H2 SO4, a strong dibasic acid corresponding to the highest oxidation state of sulfur (+ 6). Under usual conditions, sulfuric acid is a heavy, oily, colorless, and odorless liquid. In industry, mixtures of sulfuric acid both with water and sulfur trioxide are also called sulfuric acid. If the SO3: H2 O molecular ratio is less than 1, the mixture is an aqueous solution of sulfuric acid; if it is more than 1, the mixture is a solution of SO3 in sulfuric acid.
Physical and chemical properties. When existing in a concentration of 100 percent, sulfuric acid (monohydrate, SO3 · H2 O) crystallizes at 10.45°C, boils at 296.2°C, and has a density of 1.9203 g/cm3 and a heat capacity of 1.62joules/g•°K. H2 SO4 is miscible with water and SO3 in all proportions, forming the compounds
H2 SO4·4H2 O (melting point [mp] – 28.36°C)
H2 SO4·3H2 O (mp -36.31 °C)
H2 SO4·2H2 O (mp -39.60°C)
H2 SO4·H2 O (mp 8.48°C)
The compounds H2 SO4·SO3 (H2 S2 O7, disulfuric, or pyrosulfuric acid, mp 35.15°C) and H2 SO4·2SO3 (H2 S3 O10, trisulfuric acid, mp 1.20°C) are also formed. Only water vapor is given off into the vapor phase upon heating and boiling aqueous solutions of sulfuric acid containing up to 70 percent H2 S04. Vapors of sulfuric acid are formed above more concentrated solutions. A solution of 98.3 percent H2 SO4 (azeotropic mixture) upon boiling (336.5°C) is completely distilled. Sulfuric acid containing more than 98.3 percent H2 SO4 releases vapors of SO3 upon heating.
Concentrated sulfuric acid is a strong oxidizing agent. It oxidizes HI and HBr to the free halogens, and upon heating, it will oxidize all metals except Au and the platinum metals (with the exception of Pd). In the cold, concentrated sulfuric acid passi-vates many metals, including Pb, Cr, Ni, steel, and pig iron. Dilute sulfuric acid reacts with all the metals (except Pb) that are above hydrogen in the electromotive force series, for example,
Zn + H2 SO4 = ZnSO4 + H2
As a strong acid, sulfuric acid displaces weaker acids from their salts, for example, boric acid from borax:
Na2 B4 O7 + H2 SO4 + 5H2 O = Na2 SO4 + 4H3 BO3
Upon heating, it displaces more volatile acids, for example,
NaNO3 + H2 SO4 = NaHSO4 + HNO3
Sulfuric acid removes water that is chemically bound to organic compounds containing OH, or hydroxyl, groups. The dehydration of ethyl alcohol in the presence of concentrated sulfuric acid results in the formation of ethylene or diethyl ether. The charring of sugar, cellulose, starch, and other carbohydrates upon contact with sulfuric acid also derives from the dehydration of these substances. As a dibasic acid, sulfuric acid forms two types of salts: sulfates and bisulfates.
Production. The first descriptions of oil of vitriol, that is, concentrated sulfuric acid, were given by the Italian scientist V. Biringuccio in 1540 and the German alchemist whose works were published under the name of Basilius Valentinus in the late 16th and early 17th centuries. By 1690, the French chemists N. Lemery and N. Lefebvre had laid the basis for the first industrial method of obtaining sulfuric acid, a method applied in England in 1740. According to this method, a mixture of sulfur and saltpeter was burned in a ladle suspended in a glass jar containing a certain amount of water. The SO3 generated reacted with the water to form sulfuric acid. In 1746, J. Roebuck in Birmingham replaced the glass jars with chambers made of sheet lead, thus laying the basis for the chamber process for the production of sulfuric acid. Continuous improvement in the process for the production of sulfuric acid in Great Britain and France resulted in 1908 in the first tower system. In the USSR, the first tower installation went into operation in 1926 at the Po-levskoi Metallurgical Plant in the Urals.
The raw material for the production of sulfuric acid can be sulfur, pyrite (FeS2), or exhaust gases containing SO2 from furnaces for the oxidative roasting of the sulfide ores of Cu, Pb, Zn, and other metals. In the USSR, most sulfuric acid is obtained from pyrite. Here, the FeS2 is burned in furnaces in the state of a fluidized bed, a state achieved by blowing a rapid stream of air through a layer of finely ground pyrite. The gaseous mixture obtained contains SO2, O2, N2, impurities of SO3, and vapors of H2 O, As2 O3, and SiO2 and holds considerable cinder dust, which is removed from the gas in electrostatic precipitators.
Sulfuric acid is produced from SO2 by the nitrous (tower) and contact methods. The conversion of SO2 into sulfuric acid by the nitrous method is carried out in cylindrical production towers 15 m and more in height and filled with a packing of ceramic rings. Nitrous vitriol—a mixture of dilute sulfuric acid and ni-trosylsulfuric acid (NOOSO3 H) is sprayed from above into a rising stream of gases. The nitrosylsulfuric acid is produced by the reaction
N2 O3 + 2H2 SO4 = 2NOOSO3 H + H2 O
The oxidation of SO2 by nitrogen oxides occurs in solution after the absorption of SO2 by nitrous vitriol. The mixture is hydro-lyzed by water:
NOOSO3 H + H2 O = H2 SO4 + HNO2
The sulfur dioxide gas entering the tower reacts with water to form sulfurous acid:
SO2 + H2 O = H2 SO3
The reaction of HNO2 with H2 SO3 results in the formation of sulfuric acid:
2HNO2 + H2 SO3 = H2 SO4 + 2NO + H2 O
The NO produced is converted into N2 O3, or, more precisely, a mixture of NO and NO2, in the oxidizing tower. The gases are then introduced into absorption towers, where they encounter sulfuric acid supplied from the top. It is here that nitrous vitriol is obtained; the mixture is then transferred to the production towers. Thus, there is continuous production and a circulation of nitrogen oxides. The inevitable losses of nitrogen oxides with the exhaust gases are balanced by the addition of HNO3.
The sulfuric acid produced by the nitrous method is of insufficient concentration and contains harmful impurities, for example. As. Production is accompanied by the release of nitrogen oxides into the atmosphere (“foxtails,” named for the color of NO2).
The principle of the contact process for the production of sulfuric acid was discovered in 1831 by P. Phillips in Great Britain. The first catalyst was platinum. In the late 19th century and early 20th, it was found that the oxidation of SO2 to SO3 could be accelerated by vanadium pentoxide (V2 O5). The studies of the Soviet scientists A. E. Adadurov, G. K. Boreskov, and F. N. Iushkevich were important in determining the action and selection of vanadium catalysts. Modern sulfuric acid plants are constructed for the operation of the contact process. Vanadium oxides with additives of SiO2, AI2 O3, K2 O, CaO, and BaO in various proportions are used as catalyst bases. All vanadium contact materials are reactive only at temperatures above ~420°C. In the contact apparatus, the gas usually passes through four or five tiers of contact material. In the production of sulfuric acid by the contact process, the gas undergoing oxidation is first subjected to removal of impurities that might poison the catalyst. As, Se, and traces of dust are removed in scrubbing towers in which there is a countercurrent trickling of sulfuric acid. The sulfuric acid mist formed from the SO3 and H2 O present in the gaseous mixture is eliminated in wet electrostatic precipitators. Vapors of H2 O are absorbed by concentrated sulfuric acid in drying towers. The mixture of SO2 and air then passes through the catalyst (contact material) and undergoes oxidation to SO3:
The sulfur trioxide is then absorbed by the water in dilute H2 SO4:
SO3 + H2 O = H2 SO4
Depending on the amount of water introduced into the process, either oleum or a solution of sulfuric acid in water is obtained.
In 1973 the production of sulfuric acid (in the monohydrate) was (in millions of tons): 14.9 in the USSR, 28.7 in the United States, 7.1 in Japan, 5.5 in the Federal Republic of Germany, 4.4 in France, 3.9 in Great Britain, 3.0 in Italy, 2.9 in Poland, 1.2 in Czechoslovakia, 1.1 in the German Democratic Republic, and 0.9 in Yugoslavia.
USE. Sulfuric acid is one of the most important products of the heavy chemical industry. The available grades include chamber acid (not less than 75 percent H2 SO4), oil of vitriol (not less than 92.5 percent), and oleum, or fuming sulfuric acid (a solution of 18.5–20 percent SO3 in H2 SO4), as well as especially pure battery acid (92–94 percent; when diluted by water to 26–31 percent, it serves as the electrolyte in lead batteries). In addition, reagent-grade sulfuric acid (92–94 percent) is produced by the contact process in quartz or platinum apparatus. The strength of sulfuric acid is determined by the density, which is measured with a hydrometer. Most of the chamber acid is used in the production of mineral fertilizers. Sulfuric acid is used in the production of, for example, phosphoric, hydrochloric, boric, and hydrofluoric acids because of its ability to displace these acids from their salts. Concentrated sulfuric acid is used in separating organosulfur compounds and unsaturated organic compounds from petroleum products. Dilute sulfuric acid is used for the removal of scale from wire and sheets before plating with tin or zinc and for the pickling of metal surfaces before plating with chromium, nickel, or copper. It is used in metallurgy for the decomposition of complex ores, in particular, those of uranium. In organic synthesis, concentrated sulfuric acid is a necessary component of nitrating mixtures and a sulfonating agent in the production of many dyes and pharmaceuticals. Owing to its high hygroscopicity, sulfuric acid is used in drying gases and in concentrating nitric acid.
Safety measures. Poisonous gases—SO2 and NO2—as well as vapors of SO3 and H2 SO4, present a danger in the production of sulfuric acid. Proper ventilation and hermetically sealed production apparatus are therefore mandatory. Since sulfuric acid causes serious burns of the skin, handling requires extreme care and protective devices (goggles, rubber gloves, aprons, boots). When diluting sulfuric acid, the acid must be poured into water in a thin stream while stirring. Pouring the water into the acid leads to spattering because of the evolution of a great amount of heat.
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The Great Soviet Encyclopedia, 3rd Edition (1970-1979). © 2010 The Gale Group, Inc. All rights reserved.