nitric acid(redirected from HNO 3)
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nitric acid,chemical compound, HNO3, colorless, highly corrosive, poisonous liquid that gives off choking red or yellow fumes in moist air. It is miscible with water in all proportions. It forms an azeotrope (constant-boiling mixture) that has the composition 68% nitric acid and 32% water and that boils at 120.5°C;. The nitric acid of commerce is typically a solution of 52% to 68% nitric acid in water. Solutions containing over 86% nitric acid are commonly called fuming nitric acid. White fuming nitric acid (WFNA) is similar to the anhydrous variety, and red fuming nitric acid (RFNA) has a reddish brown color from dissolved nitrogen oxides. When treated with hydrogen fluoride, both varieties form inhibited fuming nitric acid, which has increased corrosion resistance in metal tanks, e.g., when used as an oxidizer in liquid fuel rockets.
Nitric acid is a strong oxidizing agent. It ionizes readily in solution, forming a good conductor of electricity. It reacts with metals, oxides, and hydroxides, forming nitratenitrate,
chemical compound containing the nitrate (NO3) radical. Nitrates are salts or esters of nitric acid, HNO3, formed by replacing the hydrogen with a metal (e.g., sodium or potassium) or a radical (e.g., ammonium or ethyl).
..... Click the link for more information. salts. Chief uses of nitric acid are in the preparation of fertilizers, e.g., ammonium nitrateammonium nitrate,
chemical compound, NH4NO3, that exists as colorless, rhombohedral crystals at room temperature but changes to monoclinic crystals when heated above 32°C;. It is extremely soluble in water and soluble in alcohol and liquid ammonia.
..... Click the link for more information. , and explosives, e.g., nitroglycerin and trinitrotoluene (TNT). It is also used in the manufacture of chemicals, e.g., in making dyes, and in metallurgy, ore flotation, etching steel, photoengraving, and reprocessing of spent nuclear fuel. It is produced chiefly by oxidation of ammonia (the Ostwald process). Small amounts are produced by the treatment of sodium nitratesodium nitrate,
chemical compound, NaNO3, a colorless, odorless crystalline compound that closely resembles potassium nitrate (saltpeter or niter) in appearance and chemical properties. It is soluble in water, alcohol, and liquid ammonia.
..... Click the link for more information. with sulfuric acid. Nitric acid was known to the alchemists as aqua fortis; the name is used in commerce for impure grades of it. Aqua regiaaqua regia
[Lat.,=royal water], corrosive, fuming yellow liquid prepared by mixing one volume of concentrated nitric acid with three to four volumes of concentrated hydrochloric acid.
..... Click the link for more information. is a mixture of nitric and hydrochloric acids. Niric acid is a component of acid rainacid rain
or acid deposition,
form of precipitation (rain, snow, sleet, or hail) containing high levels of sulfuric or nitric acids (pH below 5.5–5.6).
..... Click the link for more information. .
(HNO3) a strong, monobasic acid that is a colorless liquid under ordinary conditions. It is one of the most important products of the chemical industry. Its structural formula is
Physical and chemical properties Anhydrous nitric acid has a density of 1,522 kg/m3, a melting point of -41.15°C, and a boiling point of 84°C. It is miscible with water in all proportions and forms an azeotropic mixture containing 69.2 percent nitric acid and having a boiling point of 121.8°C. The crystal hydrates HNO3·H2O, with a melting point of −37.85°C, and HNO3·3 H2O, with a melting point of −18.5°C, are known. In the absence of water, nitric acid is unstable and decomposes in the light even at ordinary temperatures, liberating oxygen (4HNO3 = 4NO2 + 2H2O + O2); the liberated nitrogen dioxide colors the nitric acid yellow and, at high N02 concentrations, red.
Nitric acid is a strong oxidizing agent. It oxidizes sulfur to sulfuric acid and phosphorus to phosphoric acid. Only gold, tantalum, and some platinum metals do not react with nitric acid. With most metals, nitric acid interacts primarily with the liberation of oxides of nitrogen: 3Cu + 8HN03 = 3Cu(N03) 2+ 2NO + 4H2O. Some metals—for example, iron, chromium, and aluminum—readily dissolve in diluted nitric acid but are resistant to the action of concentrated acid; this is explained by the formation of a protective oxide layer on the metal’s surface. This feature makes it possible to store and ship concentrated nitric acid in steel containers. A mixture of concentrated nitric acid and hydrochloric acid (1:3), called aqua regia, dissolves even gold and platinum. Organic compounds are oxidized or nitrated by nitric acid; in the latter case, the nitric acid residue N02—the nitro group—replaces hydrogen in the organic compounds. The salts of nitric acid are called nitrates; the salts with K, Na, Ga, and NH4 are also called saltpeters.
Preparation and use In the 13th century the preparation of nitric acid by heating potassium nitrate with alum, iron vitriol, and clay was described. In the middle of the 17th century J. R. Glauber proposed a method of preparing nitric acid by moderately heating (up to 150°C) potassium nitrate with concentrated sulfuric acid: KN03 + H2S04 = HN03 + KHSO4. Until the beginning of the 20th century, this method was used in industry, with potassium nitrate replaced by the cheaper, naturally occurring Chile saltpeter, NaN03.
The modern method of nitric-acid production is based on the catalytic oxidation of ammonia by atmospheric oxygen. The main stages of this process are the contact oxidation of ammonia to nitric oxide, 4NH3 + 502 = 4NO + 6H2O; the oxidation of nitric oxide to nitrogen dioxide; and the absorption of the mixture of so-called nitrous gases in water: 2NO + 02 = 2N02, 3NO2 + H2O - 2HNO3 + NO. The air and ammonia (10–12 percent) mixture is passed through a catalyst gauze heated to 750–900°C in which ternary (93 percent Pt, 3 percent Rh, 4 percent Pd) or binary (90–95 percent Pt, 5–10 percent Rh) platinum alloys serve as the catalyst. A two-stage catalyst, in which the first stage is a platinoid gauze and the second stage is a nonplatinum catalyst, are also used; this permits a 25–30 percent reduction in platinum consumption. The time of contact between the air-ammonia mixture and the catalyst must not exceed 1 msec; otherwise, the nitric oxide that has formed decomposes. The second stage of the process—the oxidation of NO to N02 and the dissolving of N02 in water—can be carried out at atmospheric pressure, with pressures of up to 1 meganewton/m2 (10 kg force/cm2) or by a combined method in which only the absorption of the nitrous gases by water occurs under pressure. Nitric acid is obtained at concentrations of 45–49 percent, or 55–58 percent if pressure is used. Through the distillation of such solutions, nitric acid of azeotropic composition can be produced. More concentrated acid (up to 100 percent) is obtained by distilling solutions of nitric acid with concentrated H2S04 or by direct synthesis—the interaction of N2O4 with water (or diluted nitric acid) and oxygen: 2N2O4 + 2H2O + O2 = 4HN03. In the USSR, 97–98 percent nitric acid is produced.
The most important uses of nitric acid are in the production of nitrogenous and combined fertilizers, explosives (trinitrotoluene and others), and organic dyes. In organic synthesis a mixture of concentrated nitric acid and sulfuric acid, called a nitrating mixture, is widely used. Nitric acid is used in the chamber process for producing sulfuric acid, in the production of phosphoric acid from phosphorus, and as an oxidizing agent for rocket fuel. In metallurgy nitric acid is used for pickling and dissolving metals as well as separating gold and silver.
Inhaling the vapors of nitric acid leads to poisoning, and contact of nitric acid (in particular, the concentrated acid) with the skin causes burns. The maximum permissible content of nitric acid in the air in industrial establishments is 50 mg/m3, calculated as N2O5. In contact with organic substances, concentrated nitric acid causes fires and explosions.
REFERENCESAtroshchenko, V. I., and S. I. Kargin. Tekhnologiia azotnoi kisloty. Moscow-Leningrad, 1949.
Miniovich, M. A. “O sovremennom sostoianii i o perspektivakh razvitiia proizvodstva razbavlennoi azotnoi kisloty.” Zhurnalpri-kladnoi khimii, 1958, vol. 31, issue 8.
Miniovich, M. A. “Azotnaia kislota.” Kratkaia khimicheskaia entsiklopediia, vol. 1. Moscow, 1961. Pages 74–79.
E. B. SHILLER