope of nitric acid and water. Technical grades are normally 68% HNO3, (approx 15 molar), while reagent grades are specified at 70% HNO3. The density of concentrated nitric acid is 1.42 g/mL. An older density scale is occasionally seen, with concentrated nitric acid specified as 42° Baumé.
White fuming nitric acid, also called 100% nitric acid or WFNA, is very close to anhydrous nitric acid. 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/mL and has the approximate concentration of 24 molar.
A commercial grade of fuming nitric acid, referred to in the trade as "strong nitric acid" contains 90% HNO3 and has a density of 1.50 g/mL. This grade is much 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 molar.
Red fuming nitric acid, or RFNA, contains substantial quantities of dissolved nitrogen dioxide (NO2) leaving the solution with a reddish-brown color. One formulation of RFNA specifies a minimum of 17% NO2, another specifies 13% NO2. Because of the dissolved nitrogen dioxide, the density of red fuming nitric acid is lower at 1.490 g/mL.
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.
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, strong nitric acid is produced by dissolving additional nitrogen dioxide in 68% nitric acid in an absorption tower. 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.
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, which is then reacted with oxygen in air to form nitrogen dioxide. This is subsequently absorbed in water to form nitric acid and nitric oxide. The nitric oxide is cycled back for reoxidation. 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.
Prior to the introduction of the Haber process for the production of ammonia in 1923, 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.
In laboratory, nitric acid can be made from copper(II) nitrate or by reacting approximately equal masses of a nitrate salt with 96% sulfuric acid (H2SO4), and distilling this mixture at nitric acid's boiling point of 83 °C until only a white crystalline mass, a metal sulfate, remains in the reaction vessel. The red fuming nitric acid obtained may be converted to the white nitric acid.
The dissolved NOx are 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.
The main use of nitric acid is for the production of fertilizers; other important uses include the production of explosives, etching and dissolution of metals, especially as a component of aqua regia for the purification and extraction of gold, and in chemical synthesis.
In elemental analysis by ICP-MS, ICP-AES, GFAA, and Flame AA, dilute nitric acid (0.5 to 5.0 %) is used as a matrix compound for determining metal traces in solutions. 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 FAA. Typically these digestions use a 50% solution of the purchased HNO3 mixed with Type 1 DI Water.
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).
Ahmed Ressam, the al-Qaeda Millenium Bomber, used nitric acid as one of the components in the explosives that he prepared to bomb Los Angeles International Airport on New Year's Eve 1999/2000; the explosives could have produced a blast 40x greater than that of a devastating car bomb.
A solution of nitric acid and alcohol, Nital, is used for etching of metals to reveal the microstructure.
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).
A mixture of concentrated nitric and sulfuric acids causes the nitration of aromatic compounds, such as benzene. Examination of the infrared spectrum of the acid mixture using a corrosive resistant diamond cell shows infrared peaks close to that expected for carbon dioxide. The species responsible for the peaks is the nitronium ion, NO+2, which like CO2, is a linear molecule. The nitronium ion is the species responsible for nitration: being positive, it is attacked by electron-rich benzene rings. This is described more fully in organic chemistry books.
Nitric acid is a powerful oxidizing agent, and the reactions of nitric acid with compounds such as cyanides, carbides, and metallic powders can be explosive. Reactions of nitric acid with many organic compounds, such as turpentine, are violent and hypergolic (i.e., self-igniting). Due to its properties it is stored away from bases and organics.
NITRIC ACID (aqua fortis), HN03, an important mineral acid. It is mentioned in the De inventione veritatis ascribed to Geber, wherein it is obtained by calcining a mixture of nitre, alum and blue vitriol. It was again described by Albert le Grand in the 13th century and by Raimon Lull, who prepared it by heating nitre and clay and called it "eau forte." Glauber devised the process in common use to-day, viz. by heating nitre with strong sulphuric acid. Its true nature was not determined until the 8th century, when A. L. Lavoisier (1776) showed that it contained oxygen, whilst in 1785 H. Cavendish determined its constitution and showed that it could be synthesized by passing a stream of electric sparks through moist air. The acid is found to exist to a slight extent in the free condition in some waters, but chiefly occurs in combination with various metals, as nitrates, principally as nitre or saltpetre, KN03, and Chile saltpetre, NaNO 3. It is formed when a stream of electric sparks is passed through moist air, and in the oxidation ',of nitrogenous matter in the presence of water.
For experimental purposes it is usually obtained by distilling potassium or sodium nitrate with concentrated sulphuric acid. The acid so obtained usually contains more or less water and some dissolved nitrogen peroxide which gives it a yellowish red colour. It may be purified by redistillation over barium and silver nitrates, followed by treatment of the distillate with a stream of ozonized air. The product so obtained is then redistilled under diminished pressure and finally distilled again from a sealed and evacuated apparatus (V. Veley and Manley, Phil. Trans., 1898, A. 291, p. 365). On the large scale it is obtained by distilling Chile saltpetre with concentrated sulphuric acid in horizontal cast iron stills, the vapours being condensed in a series of stoneware Woulfe's bottles. In practice the theoretical quantity of acid and Chile saltpetre is not used, but the charge is so regulated that the mixture of acid and neutral sodium sulphate formed in the retort remains liquid at the temperature employed, and consequently can be readily removed. Various modifications have been made in the form of the condensing apparatus, the Guttmann condenser (Jour. Soc. Chem. Ind., 1893, p. 203) being now frequently employed. This consists of a series of vertical earthenware condensing tubes through which compressed air is passed in order to reduce the quantity of nitrogen peroxide to a minimum. The temperature of the condenser is so regulated as to bring about the condensation of the nitric acid only, which runs out at the bottom of the pipe, whilst any uncondensed steam, nitrogen peroxide and other impurities pass into a Lunge tower, where they meet a descending stream of water and are condensed, giving rise to an impure acid. F. Valentiner [Eng. Pat. 610 (1892), 19192 (1895)] recommends distillation and condensation of nitric acid in a partial vacuum. For the production of nitric acid from air see Nitrogen. Fuming nitric acid consists of a solution of nitrogen peroxide in concentrated nitric acid and is prepared by distilling dry sodium nitrate with concentrated sulphuric acid.
Nitric acid is a colourless strongly fuming liquid, having a specific gravity of 1.50394 (24.2° C.) (V. Veley, Proc. Roy. Soc., 62, p. 223). It is exceedingly hygroscopic and corrosive. On distillation, the pure acid begins to boil at 78.2° C. (W. Ramsay), partial decomposition into water, oxygen and nitrogen peroxide taking place. The acid solidifies when strongly cooled, the solid melting at - 47° C. Concentrated nitric acid forms with water a constant boiling mixture which boils at 120.5° C., contains 68% of acid and possesses a specific gravity of 1.414 (15.5° C.). If a more dilute acid than this be distilled, water passes over in excess and the residue in the retort reaches the above composition and boiling point; on distillation of a stronger acid, excess of acid passes into the distillate and the boiling point rises until the values of the constant boiling mixture are reached. On the hydrates of nitric acid see V. Veley, Jour. Chem. Soc., 1903, 83, p. 1015, and F. W. Kuster, Zeit. anorg., Chem. 1904, 41, p. 1. On mixing nitric acid with water there is a rise of temperature and a contraction in volume. The acid is a powerful oxidizing agent. It attacks most metals readily, usually with production of a nitrate or hydrated oxide of the metal and one or other of the oxides of nitrogen, or occasionally with the production of ammonium salts; magnesium, however, liberates hydrogen from the very dilute acid. Its action on metals depends in most cases on the temperature, strength of the acid, and the nature of the products of reaction. Thus in the case of copper, it is found that the diluted acid acts very slowly upon the metal at first, but as the reaction proceeds the copper dissolves more rapidly up to a certain point and then the rate of solution again diminishes. This is possibly due to the accelerating action of the nitrous acid which is produc-ed in the direct action of the copper on the nitric acid and which, when a certain amount has been formed in the system, begins to decompose, thus 3HNO 2 = HN03+ 2N0+H 2 0 (V. Veley, Phil. Trans., 1891, 182, p. 279; G. 0. Higley, Amer. Chem. Jour., 18 93, 1 5, p. 71, 18 95, 1 7, p. 18, 1896, 18, p. 587). Iron when brought into contact with nitric acid under certain conditions, remains passive to the acid. Thus at 55° C. it is passive to an acid of specific gravity 1.42, and at 31° C. to an acid of specific gravity. 38. No satisfactory explanation of this passivity has been given (see J. B. Senderens, Bull. Soc. Chem., 18 9 6 , 1 5, p. 691; A. Finkelstein, Zeit. phys. Chem., 1901, 39, p. 91; W. J. Muller, ibid. 1904, 4 8, p. 577). Nitric acid is without action on gold, platinum, iridium and rhodium.
The salts of nitric acid, known as nitrates, are mostly readily soluble in water and crystallize well. They are all decomposed when heated to a sufficiently high temperature, with evolution for the most part of oxygen and nitrogen peroxide, leaving a residue of oxide of the metal. They may be recognized by the fact that on the addition of a solution of ferrous sulphate, followed by that of concentrated sulphuric acid (the mixture being kept quite cold), the ferrous sulphate solution becomes of a deep brown colour, owing to the reducing action of the ferrous sulphate on the nitric acid which is liberated by the action of the sulphuric acid on the nitrate. As an alternative method the nitrate may be warmed with some fragments of copper and sulphuric acid which has been diluted with its own volume of water, when characteristic brown vapours will be seen.
In medicine, nitric acid is used externally in a pure state as a caustic to destroy chancres, warts and phagadenic ulcers; and diluted preparations are employed in the treatment of dyspepsia, &c. Poisoning by strong nitric acid produces a widespread gastroenteritis, burning pain in the oesophagus and abdomen and bloody diarrhoea. There may also be blood in the urine. Death occurs from collapse or from secondary destructive changes in the intestinal canal. Characteristic yellow staining of the skin round the mouth from the formation of xanthoproteic acid serves to distinguish it from poisoning by other acids. The antidotes are mild alkalis, together with the use of opium to relieve pain.